Laboratory Manual for General, Organic, and Biological Chemistry [3rd ed.] 9780321811851, 0321811852

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Laboratory Manual for General, Organic, and Biological Chemistry [3rd ed.]
 9780321811851, 0321811852

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Table of contents :
Preface To the Student Using This Laboratory Manual Working Safely in the Laboratory Commitment to Safety in the Laboratory A Visual Guide to Laboratory Equipment Graphing Experimental Data Using the Laboratory Burner Using a Pipet 1 Measurement and Significant Figures A. Measuring Length B. Measuring Volume C. Measuring Mass 2 Conversion Factors and Problem Solving A. Rounding Off B. Significant Figures in Calculations C. Equalities and Conversion Factors D. Problem Solving Using Conversion Factors 3 Density and Specific Gravity A. Density of a Liquid B. Specific Gravity C. Density of a Solid D. Graphing Mass and Volume 4 Temperature and Specific Heat A. Temperature B. Specific Heat of a Metal C. Energy and Nutrition D. Energy Values for Foods 5 Energy and Matter A. A Heating Curve for Water B. Graphing a Cooling Curve for Salol C. Energy in Changes of State: Heat of Fusion 6 Atoms and Elements A. Elements and Symbols B. The Periodic Table C. The Atoms D. Isotopes and Atomic Mass 7 Electronic Configuration and Periodic Properties A. Flame Tests B. Electron Configurations C. Atomic Radius 8 Nuclear Chemistry A. Nuclear Equations B. Radiation Measurement C. Radiation Levels from Radioactive Sources D. Effect of Shielding on Radiation Level E. Effect of Time on Radiation Level F. Effect of Distance on Radiation Level 9 Compounds and Their Bonds A. Ions: Transfer of Electrons B. Ionic Compounds and Formulas C. Metals in Ionic Compounds with Variable Charge D. Polyatomic Ions E. Molecular Compounds F. Electron-Dot Formulas and Molecular Shape 10 Chemical Reactions and Equations A. Magnesium and Oxygen B. Zinc and Copper (II) Sulfate C. Reactions of Metals and HCl D. Reactions of Ionic Compounds E. Sodium Carbonate and HCl F. Hydrogen Perioxide 11 Moles and Chemical Formulas A. Finding the Simplest Formula B. Formula of a Hydrate 114 12 Gas Laws A. Boyle's Law B. Charles's Law 13 Dalton's Law of Partial Pressures A. Partial Pressures of Oxygen and Nitrogen in Air B. Carbon Dioxide in the Atmosphere C. Carbon Dioxide in the Expired Air 14 Solutions, Electrolytes, and Concentration A. Polarity of Solutes and Solvents B. Electrolytes and Nonelectrolytes C. Electrolytes in Body Fluids D. Concentration of a Sodium Chloride Solution 15 Soluble and Insoluble Salts A. Soluble and Insoluble Salts B. Solubility of KNO3 C. Testing the hardness of Water D. Purification of Water 16 Testing for Cations and Anions A. Flames Tests for K+ , Ca2+ , and Na+ Ions B. Tests for Ammonium Ion, NH4+ , and Iron(III) Ion, Fe3+ C. Tests for Negative Ions (Anions) D. Writing the Formula of Your Unknown Salt E. Testing Consumer Products for Some Cations and Anions 17 Properties of Solutions A. Identification Tests B. Osmosis and Dialysis C. Filtration 18 Reaction Rates and Chemical Equilibrium A. Factors That Affect the Rate of a Reaction B. Chemical Equilibrium: Reversible Reactions C. Changing Equilibrium Conditions: Le Ch (R)telier's Principle 19 Acids, Bases, pH and Buffers A. Reference Colors for pH Using Red Cabbage Indicator B. Measuring pH C. Effect of Buffers on pH 20 Acid-Base Titration A. Acetic Acid in Vinegar B. Titration of an Antacid 21 Organic Compounds: Alkanes A. Comparison of Organic and Inorganic Compounds B. Alkanes C. Functional Groups 22 Reactions of Unsaturated Hydrocarbons A. Types of Unsaturated Hydrocarbons B. Addition Reaction: Bromine Test C. Oxidation: Potassium Permanganate (KMnO4 ) Test D. Identification of Unknown 23 Alcohols and Phenols A. Structures of Alcohols and Phenol B. Properties of Alcohols and Phenol C. Oxidation of Alcohols D. Ferric Chloride Test E. Identification of Unknown 24 Aldehydes and Ketones A. Structures of Some Aldehydes and Ketones B. Odor of Aldehydes and Ketones C. Solubility, Iodoform Test, and Benedict's 25 Carboxylic Acids and Esters A. Carboxylic Acids and Their Salts B. Esters C. Saponificiation 26 Aspirin and Other Analgesics A. Preparation of Aspirin B. Testing Aspirin Products C. Analysis of Analgesics 27 Amines and Amides A. Structure, Classification, and Solubility of Amines B. Neutralization of Amines with Acid C. Amides D. Hydrolysis of an Amide 28 Synthesis of Acetaminophen Synthesis of Acetaminophen 29 Types of Carbohydrates A. Monosaccharides B. Disaccharides C. Polysaccharides 30 Tests for Carbohydrates A. Benedict's Test for Reducing Sugars B. Seliwanoff's Test for Ketoses C. Fermentation Test D. Iodine Test for Polysaccharides E. Hydrolysis of Disaccharides and Polysaccharides F. Testing Foods for Carbohydrates 31 Lipids A. Physical Properties of Lipids and Fatty Acids B. Triacylglycerols C. Bromine Test for Unsaturation D. Preparation of Hand Lotion 32 Saponification and Soaps A. Saponification: Preparation of Soap B. Properties of Soap and Detergents 33 Amino Acids A. Amino Acids B. Chromatography of Amino Acids 34 Peptides and Proteins A. Peptides B. Proteins C. Denaturation of Proteins D. Isolation of Casein (Milk Protein) E. Color Tests for Proteins 35 Enzymes A. Effect of Enzyme Concentration B. Effect of Temperature C. Effect of pH D. Inhibition of Enzyme Activity

Citation preview

Representative elements Alkali Alkaline metals earth metals

Period number

1 Group 1A

2 3 4 5 6 7

13 14 15 16 17 Group Group Group Group Group 3A 4A 5A 6A 7A

18 Group 8A

1.008

2 Group 2A

3

4

5

6

7

8

9

10

Li

Be

B

C

N

O

F

Ne

6.941

9.012

10.81

12.01

14.01

16.00

19.00

20.18

11

12

13

14

15

16

17

18

1

1

Halogens Noble gases

H

Transition elements

22.99

24.31

3 3B

19

20

21

39.10

40.08

37

38

Na Mg K

Ca Sc

5 5B

6 6B

7 7B

8

9 8B

10

11 1B

12 2B

Al

Si

P

S

Cl

Ar

26.98

28.09

30.97

32.07

35.45

39.95

22

23

24

25

26

27

28

29

30

31

32

33

34

35

Cr Mn Fe Co Ni Cu Zn Ga Ge As

Br

Kr

74.92

78.96

79.90

83.80

51

52

53

54

Sn Sb

Te

I

Xe 131.3

V

44.96

47.87

50.94

52.00

54.94

55.85

58.93

58.69

63.55

65.41

69.72

72.64

39

40

41

42

43

44

45

46

47

48

49

50

Y

85.47

87.62

88.91

91.22

92.91

95.94

1992

101.1

102.9

106.4

107.9

112.4

114.8

118.7

121.8

127.6

126.9

55

56

57*

72

73

74

75

76

77

78

79

80

81

82

83

84

85

Cs Ba

La

Hf

Ta

W

Re Os

Ir

Pt

Au Hg

Tl

Pb

Bi

132.9

137.3

138.9

178.5

180.9

183.8

186.2

190.2

192.2

195.1

197.0

200.6

204.4

87

88

89†

104

105

106

107

108

109

110

111

112

113

207.2 114

209.0 115

Ra Ac

In

Rf Db Sg Bh Hs Mt Ds Rg Cn —

86

Po At

Rn

12102 117

12222 118

12092 116

Fl



Lv





12662

12642

12652

12662

12712

12722

12852

12842

12892

12882

12932

12932

12942

59

60

61

62

63

64

65

66

67

68

69

70

71

12622

*Lanthanides

Ce

Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

140.1

140.9

144.2

11452

150.4

152.0

157.3

158.9

162.5

164.9

167.3

168.9

173.0

175.0

90

91

92

93

94

95

96

97

Np Pu Am Cm Bk

98

Cf

99

100

101

102

Es Fm Md No Lr

103

232.0

231.0

12372

12442

12432

12472

12472

12512

12522

12572

12582

12592

12622

58

†Actinides

Metals

Ru Rh Pd Ag Cd

12612

12262

12272

Zr Nb Mo Tc

36

Se

Ti

Sr

Fr

4.003

4 4B

Rb

12232

2

He

Th Pa

Metalloids

U 238.0

Nonmetals

LABORATORY MANUAL for General, Organic, and Biological Chemistry THIRD EDITION

TIMBERLAKE

Boston Columbus Indianapolis New York San Francisco Upper Saddle River Amsterdam Cape Town Dubai London Madrid Milan Munich Paris Montréal Toronto Delhi Mexico City São Paulo Sydney Hong Kong Seoul Singapore Taipei Tokyo

Editor in Chief: Adam Jaworski Executive Editor: Jeanne Zalesky Senior Marketing Manager: Jonathan Cottrell Project Editor: Jessica Moro Assistant Editor: Lisa R. Pierce Editorial Assistant: Lisa Tarabokjia Marketing Assistant: Nicola Houston Executive Editorial Media Producer: Deb Perry Managing Editor, Chemistry and Geosciences: Gina M. Cheselka Production Project Manager: Wendy Perez Full Service/Composition: PreMediaGlobal Project Manager, Full Service: Revathi Viswanathan Photo Manager: Maya Melenchuck Photo Permissions Specialist: Eric Schrader Text Permissions Manager: Alison Bruckner Text Permissions Specialist: Danielle Simon, Creative Compliance Operations Specialist: Jeffrey Sargent Cover Designer: Seventeenth Street Studios Cover Photograph: shoot4u/Fotolia Credits and acknowledgments borrowed from other sources and reproduced, with permission, in this textbook appear on the appropriate page within the text. Copyright © 2014, 2011, 2007, 2002 by Pearson Education, Inc. All rights reserved. Manufactured in the United States of America. This publication is protected by Copyright, and permission should be obtained from the publisher prior to any prohibited reproduction, storage in a retrieval system, or transmission in any form or by any means: electronic, mechanical, photocopying, recording, or likewise. To obtain permission(s) to use material from this work, please submit a written request to Pearson Education, Inc., Permissions Department, 1 Lake Street, Department 1G, Upper Saddle River, NJ 07458. Many of the designations used by manufacturers and sellers to distinguish their products are claimed as trademarks. Where those designations appear in this book, and the publisher was aware of a trademark claim, the designations have been printed in initial caps or all caps. The author and the publisher of this book have used their best efforts in preparing this book. These efforts include the development, research, and testing of the theories and programs to determine their effectiveness. The author and publisher make no warranty of any kind, expressed or implied, with regard to these programs or the documentation contained in this book. The author and publisher shall not be liable in any event for incidental or consequential damages in connection with, or arising out of, the furnishing, performance, or use of these programs.

ISBN-10: 0-321-81185-2; ISBN-13: 978-0-321-81185-1 www.pearsonhighered.com 1 2 3 4 5 6 7 8 9 10—VHO—16 15 14 13 12

Contents Preface New to This Edition

ix

To the Student

x

Using This Laboratory Manual

xi

Working Safely in the Laboratory

xii

Commitment to Safety in the Laboratory

xvii

Laboratory Equipment

xviii

Graphing Experimental Data

xxi

Using the Laboratory Burner

xxiii

Using a Pipet 1

2

3

4

5

viii

xxv

Chemistry and Measurement

1

A. Measuring Length

6

B. Measuring Volume

6

C. Measuring Mass

6

Conversion Factors and Problem Solving

11

A. Rounding Off

15

B. Significant Figures in Calculations

15

C. Equalities and Conversion Factors

16

D. Problem Solving Using Conversion Factors

16

Density and Specific Gravity

23

A. Density of Liquids

25

B. Specific Gravity

25

C. Density of Solids

26

D. Graphing Mass and Volume

27

Temperature and Specific Heat

35

A. Temperature

38

B. Specific Heat

38

C. Energy and Nutrition

39

D. Energy Values for Foods

39

Energy and Matter

47

A. A Heating Curve for Water

49

B. A Cooling Curve

49

C. Heat of Fusion

51

Copyright © 2014 Pearson Education, Inc.

iii

iv

6

7

8

9

10

11

12

Laboratory Manual for General, Organic, and Biological Chemistry

Atoms and Elements

61

A. Physical Properties of Elements

66

B. The Periodic Table

66

C. The Atom

66

D. Isotopes and Atomic Mass

66

Electron Configuration and Periodic Properties

73

A. Flame Tests

76

B. Electron Configurations

77

C. Atomic Radius

77

Nuclear Chemistry

85

A. Nuclear Equations

87

B. Radiation Measurement

87

C. Radiation Levels from Radioactive Sources

87

D. Effect of Shielding on Radiation Level

87

E. Effect of Time on Radiation Level

88

F. Effect of Distance on Radiation Level

88

Compounds and Their Bonds

95

A. Ions: Transfer of Electrons

101

B. Ionic Compounds and Formulas

101

C. Metals in Ionic Compounds with Variable Charge

101

D. Polyatomic Ions

101

E. Molecular Compounds: Sharing Electrons

102

F. Electron-Dot Formulas and Shape

102

Chemical Reactions and Equations

109

A. Magnesium and Oxygen

112

B. Zinc and Copper(II) Sulfate

112

C. Reactions of Metals and HCl

112

D. Reactions of Ionic Compounds

113

E. Sodium Carbonate and HCl

113

F. Hydrogen Peroxide

113

Moles and Chemical Formulas

121

A. Finding the Simplest Formula

123

B. Formula of a Hydrate

124

Gas Laws

133

A. Boyle’s Law

135

B. Charles’s Law

135

Contents

13

14

15

16

17

18

19

20

v

Dalton’s Law of Partial Pressures

143

A. Partial Pressures of Oxygen and Nitrogen in Air

145

B. Carbon Dioxide in the Atmosphere

146

C. Carbon Dioxide in Exhaled Air

148

Solutions, Electrolytes, and Concentration

155

A. Polarity of Solutes and Solvents

158

B. Electrolytes and Nonelectrolytes

158

C. Electrolytes in Body Fluids

158

D. Concentration of a Sodium Chloride Solution

159

Soluble and Insoluble Salts A. Soluble and Insoluble Salts

167 170

B. Solubility of KNO3

170

C. Testing the Hardness of Water

171

D. Purification of Water

171

Testing for Cations and Anions A. Flames Tests for K + , Ca 2+ , and Na + Ions

179 181

B. Tests for Ammonium Ion, NH 4+ , and Iron(III) Ion, Fe3+

182

C. Tests for Negative Ions (Anions)

182

D. Writing the Formula of Your Unknown Salt

183

E. Testing Consumer Products for Some Cations and Anions

184

Properties of Solutions

191

A. Identification Tests

194

B. Osmosis and Dialysis

194

C. Filtration

196

Reaction Rates and Chemical Equilibrium

201

A. Factors That Affect the Rate of a Reaction

204

B. Chemical Equilibrium: Reversible Reactions

205

C. Changing Equilibrium Conditions: Le Châtelier’s Principle

206

Acids, Bases, pH, and Buffers

213

A. Reference Colors for pH Using Red Cabbage Indicator

216

B. Measuring pH

216

C. Effect of Buffers on pH

216

Acid-Base Titration

225

A. Concentration of Acetic Acid in Vinegar

227

B. Titration of an Antacid

229

vi

21

22

23

24

25

26

27

28

Laboratory Manual for General, Organic, and Biological Chemistry

Organic Compounds: Alkanes

235

A. Comparison of Organic and Inorganic Compounds

241

B. Alkanes

241

C. Functional Groups

242

Reactions of Unsaturated Hydrocarbons

251

A. Types of Unsaturated Hydrocarbons

253

B. Addition Reaction: Bromine Test

253

C. Oxidation: Potassium Permanganate Test

253

D. Identification of an Unknown

253

Alcohols and Phenols

259

A. Structures of Alcohols and Phenol

262

B. Properties of Alcohols and Phenol

262

C. Oxidation of Alcohols

262

D. Iron(III) Chloride Test

262

E. Identification of an Unknown

262

Aldehydes and Ketones

269

A. Structures of Some Aldehydes and Ketones

271

B. Properties of Aldehydes and Ketones

271

C. Solubility, Iodoform Test, and Benedict’s Test

271

Carboxylic Acids and Esters

279

A. Carboxylic Acids and Their Salts

282

B. Esters

282

C. Saponificiation

283

Aspirin and Other Analgesics

291

A. Preparation of Aspirin

294

B. Testing Aspirin Products

295

C. Analysis of Analgesics

295

Amines and Amides

303

A. Structure, Classification, and Solubility of Amines

306

B. Neutralization of Amines with Acid

306

C. Amides

306

D. Hydrolysis of an Amide

306

Synthesis of Acetaminophen

315

Synthesis of Acetaminophen

317

Contents

29

30

31

32

33

34

35

vii

Types of Carbohydrates

323

A. Monosaccharides

329

B. Disaccharides

329

C. Polysaccharides

329

Tests for Carbohydrates

337

A. Benedict’s Test for Reducing Sugars

339

B. Seliwanoff’s Test for Ketoses

339

C. Fermentation Test

339

D. Iodine Test for Polysaccharides

340

E. Hydrolysis of Disaccharides and Polysaccharides

340

F. Testing Foods for Carbohydrates

341

Lipids

349

A. Physical Properties of Lipids and Fatty Acids

354

B. Triacylglycerols

354

C. Bromine Test for Unsaturation

354

D. Preparation of Hand Lotion

354

Saponification and Soaps

363

A. Saponification: Preparation of Soap

366

B. Properties of Soaps and Detergents

367

Amino Acids

373

A. Amino Acids

376

B. Chromatography of Amino Acids

376

Peptides and Proteins

383

A. Peptides

386

B. Denaturation of Proteins

386

C. Isolation of Casein (Milk Protein)

386

D. Color Tests for Proteins

387

Enzymes

395

A. Effect of Enzyme Concentration

398

B. Effect of Temperature

398

C. Effect of pH

399

D. Inhibition of Enzyme Activity

399

Preface In the process of writing the lab manual for General, Organic, and Biological Chemistry, I have developed experiments that illustrate each of the chemical principles we discuss in our chemistry classes. I have also taken care to make each experiment workable as well as provide critical-thinking tools. The main organizational concepts of this laboratory manual are as follows: ● Experiments relate to basic concepts of chemistry. Experiments are designed to illustrate the chemical principles we discuss in our classes. ● Experiments are flexible. Each experiment includes a flexible group of sections that can be taught in any order, which allows instructors to select sections that fit into their weekly laboratory schedule. Lab times and comments are given for each. ● Safety. A detailed safety section in this preface includes a safety quiz. The aim here is to highlight the safety and equipment preparation on the first day of lab. In addition, each lab section contains instructions on safe behavior in the laboratory specific to that section. Students are reminded to wear goggles for every lab session. Some experiments are recommended as instructor demonstrations. ● Experiment format provides clear instructions and evaluation. Each lab section begins with a set of laboratory goals, a discussion of the chemical concepts, and examples of calculations. The report sheets begin with pre-lab questions to prepare students for laboratory work. Students obtain data, draw graphs, make calculations, and write conclusions about their results. Each lab contains questions and problems that require the student to discuss the experiment, make additional calculations, and use critical thinking to apply concepts to real life.

I hope that this laboratory manual will help you in your chemistry instruction and that students will find they learn more chemistry by participating in the laboratory experience. Karen C. Timberlake Los Angeles Valley College Valley Glen, CA 91401

Copyright © 2014 Pearson Education, Inc.

viii

New to This Edition ● Each experiment now begins with the Laboratory Goals and Lab Information. ● The Chemical Concepts (formerly called ‘Discussion’) now include a complete discussion of the chemistry topics relevant to the experiment. ● Art has been added to the Chemical Concepts, including figures from the text that give visual representations of the concepts. ● The diagrams of Laboratory Equipment in the Preface were updated and new illustrations were added. ● Section titled Using a Pipet was added to the Preface. ● Layout of the Experimental Procedures section was revised. Changes include: ● ● ●

An overview statement with each procedure to indicate what students will be doing. Procedures rewritten for clarity and consolidated for better flow of laboratory work. Removal of the boxed-type framework and procedures keyed to the Report Sheet.

● All experiments were revised for this edition to create a more efficient lab experience for students. ● A Correlation Guide that matches each experiment lab to the corresponding chapters in all Timberlake texts is placed on the Instructor Resource Center at www.pearsonhighered.com/ chemistry. ● The Appendices, which include materials needed for each lab for 20 students, have been moved to the Instructor Resource Center at www.pearsonhighered.com/chemistry. ● Inclusion of both new and rewritten Pre-lab questions for clarity and stronger assessment of concept comprehension. ● Increased font size for the Pre-lab questions and Report Sheets to improve legibility.

Copyright © 2014 Pearson Education, Inc.

ix

To the Student Here you are in a chemistry laboratory with your laboratory book in front of you. Perhaps you have already been assigned a laboratory drawer, full of glassware and equipment you may never have seen before. Looking around the laboratory, you see bottles of chemical compounds, balances, burners, and other equipment that you are going to use. This may very well be your first experience with experimental procedures. At this point you might have some questions about what is expected of you. This laboratory manual is written with those considerations in mind. The activities in this manual were written specifically to parallel the topics you are learning in your chemistry class. Many of the laboratory activities include materials that will be familiar to you, such as household products, diet drinks, cabbage juice, antacids, and aspirin. In this way, chemical topics are related to the real world and to your own science experience. Some of the labs teach basic skills; others encourage you to extend your scientific curiosity beyond the lab. It is important to realize that the value of the laboratory experience depends on you investing time and effort in it. Only then will you find that the laboratory can be a valuable learning experience and an integral part of the chemistry class. The laboratory gives you an opportunity to go beyond the lectures and words in your textbook and experience the scientific process from which conclusions and theories concerning chemical behavior are drawn. In some experiments, the concepts are correlated with health and biological concepts. Chemistry is not an inanimate science, but one that helps us to understand the behavior of living systems. A section on laboratory safety is included in the preface and followed by a detailed Safety Quiz, which highlights safety and equipment. Each lab also includes reminders of safe behavior and the author specifically recommends certain experiments for instructor demonstrations.

Copyright © 2014 Pearson Education, Inc.

x

Using This Laboratory Manual ● Each experiment begins with a clear set of laboratory goals to give you an overview of the topics you will be studying in that experiment. ● There are Pre-Lab Study Questions on easy-to-remove pages at the beginning of each laboratory report section. These questions should be completed before you come to lab. Your laboratory instructor may require that you hand in the Pre-Lab Study Questions before you begin your laboratory work. ● Each experiment is correlated to concepts you are currently learning in your chemistry class. Your instructor will indicate which activities you are to do. You will find a list of the materials needed at the beginning of each of the experiments. ● When you are ready to begin a lab, remove the Report Sheets and place them next to the Laboratory Procedures for that experiment. ● The Experimental Procedures give instructions for you to complete each experiment successfully. The steps are numbered to guide you as you proceed to measure carefully, report your data, and complete calculations. ● Questions throughout the report sheet are designed to test your understanding of the chemical concepts from the experiment.

It is my hope that the laboratory experience will help illustrate the concepts you are learning in your chemistry class. The experimental process can help make chemistry a real and exciting part of your life and provide you with skills necessary for your future.

Copyright © 2014 Pearson Education, Inc.

xi

Working Safely in the Laboratory The chemistry laboratory with its equipment, glassware, and chemicals is a place where accidents can occur. Precautions must be taken by every student to ensure the safety of everyone working in the laboratory. By following the rules for handling chemicals safely and carrying out only the approved procedures, you will help to create a safe environment in the laboratory. After you have read the following sections, complete the safety quiz and the questions on laboratory equipment. Then sign and submit the lab safety commitment to your instructor.

A. Preparing for Laboratory Work Pre-read Before you come to the laboratory, read the discussion of and directions for the experiment you will be doing that day. Make sure you know what the experiment is about before you start the actual work. If you have a question, ask your instructor to clarify the procedures. Do assigned work only Do only the experiments that have been assigned by your instructor. No unauthorized experiments are to be carried out in the laboratory. Experiments are done at assigned times, unless you have an open lab situation. Your instructor must approve any change in procedure.

Do not work alone in a laboratory. Safety awareness Learn the location and use of the emergency eyewash fountains, the emergency shower, fire blanket, fire extinguishers, and exits. Memorize their locations in the laboratory. Be aware of other students in the lab carrying chemicals to their desk or to a balance. Safety goggles must be worn all the time when you are in the lab In the Experimental Procedures for each lab, you will see the following reminder to wear your safety goggles:

EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

The particular type of goggles you should wear depends on state law, which usually requires industrialquality eye protection. Contact lenses may be worn in the lab if needed for therapeutic reasons, provided that safety goggles are worn over the contact lenses. Contact lenses without goggles are dangerous because splashed chemicals make them difficult to remove. If chemicals accumulate under a lens, permanent eye damage can result. If a chemical should splash into your eyes, flood the eyes with water at the eyewash fountain. Continue to rinse with water for at least 10 minutes. Wear protective clothing Wear sensible clothing in the laboratory. Loose sleeves, shorts, or open-toed shoes can be dangerous. A lab coat is useful in protecting clothes and covering arms. Wear shoes that cover your feet to prevent glass cuts; wear long pants and long-sleeved shirts to protect skin. Long hair must be tied back so it does not fall into chemicals or a flame from a Bunsen burner. No food or drink is allowed at any time in the laboratory Do not let your friends or children visit while you are working in the lab; have them wait outside. Prepare your work area Before you begin a lab, clear the lab bench or work area of all your personal items, such as backpacks, books, sweaters, and coats. Find a storage place in the lab for them. All you will need is your laboratory manual, a calculator, a pen or pencil, the text, and equipment from your lab drawer.

Copyright © 2014 Pearson Education, Inc.

xii

Working Safely in the Laboratory

xiii

B. Handling Chemicals Safely Check labels twice Be sure you take the correct chemical. DOUBLE-CHECK THE LABEL on the bottle before you remove a chemical from its container because some chemical names are similar. For example, sodium sulfate (Na 2SO 4 ) could be mistaken for sodium sulfite (Na 2SO3 ) if the label is not read carefully. Use small amounts of chemicals Pour or transfer a chemical into a small, clean container (beaker, test tube, flask, etc.) available in your lab drawer. To avoid contamination of the chemical reagents, never insert droppers, pipets, or spatulas into the reagent bottles. Take only the quantity of chemical you need for the experiment. Do not keep a reagent bottle at your desk; return it to its proper location in the laboratory. Immediately label the container that you transfer the chemical to. Many containers have etched sections on which you can write in pencil. If not, use tape or a marking pencil. Do not return chemicals to the original containers To avoid contamination of chemical reagent bottles, dispose of used chemicals according to your instructor’s instructions. Never return unused chemicals to reagent bottles. Some liquids and water-soluble compounds may be washed down the sink with plenty of water, but check with your instructor first. Dispose of organic compounds in specially marked containers in the hoods. Do not taste chemicals; smell a chemical cautiously Never use any equipment such as a beaker to drink from. When required to note the odor of a chemical, first take a deep breath of fresh air while you use your hand to fan some vapors toward your nose and note the odor. Do not inhale the fumes directly. If a compound gives off an irritating vapor, use it in the fume hood to avoid exposure to it. Do not shake laboratory thermometers Laboratory thermometers respond quickly to the temperature of their environment. Shaking a thermometer is unnecessary and can cause breakage. Liquid spills Clean up spills of water or liquids at your work area or floor immediately. Small spills of liquid chemicals can be cleaned up with a paper towel. Large chemical spills must be treated with absorbent material such as cat litter. Place the contaminated material in the proper waste container and label it. If a liquid chemical is spilled on the skin, flood immediately with water for at least 10 minutes. Any clothing soaked with a chemical must be removed immediately, because it can continue to damage the skin. Mercury spills The cleanup of mercury requires special attention. Mercury spills may occur from broken thermometers. Notify your instructor immediately of any mercury spills so that special methods can be used to clean up the mercury. Place any free mercury and mercury cleanup material in special containers for mercury only. Laboratory accidents Always notify your instructor of any chemical spill or accident in the laboratory. Broken glass can be swept up with a brush and dust pan and placed in a specially labeled container for broken glass. Cuts are the most common injuries in a lab. If a cut should occur, wash, elevate, and apply pressure if necessary. Clean up Wash glassware as you work. Begin your final cleanup 15 min before the end of the laboratory session. Return any borrowed equipment to the stockroom. Be sure that you always turn off the gas and water at your work area. Make sure you leave a clean desk. Wash your hands before you leave the laboratory.

C. Heating Chemicals Safely Heat only heat-resistant glassware Only glassware marked Pyrex or Kimax can be heated; other glassware may shatter at high temperatures. To heat a substance in a test tube, use a test tube holder. Holding the test tube at an angle, move it continuously through the flame. Never point the open end of the test tube at anyone or look directly into it. A hot piece of iron or glass looks the same as it does at room temperature. Place a hot object on a tile or a wire screen to cool. Flammable liquids Never heat a flammable liquid over an open flame. If heating is necessary, your instructor will indicate the use of a steam bath or a hot plate.

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Never heat a closed container When a closed system is heated, it can explode as pressure inside builds. Fire Small fires can be extinguished by covering them with a watch glass. If a larger fire is involved, use a fire extinguisher to douse the flames. Do not direct a fire extinguisher at other people in the laboratory. Shut off gas burners in the laboratory. When working in a lab, tie long hair back away from the face. If another student’s clothing or hair catches on fire, get the student to the floor and roll him or her into a fire blanket. You can also place the student under a nearby safety shower to extinguish flames. Cold water or ice may be applied to small burns.

D. Waste Disposal Chemical wastes are produced when you work in the laboratory. Although you will use small quantities of materials, some waste products are unavoidable. To dispose of these waste products safely, you need to know some general rules for chemical waste disposal. Metals Metals should be placed in a designated container to be recycled. Nonhazardous chemical wastes Substances such as sodium chloride (NaCl) that are soluble in water and are not hazardous may be emptied into the sink and rinsed down the drain with running water. If the waste is a solid, dissolve it in water before disposal or simply throw into the regular trash. Hazardous chemical wastes If a substance is hazardous or not soluble in water, it must be placed in a container that is labeled for waste disposal. Your instructor will inform you if chemical wastes are hazardous and identify the proper waste containers. If you are not sure about the proper disposal of a substance, ask your instructor. The labels on a waste container should indicate if the contents are hazardous, the name of the chemical waste, and the date that the container was placed in the lab. Hazard rating The general hazards of a chemical are presented in a spatial arrangement of numbers with the flammability rating at the twelve o’clock position, the reactivity rating at three o’clock, and the health rating at nine o’clock. At the six o’clock position, information may be given on the reactivity of the substance with water. If there is unusual reactivity with water, the symbol W (do not mix with water) is shown. In the laboratory, you may see these ratings in color with blue for health hazard, red for flammability, and yellow for reactivity hazards.

A chemical is assigned a relative hazard rating that ranges from 1 (little hazard) to 4 (extreme hazard). The health hazard indicates the likelihood that a material will cause injury due to exposure by contact, inhalation, or ingestion. The flammability hazard indicates the potential for burning. The reactivity hazard indicates the instability of the material by itself or with water with subsequent release of energy. Special hazards may be included such as W for reactivity with water or OX for oxidizing properties.

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E. Safety Quiz This quiz will test you on the preceding safety discussion. Circle the correct answer(s) in each of the following questions and check your answers. 1.

Approved eye protection is to be worn a. for certain experiments b. only for hazardous experiments c. all the time

2.

Eating in the laboratory is a. not permitted b. allowed at lunch time c. all right if you are careful

3.

If you need to smell a chemical, you should a. inhale deeply over the test tube b. take a breath of air and fan the vapors toward you c. put some of the chemical in your hand, and smell it

4.

When heating liquids in a test tube, you should a. move the tube back and forth through the flame b. look directly into the open end of the test tube to see what is happening c. direct the open end of the tube away from other students

5.

Unauthorized experiments are a. all right as long as they don’t seem hazardous b. all right as long as no one finds out c. not allowed

6.

If a chemical is spilled on your skin, you should a. wait to see if it stings b. flood the area with water for 10 minutes c. add another chemical to absorb it

7.

When taking liquids from a reagent bottle, a. insert a dropper b. pour the reagent into a small container c. put back what you don’t use

8.

In the laboratory, open-toed shoes and shorts are a. okay if the weather is hot b. all right if you wear a lab apron c. dangerous and should not be worn

9.

When is it all right to taste a chemical used in the lab? a. never b. when the chemical is not hazardous c. when you use a clean beaker

10.

After you use a reagent bottle, a. keep it at your desk in case you need more b. return it to its proper location c. play a joke on your friends and hide it

11.

Before starting an experiment, a. read the entire procedure b. ask your lab partner how to do the experiment c. skip to the laboratory report and try to figure out what to do

12.

Working alone in the laboratory without supervision is a. all right if the experiment is not too hazardous b. not allowed c. allowed if you are sure you can complete the experiment without help

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13.

You should wash your hands a. only if they are dirty b. before eating lunch in the lab c. before you leave the lab

14.

Personal items (books, clothes, etc.) should be a. kept on your lab bench b. left outside c. stored out of the way, not on the lab bench

15.

When you have taken too much of a chemical, you should a. return the excess to the reagent bottle b. store it in your lab locker for future use c. discard it using proper disposal procedures

16.

In the lab, you should wear a. practical, protective clothing b. something fashionable c. shorts and loose-sleeved shirts

17.

If a chemical is spilled on the table, a. clean it up right away b. let the stockroom help clean it up c. use appropriate adsorbent if necessary

18.

If mercury is spilled, a. pick it up with a dropper b. call your instructor c. push it under the table where no one can see it

19.

If a student’s hair or shirt catches on fire, a. use the safety shower to extinguish the flames b. get the student to the floor and roll c. roll the student in a fire blanket

20.

Hazardous waste should be a. placed in a special waste container b. washed down the drain c. placed in the wastebasket

Answer Key to Safety Quiz Question

Answer

Question

Answer

Question

Answer

1 2 3 4 5 6 7

C A B A, C C B B

8 9 10 11 12 13 14

C A B A B C C

15 16 17 18 19 20

C A A, C B A, B, C A

Commitment to Safety in the Laboratory I have read the laboratory preparation and safety procedures. I agree to comply with the safety rules by carrying out the following procedures. I will _____ Read laboratory instructions ahead of lab time. _____ Know the locations of eyewash fountains, fire extinguishers, safety showers, and exits. _____ Wear safety goggles or safety glasses in the laboratory at all times. _____ Use the proper equipment for laboratory procedures. _____ Never perform any unauthorized experiments or work alone in the laboratory. _____ Remember what is still hot if I have used the Bunsen burner or hot plate. _____ Immediately clean up chemical spills and broken glass. _____ Immediately inform the instructor of a chemical spill or accident in the laboratory. _____ Never eat or drink in the laboratory. _____ Wear sensible clothing and closed shoes, and if my hair is long, I will tie it back in the laboratory. _____ Read the labels on reagent bottles carefully, remove only small amounts of reagent with the proper tools, and never return unused chemical to the bottle or insert droppers or any other equipment into the bottle. _____ Dispose of broken glass and waste chemicals in the appropriate waste container. _____ Never leave an experiment when substances are heating or reacting. _____ Wash my hands and leave a clean work area when I leave the lab. ____________________________________________________________ Signature ____________________________________________________________ Printed name ___________________________________ _____________________

Laboratory class and section

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Laboratory Equipment

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Laboratory Equipment

Laboratory Equipment

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A Quiz on Laboratory Equipment Look for each of the items shown in the pictures of laboratory equipment on the previous pages. This quiz will help you learn some of the common laboratory equipment.

Q1.

Match each of the above pieces of equipment with its name: a._____ Bunsen burner b._____ test tube holder c._____ beaker d._____ crucible e._____ funnel

Q2.

Match each of the above pieces of equipment with its use in the laboratory: a. _____ used to hold liquids and to carry out reactions b. _____ used to pick up a crucible c. _____ used to support a beaker on an iron ring during heating d. _____ used to transfer small amounts of a solid substance e. _____ used to measure the volume of a liquid

Answers: Q1 Q2

a. a.

3 4

b. 2 b. 3

c. c.

5 2

d. 4 d. 5

e. e.

1 1

Graphing Experimental Data When you have determined a group of experimental quantities, you can prepare a graph that gives a pictorial representation of the data. First, you will prepare a data table and then you will follow a series of steps to construct a graph.

Preparing a Data Table Prepare a data table from measurements. Suppose we measured the distance traveled by a bicycle rider in a given time. Table 1 is a data table prepared by listing the two variables—time and distance— that we measured. TABLE 1 Time and Distance Measurement Time (h)

Distance (km)

1

5

3

14

4

20

6

30

7

33

8

40

9

46

10

50

Constructing the Graph Draw vertical and horizontal axes Draw vertical and horizontal axes on graph paper. The lines should be set in to leave a margin for numbers and labels, but the graph should cover most of the graph paper. Place a title at the top of the graph. The title should describe the quantities that will be placed on the axes. Label each axis The label for each axis reflects the measurement listed in the data table. On our sample graph, the labels are time (hr) for the horizontal axis and distance (km) for the vertical axis. Apply constant scales On each axis apply a scale of equal intervals that includes the full range of data points (low to high) you have in the data table. The intervals on a scale must be equally spaced and fit on the line you have drawn. Do not exceed the graph lines. Use intervals on each axis that are convenient counting units (2, 4, 6, 8, etc. or 5, 10, 15, etc.). The interval size on one axis does not need to match the size of the intervals on the other axis.

For our sample graph, we used a scale for a distance range of 0 km to 50 km. Each graph division represents 5 km. (You only have to number a few lines in order to interpret the scale. It gets too crowded with numbers if every line is marked.) Every two divisions on the time scale represent a time interval of 1 h within the time span (10 h) of the bicycle ride (see Figure 1).

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▲ FIGURE 1 Marking equal intervals for distance and time on the axes Plot the data points Plot the points for each pair of measurements on the data table. Follow a measurement on the vertical axis across until it meets a line that would be drawn from the corresponding measured value on the horizontal axis.

For example, at 4 hours, the rider has traveled 20 km. On the graph, find 20 km on the distance scale, and 4 h on the time scale. Then follow the perpendicular lines to where they intersect. That is a point on the graph. Plotting each data pair will show the relationship between distance and time. Draw a smooth line or curve that best fits the data points. Some points may not fit exactly on the line or curve you draw. That occurs when error is associated with the measurements or when the data are affected by other variables, such as terrain and energy level of the bicycle rider (see Figure 2).

▲ FIGURE 2 A completed graph with data points connected in a smooth line

Using the Laboratory Burner EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

Materials: Bunsen burner, striker or matches

In the laboratory, substances are often heated with a Bunsen burner, shown below. The burner consists of a metal tube and base connected to a gas source. The flow of gas is controlled by adjusting the gas lever at the bench or by turning the wheel at the base of the burner. The amount of air that enters the burner is adjusted by twisting the tube to open or close the air vents. The gas and air mixture is ignited at the top of the tube using a match or a striker. Make sure the gas valves are tightly closed after using the Bunsen burner.

1. Before you light the burner, practice the following: a. Open and close the gas lever at the lab bench. b. Open and close the gas needle valve (wheel) at the base of the burner. c. Open and close the air vents. With the air vents closed, ignite the burner with a striker or match. Your instructor may demonstrate the use of the striker. Turn on the gas and hold the flame or spark at the top rim of the burner tube. 2. If the flame is yellow and sooty, the gas mixture does not have an adequate supply of oxygen. Open the air vents until the color of the flame changes to blue. Adjust the gas flow until you have a flame that is 6–8 cm high with two distinct parts, an inner cone and an outer flame. 3. The hottest part of a flame is at the tip of the inner blue flame. For the most effective heating, be sure that the tip of the inner flame is placed just under the substance you heat. Remember what you heated: Hot metal and glass items do not look hot!

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Questions Q3

What is the color of the flame with the air vent closed? ________________________________ Open? _______________________________________________________________________

Q4

How do you control the height of the flame?_________________________________________ ____________________________________________________________________________ ____________________________________________________________________________

Q5

Where is the hottest part of a flame?_______________________________________________

Using a Pipet EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Using a Standard Pipet Bulb Caution: Squeeze the pipet over a sink to ensure no foreign liquid is in the bulb 1. Hold the pipet bulb over a sink and squeeze to expel any liquid that it may contain. Squeeze the bulb about halfway closed. Place the pipet bulb on the upper end of the pipet. 2. Place the tapered end of the pipet into the liquid. Release the pressure on the bulb while maintaining a tight seal with the top of the pipet. As the bulb inflates, liquid will move into the pipet. The level of liquid should rise above the volume mark, but should not go into the bulb. 3. Remove the bulb and quickly place your index finger over the mouth of the pipet. Adjust the pressure of your finger to slowly drain the liquid until it is level with the volume mark. Touch the tip to the inside of the container to remove the last drop of liquid hanging from the pipet tip. 4. With your finger still over the mouth of the pipet, take the pipet out of the liquid and move it over the evaporating dish. Lift your finger off the pipet to let the liquid flow out, again touching the tip of the pipet to the inside of the container. Some liquid should remain in the tip. Do not shake or try to blow out the liquid that remains.

▲ FIGURE 3 Using a pipet

B. Using a Three-Way Safety Pipet Bulb Caution: Squeeze the pipet over a sink to ensure no foreign liquid is in the bulb

1.

2.

3. 4.

Hold the pipet bulb over a sink and squeeze to expel any liquid that it may contain. Pinch the upper valve on the large bulb and squeeze the large bulb until the air is expelled. Now, place the three-way safety pipet bulb over the mouth of the pipet Place the end of the pipet in the liquid. Press on the lower valve of the large bulb valve until level of liquid is above the volume mark, but not into the bulb. Squeeze the valve on the small bulb until the bottom of the meniscus of the liquid is level with the volume mark. Touch the tip to the inside of the container to remove any liquid that may have formed a drop from the tip. Remove the tip of the pipet from the liquid and move the bulb and pipet over the second container. Squeeze the valve on the small bulb until all the liquid flows into the evaporating dish touching the tip of the pipet to the container to dislodge the hanging drop. Some liquid should remain in the tip of the pipet: Do not shake or blow out the liquid that remains.

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Chemistry and Measurement 1 LABORATORY GOALS • Identify metric units used in measurement, such as gram, meter, centimeter, and milliliter. • Obtain a correct measurement using a meter stick, a balance, and a graduated cylinder. • State the correct number of significant figures in a measurement.

LAB INFORMATION Time: Comments:

Related Topics:

2h Tear out the report sheets and place them beside the experimental procedures as you work. Determine the markings on each measuring tool before you measure. Record all the numbers for a measurement, including the estimated digit. Write a unit of measurement after each measured number. Significant figures, measured and exact numbers, metric prefixes.

CHEMICAL CONCEPTS Scientists and allied health personnel carry out laboratory procedures, take measurements, and report results accurately. The system of measurement used in science, hospitals, and clinics is the metric system. The metric system is a decimal system in which measurements of each type are related by factors of 10. You use a decimal system when you change U.S. money. For example, 1 dime is the same as 10 cents or one cent is 1/10 of a dime. A dime and a cent are related by a factor of 10.

Metric System The metric system has one standard unit for each type of measurement. For example, the standard metric unit of length is the meter, whereas the U.S. system of measurement uses many units of length such as inch, foot, yard, and mile. Most of the rest of the world uses the metric and the updated SI (International System of Units) systems only. The most common units are listed in Table 1.1. TABLE 1.1 Common Metric and SI Units of Measurement Measurement

Metric (Symbol)

SI (Symbol)

Length

meter (m)

meter (m)

Volume

liter (L)

cubic meter (m3 )

Mass

gram (g)

kilogram (kg)

Temperature

degrees Celsius (°C)

kelvins (K)

Time

second (s)

second (s)

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A unit is always included when reporting a measurement. For example, 5.0 m indicates a quantity and unit in this measurement of length. Without the unit, we would not know the units used to obtain the number 5.0. It could be 5.0 ft, 5.0 km, or 5.0 in. Thus, a unit is required to complete the measurement reported.

Prefixes For larger and smaller measurements, prefixes are attached to the standard unit. Some prefixes such as kilo are used for larger quantities; other prefixes such as milli are used for smaller quantities. The most common prefixes are listed in Table 1.2. TABLE 1.2 Common Prefixes in the Metric System Prefix

Symbol

Value

kilo

k

1000

deci

d

0.1 (1/10)

centi

c

0.01 (1/100)

milli

m

0.001 (1/1000)

A. Measuring Length The standard unit of length in the metric system is the meter (m). Using an appropriate prefix, you can indicate a length that is greater or less than a meter (see Table 1.3). Kilometers are used in most countries for measuring the distance between cities, whereas centimeters or millimeters are used for small lengths. TABLE 1.3 Some Metric Units Used to Measure Length Length

Value

1 km

1000 m or 103 m

1 dm

0.1 m (1/10 m) or 10−1 m

1 cm

0.01 m (1/100 m) or 10−2 m

1 mm

0.001 m (1/1000 m) or 10−3 m

A meterstick is divided into 100 cm, as seen in Figure 1.1. The smallest lines indicate centimeters. That means that each measurement you make can be certain to the centimeter. The final digit in a measurement is obtained by estimating between the smallest marked lines. For example, the shorter line reaches the 44-cm mark and is about halfway to 45 cm. We might report its length as 44.5 cm. The last digit (0.5) is the estimated digit. If the line appears to end at a centimeter mark, then the estimated digit is 0.0 cm. The longer line in Figure 1.1 appears to end right at the 67-cm line, which is indicated by reporting its length as 67.0 cm (see Sample Problem 1.3).

▲ FIGURE 1.1 A meterstick divided into centimeters (cm)

Chemistry and Measurement

3

SAMPLE PROBLEM 1.1 What is the estimated digit in each of the following measured masses? a. beaker 42.18 g

b. pencil 11.6 g

SOLUTION: a. hundredths place (0.08 g)

b. tenths place (0.6 g)

B. Measuring Volume The volume of a substance measures the space it occupies. In the metric system, the unit for volume is the liter (L). Prefixes are used to express smaller volumes such as deciliters (dL), centiliters (cL), or milliliters (mL). One cubic centimeter (cm3 or cc) is equal to 1 mL. The terms mL and cc are used interchangeably (see Table 1.4). TABLE 1.4 Some Metric Units Used to Measure Volume Unit of Volume

Value

1 dL

0.1 L (1/10 L) or 10−1 L

1 cL

0.01 L (1/100) or 10−2 L

1 mL

0.001 L (1/1000 L) or 10−3 L

In the laboratory, the volume of a liquid can be measured in a graduated cylinder (see Figure 1.2). Set the cylinder on a level surface and bring your eyes even with the liquid level. Notice that the water level is not a straight line but curves downward in the center. This curve, called a meniscus, is read at its lowest point (center) to obtain the correct volume measurement for the liquid. Your eyes should be aligned with the bottom of the meniscus in order to avoid error in making the reading. In this graduated cylinder, the volume of the liquid is 42.0 mL.

▲ FIGURE 1.2 Reading a volume of 42.0 mL in a graduated cylinder On large cylinders, the lines may represent volumes of 2 mL, 5 mL, or 10 mL. On a 250-mL cylinder, the marked lines usually represent 5 mL. On a 1000-mL cylinder, each line may be 10 mL. Then your precision on a measurement will be to the milliliter or mL.

Volume of a Solid by Displacement When an object is submerged in water, it displaces its own volume of water, causing the water level to rise. The volume of the object is the difference in the water level before and after the object is submerged (see Figure 1.3).

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▲ FIGURE 1.3 Using volume displacement to determine the volume of a solid

C. Measuring Mass The mass of an object indicates the amount of matter present in that object. In the metric system, the unit of mass is the gram (g). A larger unit, the kilogram (kg), is used for larger objects, for example, in measuring a patient’s weight in a hospital. A smaller unit of mass, the milligram (mg), is often used in the laboratory (see Table 1.5). TABLE 1.5 Some Metric Units Used to Measure Mass Mass

Value

1 kg

1000 g or 103 g

1 mg

1/1000 g (0.001 g) or 10−3 g

Measured and Exact Numbers When we measure the length, volume, or mass of an object, the numbers we report are called measured numbers. Suppose you stepped on a scale this morning and saw that you weighed 145 lb. The scale is a measuring tool and your weight is a measured number. Each time we use a measuring tool to determine a quantity, the result is a measured number. Exact numbers are obtained when we count objects. Suppose you counted 5 beakers in your laboratory drawer. The number 5 is an exact number. You did not use a measuring tool to obtain the number. Exact numbers are also found in the numbers that define a relationship between two metric units or between two U.S. system units. For example, the numbers in the following definitions are exact: 1 meter is equal to 100 cm; 1 ft is equal to 12 in. (see Sample Problem 1.1).

SAMPLE PROBLEM 1.2 Describe each of the following as a measured or exact number: a. 14 in.

b. 14 pencils

c. 60 min in 1 h

d. 7.5 kg

SOLUTION: a. measured

b. exact (counted)

c. exact (definition)

d. measured

Chemistry and Measurement

5

Significant Figures In measured numbers, all the reported figures are called significant figures. The first significant figure is the first nonzero digit. The last significant figure is always the estimated digit. Zeros between other digits or to the right of the decimal point in a decimal number are counted as significant figures. However, zeros to the left of nonzero numbers are not significant; they are placeholders. Zeros are not significant in large numbers with no decimal points; they are placeholders needed to express the magnitude of the number. When a number is written in scientific notation, all the figures in the coefficient are significant. Examples of counting significant figures in measured numbers are in Table 1.6 and Sample Problem 1.3. TABLE 1.6 Examples of Counting Significant Figures Measurement

Number of Significant Figures

Reason

455.2 cm

4

All nonzero digits are significant.

0.80 m

2

A zero following a decimal number is significant.

50.2 L

3

A zero between nonzero digits is significant.

0.0005 lb

1

Zeros to the left of nonzero numbers are not significant.

25 000 ft

2

Placeholder zeros are not significant.

SAMPLE PROBLEM 1.3 State the number of significant figures in each of the following measured numbers: a. 0.00580

b. 132.08 g

SOLUTION: a. Three significant figures. The zeros after the decimal point are placeholder zeros, but the zero following nonzero digits is significant. b. Five significant figures. The zero between nonzero digits is significant.

When you use a meterstick, or read the volume in a graduated cylinder, the measurement must be reported as precisely as possible. The number of significant figures you can report depends on the lines marked on the measuring tool you use. For example, on a 50-mL graduated cylinder, the small lines represent a 1-mL volume. If the liquid level is between 21 mL and 22 mL, you know you can report 21 mL for certain. However, you can add one more digit (the last digit) to your reported value by estimating between the 1-mL lines. For example, if the volume level was halfway between the 21-mL and 22-mL lines, you would report the volume as 21.5 mL. If the volume level is exactly on the 21-mL line, you indicate this precision by adding a significant zero to give a measured volume of 21.0 mL.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Measuring Length Materials: Meterstick, string 1. 2.

3. 4. 5.

Observe the marked lines on a meterstick. Identify the smallest lines on the meterstick. Use the meterstick to make the length measurements (cm) indicated on the report sheet. String may be used to determine the distance around your wrist. Include the estimated digit in each measurement. Indicate the estimated digit in each measurement. Indicate the number of significant figures in each measurement. Measure the length of the line drawn on the report sheet including the estimated value. Compare your value with those of two other students.

B. Measuring Volume Materials: Display of graduated cylinders with liquids, 10-mL, 50-mL, and 100-mL (or larger) graduated cylinders, solid object, thread Volume of a liquid 1.

Read the volume of each of the liquids in the display of graduated cylinders. Be as precise as you can. Be sure to estimate between the smallest markings to obtain the estimated digit.

Volume of a solid by displacement 2. 3. 4.

Obtain a graduated cylinder that will fit the solid object. Place water in the graduated cylinder until it is about half full. Record the volume, in milliliters, of water. Tie a piece of thread around the solid object. Slowly lower the solid object until it is completely submerged. Record the new volume of the water. Calculate the volume, in milliliters, displaced by the solid.

C. Measuring Mass Materials: Balance, objects to weigh (beaker, rubber stopper, evaporating dish), unknown mass Your instructor will show you how to use a laboratory balance. Be sure the reading is 0.00 when the balance pan is empty. 1. 2. 3. 4.

Separately place each object on the balance pan and record its mass. Obtain the object of unknown mass and record its code number. Place it on the balance pan and record its mass. Obtain the mass of the object from your instructor and compare to your experimental value. Determine the correct number of significant figures for each mass obtained.

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

1

1. What are the standard units of length, mass, volume, and temperature in the metric system?

2. Why is the metric (SI) system called a decimal system of measurement?

3. What is the purpose of using prefixes in the metric system?

4. Fill in the blank lines below with the unit name, abbreviation, and property measured. Unit Name

Abbreviation

Property Measured

L centimeter km mg 5. Identify each of the following as a measured number or an exact number:

5 books

9.25 L

0.035 kg

100 cm in 1 m

12 beakers

59.067 g

1000 mL = 1 L

8.03 mL

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Date

Name

Section

Team

Instructor

REPORT SHEET

Chemistry and Measurement

LAB

1

A. Measuring Length 1. What units are represented by the numbers marked on the meterstick? What do the small lines marked on the meterstick represent? Complete the following statements: There are There are There are

centimeters (cm) in 1 meter (m). millimeters (mm) in 1 meter (m). millimeters (mm) in 1 centimeter (cm).

Item

2. Length

3. Estimated Digit

4. Number of Significant Figures

Width of little fingernail Distance around wrist Length of your shoe 5. Length of line Your measurement Other students’ values How does your value of the line length compare to those of other students?

Questions and Problems Q1 What digits in the measurements for the line by other students are the same as yours and which are different?

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B. Measuring Volume Volume of a liquid (include units for every measurement) Cylinder 1

Cylinder 2

Cylinder 3

1. Volume (mL) Volume of a solid by displacement 2. Initial volume of water 3. Volume of water and submerged solid 4. Volume of solid (3 − 2)

C. Measuring Mass Item

1. Mass

4. Number of Significant Figures

1. Beaker Stopper Evaporating dish 2. Unknown # 3. Actual mass of unknown Questions and Problems Q2 State the number of significant figures in each of the following measurements: 4.5 m

204.52 g

0.0004 L

625.000 mm

805 lb

34.80 km

Q3 Indicate the estimated digit in each of the following measurements: 1.5 cm

4500 mi

0.0782 in.

42.50 g

48.231 g

8.07 lb

Conversion Factors and Problem Solving 2 LABORATORY GOALS • Round off a calculated answer to the correct number of significant figures or decimal places. • Determine the area of a rectangle and the volume of a solid by direct measurement. • Determine metric, U.S. system, and metric−U.S. system equalities and their conversion factors. • Use conversion factors in calculations to convert units of length, volume, and mass.

LAB INFORMATION Time: Comments:

Related Topics:

3h Tear out the report sheets and place them beside the procedures. Identify the smallest unit of measurement on each measuring tool you use. Include an estimated digit for each measurement. Round off the calculator answers to the correct number of significant figures. Conversion factors, significant figures or decimal places, density

CHEMICAL CONCEPTS Every day, you make some measurements, such as weighing yourself or checking the temperature. Such measurements are probably in units of the U.S. system. In the laboratory, you will make measurements too, and perform calculations. Most of these measurements, or measured numbers, will use units from the metric system (see Figure 2.1).

▲ FIGURE 2.1 There are many measurements in everyday life. When you use measured numbers in calculations, the answers that you report must reflect the precision of the original measurements. Thus, it is often necessary to adjust the results you see on the calculator display. Every time you use your calculator, you will need to assess the mathematical operations, count significant figures, and your calculator can not do this for you!

A. Rounding Off Usually there are more digits in a calculator display than there are significant digits in the measured numbers used in the calculation. Therefore, we adjust the calculator result by rounding off. If the first digit to be dropped is less than 5, it and all following digits are dropped. If the first digit to be dropped is 5 or greater, all the following digits are dropped and the value of the last retained digit is increased by 1. When you round a large number, the correct magnitude is retained by replacing the dropped digits with placeholder zeros (see Sample Problem 2.l). When a whole number appears in the calculator display, significant zeros may need to be added.

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SAMPLE PROBLEM 2.1 Round off each of the following to report an answer with three significant figures: a. 75.6243 m

b. 0.528392 L

c. 3. 8765 × 104 s

d. 4 kg

SOLUTION: a. 75.6 m

b. 0.528 L

c. 3.88 × 104 s

d. 4.00 kg

B. Significant Figures in Calculations When you measure the mass of a sample, the volume of a liquid, or the temperature of a solution, the measured numbers you obtain are often used in mathematical calculations. The answers you report depend on the number of significant figures in those measured numbers. Multiplication/Division When multiplying or dividing, the final answer is written so that it has the same number of significant figures as the measurement with the fewest significant figures (SFs) (see Sample Problem 2.2).

SAMPLE PROBLEM 2.2 Solve:

0.025 × 4.620 = 3.44

SOLUTION:

On your calculator, enter the number and the operation for multiplication or division as follows: Two SFs Four SFs 0.025 × 4.620

Three SFs ÷ 3.44

=

0.033575581 calculator display

=

0.034

answer rounded off to two SFS

Addition/Subtraction When adding or subtracting, the final answer is written so that it has the same number of decimal places as the measurement having the fewest decimal places (see Sample Problem 2.3).

SAMPLE PROBLEM 2.3 Add: 2.11 + 104.056 + 0.1205 SOLUTION:

2.11 104.056 + 0.1205

two decimal places three decimal places four decimal places

106.2865

calculator display

106.29

answer rounded off to two decimal places

C. Equalities and Conversion Factors Metric Conversion Factors If a quantity is expressed in two different metric units, a metric equality can be stated. For example, the length of 1 m is the same as 1000 mm, which gives the equality 1 m = 1000 mm. The ratio of the two values is called a conversion factor.

Conversion Factors and Problem Solving

Metric equality:

1 m = 1000 mm

Conversion factors:

1m 1000 mm and 1000 mm 1m

13

Two conversion factors are always possible for any equality. Metric conversion factors, which are from definitions, are exact and do not limit the number of significant figures in the answer. Metric—U.S. System Conversion Factors A metric—U.S. system equality gives the relationship between a metric unit and a U.S. unit. For example, 454 g is the same mass as 1 lb. Usually one value is measured and the other value is exact. Then 454 g is a measured number with three SFs, while the 1 lb is exact. In the equality 1 in. = 2.54 cm, both numbers have been defined as exact (see Figure 2.2).

▲ FIGURE 2.2 Comparing centimeters and inches See Table 2.1 for some common equalities and conversion factors for length, mass, and volume from the metric system and the metric—U.S. systems of measurement. TABLE 2.1 Common Equalities and Their Conversion Factors Metric—Metric Equality

Conversion Factors

Length

1 m = 100 cm

Mass

1 kg = 1000 g

Volume

1 L = 1000 mL

1m 100 cm and 100 cm 1m 1 kg 1000 g and 1000 g 1 kg

Metric—U.S. Equality

1 in. = 2.54 cm 1 lb = 454 g

1L 1000 mL and 1 qt = 946 mL 1000 mL 1L

Conversion Factors

2.54 cm 1 in. and 1 in. 2.54 cm 1 lb 454 g and 454 g 1 lb 946 mL 1 qt and 1 qt 946 mL

D. Problem Solving Using Conversion Factors The process of problem solving in chemistry often requires the conversion of a given quantity with one unit to the needed quantity with a different unit (see Sample Problem 2.3).

SAMPLE PROBLEM 2.3 If a melon has a mass of 546 g, what is its weight in pounds? SOLUTION: Step 1 State the given and needed quantities. Given: 546 g

Need: weight in pounds

14

Laboratory Manual for General, Organic, and Biological Chemistry Step 2 Write a plan to convert the given unit to the needed unit.

grams

Metric—U.S. factor

pounds

Step 3 State the equalities and conversion factors. Equality

Conversion Factors 1 lb 454 g and 454 g 1 lb

1 lb = 454 g

Step 4 Set up the problem to cancel units and calculate the answer.

Note that the conversion factor selected has grams in the denominator to cancel grams of the given unit in the numerator. Always begin the set up by writing the giving information, not the conversion factor. Use the conversion factor to convert into the desired unit. Exact 1 lb 546 g × = 1.20 lb 454 g Three SFs

Three SFs Three SFs

Conversion Factors and Problem Solving

EXPERIMENTAL PROCEDURES

15

GOGGLES REQUIRED!

A. Rounding Off A student rounded off some numbers to three significant figures. In a few cases, significant zeros were added. 1. 2.

Determine whether the rounding was done correctly. If it is incorrect, write the correctly rounded number.

B. Significant Figures in Calculations B.1 Multiplication and Division

Solve the multiplication and division problems. Report your answers with the correct number of significant figures. B.2 Addition and Subtraction

Solve the addition and subtraction problems. Report your answers with the correct number of decimal places. B.3 Area Materials: Meterstick or metric ruler 1. 2. 3. 4.

Use a metric ruler to measure the length (cm) and width (cm) of the rectangle drawn on the report sheet. Calculate the area (cm 2 ) of the rectangle using your measurements and the formula Area = L × W. Obtain a second set of measurements from another student and record. Calculate the area using the measurements from the other student.

B.4 Volume of a Solid Materials: Meterstick or metric ruler, solid object 1. 2. 3.

Record the shape of the solid object and the dimensions to measure (see list below). Use a metric ruler to measure the dimensions of the solid in centimeters (cm). Calculate the volume of the solid object, in cm3 , using the appropriate formula from the following: Shape

Dimensions to Measure

Formulas for Volume

Cube

Length (L)

V = L3

Rectangular solid

Length (L), width (W), height (H)

V = L×W×H

Cylinder

Diameter (D), height (H)

V=

πD 2 H 3.14D 2 H = 4 4

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Laboratory Manual for General, Organic, and Biological Chemistry

C. Equalities and Conversion Factors C.1 Metric—Metric Conversion Factors for Volume Materials: 1-L graduated cylinder 1. 2.

Observe the markings on a 1-L graduated cylinder. Write an equality that states the number of milliliters in 1 L. Write two conversion factors for the equality.

C.2 Metric—U.S. System Conversion Factors for Volume Materials: 1-L graduated cylinder, 1-qt measure (or 1-pt measure) 1. 2. 3.

Using a 1-pt or 1-qt measure, measure 1 qt (or 2 pt) of water and transfer to a 1-L graduated cylinder. Record the number of milliliters in 1 qt. Write the true equality that states the number of milliliters in a quart. Write two conversion factors for the equality.

C.3 Metric—U.S. System Conversion Factors for Length Materials: Metric ruler or meterstick 1. 2. 3. 4.

Measure the vertical length of this page in inches and in centimeters. Divide the number of centimeters by the number of inches to give the relationship of cm/in. Round off to give the experimental factor for the number of centimeters in 1 in. Write the equality that states the number of milliliters in a quart. Write two conversion factors for the equality.

C.4 Metric—U.S. System Conversion Factors for Mass Materials: Commercial product with mass (weight) of contents given on label that list the contents in both metric and U.S. units. 1.

Write the name of the commercial product.

2.

Read the label on the product. Record the mass of the contents, in grams, and the weight, in ounces or pounds. (Do not weigh contents.) If the weight is given in ounces, convert it to pounds (1 lb = 16 oz). Divide the mass, in grams, of the product by its weight in pounds. This is your experimental factor for grams in one pound (g/1 lb). Write the equality that states the number of grams in 1 lb. Write two conversion factors for the equality.

3. 4. 5. 6.

D. Problem Solving Using Conversion Factors Calculating Your Metric Height Materials: Yardstick 1. 2.

Record your height in inches. Or use a yardstick to measure. Using the appropriate conversion factor, calculate your height in centimeters. Show your setup for each calculation. Height ( in .) ×

3.

2.54 cm = your height (cm) 1 in .

Using the appropriate conversion factor, calculate your height in meters. Show your setup for each calculation. Height ( cm ) ×

1m = your height (m) 100 cm

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

2

1. What are the rules for rounding off numbers?

2. How do you determine the number of significant figures for an answer obtained by multiplication or division?

3. How do you determine the number of decimal places for an answer obtained by addition or subtraction?

4. What is an equality and how is it used to write a conversion factor?

5. Write the equality and conversions factors for the relationship between miles and hours for a car traveling at 55 mi/h.

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Date

Name

Section

Team

Instructor

REPORT SHEET

Conversion Factors and Problem Solving

LAB

2

A. Rounding Off Initial Number

Student’s 1. Correct? Rounded Value (yes/no)

143.63212

144

532 800

533

0.008 583 45

0.009

8

8.00

2. Corrected (if needed)

B. Significant Figures in Calculations B.1 Multiplication and Division Perform the following multiplication and division calculations. Give a final answer with the correct number of significant figures:

0.1184 × 8.00 × 0.0345 (42.4)(15.6) 1.265 (35.56)(1.45) (4.8)(0.56) B.2 Addition and Subtraction

Perform the following addition and subtraction calculations. Give a final answer with the correct number of decimal places. 13.45 mL + 0.4552 mL 145.5 m + 86.58 m + 1045 m

245.625 g – 80.2 g 4.62 cm – 0.885 cm Copyright © 2014 Pearson Education, Inc.

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Laboratory Manual for General, Organic, and Biological Chemistry

B.3 Area

1. Your measurements

3. Another student’s measurements

Length Width Area (2., 4.) Why could two students obtain different values for the calculated area of the same rectangle?

B.4 Volume of a Solid 1. Shape of the solid

Dimensions to measure 2. Height

Width

Length Diameter (if cylinder)

3. Formula for volume of solid

Volume of the solid (Show calculations of volume, including the units.)

Conversion Factors and Problem Solving

21

C. Equalities and Conversion Factors C.1 Metric—Metric Conversion Factors for Volume 1. Equality

1L =

mL

2. Conversion factors C.2 Metric—U.S. System Conversion Factors for Volume 1. Number of milliliters in 1 qt 2. True equality

mL (Experimental factor)

1 qt =

mL

3. Conversion factors

How does your experimental factor compare to the conversion factor 946 mL/1 qt.?

C.3 Metric—U.S. System Conversion Factors for Length 1. Vertical page length (measured)

in.

Vertical page length (measured)

cm

cm/in. =

2. 3. True equality

cm/1 in. (Experimental factor)

in. =

cm

4. Conversion factors

How does your experimental factor compare to the conversion factor 2.54 cm/1 in.?

C.4 Metric—U.S. System Conversion Factors for Mass 1. Name of commercial product 2. Mass, in grams, stated on label

g

Weight, in pounds or ounces, stated on label 3. Weight in lb (Convert oz to lb if needed.) 4.

lb g/lb =

g/1 lb

1 lb =

g

(Experimental factor) 5. True equality 6. Conversion factors

How does your experimental factor compare to the standard conversion factor of 454 g/lb?

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Laboratory Manual for General, Organic, and Biological Chemistry

D. Problem Solving Using Conversion Factors Your metric height 1. Height (inches)

in.

2. Height in centimeters

cm (Show your calculations)

3. Height in meters

m (Show your calculations)

Questions and Problems Q1 A pencil is 16.2 cm long. What is its length in inches?

Q2 A person has a mass of 63.4 kg. What is that weight in pounds?

Q3 A bottle of olive oil contains 1.4 qt of olive oil. What is that volume in milliliters?

Q4 How many liters of plasma are present in 8.5 pt?

Density and Specific Gravity 3 LABORATORY GOALS • Calculate the density of a substance from measurements of its mass and volume. • Calculate the specific gravity of a liquid from its density. • Determine the specific gravity of a liquid using a hydrometer.

LAB INFORMATION Time: Comments: Related Topics:

2h Tear out the report sheets and place them beside the procedures. Round off the calculator answers to the correct number of significant figures. Dispose of liquids properly, as directed by your instructor. Mass, volume, prefixes, significant figures, density

CHEMICAL CONCEPTS A. Density of Liquids To determine the density of a liquid, you need the mass and volume of the liquid. The mass of a liquid is determined by weighing. The mass of a container is obtained, a certain volume of liquid is added, and then the combined mass is determined. Subtracting the mass of the container gives the mass of the liquid. From the mass and volume, the density is calculated. Every substance has a unique density, which is the relationship of its mass to its volume. Once you have measured the mass and volume of a substance you can calculate its density. If the mass is measured in grams and the volume in milliliters or cubic centimeters, its density will have units of g/mL or g/cm3 . Density of a substance =

mass of substance g of substance = . volume of substance mL or cm3 of substance

B. Specific Gravity The specific gravity of a liquid is a comparison of the density of that liquid with the density of water, which is 1.00 g/mL at 4 °C. Specific gravity =

denisty of a substance (g/mL) . denisty of water (1.00 g/mL)

Specific gravity is a number with no units; the units of density (g/mL) have canceled out. This is one of the few measurements in chemistry written without any units.

C. Density of Solids The volume of a solid is often determined using volume displacement, in which a solid submerged in a known amount of water displaces its own volume (see Figure 3.1).

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▲ FIGURE 3.1 The density of a solid can be determined by volume displacement, in which a submerged object displaces a volume of water equal to its own volume.

Density is a unique property of every substance and can be used to identify a substance. Some solid objects have a density that is less than that of water, and so they float in water. Other solid objects have a density that is greater than that of water, and so they sink in water (see Figure 3.2).

▲ FIGURE 3.2 Objects that sink in water are more dense than water; objects that float are less dense.

D. Graphing Mass and Volume A graph can be plotted to show the relationship between the mass and volume of metal samples (see Graphing Experimental Data at the beginning of this lab manual). After samples of a metal are weighed and their volume determined, a graph is drawn by plotting the mass (g) of the metal pieces and their volume (mL). The slope of the line is used to determine the density of the metal.

Density and Specific Gravity

EXPERIMENTAL PROCEDURES

25

GOGGLES REQUIRED!

A. Density of Liquids A.1 Density of Water Materials: 50-mL graduated cylinder, water, 100-mL or 250-mL beaker 1. 2. 3. 4. 5.

Place about 20 mL of water into a 50-mL graduated cylinder. Record the volume with the correct number of significant figures. Weigh a small dry beaker. Record its mass. Pour the water into the beaker and weigh. Record the combined mass. Determine the mass of the water by subtracting the mass of the beaker from the mass of the beaker plus the water. Calculate the density of the water sample by dividing its mass (g) by its volume (mL).

A.2 Density of an Unknown Liquid Materials: 50-mL graduated cylinder, liquid sample, 100-mL or 250-mL beaker 1. 2. 3. 4. 5.

Place about 20 mL of the liquid sample into a 50-mL graduated cylinder. Record the type of liquid and its volume, in milliliters, with the correct number of significant figures. Weigh a small dry beaker. Record its mass. Pour the liquid sample into the beaker and weigh. Record the mass with the correct number of significant figures. Determine the mass of the liquid sample by subtracting the mass of the beaker from the mass of the beaker with the liquid sample. Calculate the density of the liquid sample by dividing its mass (g) by its volume (mL). Dispose of the liquid sample according to your instructor’s directions.

B. Specific Gravity Materials: Display of graduated cylinders with hydrometers in water and liquid samples 1. 2. 3. 4.

Calculate the specific gravity of water by dividing its density from A.1 by the density of water (1.00 g/mL). Calculate the specific gravity of the liquid sample you used in A.2 by dividing its density by the density of water (1.00 g/mL). Read the hydrometer floating in a graduated cylinder containing water that you used in B.1. (see Figure 3.3). Record its specific gravity as a decimal number. Read the hydrometer floating in a graduated cylinder containing the liquid sample that you used in B.2. Record its specific gravity as a decimal number.

Using a hydrometer The hydrometer is an instrument used to determine the specific gravity of a fluid. A hydrometer is placed in a liquid and spun slowly to keep it from sticking to the sides of the container. The scale on the hydrometer is read at the lowest (center) point of the meniscus of the fluid (see Figure 3.3). Some hydrometers use a comma to represent a decimal point, while other hydrometers don’t use any markings. Thus, the value of 1,000 or 1000 on a hydrometer is read as 1.000.

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Laboratory Manual for General, Organic, and Biological Chemistry

◄ FIGURE 3.3 Measuring specific gravity using a hydrometer

C. Density of Solids Materials: Metal object, string or thread, graduated cylinder 1. 2. 3. 4. 5. 6.

Obtain a metal object. Weigh it and record the mass, in grams. Obtain a graduated cylinder that is large enough to hold the solid metal object. Add water to the cylinder until it is about full. Record the water level, in milliliters. Attach a string or thread to the object. Lower it slowly into the water until it is submerged. Record the final water level, in milliliters. Calculate the volume, in milliliters, of the object. Calculate the density (g/mL) of the metal object. If your metal object is made of one of the metals in Table 3.1, use your density from step 5 to identify the metal. TABLE 3.1 Density Values of Some Metals Substance

Density (g/mL)

Aluminum

2.7

Brass

8.4

Copper

8.9

Iron

7.9

Lead

11.3

Nickel

8.9

Tin

7.3

Zinc

7.1

Density and Specific Gravity

27

D. Graphing Mass and Volume Materials: Metal pieces such as aluminum, copper, zinc, or pennies (pre-1980 or post-1980), 50-mL graduated cylinder In this graphing activity, volume and mass of five different samples of the same metal will be measured and used to prepare a graph. The density (g/mL) will be visually represented on a graph. 1. 2. 3. 4. 5. 6. 7.

Place about 20–25 mL of water in a 50-mL graduated cylinder. Carefully record this initial volume of water. Place the cylinder and water on a top-loading balance. Record this initial mass. Add two or three pieces of metal or pennies. Indicate the type of metal pieces. Record total mass. Record the new level of the water. Add a few more metal pieces to the cylinder and repeat step 4. Add a few more metal pieces to the cylinder and repeat step 4. Add a few more metal pieces to the cylinder and repeat step 4.

Calculations 8. Calculate the mass of the metal pieces or pennies by subtracting the initial mass of the cylinder and water from each combined masses of cylinder, water, and pieces of metal. 9. Calculate the volume displaced by the metal pieces by subtracting the initial volume of water from each of the new water levels. 10. Calculate the density (g/mL) for the metal in each experiment. Drawing the Graph 11. Label the axes of the graph with equal intervals that include the highest values. The origin of each axis is 0.0. Draw the graph by plotting the mass (g) of the metal pieces or pennies on the vertical axis and the volume (mL) of the metal pieces or pennies on the horizontal axis. Place a point at (0,0) for a mass of 0.00 g and a volume of 0.00 mL. Use a ruler to draw a straight line through the points you have marked on the graph. If some of the points are not on the line, draw the line between them so that as many points appear above the line as below the line. Slope 12. The slope of the line on the graph represents the density of the metal. Mark two places on the line and use these in the next step. Divide the difference between the two mass values by the difference of the two values for volume. See equation below. Mass (2) − Mass (1) g = = density of metal or pennies Volume (2) − Volume (1) mL

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Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

3

1. What property of oil makes it float on water?

2. Why would heating the gas in a hot air balloon make the balloon rise?

3. What is the difference between density and specific gravity?

4. An object has a mass of 18.4 g and a volume of 11.2 mL. a. What is the density of the object?

b. What is the specific gravity of the object?

c. Will the object sink or float in water?

5. A 15.0 g-sample of a liquid has a density of 0.875 g/mL. a. What is the volume, in milliliters, of the liquid? Use conversion factors to calculate answer.

b. What mass, in grams, of the liquid is in 34.6 mL of the liquid?

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Date

Name

Section

Team

Instructor

REPORT SHEET

Density and Specific Gravity

LAB

3

A. Density of Liquids A.1 Density of Water 1. Volume (mL) 2. Mass of beaker 3. Mass of beaker + liquid 4. Mass of liquid 5. Density (Show calculations)

A.2 Density of an Unknown Liquid 1. Type of liquid sample Volume of liquid 2. Mass of beaker 3. Mass of beaker + liquid 4. Mass of liquid 5. Density (Show calculations)

B. Specific Gravity 1. Specific gravity of water 2. Specific gravity of liquid sample 3. Specific gravity of water using a hydrometer 4. Specific gravity of liquid using a hydrometer How do the calculated specific gravity of water and the liquid sample (1, 2) compare to the hydrometer readings (3, 4) for water and the liquid sample?

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C. Density of Solids 1. Mass of the solid 2. Initial water level 3. Final water level with solid 4. Volume of solid (3 − 2) 5. Density of the solid

g/mL

(Show calculations) 6. Type of metal

D. Graphing Mass and Volume 1. Initial volume of water

mL

2. Initial mass of cylinder + water

g

3. Type of metal Total Mass of Cylinder, Water, and Metal Pieces

Total Volume of Water and Metal Pieces

4.

g

4.

mL

5.

g

5.

mL

6.

g

6.

mL

7.

g

7.

mL

Calculations 8. Mass of Metal Pieces

9. Volume of Metal Pieces

10. Density (8/9)

(4–2)

g

(4–1)

mL

g/mL

(5–2)

g

(5–1)

mL

g/mL

(6–2)

g

(6–1)

mL

g/mL

(7–2)

g

(7–1)

mL

g/mL

Density and Specific Gravity

33

11. Drawing the Graph

12. Density of the metal =

Mass (2) − Mass (1) = Volume (2) − Volume (1)

g/mL

Questions and Problems Q1 A metal object has a mass of 8.37 g. When it was placed in a graduated cylinder containing 20.0 mL of water, the water level rose to 23.1 mL. What is the density and identity of the metal? (see Table 3.1)

Q2 What is the mass of a solution that has a density of 0.775 g/mL and a volume of 50.0 mL?

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Laboratory Manual for General, Organic, and Biological Chemistry

Q3 What is the volume of a solution that has a specific gravity of 1.2 and a mass of 185 g?

Q4 A pump delivers the following volume of a solution over 4 h. Volume (mL)

Time (h)

0

0

50

1.0

100

2.0

125

2.5

150

3.0

200

4.0

Prepare a graph to represent the data.

Temperature and Specific Heat 4 LABORATORY GOALS • • • • • •

Convert a Celsius temperature to Fahrenheit and to Kelvin temperatures. Distinguish between a calorie, joule, kilocalorie, and nutritional Calorie. Use the specific heat of water to calculate heat lost or gained. Calculate the specific heat of a metal object in cal/g °C and J/g °C. Calculate the energy values of foods in kcal/g. Use nutritional data to determine the kilocalories and kilojoules in one serving.

LAB INFORMATION Time: Comments: Related Topics:

2–3 h Tear out the lab report sheets and place them beside the experimental procedures. Be careful with boiling water. Specific heat, measuring heat energy, calculating heat in calories and joules, nutritional Calorie

CHEMICAL CONCEPTS A. Temperature Temperature measures the intensity of heat in a substance. On the Celsius scale, water freezes at 0 °C; on the Fahrenheit scale, water freezes at 32 °F (see Figure 4.1).

▲ FIGURE 4.1 The Fahrenheit, Celsius, and Kelvin temperatures use the freezing point and boiling point of water as reference points.

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A Celsius temperature is converted to its corresponding Fahrenheit temperature by using the following equation: TF = 1.8(TC ) + 32

When the Fahrenheit temperature is known, the Celsius temperature is determined by rearranging the equation. TC =

(TF − 32) 1.8

A Celsius temperature can be converted to a Kelvin temperature by using the following equation: TK = TC + 273

B. Specific Heat Every substance has the capacity to absorb heat. When heat is added to a substance, the temperature of that substance increases. Different substances vary in the amount of heat required to raise their temperatures by 1 °C. Specific heat is the amount of heat in calories or joules that raises the temperature of 1 g of a substance by 1 °C (see Table 4.1). Specific heat =

amount of heat (cal or J) mass(g) × change in temperature ΔT (°C)

TABLE 4.1 Specific Heat Values for Selected Substances Substance

Specific Heat J/g °C

cal/g °C

Water, H 2 O (liquid)

4.184

1.00

Iron

0.452

0.108

Copper

0.385

0.0920

Aluminum

0.897

0.214

Lead

0.128

0.306

Zinc

0.387

0.0925

For the experiment, a Styrofoam cup is used as a calorimeter (see Figure 4.2). A measured amount of water is placed in the cup and its temperature is determined. A piece of metal is heated to the temperature of boiling water (about 100 °C) and quickly transferred to the water in the calorimeter. The heat lost by the hot metal warms the water in the cup. As the temperature of the water increases, the temperature of the metal decreases until both the water and the metal reach the same final temperature. By measuring the increase in the temperature of the water, the amount of heat (calories or joules) given off by the hot metal can be calculated.

▲ FIGURE 4.2 A calorimeter consists of a Styrofoam cup, a cover, water, and a thermometer.

Temperature and Specific Heat

37

For example, suppose 95 g of water in the Styrofoam cup has an initial temperature of 24.5 °C. When a 30.3-g piece of metal at a temperature of 100.0 °C (boiling water) is added, the temperature of the water rises to a final temperature of 27.0 °C. The temperature increase for the water is 2.5 °C (27.0 °C – 24.5 °C). The temperature decrease of the metal is 73.0 °C (100.0 °C – 27.0 °C). Heat gain for the water

Heat (cal) = g × ΔT (°C) × SH (H 2 O) Heat (cal) = 95 g × 2.5 °C × 4.184 J/g °C = 990 J If 990 J is gained by water, then 990 J is the amount of heat lost by the piece of metal. Calculate the specific heat of the metal SH (metal) =

J 990 J = = 0.45 J/g °C Mass(g) × ΔT (°C) (30.3 g)(73.0 °C)

The metal is probably iron because the experimental result is close to the specific heat of iron.

C. Energy and Nutrition The energy values of a food are measured in Calories or kilojoules using a calorimeter. In this lab, you will determine the number of Calories in a sample of food. Burning a cheese puff or some chips releases heat that increases the temperature of water in an aluminum can. Some heat is lost to the can and the surrounding air. However, you can get an idea of how nutritionists determine the energy values of foods. Energy values are usually reported on U.S. food packages in kilocalories (Calories). Energy values are given in kcal (Cal)/g. However, the international unit, kilojoule (kJ), is becoming more prevalent. For example, a baked potato has an energy value of 110 Calories, which is 110 kcal or 460 kJ. A typical diet of 2100 Cal (kcal) is the same as an 8800 kJ diet. Energy Values in Nutrition 1 Cal = 1 kcal = 1000 cal

1 Cal = 4.184 kJ

D. Energy Values for Foods Our diets contain foods that provide us with energy. We need energy to make our muscles work, to breathe, to synthesize molecules such as protein and fats, and to repair tissues. A typical diet required by a 25-year-old woman is 2000–2500 kcal. Nutritionists established the energy values for the three food types: carbohydrates, 17 kJ/g (4 kcal/g); fats, 38 kJ/g (9 kcal/g); and proteins, 17 kJ/g (4 kcal/g). For example, a candy that contains 12 g of carbohydrate will provide 200 kJ or 50 Cal. The energy values are usually rounded off to the tens place. 12 g carbohydrate ×

17 kJ = 204 kJ = 200 kJ 1 g carbohydrate

12 g carbohydrate ×

4 kcal = 48 kcal = 50 kcal = 50 Cal 1 g carbohydrate

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Temperature Materials: Thermometer (°C), a 150- or 250-mL beaker, ice, rock salt Caution: 1.

2. 3. 4. 5.

Never shake down a laboratory thermometer. Shaking a laboratory thermometer can cause breakage and serious accidents.

Observe the markings on a thermometer. Indicate the lowest and highest temperatures that can be read using that thermometer. On most thermometers, you can estimate the tenths of a degree (0.1 °C) To measure the temperature of a liquid, place the bulb of the thermometer in the center of the solution. Keep it immersed while you read the temperature scale. Wait until the temperature becomes constant before you record the temperature. Fill a 250-mL beaker about 1/3 full of water. Record its temperature. Fill a 250-mL beaker about 1/3 full of water and add an equal amount of ice. After about 5 min, record the temperature of the ice-water mixture. Keep the mixture for step 4. Add 20–30 g of rock salt to the ice-water mixture from step 3. Stir and allow a few minutes for the temperature to change. Record the temperature of the salt-ice-water. Convert the Celsius temperatures to °F and K.

B. Specific Heat Materials: Thermometer, Bunsen burner, ring stand, iron ring, wire screen, 400-mL beaker, balance, calorimeter (Styrofoam cup and cover), stirring rod, metal object, and string Fill the 400-mL beaker about two-thirds full of water. Place the beaker and water on a wire screen set on an iron ring so that the ring is about 6 cm above the top of the Bunsen burner. Light your Bunsen burner and begin heating the water in the beaker. 1.

Obtain a metal object. Record the identification numbers on the metal.

2.

Weigh the metal object and record its mass. Tie a length of string or fishing line to the metal object and gently lower it into the water bath. Allow the water bath containing the metal object to boil for 10 min.

3.

Obtain a Styrofoam cup and weigh it. Record its mass.

4.

Add about 50 mL of water to the cup (calorimeter) making sure there is enough water to cover the metal object. If the metal object is large, add more water. Weigh the cup and water. Record.

5.

Measure the temperature of the boiling water bath after it has boiled for 10 min or more. Record.

6.

Measure the temperature of the water in the cup (calorimeter). Record. Carefully remove the hot metal object from the boiling water bath and transfer quickly to the colder water in the cup (calorimeter). Immediately place the cover on the calorimeter, and stir gently.

7.

Using a thermometer, record the highest temperature reached by the water in the cup (calorimeter). This will be the final temperature of both the water and the metal object. If you are going to run a second trial with the same metal piece, repeat steps 3 to 7.

Calculations 8.

Calculate the temperature change for the water in the Styrofoam calorimeter.

9.

Calculate the mass of the water in the cup (calorimeter).

10.

Calculate the heat, in joules, gained by the water. Heat (J) = g H 2 O (9) × ΔT (8) × 4.184 J/g °C

11.

State the heat, in joules, lost by the metal object.

12.

Calculate the temperature change for the metal.

Temperature and Specific Heat

13.

Calculate the specific heat of the metal. SH of metal =

14.

39

heat (J)(11) mass (metal) (2) × ΔT (12)

Use Table 4.1 to identify the type of metal in the object. Your instructor may add other values for identification.

C. Energy and Nutrition (This may be an instructor demonstration) Materials: Aluminum can, food sample such as chips or cheese puffs with a nutrition label, thermometer, ring stand, two iron rings, a clamp, wire screen, Bunsen burner 1. Record the type of food used. 2. Weigh and record the mass of an aluminum can. 3. Pour about 100 mL of water into the can. Weigh the can and water. Record the total mass, in grams. Set up the apparatus in the hood: Place the can in an iron ring that holds it securely upright. Attach a second iron ring covered with a wire screen a short distance below the aluminum can. Suspend a thermometer from a clamp so that the bulb is below the water level in the aluminum can. 4. Weigh the food sample and record. 5. Using a thermometer, record the initial temperature of the water in the can. Place the food sample on the wire screen below the aluminum can. Ignite the food sample using a match or Bunsen burner. Remove the heat source immediately and let the food sample burn. 6. Using a thermometer, record the highest temperature of the water in the can. 7. Weigh any remaining ash or food sample. Record. Calculations 8. 9. 10.

Calculate the mass, in grams, for the water in the aluminum can. Calculate the temperature change of the water after it is heated by the burning food. Calculate the heat gain of the water, in calories. Heat (cal) = g water (8) × ΔT (9) × 1.00 cal/g °C

11.

Write the calories given off by the burning food, which is equal to the heat gained by the water.

12.

Heat (cal) lost by burning food = Heat (cal) gained by water (10) Calculate the heat given off by the burning food sample in kilocalories (Cal).

13. 14.

Kilocalories = heat (cal) (11) × 1 kcal/1000 cal Calculate the mass of food sample that burned by subtracting the mass of ash and any remaining unburned food after combustion. Calculate the energy value of the food. Energy value of food =

kcal (12) g of food burned (13)

D. Energy Values for Foods Materials: Food products with nutrition data on labels 1. 2. 3. 4. 5. 6.

Obtain a food product that has a Nutrition Facts label. Record the type of food. Record the serving size. List the grams of carbohydrate, fat, and protein in one serving of the food. From the mass of each food type, calculate the Calories (kcal) of each food type in one serving using the accepted energy values. Round the answer to the tens place. Determine the total Calories (kcal) in one serving. Compare your total to the Calories listed on the label.

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Pre-Lab Study Questions

4

1. What is the equation for converting a Fahrenheit temperature to a Celsius temperature?

2. What is meant by the term specific heat?

3. Why is a measured amount of water needed to determine the specific heat of a metal object?

4. What are the units of specific heat?

5. How is the caloric value of a food sample determined?

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REPORT SHEET

Temperature and Specific Heat

LAB

4

A. Temperature 1. Lowest temperature

___________ Highest temperature ________________ °C

°F

K

2. Tap water

_____________

_____________

_______________

3. Ice-water mixture

_____________

_____________

_______________

4. Ice-water-salt mixture _____________

_____________

_______________

Questions and Problems Q1 Calculate each of the following temperatures conversions: a. 25 °C to °F b. 18 °F to °C c. 45 °C to K Q2 A recipe calls for a baking temperature of 205 °C. What temperature in °F should be set on the oven?

B. Specific Heat 1. Identification Number ______________ First run

Second run

2. Mass of metal

________________ ________________

3. Mass of cup calorimeter

________________ ________________

4. Mass of cup calorimeter + water

________________ ________________

5. Temperature of boiling water bath

________________ ________________

6. Initial temperature of water

________________ ________________

7. Final temperature of water and metal ________________ ________________

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Calculations

First run

Second run

8. ∆T (°C) for water (7−6)

_________________

________________

9. Mass of water (4−3)

_________________

________________

10. Heat (J) gained by water (Show calculations.)

_________________

________________

11. Heat (J) lost by the metal

_________________

________________

Heat (J) lost by metal = Heat (J) gained by water (10) _________________

________________

13. Specific heat of the metal (J/g °C) _________________

________________

12. ∆T (°C) for metal (5−7) (Show calculations.)

14. Identity of metal ________________________ Questions and Problems Q3 Why did you need to transfer the metal quickly from the hot water bath to the water in the cup calorimeter?

Q4 How many calories are required to raise the temperature of 225 g of water from 42 °C to 75 °C?

Q5 A metal object with a mass of 19 g is heated to 96 °C and then transferred to a calorimeter containing 75 g of water at 18 °C. If the water and metal object reach a final temperature of 22 °C, what is the specific heat of this metal object?

Temperature and Specific Heat

45

C. Energy and Nutrition 1. Type of food sample

________________________

2. Mass of aluminum can

________________________

3. Mass of aluminum can and water

________________________

4. Mass of food

________________________

5. Initial temperature of water

________________________

6. Final temperature of water

________________________

7. Mass of ash or food remaining

________________________

Calculations 8. Mass of water (3 − 2)

________________________

9. Temperature change for water (6 − 5)

________________________

10. Heat gained by water (Show calculations.)

________________________

11. Heat (cal) given off by burning of food sample

________________________

12. Heat (kcal) given off by burning of food sample ________________________ 13. Mass of food undergoing combustion (4−7)

________________________

14. Energy value (kcal/g) of food sample (Show calculations.)

________________________

Questions and Problems Q6 A 0.25 g sample of a pretzel is burned. The heat it gives off is used to heat 50. g of water from 18 °C to 42 °C. What is the energy value of the pretzel, in kcal/g?

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D. Energy Values for Foods 1. Name of food product _________________________ _________________________ 2. Serving size 3. Mass of food types in one serving Carbohydrate ___________ g

Fat ___________ g

Protein ___________ g

4. Calculations for kcal per serving (Show calculations.) ____________________ kcal (Cal) Carbohydrate

Fat

____________________ kcal (Cal)

Protein

____________________ kcal (Cal)

5. Total Calories (kcal) per serving

____________________ kcal (Cal)

6. Calories (for one serving) listed on the label ____________________ Cal Questions and Problems Q7 What percent (%) of the total Calories in your food product is from fat?

What percent (%) of the total Calories in your food product is from carbohydrate?

What percent (%) of the total Calories in your food product is from protein?

Q8 How does your calculated number of Calories compare to the Calories listed on the label of the food product?

Energy and Matter 5 LABORATORY GOALS • • • •

Prepare a heating curve and a cooling curve. Use the specific heat of water to calculate heat lost or gained. Calculate the heat of fusion for water. Determine the freezing point of Salol.

LAB INFORMATION Time: Comments: Related Topics:

2h Tear out the report sheets and place them beside the matching procedures. Be careful with boiling water. Use mitts or beaker tongs to move hot beakers, or let them cool. Changes of state, heating and cooling curves, heat of fusion, heat of vaporization

CHEMICAL CONCEPTS A. A Heating Curve for Water The temperature of a substance indicates the kinetic energy (energy of motion) of its molecules. When water molecules gain heat energy, they move faster and the temperature rises. Eventually, the water molecules gain sufficient energy to separate from the other liquid molecules and form a gas. The liquid changes to a gas in a change of state called boiling. When a liquid boils, forming a gas, a horizontal line (plateau) appears on the graph, as shown in Figure 5.1. This constant temperature is called its boiling point.

▲ FIGURE 5.1 A heating curve illustrates the change in temperature and changes in state as heat is added.

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B. A Cooling Curve When a liquid cools, its particles move more slowly. Eventually, the attractions between the particles cause a solid to form (freezing). There is no change in the temperature while the liquid is freezing. When the liquid reaches its freezing point, its temperature becomes constant. On the graph in Figure 5.2, the drop in temperature is shown. As the liquid cools, the temperature may drop temporarily below its freezing point. This condition is called supercooling. The horizontal line or plateau on the cooling curve indicates the freezing point of the substance as it changes from a liquid to a solid.

▲ FIGURE 5.2 A typical cooling curve showing supercooling and the freezing point plateau.

ENERGY IN CHANGES OF STATE C. Heat of Fusion Changing state from solid to liquid (melting) requires energy. Ice melts at a constant temperature of 0 °C. At that melting point, the amount of heat required to melt 1 g of ice is called the heat of fusion. For water, the energy needed to melt 1 g of ice (at 0 °C) is 334. joules. That is also the amount of heat released when liquid water freezes to solid ice at 0 °C. Melting (0 °C): H 2 O(s ) + heat fusion (334 J/g) → H 2 O(l ) Freezing (0 °C): H 2 O(l ) → H 2 O(s ) + heat fusion (334 J/g)

In this experiment, ice will be added to a sample of water. From the temperature change, the amount of heat lost by the water sample can be calculated. This is also the amount of heat needed to melt the ice. Heat (J) lost by water = g water × ΔT × 4.184 J/g °C = heat (J) gained to melt ice By measuring the amount of ice that melted, the heat of fusion, in joules/gram, can be calculated. Heat of fusion (J/g) =

heat (J) needed to melt ice grams of ice

Heat of Vaporization A similar situation occurs for a substance that changes from a liquid to a gas (vapor). Water boils at 100 °C, its boiling point. At that temperature, the energy required to convert liquid to gas is called the heat of vaporization. For water, the energy required to vaporize 1 g of water at 100 °C is 2260 J. This is also the amount of heat released when 1 gram of steam condenses to liquid at 100 °C. Boiling (100 °C):

H 2 O(l ) + heat vaporization (2260 J/g) → H 2 O(g )

Condensation (100 °C): H 2 O(g) → H 2 O(l ) + heat vaporization (2260 J/g)

Energy and Matter

EXPERIMENTAL PROCEDURES

49

GOGGLES REQUIRED!

A. A Heating Curve for Water Materials: Beaker (250- or 400-mL), Bunsen burner (or hot plate), ring stand, graduated cylinder, iron ring, wire gauze, clamp, thermometer, timer Place the beaker on a hot plate or on a wire screen placed on an iron ring above a Bunsen burner (see Figure 5.3). The height of the iron ring should be about 3–5 cm above the burner. Place the thermometer securely in a clamp. Adjust the thermometer so that the bulb is in the center portion of the liquid. Do not let the thermometer rest on the side or bottom of the beaker.

▲ FIGURE 5.3 Setup for heating water with (left) a hot plate or (right) a Bunsen burner. 1. 2. 3.

Using a graduated cylinder, pour 100 mL of cool water into a 250-mL beaker. Record the volume with the correct number of significant figures. Using a thermometer, measure and record the initial temperature of the water. Light the burner (or turn on hot plate). Using a timer or a watch with a second hand, record the temperature of the water at 1-min intervals. You may need to add more lines on the data table for time and temperature. Eventually the water will come to a full boil. The early appearance of small bubbles of escaping gas does not indicate boiling. When the water is boiling, the temperature will become constant. Record the boiling temperature at 1-min intervals for another 4–5 min. Turn off the Bunsen burner or the hot plate.

Drawing a Heating Curve for Water 4.

5. 6. 7. 8.

Prepare a heating curve for water by graphing the temperature versus the time. (Review graphing in the preface of this lab manual) (see Figure 5.4). Label the parts of the graph that represent the liquid state and boiling. The plateau (flat part of graph) indicates the boiling point of the water. Record its value. Use the graph to determine the temperature change from the initial temperature of the cool water to the boiling point (plateau). Using the measured volume of water and its density (1.00 g/mL), calculate the mass of the water. Use the heat equation to calculate the joules required to heat the water to boiling. joules = mass (7) × ΔT (6) × 4.184 J/g °C

B. A Cooling Curve Materials: 400-mL beaker, water, Bunsen burner (or hot plate), freezing point apparatus (large test tube containing a small amount of Salol (phenylsalicylate) and fitted with a two-hole stopper, thermometer, and wire stirrer). Salol is a solid at room temperature. Your instructor will indicate the location of the already prepared test tubes containing Salol and the stirring apparatus (see Figure 5.5). Do not try to pull out the stopper or thermometer. It is frozen in the Salol until heated.

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▲ FIGURE 5.4 Example of a graph of a heating curve. Hot Water Bath to Melt the Salol The Salol in the freezing point apparatus must be melted before this experiment can begin. Prepare a hot water bath by filling a 400-mL beaker about half full of water. Place it over a Bunsen burner (or hot plate) with an iron ring and wire screen. Place the test tube setup containing the Salol in the warm water. Let the temperature of the melted Salol go up to about 70 °C but not to the boiling temperature of water. All of the Salol should now be liquid. Turn off the Bunsen burner or hot plate and remove the test tube with the liquid Salol from the warm water. Clamp the test tube and contents to a ring stand (see Figure 5.5). Gently raise and lower the wire stirrer to mix the contents. You will have to stop stirring when the stirrer becomes frozen in the Salol. 1.

Record the initial temperature and temperatures at 1-min intervals as the temperature of the liquid Salol decreases. After solid forms, continue to make at least five more temperature readings. A constant temperature for five or more minutes indicates that the Salol has reached its freezing point. Return the freezing point apparatus and contents to your instructor.

Drawing the Cooling Curve for Salol 2.

Plot the cooling curve for Salol on the graph provided in the report page. For a review of graphing, see the preface of this lab manual. Label the liquid state, solid state, supercooling (if any), and the freezing point of the Salol.

▲ FIGURE 5.5 Stirring apparatus for cooling curve and freezing point determination of Salol.

Energy and Matter

51

C. Heat of Fusion Materials: Calorimeter (Styrofoam cup and cardboard cover), thermometer, 50- or 100-mL graduated cylinder, 100-mL beaker, ice cubes or crushed ice 1. 2. 3.

Weigh an empty Styrofoam cup. Record its mass. Add 100 mL of water to the cup and reweigh. Record the combined mass. Using a thermometer, record the initial temperature of the water (see Figure 5.6).

Add 2 or 3 ice cubes (or use crushed ice that fills a 100-mL beaker) to the water in the cup. Stir. Add ice until the temperature drops to at least 2–3 °C. If some ice is not melted, remove it immediately and discard. 4.

Record the final temperature of the water.

5.

Weigh the calorimetry cup with the initial sample of water and the melted ice. The increase in mass indicates the amount of ice that melted.

▲ FIGURE 5.6 Calorimetry setup with water, a thermometer, Styrofoam cup, and a cardboard cover. Calculations 6. 7. 8. 9. 10. 11.

Calculate the mass of water initially placed in the calorimetry cup. Calculate the temperature change (∆T) for the water. Calculate the heat, in joules, lost by the water. joules = mass (6) × ΔT (7) × 4.184 J/g °C State the number of joules that were needed to melt the ice, which is equal to the joules lost. Calculate the grams of ice that melted. Calculate your experimental value for the heat of fusion for ice, in J/g. Heat of fusion (J/g) =

heat (J) needed to melt ice (9) grams of ice (10)

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Pre-Lab Study Questions

5

1. Why is energy required for the boiling process?

2. When water at 0 °C freezes, is heat lost or gained? Explain your answer.

3. Why are there two plateaus on the heating curve for water that begins at −15 °C and ends at 120 °C?

4. How many joules are required for each of the following changes? a. heating 65 g of water from 12 °C to 76 °C

b. melting 12.8 g of ice at 0 °C

c. boiling 21.1 g of water at 100 °C

5. How many kilojoules are released when 8.2 g of water condenses at 100 °C and cools to 15 °C?

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REPORT SHEET

Energy and Matter

LAB

5

A. A Heating Curve for Water 1. Volume of water Time (min) 2. 0 (initial) 3. 1 2 __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________

Copyright © 2014 Pearson Education, Inc.

mL Temperature (° C) __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________

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4. Drawing a Heating Curve for Water

5. Boiling point of water

________________________ °C

6. Temperature change (∆T)

________________________ °C

________________________ g 7. Mass of water Volume of water (1) × 1.00 g/mL 8. Joules needed to heat water (Show calculations.)

________________________ J

Energy and Matter

B. A Cooling Curve 1. Time (min) 0 1 2 3 4 5 __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________

Temperature (° C) __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________ __________

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2. Drawing the Cooling Curve for Salol

Questions and Problems Q1 What is the freezing point of Salol?

C. Heat of Fusion 1. Mass of calorimeter

___________________________ g

2. Mass of calorimeter + water

___________________________ g

3. Initial water temperature

___________________________°C

4. Final water temperature (ice added)

___________________________°C

5. Mass of calorimeter + water + melted ice ___________________________ g

Energy and Matter

59

Calculations 6. Mass of water (2 − 1)

____________________ g

7. Temperature change (4 − 3)

____________________ °C

8. Joules lost by water (Show calculations.)

____________________ J

9. Joules needed to melt ice = joules lost by water (8) ____________________ J 10. Mass of ice that melted (5 − 2) (Show calculations.)

____________________ g

11. Heat of fusion (joules to melt 1 g of ice) (Show calculations.)

____________________ J/g

Questions and Problems Q2 When water is heated, the temperature eventually reaches a constant value and forms a plateau on the graph. What does the plateau indicate?

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Q3 175 g of water was heated from 15 °C to 88 °C. How many kilojoules were absorbed by the water?

Q4 How many kilojoules are required at 0 °C to melt an ice cube with a mass of 25 g?

Q5 How many kilojoules are required to melt 15 g of ice at 0 °C, and raise the temperature of the liquid that forms to 85 °C?

Atoms and Elements 6 LABORATORY GOALS • • • •

Write the correct symbols or names of some elements. Describe some physical properties of the elements you observe. Categorize an element as a metal, metalloid, or nonmetal from its physical properties. Given the atomic symbol, determine its mass number; atomic number; and the number of protons, neutrons, and electrons.

LAB INFORMATION Time: Comments: Related Topics:

2h Obtain a periodic table as a reference. Tear out the report sheets and place them beside the procedures. Carefully observe the physical properties of the elements in the display. Names and symbols of the elements, periodic table, atoms, subatomic particles, isotopes, atomic mass

CHEMICAL CONCEPTS Elements and Symbol Primary substances, called elements, build all the materials about you. Chemical symbols are one- or two-letter abbreviations for the names of the elements. Only the first letter of an element’s symbol is capitalized. If the symbol has a second letter, it is lowercase so that we know when a different element is indicated. If two letters are capitalized, they represent the symbols of two different elements. For example, the element cobalt has the symbol Co, whereas the two capital letters CO specify two elements, carbon (C) and oxygen (O). One-Letter Symbols

Two-Letter Symbols

C S N I

Co Si Ne Ni

carbon sulfur nitrogen iodine

cobalt silicon neon nickel

Most of the elements have symbols that use letters from their current names. However, some symbols are derived from the ancient names of elements. For example, Na, the symbol for sodium, comes from the Latin word natrium. The symbol for iron, Fe, is derived from the Latin name ferrum. Table 6.1 lists the names and symbols of some common elements. Learning their names and symbols will greatly help your learning of chemistry.

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TABLE 6.1 Names and Symbols of Some Common Elements Name*

Symbol

Name*

Symbol

Name*

Symbol

Aluminum Argon Arsenic Barium Boron Bromine Cadmium Calcium Carbon Chlorine

Al Ar As Ba B Br Cd Ca C Cl

Au He H I Fe Pb Li Mg Mn Hg

Phosphorus Platinum Potassium (kalium) Radium Silicon Silver (argentum) Sodium (natrium) Strontium Sulfur Tin (stannum)

P Pt K Ra Si Ag Na Sr S Sn

Chromium Cobalt Copper (cuprum) Fluorine

Cr Co Cu F

Gold (aurum) Helium Hydrogen Iodine Iron (ferrum) Lead (plumbum) Lithium Magnesium Manganese Mercury (hydrargyrum) Neon Nickel Nitrogen Oxygen

Ne Ni N O

Titanium Uranium Zinc

Ti U Zn

*Names given in parentheses are ancient Latin or Greek words from which the symbols are derived.

Metals are elements that are usually shiny or have a metallic luster. They are usually good conductors of heat and electricity, ductile (can be drawn into a wire), and malleable (can be molded into a shape). Some metals such as sodium or calcium may have a white coating formed by reacting with oxygen in the air. If these are cut, you can see the fresh shiny metal underneath. In contrast, nonmetals are not good conductors of heat and electricity; are brittle (not ductile or malleable); and appear dull, not shiny.

The Periodic Table The periodic table, shown on the inside front cover of this lab manual and your textbook, contains information about each of the elements. On the table, the horizontal rows are periods and the vertical columns are groups. Each group contains elements that have similar physical and chemical properties. The groups are numbered across the top of the chart. Elements in Group 1A (1) are the alkali metals, elements in Group 2A (2) are the alkaline earth metals, and Group 7A (17) are the halogens (see Figure 6.1). Group 8A (18) are the noble gases, which are elements that are not very reactive compared to other elements. A dark zigzag line that looks like a staircase separates the metals on the left side from the nonmetals on the right side.

◄ FIGURE 6.1 Lithium (Li), sodium (Na), and potassium (K) are some alkali metals from Group 1A (1).

Atoms and Elements

63

Except for aluminum, the elements located along the heavy zigzag line are metalloids: B, Si, Ge, As, Sb, Te, Po, and At (see Figure 6.2). Metalloids exhibit some properties that are typical of the metals and other properties that are characteristic of the nonmetals. For example, metalloids are better conductors of heat and electricity than the nonmetals, but they are not as good conductors as the metals.

▲ FIGURE 6.2 The metalloids that border the heavy zigzag line have characteristics of both metals and nonmetals.

The Atom All the elements listed on the periodic table are made up of atoms. An atom is the smallest particle of an element. If you could divide a piece of aluminum foil into smaller and smaller pieces, you would eventually have a piece so small that you could not divide it further. Then you would have a single atom of aluminum (see Figure 6.3). There are different kinds of atoms for each of the elements.

▲ FIGURE 6.3 Aluminum foil consists of atoms of aluminum.

Atomic Number and Mass Number Atoms are made up of smaller bits of matter called subatomic particles. Protons are positively charged particles, electrons are negatively charged, and neutrons are neutral (no charge). In an atom, the protons and neutrons are tightly packed in the tiny center called the nucleus. Most of the atom is empty space, which contains fast-moving electrons. Electrons are so small that their mass is considered to be negligible compared to the mass of the proton or neutron. The atomic number is equal to the number of protons. The mass number of an atom is the number of protons plus the number of neutrons (see Figure 6.4). atomic number

=

number of protons (p + )

mass number

=

sum of the number of protons and neutrons (p + + n0 )

number of neutrons =

mass number (p + + n0 ) − atomic number (p + )

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◄ FIGURE 6.4 All atoms of carbon (atomic number 6) contain six protons in the nucleus, some neutrons, and six electrons outside the nucleus.

Table 6.2 illustrates these relationships between atomic number; mass number; and the number of protons, neutrons, and electrons in examples of single atoms for different elements. TABLE 6.2 Composition of Some Atoms of Different Elements Element

Symbol

Atomic Number

Mass Number

Number of Protons

Number of Neutrons

Number of Electrons

Hydrogen Nitrogen Oxygen Chlorine

H N O Cl

1 7 8 17

1 14 16 37

1 7 8 17

0 7 8 20

1 7 8 17

Isotopes and Atomic Mass Isotopes are atoms of the same element that differ in the number of neutrons. This means that isotopes of an element have the same number of protons, but different mass numbers. To distinguish between the different isotopes of an element, we write an atomic symbol for a particular isotope with its mass number in the upper left corner and its atomic number in the lower left corner (see Figure 6.5).

▲ FIGURE 6.5 The atomic symbol for an isotope of magnesium with 12 neutrons is Mg−24. All atoms of the element magnesium (Mg) have an atomic number of 12. Thus, every magnesium atom always has 12 protons. However, some of the magnesium atoms have 12 neutrons, others have 13 neutrons, and still others have 14 neutrons (see Figure 6.6). These different numbers of neutrons give the magnesium atoms different mass numbers but do not change their chemical behavior.

Atoms and Elements

65

◄ FIGURE 6.6 The nuclei of the three naturally occurring magnesium isotopes have the same number of protons, but different numbers of neutrons.

Calculating the Atomic Mass of an Element To calculate the atomic mass of an element, we need to know the percent abundance of each isotope and its mass, which are determined experimentally. For example, a sample of naturally occurring chlorine 35 37 Cl atoms and 24.24% of 17 Cl atoms. The atomic mass is a weighted average consists of 75.76% of 17 because it is calculated from the percent abundance of each isotope and its mass: the isotope mass of 34.97 amu, and the isotope

37 17 Cl

has a mass of 36.97 amu.

Isotope

Mass (amu) × Abundance (%)

35 17 Cl

= Contribution to Average Cl Atom

34.97

×

75.76 100

= 26.49 amu

37 17 Cl

36.97

×

24.24 100

= 8.962 amu

Atomic mass of Cl

= 35.45 amu (weighted average mass)

35 17 Cl

has a

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Physical Properties of Elements Materials: A display of elements 1. 2. 3.

Write the symbol and atomic number for each element listed. Observe the elements in the laboratory display. Describe their color and luster (shininess). From your observations, identify each element as a metal (M), metalloid (ML), or a nonmetal (NM).

B. The Periodic Table Materials: Periodic table, colored pencils, display of elements 1. 2. 3. 4. 5.

On the partial periodic table in the report sheet, write the symbols and atomic numbers (above the symbols) of the 14 elements you observed in part A. Write the period numbers on the left side of the table. Outline and label the columns that contain the alkali metals, alkaline earths, halogens, and noble gases. Outline and label the sections that contain the transition elements. Draw a dark, heavy line to separate the metals and nonmetals.

C. The Atom Complete the table for each of the neutral atoms with the symbol, atomic number, mass number, number of protons, neutrons, and electrons.

D. Isotopes and Atomic Mass 1. 2.

Complete the information for each of the isotopes of silver: the atomic symbol and the number of protons, neutrons, and electrons. Silver consists of two naturally occurring isotopes. Calculate the atomic mass for silver using the percent abundance of each of the isotopes and their isotopic masses.

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

6

1. Describe the periodic table.

2. Where are the alkali metals and the halogens located on the periodic table?

3. On the following list of elements, circle the symbols of the transition elements and underline the symbols of the halogens: Mg Cu Br Ag Ni Cl Fe F

4. Complete the list of names of elements and symbols: Name of Element

Symbol

Name of Element

Symbol

Potassium

Na

Sulfur

P

Nitrogen

Fe

Magnesium

Cl

Copper

Ag

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Date

Name

Section

Team

Instructor

REPORT SHEET

Atoms and Elements

LAB

6

A. Physical Properties of Elements Element

1. Symbol 1. Atomic 2. Color Number

2. Luster 3. Metal/Metalloid/ Nonmetal

Aluminum

_________ ________ _________ _________ ______________

Carbon

_________ ________ _________ _________ ______________

Copper

_________ ________ _________ _________ ______________

Iron

_________ ________ _________ _________ ______________

Magnesium _________ ________ _________ _________ ______________ Nickel

_________ ________ _________ _________ ______________

Nitrogen

_________ ________ _________ _________ ______________

Oxygen

_________ ________ _________ _________ ______________

Phosphorus

_________ ________ _________ _________ ______________

Silicon

_________ ________ _________ _________ ______________

Silver

_________ ________ _________ _________ ______________

Sulfur

_________ ________ _________ _________ ______________

Tin

_________ ________ _________ _________ ______________

Zinc

_________ ________ _________ _________ ______________

B. The Periodic Table

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Questions and Problems Q1 From their positions on the periodic table, categorize each of the following elements as a metal (M), a metalloid (ML), or a nonmetal (NM): Na ___________ S ____________ Cu _____________ F

_____________

Fe ___________ C ____________ As _____________ Ca _____________ Q2 Give the symbol of each of the following elements: a. noble gas in Period 2

____________

c. alkali metal in Period 3 ___________ e. alkali metal in Period 4 ___________

b. halogen in Period 2 ___________ d. halogen in Period 3 ___________ f. metalloid in Period 2 ___________

Q3 Complete the following table for the elements that are listed. If you have a full display of elements, check to see if your predictions are correct. Element

Metal/Metalloid/ Nonmetal

Prediction: Shiny/Dull

Correct? Yes/No

Chromium Gold Lead Cadmium Silicon

C. The Atom Name of Element

Symbol of Element

Atomic Number

Mass Number

Protons

Neutrons Electrons 30

Fe 27

13 19

Bromine

80 Au

197

20

Atoms and Elements

71

D. Isotopes and Atomic Mass 1.

Nuclear Symbol

Protons

Neutrons

Electrons

Nuclear Symbol

Isotopic Mass

Percent Abundance

107 47Ag

106.9

51.84%

109 47Ag

108.9

48.15%

107 47Ag 109 47Ag

2.

(Show calculations here)

Questions and Problems Q4 A neutral atom has a mass number of 80 and has 45 neutrons. Write its atomic symbol.

Q5 An atom has two more protons and two more electrons than the atom in Question 4. What is its atomic symbol?

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Electron Configuration and Periodic Properties 7 LABORATORY GOALS • • • • •

Describe the color of a flame produced by an element. Use the color of a flame to identify an element. Write the electron configuration for an element. Plot a graph of atomic radius versus atomic number. Interpret the trends in atomic radii within a family and a period.

LAB INFORMATION Time: Comments: Related Topics:

1½ h Obtain a periodic table or use the one on the inside cover of your textbook. Tear out the report sheets and place them beside the procedures. Electrons and protons, energy levels, and electron arrangement

CHEMICAL CONCEPTS Electron Energy Levels The chemical properties of an element strongly depends on the arrangement of the electrons. Every electron has a specific energy known as its energy level (n), starting with the lowest energy level n = 1 up to the much higher energy level n = 7. Electrons in the lower energy levels are usually closer to the nucleus, while electrons in the higher energy levels are farther away. Each of the energy levels consists of one or more sublevels, in which electrons with identical energy are found. The sublevels are identified by the letters s, p, d, and f (see Figure 7.1). For example, the first energy level (n = 1) has one sublevel, 1s. The second energy level (n = 2) has two sublevels, 2s and 2p. The third energy level (n = 3) has three sublevels, 3s, 3p, and 3d. The fourth energy level (n = 4) has four sublevels: 4s, 4p, 4d, and 4f. Within each energy level, the s sublevel has the lowest energy, the p sublevel has the next lowest energy, followed by the d sublevel and finally the f sublevel.

▲ FIGURE 7.1 The number of sublevels in an energy level is the same as the energy-level value of n.

A. Flame Tests When electrons absorb specific amounts of energy, they can attain higher energy levels. When they drop to lower, more stable energy levels, energy is emitted. If the energy released corresponds to the energy of visible light, the emission produces a color that can be detected by the human eye. When heated, many of the elements in Groups 1A (1) and 2A (2) produce colorful flames. Each element produces a characteristic color. Copyright © 2014 Pearson Education, Inc.

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B. Electron Configurations In an electron configuration, electrons are arranged in orbitals starting with the sublevels that have the lowest energy. The number of electrons in each sublevel is written as a superscript.

The electron arrangement of an element can be determined from its position in the periodic table. The electron configuration can be written by following the sublevel blocks across the periodic table, starting with period 1. The s block is formed by Groups 1A (1) and 2A (2). The p block includes the elements in Groups 3A (13) to 8A (18). The period number gives the particular energy level of each p sublevel, beginning with 2p (see Figure 7.2). Li 1s 22s1

O 1s 22s 22p 4

Ne 1s 22s 22p 6

Na 1s 22s 22p 63s1

S 1s 22s 22p 63s 23p 4

Ar 1s 22s 22p 63s 23p 6

On the periodic table, the 4s block fills next. K 1s 22s 22p 63s 23p 64s1

Ca 1s 22s 22p 63s 23p 64s 2

The d block begins with atomic number 21 and includes 10 transition metals. The energy level of each d block is one less than its period number. Sc 1s 22s 22p 63s 23p 64s 23d 1

Fe 1s 22s 22p 63s 23p 64s 23d 6

Ga 1s 22s 22p 63s 23p 64s 23d 104p1

As 1s 22s 22p 63s 23p 64s 23d 104p3

Br 1s 22s 22p 63s 23p 64s 23d 104p5

Kr 1s 22s 22p 63s 23p 64s 23d 104p 6

Zn 1s 22s 22p 63s 23p 64s 23d 10

The f block, which has a maximum of 14 electrons, follows the 6s block. The energy level of each f block is two less than the corresponding period number.

▲ FIGURE 7.2 Electron configuration follows the order of sublevels on the periodic table.

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75

C. Atomic Radius The size of an atom is determined by its atomic radius, which is the distance of the valence electrons from the nucleus. For each group of representative elements, the atomic size increases going from the top to the bottom because the outermost electrons in each energy level are farther from the nucleus. For example, in Group 1A (1), Li has a valence electron in energy level 2; Na has a valence electron in energy level 3; and K has a valence electron in energy level 4. This means that a K atom is larger than a Na atom, and a Na atom is larger than a Li atom. The atomic radius of representative elements is affected by the attractive forces of the protons in the nucleus on the electrons in the outermost level. For the elements going across a period, the increase in the number of protons in the nucleus increases the positive charge of the nucleus. As a result, the electrons are pulled closer to the nucleus, which means that the atomic sizes of representative elements decreases going from left to right across a period.

Periodic Properties The electron configurations of atoms are an important factor in the physical and chemical properties of the elements. Now we will look at properties of atoms, including the valence electrons, atomic size, ionization energy, and metallic character. Known as periodic properties, each type of property increases or decreases across a period, and then the trend is repeated again in each subsequent period. In an electron configuration, the electrons in the highest, or outermost, energy level are called the valence electrons. The valence electrons determine the chemical properties of the elements. The group numbers of the representative elements give the number of valence electrons in elements of each group (vertical column). For example, all the elements in Group 1A (1) have one valence electron in an s orbital. All the elements in Group 2A (2) have two (2) valence electrons in an s orbital. Group 3A (13) has three valence electrons in s and p orbitals, Group 4A (14) has four valence electrons, Group 5A (15) has five valence electrons, Group 6A (18) has six valence electrons, halogens in Group 7A (17) have seven valence electrons, and Group 8 (18) (except for He) has eight valence electrons. The similarities of behavior occur periodically as the number of valence electrons is repeated.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Flame Tests Materials: Bunsen burner, spot plate, flame-test (nichrome) wire, small graduated cylinder, 100-mL beaker, 1 M HCl, 0.1 M solutions (dropper bottles): CaCl2 , KCl, BaCl2 , SrCl2 , CuCl2 , NaCl, and unknown solutions

Obtain a spot plate and flame-test wire. Rinse the spot plate in distilled water. Pour 10 mL of 1 M HCl into a 100-mL beaker. CAUTION 1 M HCl is corrosive! Be careful when you use it. Wash off any HCl spills on the skin with tap water for 10 minutes. 1.

Place 6–8 drops of each test solution in separate indentations of the spot plate. Label the spot plate diagram in the laboratory report to match the types of solutions. Be careful not to mix the different solutions. Adjust the flame of a Bunsen burner until it is nearly colorless with a blue inner cone. Clean the test wire by dipping the loop in the HCl in the beaker and placing it in the flame of the Bunsen burner. (If you see a strong color in the flame while heating the wire, dip it in the HCl again. Repeat until the color is gone.)

Observing Flame Colors 2.

Dip the cleaned wire in one of the solutions on the spot plate. Make sure that a thin film of the solution adheres to the loop (see Figure 7.3). Move the loop of the wire to the tip of the inner blue flame. Record the color of each solution. Clean the wire in HCl after each test. Repeat the flame test with the other solutions. Note: The color of potassium in the KCl flame is short-lived. Be sure to observe the color of the flame from the KCl solution as soon as you put the wire in the flame.

▲ FIGURE 7.3 Using a flame-test wire to test for flame color.

Identifying an Unknown Solution Obtain one or more unknown solutions from your instructor and record their code letters. 3. Place 6–8 drops of each unknown solution in a clean spot plate. Repeat the flame-test procedure with each unknown solution and record the color of its flame. 4. From your list of colors for elements from step 2, identify the element in each of the unknown solution. You may wish to recheck the flame color of the known solution that best matches the flame color of an unknown. For example, if you think your unknown is KCl, recheck the color of the KCl solution to confirm.

Electron Configuration and Periodic Properties

77

B. Electron Configurations 1. 2. 3.

Write the electron configuration of each atom listed on the laboratory report. Indicate the number of valence electrons. Determine the group number for each element.

C. Atomic Radius The atomic radii for elements with atomic numbers 1–25 are listed in Table 7.1. On the graph, plot the atomic radius of each element versus the atomic number of that element. Be sure to connect all the points. Use the completed graph to answer questions in the report sheet. TABLE 7.1 Atomic Radii for the Elements with Atomic Numbers 1–25 Element

Symbol

Atomic Number

Atomic Radius (pm*)

first period hydrogen helium

H He

1 2

37 50

second period lithium beryllium boron carbon nitrogen oxygen fluorine neon

Li Be B C N O F Ne

3 4 5 6 7 8 9 10

152 111 88 77 70 66 64 70

third period sodium magnesium aluminum silicon phosphorus sulfur chlorine argon

Na Mg Al Si P S Cl Ar

11 12 13 14 15 16 17 18

186 160 143 117 110 104 99 94

fourth period potassium calcium scandium titanium vanadium chromium manganese

K Ca Sc Ti V Cr Mn

19 20 21 22 23 24 25

231 197 160 150 135 125 125

*(picometer = 10−12 m)

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Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

7

1. Why does a sodium street lamp give off a yellow color, whereas a neon light gives off a red color?

2. Describe the energy levels of electrons in an atom.

3. Why do some elements produce colorful flames?

4. How can you identify an unknown element using a flame test?

5. Write the electron configuration for each of the following: a. lithium

b. sodium

c. potassium

6. Why do the electron configurations in Question 5 all end with the same sublevel notation?

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Date

Name

Section

Team

Instructor

REPORT SHEET

Electron Configuration and Periodic Properties

LAB

7

A. Flame Tests 1.

2. Solution CaCl2

Element Ca

Color of Flame ________________________

KCl

K

________________________

BaCl2

Ba

________________________

SrCl2

Sr

________________________

CuCl2

Cu

________________________

NaCl

Na

________________________

3. Unknown Solution(s) Identification letter _______________

_______________

_______________

4. Element present _______________

_______________

_______________

Color of flame

Questions and Problems Q1 You are cooking spaghetti in water you have salted with NaCl. You notice that when the water boils over, it causes the flame of the gas burner to turn bright orange. How would you explain the appearance of a color in the flame?

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B. Electron Configurations Atom

1. Electron Configuration

2. Number of 3. Group Valence Electrons Number

O Na Ca Fe Zn Br Sr Cd Xe Cs Pb

Questions and Problems Q2 Complete the following: Number of sublevels in n = 3

Group number of carbon

Number of orbitals in the 2p sublevel Maximum number of electrons in 3d sublevel Maximum number of electrons in a 3p orbital

Sublevel being filled by element with atomic number 47 Sublevel that begins to fill after 4s 2 Number of valence electrons in As

Q3 Give the symbol of the element described by each of the following: First element that fills 3s sublevel

First element with five 3p electrons

Period 4 element in the same group as F

1s 22s 22p 63s 23p 64s 23d 104p3

Element with 3d 6

First element that completes n = 3

Element with a half-filled 5p level

Period 6 element in the same group as Mg

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83

C. Atomic Radius

Questions and Problems Q4 Describe the change in the atomic radii for the elements in Period 2, from lithium to neon.

Q5 Why does the change for the atomic radii of the elements in Period 3 from sodium to argon look similar to Period 2?

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Nuclear Chemistry 8 LABORATORY GOALS • Observe the use of a Geiger counter. • Determine the effect of shielding materials, distance, and time of radiation. • Complete a nuclear equation.

LAB INFORMATION Time: Comments: Related Topics:

2h Tear out the report sheets and place them beside the procedures. Follow your instructor’s directions for protection from radiation. Radioactivity, alpha particles, beta particles, gamma rays, shielding, nuclear decay, atomic symbols, nuclear equations

CHEMICAL CONCEPTS Radioactivity and Nuclear Equations Radioactivity occurs when a proton or neutron breaks down in the nucleus of an unstable atom or the particles in the nucleus are rearranged. Then a particle of energy called nuclear radiation is emitted from the nucleus. The nucleus has undergone nuclear decay. In nuclear reactions, the most typical kinds of radiation include alpha particles (α ), beta particles ( β ), positrons ( β + ), and gamma rays (γ ). alpha decay

147 143 4 62 Sm → 60 Nd + 2 He

alpha particle (α )

= 42 He

beta decay

40 20 Ca

beta particle ( β )

=

0 −1e

positron emission

49 25 Mn



positron ( β + )

=

0 +1e

gamma emission

167 68 Er

0 → 167 68 Er + 0 γ



40 0 21Sc + −1e 49 0 24 Cr + +1e

gamma ray (γ )

Radiation Levels from Radioactive Sources If radiation passes through the cells of the body, the cells may be damaged. You can protect yourself by using shielding materials, by limiting the amount of time near radiation sources, and by keeping a reasonable distance from the radioactive source (see Figure 8.1). In a process called radioactive decay, a nucleus spontaneously breaks down by emitting radiation or the particles described above. In a nuclear equation, the sum of the mass numbers and atomic numbers of the radioactive nucleus on the left of the arrow must equal the sum of the mass numbers and atomic numbers written on the right of the arrow (see Figure 8.2).

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◄ FIGURE 8.1 Heavy alpha particles are shielded by clothing and skin, beta particles are shielded by lab coats and gloves, whereas gamma rays require dense shielding such as lead.

To detect radiation, a device such as a Geiger counter is used. Radiation passes through the gas held within the tube, producing ion pairs. These charged particles emit bursts of electrical current that are converted to flashes of light and audible clicks on the Geiger counter.

▲ FIGURE 8.2 In alpha decay, the mass number of the new nucleus decreases by 4 and its atomic number decreases by 2.

Nuclear Chemistry

EXPERIMENTAL PROCEDURES

87

GOGGLES REQUIRED!

A. Nuclear Equations Complete and balance each of the nuclear equations in the report sheet.

B. Radiation Measurement (Experiments B, C, D, E, and F will be done as a demonstration.) Materials: Geiger counter The level of radiation that occurs naturally is called background radiation. Remove all sources of radiation near the counter. 1.

2.

Count the background radiation by operating the Geiger counter for 1 min. Record the number of counts. Measure the background count for two more 1-minute intervals. Record each. Total the number of counts in steps 1 and 2 and divide by 3. This value represents the average background level of radiation in counts/min (cpm). For the rest of this experiment, this background radiation is subtracted from other radiation counts to give the amount of radiation emitted from each radioactive source.

C. Radiation Levels from Radioactive Sources Materials: Geiger counter, meterstick, 3–4 radiation sources to test for radiation levels: Fiestaware™, minerals, old lantern mantles containing thorium compounds, camera lenses, old watches with radium-painted numbers on dials, smoke detectors containing Am-241, some foods such as table salt substitute (KCl), cream of tartar, instant tea, instant coffee, dry seaweed 1. 2. 3. 4. 5.

Place a radiation source at a distance of 10–20 cm from the detection tube. Record the type of radiation source. Count the radiation emitted for 1 min. Record. Enter the average background count (cpm) from part B2. Subtract the background count to obtain the radiation emitted by the source alone. Repeat steps 1–4 for other radiation sources.

D. Effect of Shielding on Radiation Level Materials: Geiger counter, meterstick, one of the stronger radioactive sources used in part B, samples of shielding materials: lead, paper, glass, cardboard, etc. 1.

2. 3. 4.

Place the radiation source at the same distance from the detection tube as in part B. Place one of the shielding materials between the Geiger counter detection tube and the radioactive source (see Figure 8.3). Record the type of shielding used. Record the counts from the Geiger counter for 1 min. Repeat steps 1 and 2 for other types of shielding. Enter the average background count (cpm) from part B2. For each type of shielding, subtract the background count to determine the radiation in counts/min.

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▲ FIGURE 8.3 Measuring radiation using shielding.

E. Effect of Time on Radiation Level The more time you spend near a radioactive source, the greater the amount of radiation you receive. 1.

2. 3.

Place one of the stronger radiation sources the same distance from the counter as in D1. Record the counts emitted for 1 min. Using the same radioactive source at the same distance as in part D1, record the number of counts emitted for 2 min. Using the same radioactive source at the same distance as in part D1, record the number of counts emitted for 3 min. Enter the average background count (cpm) from part B2. For 2 min and 3 min, enter a background that is two times (2×) and three times (3×) the background count at 1 min. For each amount of time (1 min, 2 min, 3 min) subtract background radiation to obtain the radiation count from the source.

F. Effect of Distance on Radiation Level By doubling your distance from a radioactive source, you receive one-fourth (1/4) the intensity of the radiation (see Figure 8.4). 1. 2. 3. 4. 5.

Place a strong radioactive source at a distance of 100 cm from the detection tube. Record the counts for 1 min. Repeat step 1 for distances of 75 cm, 50 cm, and 25 cm. Enter the average background count (cpm) from part B2. For each distance, subtract the background radiation to obtain the radiation for each source. Graph the radiation level (cpm) versus the distance (cm) of the radiation source. (Review the instructions for preparing a graph found in the preface).

▲ FIGURE 8.4 The effect of radiation lessens as the distance from the source increases.

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

8

1. In what part of the atom do alpha or beta particles originate?

2. Why is protection from radiation needed?

3. What are some medical uses of radiation?

4. Balance each of the following nuclear equations: a.

234 90 Th

⎯⎯ → ____________ +

0 −1e

→ ____________ + 42 He b. 227 87 Fr ⎯⎯

c.

63 29 Cu

+ ____________ ⎯⎯ →

64 29 Cu

+ 11H

5. Calculate each of the following involving half-lives: a. The half-life of Na-24 is 15 h. If the initial activity of the sodium-24 is 6.4 mCi, what is its activity after 60. h.

b. If a sample of I-123 with an activity of 4 kBq decays to an activity of 1 kBq in 26.4 h, what is its half-life?

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Date

Name

Section

Team

Instructor

REPORT SHEET

Nuclear Chemistry

LAB

8

A. Nuclear Equations Complete the nuclear equations by filling in the correct symbols: 131 53 I

⎯⎯ → −01e + ____________

96 40 Zr

+ ____________ ⎯⎯ → 10 n +

99 42 Mo

27 13 Al

24 + ____________ ⎯⎯ → 11 Na +

4 2 He

B. Radiation Measurement 1. Time (min)

Counts

1

______________

1

______________

1

______________

2. Average background count

Total counts ______________ counts/min (cpm) = 3

C. Radiation Levels from Radioactive Sources 1. Radiation Source

2. Counts/min (cpm)

3. Average Background (cpm)

4. Radiation from Source (cpm)

Questions and Problems Q1 Which item was the most radioactive?

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D. Effect of Shielding on Radiation Level 1. Shielding

2. Counts/min 3. Average 4. Radiation from Background (cpm) Source (cpm)

No shielding (C2)

Questions and Problems Q2 Which type of shielding provided the most protection from radiation? Why?

Q3 List shielding materials adequate for protection from the following: a. alpha particles b. beta particles c. gamma rays

E. Effect of Time on Radiation Level Time (min)

1. Counts

2. Average 3. Radiation from Background (cpm) Source (cpm)

1 2 3

F. Effect of Distance on Radiation Level 1. Distance

100 cm 75 cm 50 cm 25 cm

2. Counts/min 3. Average 4. Radiation from Background (cpm) Source (cpm)

Nuclear Chemistry

93

Q4 From the data in part F, calculate the ratio of radiation at 50 cm compared to radiation at 100 cm. counts/min at 50 cm Ratio = = counts/min at 100 cm Q5 Using your graph in part F, estimate the radiation levels at the following distances: Estimated cpm at 20 cm =

Estimated cpm at 40 cm = Estimated cpm at 60 cm = Estimated cpm at 80 cm = What happens to the amount of radiation as distance from the radiation source increases? Questions and Problems Q6 Write the symbols for the following types of radiation: a. alpha particle

b. beta particle

c. gamma ray

Q7 The bacteria in some foods are sterilized by placing the food near a source of ionizing radiation. Does that mean that the food becomes radioactive? Explain your answer.

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Compounds and Their Bonds 9 LABORATORY GOALS • • • • •

Write the electron-dot symbol for an atom and an ion. Write a correct formula and name of an ionic compound. Write a correct formula and name of an ionic compound containing a polyatomic ion. Write a correct formula and name of a molecular compound. Use the electron-dot formula of a compound to predict its shape and polarity.

LAB INFORMATION Time: Comments: Related Topics:

2–3 h Tear out the report sheets and place them beside the matching procedures. Ions, ionic bonds, electron-dot symbols, electron-dot formulas, ionic compounds, covalent bonds, molecular compounds

CHEMICAL CONCEPTS A. Ions: Transfer of Electrons When atoms of metals in Groups 1A (1), 2A (2), or 3A (13) react with atoms of nonmetals in Groups 5A (15), 6A (16), or 7A (17), the metals lose electrons and the nonmetals gain electrons in their valence shells. We can predict the number of electrons lost or gained by looking at the electron configurations of the atoms. For example, magnesium, which has an electron configuration 1s 22s 22p 63s 2 , has two valence electrons. It loses those two electrons to attain an octet with an electron arrangement of 1s 22s 22p 6 . The result is a magnesium ion with a charge of 2+. As a positive ion, it keeps the same name as the element.

When metals combine with nonmetals (5, 6, or 7 valence electrons), nonmetals gain electrons to obtain a stable electron configuration. For example, chlorine, which has an electron configuration 1s 22s 22p 63s 23p5 , has seven valence electrons. By gaining one valence electron, it becomes stable with an electron configuration of 1s 22s 22p 63s 23p 6 . The result is a negatively charged chloride ion with a charge of 1–. In the name of a binary compound with two different elements, the name of the negative ion ends in ide.

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B. Ionic Compounds and Formulas A compound consists of two or more different elements that are chemically combined. Most atoms combine by forming stable electron configurations. For example, the attractions between the positively charged Mg 2 + ions and the negatively charged Cl− ions are called ionic bonds.

The group number 1A–8A on the periodic table can be used to determine the ionic charges of the representative elements. Group number

1A (1) 2A (2)

3A (13)

4A (14)

5A (15)

6A (16) 7A (17)

8A (18)

Valence electrons

1 e−

2 e−

3 e−

4 e−

5 e−

6 e−

7 e−

8 e−

Electron change

lose 1

lose 2

lose 3

none

gain 3

gain 2

gain 1

no change

1+

2+

3+

none

3–

2–

1–

none

Ionic charge

To write an ionic formula, we use charge balance to determine the smallest number of positive and negative ions that give an overall charge of zero. For example, an overall charge of zero is obtained by using two Cl− ions to match the charge of the Mg 2 + ion. 1 (2+) + 2(−1) = 0

The number of each type of ion needed for charge balance gives the subscripts in the formula for the compound MgCl2 . (The subscript 1 for Mg is understood.) In any ionic formula, only the symbols and subscripts are written, not their ionic charges. Mg 2 + + 2 Cl− = MgCl2 = 0

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C. Metals in Ionic Compounds with Variable Charge Most of the transition metals can form more than one kind of positive ion. For example, iron forms two ions, Fe2+ and Fe3+ . To distinguish between the two ions, a Roman numeral that gives the ionic charge is written after the element name. The Roman numeral is always included within the parentheses in the names of compounds that form two or more positive ions. It is never included in the chemical formula. (see Table 9.1). TABLE 9.1 Some Ions of the Transition Elements Ion

Names

Compound

Names

Fe2+

Iron(II) ion

FeCl2

Iron(II) chloride

Fe3+

Iron(III) ion

FeCl3

Iron(III) chloride

Copper(I) ion

CuCl

Copper(I) chloride

Copper(II) ion

CuCl2

Copper(II) chloride

Cu

+

Cu 2+

Among the transition metals, a few elements (zinc, silver, and cadmium) form only a single type of ion. Thus, they are not variable and we do not need to write a Roman numeral as part of their names. Zn 2+ zinc ion

Ag + silver ion

Cd 2 + cadmium ion

On the other hand, there are two metals (tin and lead) in Group 4A (14) that have variable charges and require Roman numerals. Sn 2+ tin(II) ion and Sn 4+ tin(IV) ion

Pb 2+ lead(II) ion and Pb 4+ lead(IV)

D. Polyatomic Ions A compound that consists of three or more kinds of atoms will contain a polyatomic ion. A polyatomic ion is a group of covalently bonded atoms with an overall charge. That charge, which is usually negative, is the result of adding electrons to complete octets. Some examples of polyatomic ions are given in Table 9.2. The most common ion is named by replacing the ending of the name for the nonmetal with ate. The ite ending has one oxygen less than the most common form of the ion. Ammonium ion, NH 4+ , is positive because its group of atoms lost one electron. TABLE 9.2 Some Polyatomic Ions* Polyatomic Ion

Name

NH 4+

ammonium

OH −

hydroxide

NO 3 −

nitrate

ClO 3



chlorate

Polyatomic Ion with One Less Oxygen

Name

NO 2 −

nitrite

ClO 2



ClO− CO 3 2−

carbonate

HCO3−

hydrogen carbonate (bicarbonate)

SO 42 −

sulfate

HSO 4



PO4 3−

hypochlorite

SO32 − −

hydrogen sulfate (bisulfate)

HSO3

phosphate

PO33−

*Polyatomic ions and names in bold are the most common ions.

chlorite

sulfite hydrogen sulfite (bisulfite) phosphite

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To write a formula with a polyatomic ion, we determine the number of each type of ion needed for charge balance just as we did with the simple ions. When two or more polyatomic ions are needed, the formula of the polyatomic ion is enclosed in parentheses and the subscript placed outside. For example, the formula of calcium nitrate, which contains the ions Ca 2+ and NO3− , is Ca(NO3 ) 2 . Subscript for charge balance Ca(NO3 ) 2 Calcium nitrate Parentheses enclose polyatomic ion

E. Molecular Compounds: Sharing Electrons In a molecular compound, octets are achieved by sharing electrons between two nonmetals in Groups 4A (14), 5A (15), 6A (16), or 7A (17). For example, nitrogen in Group 5A (15) has five valence electrons, one electron pair and three single electrons. A molecule of NH3 forms when the three unpaired electrons are each shared with the unpaired electrons of three H atoms. The sharing of one pair of electrons is called a single bond.

▲ A molecule of NH3 contains three single bonds between the N atom and three H atoms. The electron-dot formula of a molecular compound is drawn by sharing the valence electrons until each atom has a complete octet. For example, in water (H 2 O), the O atom shares two unpaired electrons with two H atoms. The O atom now has an octet and both H atoms are stable with two valence electrons. Electron−Dot Formula for H2 O

Sometimes, more than one pair of electrons is shared between the central atom and its bonded atoms. For example, we use the following steps to draw the electron-dot formula for CO 2 , in which C is the central atom:

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1.

The arrangement of the atoms is O C O.

2.

The C atom has 4 valence electrons and each of the O atoms has 6 valence electrons. Thus there are 16 (4 + 6 + 6) valence electrons that can be used to form octets in CO 2 .

3.

Attach each O atom to the central C atoms using a pair of electrons.

4.

Place the remaining 12 electrons (16 – 4) around the C and O atoms.

Because the octet for the C atom is not complete, one electron pair from each O atom is shared with the C atom. This gives two double bonds in the CO 2 molecule by using 16 valence electrons.

Naming Molecular Compounds

Molecular compounds with two nonmetals are named by using prefixes, which state the number of atoms of each element in the compound (see Table 9.3). The first nonmetal is named by its element name; the second ends in ide. The first eight prefixes are mono (1), di (2), tri (3), tetra (4), penta (5), hexa (6), hepta (7), and octa (8). Usually the prefix mono is not shown, with the exception of carbon monoxide. TABLE 9.3 Some Formulas and Names of Molecular Compounds Formula

Name

CO

carbon monoxide

CO 2

carbon dioxide

PCl3

phosphorus trichloride

N 2O4

dinitrogen tetroxide (drop a in a double vowel)

SCl6

sulfur hexachloride

F. Electron-Dot Formulas and Shape The shape of a molecule or an ion can be described as linear, bent, trigonal planar, tetrahedral, or trigonal pyramidal. The valence shell electron-pair repulsion (VSEPR) model indicates that the bond angles in a molecule or ion are determined when the valence electrons in bonds and lone pairs are as far apart as possible. Counting the shared pairs and lone pairs determines the electron-group arrangement. Only the atoms bonded to the central atom determine the shape (see Table 9.4). Polar and Nonpolar Molecules

When the electronegativity difference is 0.5 to 1.8 between the central atom and an attached atom, the bond is polar. If the polar bonds in a molecule are symmetrical and the dipoles cancel, the molecule itself is nonpolar. When the dipoles do not cancel, the molecule or ion is polar.

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TABLE 9.4 Shapes for a Central Atom with Two, Three, and Four Bonded Atoms Electron Electron-Group Bonded Groups Arrangement Atoms

Lone Pairs

Bond Shape Angle

Example Three-Dimensional Model BeCl2

2

Linear

2

0

180°

Linear

3

Trigonal Planar

3

0

120°

Trigonal planar

BF3

2

1

120°

Bent

SO 2

4

0

109°

Tetrahedral

CH 4

3

1

109°

Trigonal pyramidal

NH3

2

2

109°

Bent

H2O

4

Tetrahedral

Compounds and Their Bonds

EXPERIMENTAL PROCEDURES

101

GOGGLES REQUIRED!

A. Ions: Transfer of Electrons For each atom and ion, complete the following in the table: 1. 2. 3. 4. 5. 6. 7.

electron configuration electron-dot symbol number of electrons lost or gained electron configuration for each ion ionic charge symbol, including charge name of the ion

B. Ionic Compounds and Formulas Materials: Display of compounds, Internet, reference manuals 1.

Physical properties

From the display of compounds, describe the appearance of sodium chloride, NaCl. Using the Internet or a reference manual, record the density and the melting points of sodium chloride. 2.

Formulas of ionic compounds Using the periodic table, write the positive and negative ions and the correct formula of each compound.

3.

Names of ionic compounds Using the formula of each ionic compound, write the positive and negative ions and the correct name.

C. Metals in Ionic Compounds with Variable Charge Materials: Display of compounds, Internet, reference manuals 1.

Physical properties

From the display of compounds, describe the appearance of iron(III) chloride, FeCl3 . Using the Internet or a reference manual, record the density and the melting points of iron(III) chloride. 2.

Formulas of ionic compounds From the name, write the positive and negative ions and the correct formula of each compound.

3.

Names of ionic compounds From the formula of each compound, write the positive and negative ions, and the name of the formula. Be sure to indicate the ionic charge as a Roman numeral if the transition metal has a variable valence.

D. Polyatomic Ions Materials: Display of compounds, Internet, reference manuals 1.

Physical properties

From the display of compounds, describe the appearance of potassium carbonate, K 2 CO3 . Using the Internet or a reference manual, record the density and the melting points of potassium carbonate.

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2.

Formulas of ionic compounds From the name, write the positive ion and negative polyatomic ion and the correct formula of each compound. Be sure to use parentheses when two or more polyatomic ions are needed.

3.

Names of ionic compounds From the formula of each ionic compound, write the positive and negative ions (polyatomic) and the name of the compound. Be sure to indicate the ionic charge as a Roman numeral if the transition metal has a variable valence.

E. Molecular Compounds: Sharing Electrons Materials: Display of compounds, Internet, reference manuals 1.

Physical properties

From the display of compounds, describe the appearance of water, H 2 O. Using the Internet or a reference manual, record the density and the melting points of water. 2.

Formulas of molecular compounds From the name, write the correct formula of the molecular compound using prefixes as subscripts.

3.

Names of molecular compounds Name each molecular compound, using prefixes to indicate two or more atoms of an element.

F. Electron-Dot Formulas and Shape Materials: Molecular model kit

Obtain a molecular model kit and build a model of each of the molecules in the report sheet, then complete the following for each: 1. 2. 3. 4. 5. 6.

Draw the electron-dot formula. Count the electron groups around the central atom. Use VSEPR to determine the electron-group arrangement. Count the number of atoms bonded to the central atom. Identify the molecular shape. Indicate if the molecules listed would be polar or nonpolar.

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

9

1. Where are the valence electrons in an atom?

2. How are positive and negative ions formed?

3. Why are electrons shared in molecular compounds?

4. How do the names of molecular compounds differ from the names of ionic compounds?

5. What are polyatomic ions?

6. How does the number of bonded atoms around a central atom determine its shape? Include examples in your answer.

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REPORT SHEET

LAB

Compounds and Their Bonds

9

A. Ions: Transfer of Electrons Element

5. Ionic 6. Symbol 7. Name 2. Electron- 3. Loss or 4. Electron Atomic 1. Electron of Ion Configuration Charge of Ion Gain of Dot Number ConfiguraElectrons of Ion tion of Atom Symbol

Sodium

11

Nitrogen

7

1s 22s 22 p 63s1

Na •

lose 1 e−

1s 22 s 22 p 6

Na +

1+

Sodium

Aluminum 13 Chlorine

17

Calcium

20

Oxygen

8

B. Ionic Compounds and Formulas 1. Physical properties Compound

Appearance

Density

Melting Point

Sodium chloride, NaCl 2. Formulas of ionic compounds Name

Positive Ion

Negative Ion

Formula

Lithium iodide

Li +

I−

LiI

Aluminum oxide Calcium sulfide Magnesium bromide Potassium nitride Sodium fluoride

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3. Names of ionic compounds Formula K 2S

Positive Ion K+

Negative Ion 2−

S

Name Potassium sulfide

BaF2 MgO Na 3 N AlCl3 Mg 3P2

C. Metals in Ionic Compounds with Variable Charge 1. Physical properties Compound Appearance Iron(III) chloride, FeCl3

Density

Melting Point

2. Formulas of ionic compounds Name

Positive Ion

Negative Ion

Iron(III) chloride

Fe3+

Cl−

Formula FeCl3

Negative ion

Name

Iron(II) oxide Copper(I) sulfide Copper(II) nitride Zinc oxide Silver sulfide 3. Names of ionic compounds Formula Cu 2S Fe 2O3 CuCl2 FeS Ag 2O FeBr2

Positive ion Cu

+

2−

S

Copper(I) sulfide

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D. Ionic Compounds with Polyatomic Ions 1. Physical properties Compound K 2CO3

Appearance

Density

Melting Point

2. Formulas of ionic compounds Name

Positive Ion

Negative Ion

Formula

Potassium carbonate

K+

CO32−

K 2CO3

Negative ion

Name

Sodium nitrate Calcium bicarbonate Chromium (III) hydroxide Lithium phosphate Potassium sulfate 3. Names of ionic compounds Formula

Positive ion Ca 2+

CaSO 4

SO 4

2−

Calcium sulfate

Al(NO3 )3 Na 2CO3 MgSO3 Cu(OH) 2 Mg 3 (PO 4 ) 2

E. Molecular Compounds 1. Physical properties Compound Water, H 2O

Appearance

Density

Melting Point

2. Formulas of molecular compounds Name

Formula

Name

Dinitrogen pentoxide

Dinitrogen trisulfide

Silicon tetrachloride

Oxygen difluoride

Phosphorus tribromide

Iodine heptafluoride

Formula

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3. Names of molecular compounds Formula ClF5

Name

Formula SF6

CS2

N 2 O3

PCl5

SeF6

Name

F. Electron-Dot Formulas and Shape Formula 1. Electron- Dot 2. Number of 3. Electron4. Number 5. Shape 6. Polar or Formula Electron Group of Bonded Nonpolar? Groups Arrangement Atoms

H2O

SF2

NI3

SiBr4

SO3

CO 2

Chemical Reactions and Equations 10 LABORATORY GOALS • • • •

Observe physical and chemical properties associated with chemical changes. Give evidence for the occurrence of a chemical reaction. Write a balanced equation for a chemical reaction. Identify a reaction as a combination, decomposition, replacement, or combustion reaction.

LAB INFORMATION Time: Comments:

Related Topics:

2–2½ h Read all the directions and safety instructions carefully. Match the labels on bottles and containers with the names of the substances. Label your containers with the formulas of the chemicals you place in them. A Bunsen burner is a potential hazard. Keep your work area clear of books, papers, backpacks, and other flammable items. Be sure that long hair is tied back. Tear out the report sheets and place them beside the matching procedures. Chemical change, chemical equation, balancing chemical equations

CHEMICAL CONCEPTS When a substance undergoes a physical change, it changes its appearance but not its composition. For example, when silver (Ag) melts and forms liquid silver (Ag), it undergoes a physical change from solid to liquid. In a chemical change, a substance is converted to one or more new substances with different properties. For example, when silver (Ag), which is a shiny substance, becomes tarnished it forms a dull-gray silver sulfide (Ag 2S), a new substance with different properties. Evidence of this and other chemical reactions is observed by the formation of bubbles, a solid, a change in color or a change in heat energy. Silver experienced the formation of a solid, the silver sulfide (Ag 2S). (see Table 10.1).

A chemical change produces new substances.

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TABLE 10.1 Evidence of Chemical Change Formation of a gas (bubbles) Formation of a solid (precipitates) Change in color Heat produced or absorbed

Chemical Equations In a chemical reaction, atoms in the reactants are rearranged to produce new combinations of atoms in the products. In an equation for a chemical reaction, the reactants are written on the left and the products on the right. An arrow between them indicates that a chemical reaction takes place. Reactants ⎯⎯ → Products

For example, we write the equation for the reaction of carbon and oxygen, which forms carbon dioxide, as Δ

C( s ) + O 2 (g ) ⎯⎯ → CO 2 (g )

The formula of each substance is followed by its physical state (s, l, g) in parentheses. If heat is required for the reaction, a triangle, which is a symbol for heat, is written over the arrow (see Table 10.2). TABLE 10.2 Some Symbols Used in Writing Equations Symbol

Meaning

+

Separates two or more formulas

⎯⎯ →

Reacts to form products

Δ

⎯⎯ →

Reactants are heated

(s)

Solid

(l )

Liquid

(g)

Gas or vapor

(aq )

Aqueous

Balancing Chemical Equations In a balanced equation, the total number of atoms of each element in the reactants is equal to the total number of atoms in the products. This balance is achieved by writing a coefficient in front of a formula containing that particular element. For example, we balance the equation for the reaction of hydrogen and oxygen as follows: H 2 ( g ) + O 2 (g ) ⎯⎯ → H 2 O(g )

Unbalanced equation

H 2 ( g ) + O 2 (g ) ⎯⎯ → 2H 2 O(g )

A coefficient of 2 in front of H 2 O balances the O atoms.

2H 2 ( g ) + O 2 (g ) ⎯⎯ → 2H 2 O(g )

A coefficient of 2 in front of H 2 balances the H atoms.

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111

If we count the number of atoms of H and O in the reactants, we find they are equal to the number of H and O atoms in the products. We say the equation is balanced.

Types of Reactions There are many different chemical reactions, but most can be classified into the types of reactions shown in Table 10.3. TABLE 10.3 Common Types of Chemical Reactions Type of Reaction

Description

Combination

Example Equation Δ

Elements or simple compounds form a more complex product. Decomposition A reacting substance is split into simpler substances. Single replacement One element takes the place of another element in a compound. Double replacement Elements in two compounds switch places.

Cu(s ) + S(s ) ⎯⎯ → CuS(s )

Combustion

CH 4 (g ) + 2O 2 (g ) ⎯⎯ → CO 2 (g ) + 2H 2 O(g )

Hydrocarbon fuel and oxygen form carbon dioxide and water.

Δ

CaCO3 (s ) ⎯⎯ → CaO(s ) + CO 2 (g )

Mg(s ) + 2HCl(aq ) ⎯⎯ → MgCl2 (aq ) + H 2 (g ) → AgNO3 (aq ) + NaCl(aq ) ⎯⎯ AgCl( s) + NaNO3 (aq ) Δ

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Magnesium and Oxygen Materials: Magnesium ribbon (2–3 cm long), tongs, Bunsen burner 1. 2.

3. 4.

Obtain a small strip (2–3 cm) of magnesium ribbon. Record its initial appearance. Using a pair of tongs to hold the end of the magnesium ribbon, ignite it by carefully placing it into the tip of the inner blue cone of the flame from the Bunsen burner. As soon as the magnesium ribbon ignites, remove it from the flame. Shield your eyes as the ribbon burns and do not look directly at the flame. Record your observations of chemical changes for the reaction. Balance the equation given for the reaction. Use 1 as a coefficient when one unit of that substance is required. Identify the type of reaction.

For experiments B–F, use small quantities. For solids, use the amount of compound that fits on the tip of a spatula or a small scoop. Place 3 mL of water in a test tube. Use this same volume as a reference volume when you need 3 mL of a solution. Do not dip droppers or stirring rods into reagent bottles. They may contaminate a reagent for the entire class. Discard unused chemicals as indicated by your instructor.

B. Zinc and Copper(II) Sulfate Materials: Two test tubes, test tube rack, 1 M CuSO 4 (aq ), Zn(s), small graduated cylinder 1.

2.

3. 4.

Pour 3 mL of 1 M CuSO 4 (aq) into a test tube by matching the reference volume. Obtain a small piece of zinc(s) metal. Describe the initial appearance of the CuSO 4 (aq ) solution and the zinc(s) metal. Add the Zn metal to the test tube containing the CuSO 4 (aq ) solution. Place the test tube in your test tube rack. After 30 min has elapsed, observe any color change of the CuSO 4 solution and the piece of Zn metal. Balance the equation given for the reaction. Identify the type of chemical reaction that has occurred.

Pour the CuSO 4 (aq ) solutions into a waste container or dispose of as directed by your instructor. Rinse the piece of zinc with water and place it in a recycling container or as directed by your instructor.

C. Reactions of Metals and HCl Materials: Three test tubes, test tube rack, small pieces of Zn(s), Cu(s), and Mg(s) metal, 1 M HCl Caution: HCl is a corrosive acid. Handle carefully! 1. 2.

3. 4.

Obtain a small piece each of zinc, copper, and magnesium. Describe the appearance of each. Place 3 mL of 1 M HCl (match your reference volume) in each of three test tubes. Carefully add a metal piece to the HCl solution in each of the three test tubes. Observe the reaction and record any evidence of chemical change. Carefully feel the test tubes to detect any change in the heat energy. Balance the equation for each metal that gave a chemical reaction. If there was no reaction, cross out the products and write NR. Identify the type of reaction for each chemical reaction that occurred.

Carefully pour off the acid and follow with large quantities of water to dilute. Dispose of as directed by your instructor. Rinse the metal pieces with water and place them in a recycling container or as directed by your instructor.

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113

D. Reactions of Ionic Compounds Materials: Four test tubes, test tube rack, 0.1 M solutions: CaCl2 (aq), Na 3 PO 4 (aq ), FeCl3 (aq), KSCN(aq) D1 Reaction of CaCl 2 and Na 3 PO4 1.

Place 3 mL each of 0.1 M CaCl2 (aq) and 0.1 M Na 3 PO 4 (aq) into separate test tubes. Record the appearance of each solution.

2. 3. 4.

Pour the contents of one test tube into the other and describe any evidence of a chemical reaction. Balance the equation. Identify the type of chemical reaction.

Dispose of the solutions and solids as directed by your instructor. D2 Reaction of FeCl 3 and KSCN 1.

Place 3 mL each of 0.1 M FeCl3 (aq) and 0.1 M KSCN(aq) into separate test tubes. Record the appearance of each solution.

2. 3. 4.

Pour the contents of one test tube into the other and describe any evidence of a chemical reaction. Balance the chemical equation. Identify the type of chemical reaction.

Dispose of the solution as directed by your instructor.

E. Sodium Carbonate and HCl Materials: Test tubes, test tube rack, 1 M HCl(aq), Na 2 CO3 ( s ), matches or wood splints 1. 2. 3. 4. 5.

Place 3 mL of 1 M HCl(aq) in a test tube. Record its appearance. Obtain a small amount of Na 2 CO3 ( s) (about the size of a pea). Record its appearance. Carefully add the Na 2 CO3 ( s) to the test tube containing 1 M HCl(aq). Record any evidence of a chemical reaction. Light a match or wood splint and insert it inside the neck of the test tube. Record whether the flame goes out or stays lighted. Balance the chemical equation. Identify the type of chemical reaction.

Dispose of the solutions and solids as directed by your instructor.

F. Hydrogen Peroxide Materials: Test tubes, test tube rack, 3% H 2 O 2 , 0.1 M KI(aq) 1.

Place 3 mL of 3% H 2 O 2 (aq ) in a test tube. Record its appearance.

2.

Place 3 mL of 0.1 M KI in a second test tube. Pour the KI solution (catalyst) into the first test tube. Record any evidence of a chemical reaction.

3.

Balance the chemical equation. Note that KI is included in the equation above the reaction arrow because a catalyst only speeds up a reaction, it does not change the quantities of reactants. Identify the type of chemical reaction.

4.

Dispose of the solutions as directed by your instructor.

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Pre-Lab Study Questions

10

1. Why are burning candles and rusting nails examples of chemical change?

2. What is included in a chemical equation?

3. How does a combination reaction differ from a decomposition reaction?

4. Balance each of the following reactions, and identify the type of reaction: Unbalanced Equation

Type of Reaction

Δ a. Al(s ) + Fe 2 O3 (s ) ⎯⎯ → Al2 O3 + Fe(l ) Δ b. KClO3 (s ) ⎯⎯ → KCl(s ) + O 2 (g ) Δ c. Li(s ) + Cl2 (g ) ⎯⎯ → LiCl(s ) Δ d. C2 H 4 (g ) + O 2 (g ) ⎯⎯ → CO 2 (g ) + H 2 O(g )

e. CrCl3 (aq ) + H 2S(g ) ⎯⎯ → Cr2S3 (s ) + HCl(aq ) 5. Complete and balance each of the following reactions: Reactants

Type of Reaction

a. K(s ) + N 2 (g ) ⎯⎯ →

Combination Δ

b. C5 H12 (g ) + O 2 (g ) ⎯⎯ →

Combustion

c. Al(s ) + CuSO 4 (aq ) ⎯⎯ →

Single Replacement

d. CoCl3 (aq ) + AgNO3 (aq ) ⎯⎯ →

Double Replacement

Δ

e. MgCO3 (s ) ⎯⎯ →

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REPORT SHEET

Chemical Reactions and Equations

LAB

10

A. Magnesium and Oxygen 1. Initial appearance of Mg

_________________________________________

2. Evidence of chemical reaction _________________________________________ _________________________________________ _________________________________________ 3. Balance:

_____ Mg(s) + _____ O 2 (g ) ⎯⎯ → ____ MgO(s)

4. Type of chemical reaction:

__________________________________________

B. Zinc and Copper(II) Sulfate Time

CuSO4 (aq ) Zn(s) Appearance Appearance

Evidence of a Chemical Reaction

1. initial 2. after 30 min → ___ Cu(s ) + ___ ZnSO4 (aq ) 3. ___ Zn(s ) + ___ CuSO 4 (aq ) ⎯⎯ 4. Type of chemical reaction: _______________________________________

C. Reactions of Metals and HCl Metal

1. Appearance of Metals

2. Evidence of a Chemical Reaction

Zn(s) Cu(s) Mg(s)

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3. ____ Zn(s ) + ____ HCl(aq)

⎯⎯ →

____ ZnCl2 (aq ) + ____ H 2 (g )

____ Cu(s) + ____ HCl(aq)

⎯⎯ →

____ CuCl2 (aq ) + ____ H 2 (g )

____ Mg(s) + ____ HCl(aq )

⎯⎯ →

____ MgCl2 (aq ) + ____ H 2 (g )

4. Type of chemical reaction: Zn __________________________________________________ Cu __________________________________________________ Mg __________________________________________________

D. Reactions of Ionic Compounds D1 Reaction of CaCl 2 and Na 3 PO4 Reactants

1. Appearance of Solutions

2. Evidence of a Chemical Reaction

CaCl2 (aq ) Na 3PO 4 (aq) 3. ____ CaCl2 (aq) + ____ Na 3PO 4 (aq ) ⎯⎯ → ____ Ca 3 (PO 4 ) 2 (s ) + ____ NaCl(aq) 4. Type of reaction: __________________________________________________ D2 Reaction of FeCl 3 and KSCN Reactants

1. Appearance of Solutions

2. Evidence of a Chemical Reaction

FeCl3 (aq)

KSCN(aq) 3. ____ FeCl3 (aq ) + ____ KSCN(aq) ⎯⎯ → ____ Fe(SCN)3 (aq) + ____ KCl(aq) 4. Type of reaction: _____________________________________________________

E. Sodium Carbonate and HCl Reactants

HCl(aq ) Na 2CO3 ( s )

1. Appearance of Reactants

2. Evidence of a Chemical Reaction

Chemical Reactions and Equations

119

3. Observation of burning match or splint ____________________________________

What caused the change in the burning match or splint?

→ CO2 (g ) + ____ H 2O(l ) + ____ NaCl(aq) 4. ____ Na 2CO3 ( s ) + ____ HCl(aq) ____ ⎯⎯

5. Type of reaction: _____________________________________________________

F. Hydrogen Peroxide Reactants

1. Appearance of Reactants

2. Evidence of a Chemical Reaction

H 2O 2 (aq) KI 3. ____ H 2O 2 (aq) ⎯⎯→ ____ H 2O(l ) + ____ O 2 (g )

4. Type of chemical reaction: ______________________________________________ Questions and Problems Q1 What evidence of a chemical reaction might you see in the following cases? Refer to Table 10.1. a. dropping an Alka-Seltzer tablet into a glass of water

________________________________________________________________ b. bleaching a stain

________________________________________________________________ c. burning a match

________________________________________________________________ d. rusting of an iron nail

________________________________________________________________ Q2 Balance the following equations: a. _____ Mg(s ) + _____ HCl(aq )

⎯⎯ → _____ H 2 (g ) + _____ MgCl2 (aq)

b. _____ Al(s ) + _____ O 2 (g )

⎯⎯ → _____ Al2O3 (s )

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c. _____ Fe2O3 (s ) + _____ H 2 O(l )

⎯⎯ → _____ Fe(OH)3 (s )

→ _____ Ca(NO3 )2 (aq ) + _____ H 2 O(l ) d. _____ Ca(OH)2 (aq) + _____ HNO3 (aq) ⎯⎯

Q3 Write a balanced equation for each of the following reactions. Write the correct formulas of the reactants and products and the states of each. a. Liquid pentane (C5 H12 ) and oxygen (O 2 ) gas react to form carbon dioxide and water.

b. Sodium and water react to form sodium hydroxide and hydrogen gas (H 2 ).

c. Iron and oxygen (O2 ) gas react to form iron(III) oxide. Q4 Classify each reaction as combination, decomposition, single replacement, double replacement, or combustion. a. Ni(s ) + F2 (g ) ⎯⎯ → NiF2 (s )

__________________

b. Fe2 O3 (s ) + 3C(s ) ⎯⎯ → 2Fe(s ) + 3CO(g )

__________________

c. CaCO3 (s ) ⎯⎯ → CaO(s ) + CO 2 (g )

__________________

d. H 2SO 4 (aq ) + 2KOH(aq) ⎯⎯ → K 2SO 4 (aq) + 2H 2O(l ) __________________ e. C2 H 4 (g ) + 3O 2 (g ) ⎯⎯ → 2CO 2 (g ) + 2H 2O(g )

__________________

Q5 Complete and balance each of the following chemical reactions by writing the correct formulas of the product(s) that form: Type of Reaction a. Zn(s ) + CuBr2 (aq )

⎯⎯ → __________ + __________ single displacement

b. H 2 (g ) + Cl2 (g )

⎯⎯ → ______________

c. MgCO3 (s )

⎯⎯ → __________ + __________ decomposition

combination

d. KCl(aq) + AgNO3 (aq) ⎯⎯ → _________ + _________ double displacement

Moles and Chemical Formulas 11 LABORATORY GOALS • Determine the simplest formula of a compound. • Calculate the percent water in a hydrate. • Determine the formula of a hydrate.

LAB INFORMATION Time: Comments:

Related Topics:

2–2½ h Tear out the report sheets and place them beside the matching procedures. Use steel wool to remove any coating on the magnesium ribbon until it is shiny. Check the crucible for cracks before you start to heat it. When you set a hot object aside to cool, remember that it is hot. Dispose of all chemicals as directed by your instructor. Formulas, moles, molar mass, calculating moles from grams, calculating grams from moles, hydrates, dehydration

CHEMICAL CONCEPTS A. Finding the Simplest Formula The simplest formula of a compound is the lowest whole-number ratio of the atoms in the formula. For example, the compound benzene, with molecular formula C6 H 6 , has the simplest formula CH. Some molecular formulas and their simplest formulas are shown in Table 11.1. TABLE 11.1 Examples of Molecular and Empirical Formulas Name

Molecular Formula

Simplest Formula

Acetylene

C2 H 2

CH

Benzene

C6 H 6

CH

Ammonia

NH3

NH3

Hydrazine

N2H4

NH 2

The simplest formula of a compound is determined by converting the number of grams of each element to moles and finding the lowest whole-number ratio to use as subscripts. For example, in an experiment it was determined that 0.040 mole of Zn had combined with 0.080 mole of Cl to form a compound. To calculate the simplest formula we proceed as follows: 1.

Divide the moles of each element by the smaller number of moles (0.040) and round to the nearest whole number.

0.080 mole Cl = 2 moles of Cl 0.040

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0.040 mole Zn = 1 mole of Zn 0.040

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Use the whole numbers as subscripts to write the formula of the compound. ZnCl2 (The subscript 1 for Zn is understood.)

B. Formula of a Hydrate A hydrate is an ionic compound that is combined with a specific number of water molecules. The number of water molecules is fixed for each hydrate, but differs from one hydrate to another. The number of water molecules is written after the ionic formula and separated by a large, raised dot. CaSO 4 • 2H 2 O

CuSO 4 • 5H 2 O

Na 2 CO3 •10H 2 O

The water molecules in the hydrate can be removed by heating. When all the water is removed, the remaining ionic compound is called an anhydrate. For example, when one mole of copper(II) sulfate pentahydrate is heated, five moles of H 2 O are removed. The water removed is also called the water of hydration. CuSO 4 • 5H 2 O hydrate

Δ

⎯⎯ →

CuSO 4 + anhydrate

5H 2 O(g ) water of hydration

The amount of water in a hydrate is experimentally determined by measuring the mass of the hydrate before heating and the mass of the anhydrate after heating. The difference in mass is due to the water of hydration that is lost. The percent water is calculated by dividing the grams of water by the mass of the hydrate and multiplying by 100%. For example, if 2.00 g of CuSO 4 i 5H 2 O is heated and the mass of the anhydrate CuSO 4 is 1.28 g, we would calculate the grams of H 2 O as the difference. 2.00 g of hydrate –1.28 g of anhydrate = 0.72 g of H 2 O in hydrate

Then the percent water in the hydrate is calculated as 0.72 g H 2 O × 100% = 36% H 2 O in hydrate 2.00 g CuSO 4

Moles and Chemical Formulas 123

EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Finding the Simplest Formula Materials: Crucible, crucible cover, crucible tongs, clay triangle, iron ring and stand, Bunsen burner, magnesium ribbon, steel wool, eyedropper, small 100- or 150-mL beaker, heat-resistant pad

In this experiment, you will heat magnesium so that it reacts with the oxygen (O 2 ) in the air and forms an oxide. The difference between the mass of the oxide compound and the initial mass of the magnesium is the mass of oxygen that combined with magnesium. When the moles of the magnesium and the oxygen are calculated, the simplest formula can be determined. 1.

Obtain a clean, dry crucible and its cover. Set the crucible, and cover, slightly offset, on a clay triangle and place on an iron ring attached to a ring stand (see Figure 11.1). Heat the crucible and cover for about one minute. Cool until they are at room temperature. Using crucible tongs, carry the crucible and cover to the balance. Do not place hot objects on a balance pan. Weigh the crucible and cover and record the mass.

◄ FIGURE 11.1 A crucible and cover, slightly offset, are heated on a clay triangle.

2.

Obtain a piece of magnesium ribbon that has a mass of 0.15–0.30 g. If there is tarnish on the ribbon, remove it by polishing the ribbon with steel wool. Describe the appearance of the magnesium ribbon after polishing.

3.

Twist the ribbon into a coil and place it at the bottom of the crucible. Weigh the crucible, cover, and magnesium ribbon and record the mass.

Heating the Magnesium Ribbon Do this part of the experiment in a fume hood:

• Place the crucible with the magnesium ribbon on the clay triangle. Keep the cover and a pair of tongs nearby. • Begin to heat the crucible making sure that the tip of the inner blue flame touches the bottom of the crucible. The bottom of the crucible will become red hot. Watch for smoke or fumes, which indicate that the magnesium and oxygen are reacting. • As soon as the magnesium bursts into flame, use the tongs to place the cover on the crucible. The cover should be slightly offset to allow oxygen to react with the magnesium. Caution: Avoid looking directly at the bright flame of the burning magnesium. • When the magnesium no longer produces smoke or a flame, remove the cover and set it on a heat-resistant surface.

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• Continue to heat the crucible strongly for another five minutes. Then turn off the burner and allow the crucible and its contents to cool to room temperature. During heating, some magnesium reacts with nitrogen in the air to form magnesium nitride. 3Mg(s ) + N 2 ( g )

Δ

⎯⎯ →

Mg3 N 2 ( s )

To remove this nitride product, carefully add 15–20 drops of water to the cooled contents. Mg3 N 2 ( s ) + 3H 2 O(l )

Δ

⎯⎯ →

3MgO( s ) + 2NH3 ( g )

Caution: Avoid breathing fumes from the crucible because ammonia may be released. 4.

Cover the crucible and heat gently for five minutes to drive off any excess water. Then heat strongly for five minutes. Allow the crucible to cool completely. Reweigh the crucible, cover, and oxide contents and record the mass. Remove the solid in the crucible and dispose of it as directed by your instructor.

Calculations 5. 6. 7. 8.

Determine the mass of the magnesium (3 – 1). Calculate the mass of the magnesium compound (4 – 1). Calculate the mass of oxygen that combined with the magnesium (6 – 5). Determine the number of moles of magnesium by dividing the mass of magnesium (5) by its molar mass. moles of Mg = g Mg ×

9.

Determine the number of moles of oxygen by dividing the mass of the oxygen (7) by its molar mass. moles of O = g O ×

10. 11.

1 mole Mg 24.31 g Mg

1 mole O 16.00 g O

Divide the moles of Mg (8) and the moles of O (9) by the smaller number of moles. Round each of the results to the nearest whole number. Using the whole number values obtained in 10 as subscripts, write the simplest formula of the magnesium compound.

B. Formula of a Hydrate Materials: Crucible, clay triangle, crucible tongs, hydrate of MgSO 4 , iron ring and stand, Bunsen burner, heat-resistant pad, laboratory balance 1.

Obtain a clean, dry crucible and its cover. Set the crucible and cover, slightly offset, on a clay triangle and place on an iron ring attached to a ring stand (see Figure 11.1). Heat the crucible and cover for about one minute. Cool until they are at room temperature. Using crucible tongs, carry the crucible and cover to the balance. Do not place hot objects on a balance pan. Weigh the crucible and cover and record the mass.

2.

Fill the crucible about 1/3 full with the hydrate of MgSO 4 . Weigh the crucible with the MgSO 4 hydrate and record the mass.

3.

Set the crucible and hydrate on a clay triangle that is set on an iron ring (see Figure 11.1). The cover should be slightly offset so that water vapor can escape. Heat gently for five minutes; increase the intensity of the flame and heat strongly for another 10 minutes. The bottom of the crucible should become a dull red color. Turn off the burner. Allow the crucible to cool to room temperature.

Moles and Chemical Formulas 125

3a.

Caution: Allow heated items to cool to room temperature. Do not place a hot object on the balance pan. Weigh the crucible and its contents. Record the mass. Repeat step 3. If more water has been lost, use this final mass in the calculations. Remove the solid in the crucible and dispose of it as directed by your instructor.

Calculations 4. 5. 6.

Calculate the mass of the hydrate (2 – 1) Calculate the mass of the anhydrate after heating (3 – 1) Calculate the mass of water lost from the hydrate sample (4 – 5)

7.

Calculate the percent H 2 O in the hydrate by dividing the mass of H 2 O lost (6) by the mass of the hydrate (4) and multiply by 100%. Percent water in hydrate =

8.

g water × 100% g hydrate

Calculate the moles of H 2 O in the hydrate by dividing the mass of the H 2 O (6) by its molar mass. Moles of water = g water ×

9.

1 mole water 18.02 g water

Calculate the moles of anhydrate by dividing the mass of the anhydrate by its molar mass. Moles of anhydrate = g anhydrate ×

10.

To determine the ratio of moles of water to 1 mole of anhydrate, divide the moles of water (8) by the moles of anhydrate (9). Round off the value for moles of H 2 O to the nearest whole number. ______ moles of water _____ moles of MgSO 4

11.

1 mole anhydrate 120.4 g anhydrate

=

______ moles of H 2 O 1 mole of MgSO 4

Complete the formula of the hydrate by writing in the number of moles of water for each mole of anhydrate.

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Date

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Pre-Lab Study Questions

11

1. What is meant by the simplest formula of a compound?

2. How does a hydrate differ from an anhydrate?

3. What happens when a hydrate is heated?

4. A hydrate of CoCl2 with a mass of 6.00 g is heated strongly. After cooling, the mass of the anhydrate is 3.27 g. a. How many grams of H 2 O were lost from the hydrate?

b. What is the % water in the hydrate?

c. What is the formula of the CoCl2 hydrate?

d. Write the equation for the dehydration of the CoCl2 hydrate.

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Date

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REPORT SHEET

Moles and Chemical Formulas

LAB

11

A. Finding the Simplest Formula 1. Mass of empty crucible + cover 2. Initial appearance of the magnesium 3. Mass of crucible + cover + magnesium 4. Mass of crucible + cover + oxide product

_______________________ g _______________________ _______________________ g _______________________ g

Calculations 5. Mass of magnesium 6. Mass of magnesium compound

_______________________ g _______________________ g _______________________ g _______________________ mole

7. Mass of oxygen in the product 8. Moles of Mg (Show calculations.)

_______________________ mole

9. Moles of O (Show calculations.)

10. Which number of moles (Mg or O) is smaller _______________________ moles of Mg = ______________ moles of Mg (rounded to a whole number)

moles O

Formula:

Mg

O

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= ______________ moles of O (rounded to a whole number)

subscripts

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Questions and Problems Q1 Using the rules for writing the formulas of ionic compounds, write the ions and the correct formula for magnesium oxide.

Q2 Write a balanced equation for the reaction of the magnesium and the oxygen (O 2 ), including their physical states.

Q3 Calculate the simplest formula for each of the following compounds: a. 0.200 mole of Al and 0.600 mole of Cl

b. 0.080 mole of Ba, 0.080 mole of S, 0.320 mole of O

Q4 When 2.50 g of copper reacts with oxygen, the copper oxide compound has a mass of 2.81 g. What is the simplest formula of the copper oxide?

B. Formula of a Hydrate 1. Mass of crucible

____________________________ g

2. Mass of crucible and hydrate

____________________________ g ____________________________ g

3. Mass of crucible and anhydrate 3a. Mass of crucible and anhydrate (second heating)

____________________________ g

Calculations

5. Mass of anhydrate

____________________________ g ____________________________ g

6. Mass of water lost

____________________________ g

4. Mass of hydrate

Moles and Chemical Formulas 131

7. Percent water (Show calculations.)

____________________________ %

8. Moles of water (Show calculations.)

____________________________ moles

9. Moles of salt (anhydrate) (Show calculations.)

____________________________ moles

10. Ratio of moles of water to moles of hydrate ____________________________ (Show calculations.)

11. Formula of hydrate

MgSO 4 •

H 2O

Questions and Problems Q5 Using the formula you obtained in B11, write a balanced equation for the dehydration of the MgSO 4 hydrate used in the experiment.

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Gas Laws 12 LABORATORY GOALS • • • •

Graph the relationship between the pressure and volume of a gas (Boyle’s law). Observe the effect of changes in temperature on the volume of a gas (Charles’s law). Graph the data for volume and temperature of a gas (Charles’s law). Use a graph of Charles’s law to predict the value of absolute zero.

LAB INFORMATION Time: Comments: Related Topics:

2h Be careful when you work around boiling water. Tear out the report sheets and place them beside the matching procedures. Pressure, volume, Boyle’s law, Charles’s law, combined gas law, Kelvin temperature

CHEMICAL CONCEPTS Pressure The pressure of a gas depends upon the number of molecules hitting the walls of a container and their force. You increase pressure when you add air to a car or bicycle tire or basketball or when you blow up a balloon. Airplanes must be pressurized so that you breathe sufficient oxygen. Scuba divers require increased air pressures in their air tanks while diving because the pressure on their bodies increases as they dive deeper. In medicine, blood pressure is measured to determine the force of the blood against the blood vessels. Common units for measuring pressure are atmospheres, millimeters of mercury, and torrs (see Table 12.1). TABLE 12.1 Units for Measuring Pressure Unit

Abbreviation

Unit Equivalent to 1 atm

Atmosphere

atm

1 atm (exact)

Millimeters of Hg

mmHg

760 mmHg (exact)

Torr

torr

760 torr (exact)

A. Boyle’s Law According to Boyle’s law, the pressure (P) of a gas varies inversely with the volume (V) of the gas, when the temperature (T) and the number of moles (n) are kept constant. Thus, the volume decreases when the pressure increases or the volume increases when the pressure decreases. For example, if the pressure of a gas is doubled, the corresponding volume is one-half its previous value.

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134

Laboratory Manual for General, Organic, and Biological Chemistry Boyle’s Law: As the volume of a gas decreases, the pressure of that gas increases.

Another way of stating Boyle’s law is that the PV product of a gas remains constant as long as the temperature and number of moles do not change. Therefore, we can write Boyle’s law as an equality of the initial pressure and volume (PV 1 1 ) with the final pressure and volume (P2V2 ). PV 1 1 = P2V2 (T, n constant)

Boyle’s Law

B. Charles’s Law In cold weather, the tires on a car or a bicycle seem to go flat. In hot weather, a full tire may burst. Such examples illustrate that changes in volume are related to changes in temperature. When the temperature increases, the kinetic energy of the molecules increases. To keep the pressure constant, the volume must expand. When the molecules slow down as the temperature decreases, the volume must decrease. According to Charles’s law, the volume of a gas changes directly with the Kelvin temperature as long as the pressure and number of moles remain constant. Charles’s Law: When the temperature (K) of a gas increases, the volume of that gas also increases to maintain constant pressure.

According to Charles’s law, the V/T ratio of a gas remains constant as long as the pressure and number of moles do not change. Therefore, we can write Charles’s law as an equality of the initial volume and temperature (V1/T1 ) with the final volume and temperature (V2 /T2 ). V1 V2 = T1 T2

(P, n constant)

Charles’s Law

Absolute Zero Absolute zero is the theoretical value for the coldest temperature that matter can attain. On the Kelvin scale, this would be 0 K, which is equivalent to −273 °C. All temperatures used in comparisons or calculations of gases must use the Kelvin temperature scale.

Gas Laws

EXPERIMENTAL PROCEDURES

135

GOGGLES REQUIRED!

A. Boyle’s Law In an experiment, the volume of a specific amount of gas was measured at different pressures while the temperature was kept constant. You will work with these experimental results, which are given in the report sheet table. Review the instructions for Graphing Experimental Data in the preface of this laboratory manual. 1. 2.

Calculate the P × V product for each sample. Remember to round off each result to the correct number of significant figures. Plot the pressure (mmHg) of the gas in each reading against the volume (mL). The data points will form a slight curve, not a straight line. Draw a smooth line through the data points you have plotted. Use the graph to answer Q1–Q3.

B. Charles’s Law Materials: 125-mL Erlenmeyer flask, 400-mL beaker, one-hole rubber stopper with a short piece of glass tubing inserted and attached to a piece of rubber tubing, water containers, thermometer, pinch clamp, buret clamp, Bunsen burner (or a hot plate), wire gauze, iron ring and stand, a graduated cylinder, boiling chips, ice In this experiment, the volume of a gas will be measured at different temperatures. Using this data, a graph of the relationship between volume and temperature will be prepared. The value of absolute zero is predicted by extending the graph line (extrapolating) to the axis where the volume of the gas would decrease to 0 mL.

Preparation for Experiment Set a 400-mL beaker on a hot plate or on an iron ring covered by wire gauze. Obtain a 125-mL Erlenmeyer flask and make sure the flask is dry inside. Place a one-hole rubber stopper and tubing in the neck of the flask. Place a buret clamp on the neck of the flask, and fasten the buret clamp to a ring stand (see Figure 12.1). Place the flask inside the beaker but do not touch the bottom of the beaker. Pour water into the beaker until it comes up to the neck of the flask, which leaves a little space at the top so the water will not boil over. Add a few boiling chips to the water.

◄ FIGURE 12.1 The air in an Erlenmeyer flask is heated in a boiling water bath.

1.

Begin heating and bring the water to a boil. Boil gently for 10 min to bring the temperature of the air in the flask to that of the boiling water. Measure the Celsius temperature of the boiling water. Convert from degrees Celsius to the corresponding Kelvin temperature.

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Turn off the burner and immediately place a pinch clamp on the rubber tubing. A pinch clamp closes off the rubber tubing

Undo the buret clamp from the ring stand. Using the buret clamp as a handle, carefully lift the flask out of the hot water and carry it over to one of the containers of cool water. Keeping the stopper end of the flask pointed downward, submerge the flask in the cool water (see Figure 12.2). Keeping the stopped end of the flask pointed downward, submerge the flask and remove the pinch clamp. Water will enter the flask as the air in the flask cools and decreases in volume.

◄ FIGURE 12.2 The air space in a heated flask decreases when placed in cool water.

Keep the flask submerged in the cool water for at least 10 min. We assume that the temperature of the air remaining in the flask is the same temperature as the cool water outside the flask. Measure the Celsius temperature of the cool water in the water bath. Convert from degrees Celsius to the corresponding Kelvin temperature. 3.

Keeping the flask upside down, carefully raise or lower the flask until the water level inside the flask is equal to the water level in the container (see Figure 12.3). This process equalizes the pressure inside the flask with the atmospheric pressure.

◄ FIGURE 12.3 The water levels inside and outside the submerged flask are equalized to give the same pressure.

While keeping the water levels equal, close the rubber tubing with the pinch clamp. Remove the flask from the water, then set it upright on the lab bench and remove the rubber stopper. Measure and record the volume of water (mL) that entered the flask. 4.

Fill the flask to the top with water and reinsert the rubber stopper. Some water will spill out. Remove the stopper and pour the water that filled the flask into a graduated cylinder. Measure and record the volume of water in the flask. This volume of water is equal to the volume of air initially in the flask.

5.

Repeat the process again, submerging the flask in a container of water with a different temperature. Or, if your instructor permits, obtain the data for other temperatures and volumes of water from other students in your class.

Calculations 6. 7. 8.

Calculate the volume of air in the flask for each temperature by subtracting the amount of water (3) that entered the flask at that temperature from the gas volume (4). For each gas sample, calculate the V/T value. According to Charles’s law, these ratios should be constant. Plot the points for each Kelvin temperature (1) and volume (4). Draw a straight line through (or close to) the points that you plotted on the graph. Extend the line to the horizontal axis where the volume is 0 mL. Theoretically, gases reach a temperature called absolute zero at a volume of 0 mL. Record the temperature where the graph line crosses the horizontal axis.

Date

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Pre-Lab Study Questions

12

1. Why is an airplane pressurized?

2. Why does a scuba diver need increased gas pressure in the air tank?

3. How does temperature affect the kinetic energy of gas molecules?

4. A balloon is filled with 652 mL of helium at a pressure of 1.00 atm. What is the new volume, in milliliters, if the pressure decreases to 738 mmHg, with T and n constant?

5. A sample of neon gas occupies 75.0 mL at 25 °C. What is its new volume, in milliliters, if the temperature decreases to −12 °C, with P and n constant?

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REPORT SHEET

LAB

Gas Laws

12

A. Boyle’s Law 1. Reading

Pressure (P)

Volume (V)

1 2 3 4 5 6

630. mmHg 690. mmHg 726 mmHg 790. mmHg 843 mmHg 914 mmHg

32.0 mL 29.2 mL 27.8 mL 25.6 mL 24.0 mL 22.2 mL

P × V (Product)

2. Graphing Pressure versus Volume: Boyle’s Law

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Questions and Problems Q1 According to your graph, what is the relationship between the pressure and the volume of a gas?

Q2 On your graph, what is the volume of the gas at a pressure of 760. mmHg?

Q3 On your graph, what is the pressure of the gas when the volume is 30.0 mL?

Q4 A sample of helium has a volume of 325 mL and a pressure of 655 mmHg. What will be the pressure, in mmHg, if the sample of helium is compressed to 125 mL (T, n constant)? (Show calculations.)

Q5 A 75.0-mL sample of oxygen has a pressure of 1.50 atm. What will be the new volume, in milliliters, if the pressure of the sample becomes 4.50 atm (T, n constant)? (Show calculations.)

Gas Laws

141

B. Charles’s Law 1.

2.

5. Measurements from other students

Boiling water bath

Cool Bath 1

Cool Bath 2

Temperature (°C) Temperature (K) (T) 3. Volume (mL) of water in cool flask

0.0 mL

4. Volume of water = air in flask 6. Volume of air remaining in cool flask (V) V 7. T

8. Graphing Volume versus Temperature for a Gas

Cool Bath 3

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Questions and Problems Q6 According to your graph, what is the predicted Kelvin temperature of absolute zero?

Q7 How does your predicted value for absolute temperature compare with the accepted value of 0 K?

Q8 According to your graph, how are temperature and volume related?

Q9 A gas with a volume of 525 mL at a temperature of −25 °C is heated to 175 °C. What is the new volume, in milliliters, of the gas if pressure and number of moles are held constant?

Q10 A gas has a volume of 2.8 L at a temperature of 27 °C. What temperature (°C) is needed to expand the volume to 15 L? (P and n are constant.)

Q11 Combined gas law problem: A balloon is filled with 500.0 mL of helium at a temperature of 27 °C and 755 mmHg. What volume, in milliliters, will it have when it reaches an altitude where the temperature is −33 °C and the pressure is 0.65 atm?

Dalton’s Law of Partial Pressures 13 LABORATORY GOALS • Measure the percentage of O 2 and N 2 in the air. • Determine the partial pressures of O 2 and N 2 in the air. • Determine the partial pressure of CO 2 in the air. • Determine the partial pressure of CO 2 in exhaled air.

LAB INFORMATION Time: Comments: Related Topics:

2 h (part A is finished during the next lab period) Tear out the report sheets and place them beside the matching procedures. Partial pressure, Dalton’s law, gas mixtures, atmospheric pressure

CHEMICAL CONCEPTS When we work with gases, we obtain the atmospheric pressure by reading a barometer in the laboratory. In a mercury barometer, air molecules strike the open surface of mercury, which pushes against the column of mercury in the vertical glass tube. When the height of the mercury column is 760 mmHg, it measures a pressure of 1 atmosphere (atm). As the atmospheric pressure changes, the height of the mercury in the tube also changes.

◄ The pressure exerted by the gases in the atmosphere is equal to the pressure of a mercury column in a closed glass tube.

Dalton’s Law of Partial Pressures In a mixture of two or more gases, the pressure exerted by each gas is called its partial pressure. The sum of the partial pressures is the total pressure of the gas mixture. Ptotal Total pressure of a gas mixture

= =

P1 + P2 + P3 + " Sum of the partial pressures of the gases in the mixture

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▲ When two gases are combined, the total pressure of the gas mixture is the sum of their partial pressures.

A. Partial Pressures of Oxygen and Nitrogen in Air The two major gases in the air are oxygen (O 2 ) and nitrogen (N 2 ). In this experiment, you will make measurements to determine the percentage of nitrogen and oxygen in the air. We remove the oxygen by reacting it with iron powder in a test tube. 4Fe(s ) + 3O 2 (g )



2Fe 2 O3 (s )

The oxygen that reacts is replaced by water that enters the test tube. After we measure the volume of the air (oxygen and nitrogen) in the test tube, we subtract the volume of the unreacted nitrogen, to obtain the volume of oxygen. Using the atmospheric pressure, the partial pressures (mmHg) and their percentages of oxygen and nitrogen in the air can be calculated.

B. Carbon Dioxide in the Atmosphere Carbon dioxide is an important greenhouse gas that absorbs radiation from the Sun. In the carbon cycle, plants utilize most of the CO 2 in the atmosphere and return O 2 to the atmosphere. On Earth, carbon dioxide is produced by volcanic activity, respiration of living organisms, and burning of biomass and fossil fuels for energy and heat. In this experiment, we will measure the partial pressure of carbon dioxide in the atmosphere by reacting the CO 2 in the air with NaOH. CO 2 (g ) + NaOH(aq )



NaHCO3 (aq)

As the CO 2 is absorbed by the NaOH, the pressure inside the sealed flask will decrease. The decrease in the pressure will be used to determine the partial pressure of CO 2 in the atmosphere. The other major gases in the atmosphere, oxygen and nitrogen, do not react with sodium hydroxide and continue to exert their respective partial pressures.

C. Carbon Dioxide in Exhaled Air When the body metabolizes nutrients, one of the end products is carbon dioxide, CO 2 . The body eliminates carbon dioxide by exhalation of air from the lungs. In this experiment, we will measure the partial pressure of carbon dioxide in exhaled air by reacting the CO 2 present in exhaled air with NaOH. CO 2 (g ) + NaOH(aq )



NaHCO3 (aq)

As the CO 2 is absorbed by the NaOH, the pressure inside the sealed flask will decrease. The decrease in the pressure will be used to determine the partial pressure of CO 2 in the exhaled air.

Dalton’s Law of Partial Pressures

EXPERIMENTAL PROCEDURES

145

GOGGLES REQUIRED!

A. Partial Pressures of Oxygen and Nitrogen in Air Materials: 250-mL beaker, large test tube, iron powder, 100-mL graduated cylinder, buret clamp or test tube holder 1.

Completely fill a large test tube with water. Pour all the water into a graduated cylinder. Measure and record the volume of water. This is the volume of the air in the test tube.

Laboratory Setup for Experiment Obtain a small scoop of iron powder and sprinkle the powder in the same empty, moist test tube. Shake the test tube so that the powder adheres to the inside of the test tube. Pour out any loose powder. Fill a 250-mL beaker about one-half full of water. Attach a buret clamp or test tube holder to the test tube. Turn the test tube with the iron powder upside down and carefully set its open end into the water (see Figure 13.1). With the test tube resting on the bottom of the beaker, adjust the handle of the buret clamp or test tube holder so it rests on the rim of the beaker. Store this apparatus in a place where it will remain undisturbed. The last part of this experiment will be completed during the next laboratory period.

◄ FIGURE 13.1 A beaker and test tube apparatus is used for the determination of the partial pressures of oxygen and nitrogen in the air.

Next Laboratory Period Carefully place the beaker and test tube apparatus on your lab desk. You will see that water has filled part of the test tube because water replaced the oxygen that reacted with the iron powder. Use a marking pen or tape to mark the water level inside the test tube (see Figure 13.2).

◄ FIGURE 13.2 A test tube is marked to indicate the volume of unreacted nitrogen.

2.

Remove the test tube from the beaker and fill it with water up to the marked line. Empty the water into a graduated cylinder. Measure and record the volume of the water. This is the volume of unreacted nitrogen that remained in the test tube after the oxygen reacted.

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Calculations 3.

Read a barometer and record the atmospheric pressure in mmHg.

4.

Calculate the volume of O 2 present in the test tube by subtracting the volume (2) of N 2 from the total volume (1) of the test tube. Volume of O 2 = volume (1) of test tube – volume (2) of N 2

5.

Calculate the percentage of O 2 in the air by dividing the volume (4) of O 2 by the volume (1) of the air sample. % O2 =

6.

Calculate the percentage of N 2 in the air by dividing the volume (2) of N 2 by the volume (1) of the air sample.

% N2 = 7.

mL of O 2 × 100% mL of air

mL of N 2 × 100% mL of air

Calculate the partial pressure of O 2 by multiplying the atmospheric pressure, in mmHg (3), by the percentage (5) of O 2 . Partial pressure of O2 =

8.

% O2 × atmospheric pressure 100

Calculate the partial pressure of N 2 by multiplying the atmospheric pressure, in mmHg (3), by the percentage (6) of N 2 . Partial pressure of N 2 =

% N2 × atmospheric pressure 100

B. Carbon Dioxide in the Atmosphere (This may be a demonstration by the instructor.) Materials: 250-mL Erlenmeyer flask to fit two-hole stopper, shell vial, 6 M NaOH, mineral oil,150-mL beaker, food coloring (optional), meterstick, glass tubing (60–75 cm), two-hole stopper with two short pieces of glass tubing, rubber tubing (1 short, 1 long), funnel, pinch clamp 1.

Read a barometer and record the atmospheric pressure, in mmHg.

Preparation of Laboratory Apparatus

Carefully lower an empty shell vial into a 250-mL Erlenmeyer flask so that the vial remains upright (see Figure 13.3). Place a funnel in the flask directly above the shell vial. Obtain a small amount of 6 M NaOH in a small beaker. Caution: Sodium hydroxide is caustic! Be sure to clean up any spill immediately. Wash any spills on skin for 10 minutes. Slowly pour the NaOH through the funnel into the vial and stop when the vial is almost full. Pour a small amount of mineral oil into the vial until the oil forms a thin layer of about 1 mm on top of the NaOH. The mineral oil on top of the sodium hydroxide prevents it from reacting with carbon dioxide. Remove the funnel.

Dalton’s Law of Partial Pressures

147

◄ FIGURE 13.3 Fill a shell vial with NaOH and mineral oil using a funnel in a flask.

Continue to prepare the apparatus as shown in Figure 13.4. Place a two-hole stopper containing two pieces of glass tubing in the flask. Make sure that the stopper is tight. One piece of glass tubing should be attached to a short piece of rubber tubing (A) and the other glass tubing should be attached to a longer piece of rubber tubing (B). Place a long piece (60–70 cm) of glass tubing in the open end of the longer piece of rubber tubing. Place the open end of the glass tubing (B) in a beaker filled with water. A few drops of food coloring can be added to the water to make it easier to see the level. Close the apparatus by attaching a pinch clamp to the rubber tubing (A). Make sure that the end of tube B stays below the water level in the beaker for the rest of the experiment.

▲ FIGURE 13.4 Flask, tubing, and vial provide a setup for carbon dioxide determination. Gently tilt the flask until the shell vial tips over and the NaOH spills out and covers the bottom of the flask. Swirl and shake the flask gently. Hold the flask around the top so that your hands do not warm the flask, because an increase in temperature will change the pressure. As the carbon dioxide in the flask reacts with NaOH, there is a drop in the pressure in the flask, and water will rise in the glass tube (B). Continue to swirl the flask until there is no further change in the water level in the glass tube. Because there is only a small amount of CO2 in the atmosphere, the water level will not change very much. 2.

The distance between the water level in the beaker and the water level in the vertical glass tube (B) is equal to the partial pressure of CO 2 , in mm H 2 O. Use a meterstick to measure this distance, in millimeters. Make sure your measurement is in millimeters (mm), not centimeters (cm).

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Calculations 3.

To calculate the PCO2 in mmHg, divide the PCO2 in mm H 2 O (2), by 13.6. (1.00 mmHg = 13.6 mm H 2 O) PCO2 in the atmosphere (mmHg) = height in mm H 2 O ×

4

1 mmHg 13.6 mm H 2 O

Calculate the percent CO 2 in the atmosphere by dividing the PCO2 (3) by the atmospheric pressure (1). PCO2 ( mmHg ) × 100% % CO 2 in the atmosphere = PAtmosphere ( mmHg )

C. Carbon Dioxide in Exhaled Air Materials: 250-mL Erlenmeyer flask to fit two-hole stopper, shell vial, 6 M NaOH, mineral oil, 150-mL beaker, food coloring (optional), meterstick, glass tubing (60–75 cm), two-hole stopper with two short pieces of glass tubing, rubber tubing (1 short, 1 long), funnel, pinch clamp, clean straws. Caution: NaOH is caustic; clean up any spills immediately. 1.

Read a barometer and record the atmospheric pressure (mm Hg).

Preparation of Laboratory Apparatus

Rinse out the flask you used in part B and prepare the same apparatus. This time, place a clean straw into the rubber tubing (A). As before, place glass tube B in the beaker of water. Take a breath of air, hold for a moment, and exhale through the straw. This will cause bubbling in the beaker of water. Cover the straw with your finger while you inhale. Caution: Take your time. Exhaling too rapidly may cause hyperventilation and make you dizzy. Stop and rest. Exhale through the straw again. Repeat the exhalation three times. Place a pinch clamp on the short piece of rubber tubing (A). Remove the straw. Now the flask is filled with exhaled air. Gently tilt the flask until the shell vial tips over and the NaOH spills out and covers the bottom of the flask. Swirl and shake the flask gently. Hold the flask around the top, so your hands do not warm the flask, because an increase in temperature will change the pressure. As the carbon dioxide in the flask reacts with NaOH, there is a drop in the pressure in the flask, and water will rise in the glass tube (B). Because of the higher partial pressure of CO 2 in exhaled air, the water level in tube B will make a dramatic rise. Keep swirling the flask until there is no further change in the water level in the long glass tube (B). 2.

The distance between the water level in the beaker and the water level in the vertical glass tube (B) is equal to the partial pressure of CO 2 , in mm H 2 O. Use a meterstick to measure this distance, in millimeters. Make sure your measurement is in millimeters (mm), not centimeters (cm).

Calculations 3.

To calculate the PCO2 in mmHg, divide the PCO2 , in mm H 2 O (2), by 13.6. (1.00 mmHg = 13.6 mm H 2 O) PCO2 in exhaled air (mmHg) = height in mm H 2 O ×

4.

1 mmHg 13.6 mm H 2 O

Calculate the percent CO 2 in exhaled air by dividing the PCO2 , (3), by the atmospheric pressure, (1). % CO 2 in exhaled air =

PCO2 ( mmHg ) PAtmosphere ( mmHg )

× 100%

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Pre-Lab Study Questions

13

1. What is meant by the term partial pressure?

2. How will you determine the partial pressure of N 2 in the atmosphere when it does not react in any of the experiments?

3. A tank contains a gas mixture of 0.852 atm of He, 685 mmHg of O 2 , and 455 mmHg of N 2 . What is the pressure, in mmHg, of the gas mixture?

4. A gas mixture contains 1.52 atm of Ne, 766 mmHg of He, and Ar. What is the partial pressure, in atmospheres, of Ar if the gas mixture has a total pressure of 3.27 atm?

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REPORT SHEET

Dalton’s Law of Partial Pressures

LAB

13

A. Partial Pressures of Oxygen and Nitrogen in Air 1. Volume of test tube

________________________ mL

2. Volume of N 2

________________________ mL

3. Atmospheric pressure

________________________ mmHg

Calculations 4. Volume of O 2

________________________ mL

5. Percent O2 (Show calculations.)

________________________ % O 2

6. Percent N 2 (Show calculations.)

________________________ % N 2

7. Partial pressure of O2 (Show calculations.)

________________________ mmHg

8. Partial pressure of N 2 (Show calculations.)

________________________ mmHg

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Questions and Problems Q1 If you lived in the mountains where the atmospheric pressure is 685 mmHg, what would the partial pressure of oxygen and of nitrogen be? Use your percent values from A.5 and A.6.

B. Carbon Dioxide in the Atmosphere 1. Atmospheric pressure

_______________________ mmHg

2. PCO2 (height of water column in tube B) _______________________ mm H 2O Calculations 3. PCO2 = mm H 2O ×

1 mmHg = 13.6 mm H 2 O

4. Percent CO 2 in the atmosphere (Show calculations.)

_______________________ mmHg _______________________ % CO 2

C. Carbon Dioxide in Exhaled Air 1. Atmospheric pressure

_______________________ mmHg

2. Height of water column

_______________________ mm H 2O

Calculations 3. PCO2 = mm H 2O ×

1 mmHg = 13.6 mm H 2O

4. Percent CO 2 (exhaled air) (Show calculations.)

_______________________ mmHg _______________________ % CO 2

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153

Questions and Problems Q2 Which would you expect to be higher, the percentage of CO 2 in the atmosphere or in exhaled air? Why?

Q3 A sample of gas contains O 2 at a pressure of 45 mmHg, N 2 at a pressure of 1.20 atm, and He at a pressure of 825 mmHg. What is the total pressure in mmHg?

Q4 A mixture of gases has a total pressure of 1650 mmHg. The gases in the mixture are helium at a pressure of 215 mmHg, nitrogen at a pressure of 0.28 atm, and oxygen. What is the partial pressure of the oxygen, in atm?

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Solutions, Electrolytes, and Concentration 14 LABORATORY GOALS • • • • • • •

Observe the solubility of a solute in polar and nonpolar solvents. Determine the effect of particle size, stirring, and temperature on the rate of solution formation. Identify an unsaturated and a saturated solution. Compare the conductivity of strong electrolytes, weak electrolytes, and nonelectrolytes. List the electrolytes and their concentrations (mEq/L) in intravenous solutions. Calculate the mass/mass percent and mass/volume percent concentrations for an NaCl solution. Calculate the molar concentration of an NaCl solution.

LAB INFORMATION Time: Comments: Related Topics:

3h Some solvents in part A of the experiments are flammable. Do not light any burners. Tear out the report sheets and place them beside the matching procedures. Your instructor will use the conductivity apparatus in a demonstration for part B. Solute, solvent, formation of solutions, polar and nonpolar solutes, electrolytes, nonelectrolytes, concentrations of solutions Dispose of all chemicals as directed by your lab instructor.

CHEMICAL CONCEPTS A. Polarity of Solutes and Solvents A solution is a homogenous mixture of two or more substances. The substance that is present in the greater amount is called the solvent. The substance that is present in the smaller amount is the solute. In many solutions, including body fluids and the oceans, water is the solvent. However, the solutes and solvents that make up solutions may be solids, liquids, or gases.

▲ A solution consists of at least one solute dispersed in a solvent.

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A solution forms when the attractive forces between the solute and the solvent are similar. A polar (or ionic) solute such as NaCl is soluble in water, a polar solvent. As the NaCl dissolves, its ions separate into Na + and Cl− . The positive Na + ions are attracted to the partially negative oxygen atoms of water. At the same time, the negative Cl− ions are pulled into the solution by their attraction to the partially positive hydrogen atoms of water. Once the ions are into the solution, they stay in solution because they are hydrated, which means that water molecules have surrounded each ion.

▲ Ions on the surface of NaCl are attracted to polar water molecules that pull the Na + and Cl− ions into solution. A weakly nonpolar solvent such as acetone is needed to dissolve a nonpolar solute. This requirement of similar electrical attraction between solute and solvent is sometimes stated as “like dissolves like.”

B. Electrolytes and Nonelectrolytes The types of substances in aqueous solutions can be identified as strong electrolytes, weak electrolytes, or nonelectrolytes by using a conductivity apparatus. Electrolytes are substances that produce ions in water. Strong electrolytes contain only ions in solution; weak electrolytes produce a few ions, but contain mostly molecules. Nonelectrolytes dissolve as molecules. When ions are present in an aqueous solution, the lightbulb of a conductivity apparatus will glow because the ions complete an electrical circuit. Weak electrolytes produce a few ions: the lightbulb will glow weakly. When a nonelectrolyte dissolves, only molecular substances are present, which do not carry current in an aqueous solution. The lightbulb in the conductivity apparatus does not glow (see Figure 14.1).

A strong electrolyte completely dissociates into ions in an aqueous solution.

A weak electrolyte forms mostly molecules and a few ions in an aqueous solution.

▲ FIGURE 14.1 A lightbulb apparatus is used to measure the conductivity.

A nonelectrolyte dissolves as molecules in an aqueous solution.

Solutions, Electrolytes, and Concentration 157

C. Electrolytes in Body Fluids Electrolytes play an important role in the maintenance of the cell. In the intracellular fluid of a cell, the major cation is potassium, K + , and the major anion is hydrogen carbonate, HCO3− . In the extracellular fluid, the major cation is sodium, Na + , and the major anion is chloride, Cl− . Other electrolytes are found in smaller quantities in both intracellular and extracellular fluids. The concentrations of electrolytes in intravenous fluids are usually expressed in milliequivalents per liter (mEq/L).

◄ An intravenous solution is used to maintain electrolyte balance in the body.

When there is a loss of fluid from the body or an imbalance of electrolytes, an intravenous solution may be administered. The type of IV solution reflects the needs of the cells in the body as determined by laboratory tests of the body fluids. Electrolytes in the body maintain and regulate fluid balance, muscle tone, and acid–base balance.

D. Concentration of a Sodium Chloride Solution The concentration of a solution is calculated from the amount of solute present in a certain amount of solution. The concentration may be expressed using different units for amount of solute and solution. A mass percent (m/m) concentration expresses the grams of solute in the grams of solution. The mass/volume (m/v) percent concentration of a solution states the grams of solute present in the milliliters of the solution. Mass percent (m/m) =

mass of solute (g ) × 100% mass of solution (g )

Mass/volume percent (m/v) =

grams of solute × 100% milliliters of solution

A molar (M) concentration gives the moles of solute in a liter of solution.

Molarity (M) =

moles of solute liters of solution

▲ A 2.0 M NaCl solution is prepared by adding water to 2.0 moles of NaCl to give a 1.00-L volume.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Polarity of Solutes and Solvents (This may be a demonstration by your instructor.) Materials: Test tubes (8), test tube rack, spatulas, stirring rods, KMnO 4 (s ), I 2 (s), sucrose(s), vegetable oil, cyclohexane Iodine ( I 2 ) can burn the skin. Handle cautiously! Caution: Cyclohexane is flammable—do not proceed if any laboratory burners are in use. Use a fume hood if possible. 1.

2.

Set up eight test tubes in a test tube rack. Polar solvent: To each of four test tubes, add 3 mL of water. Nonpolar solvent: To each of four more test tubes, add 3 mL of cyclohexane. To one test tube containing water and one test tube containing cyclohexane add a few crystals (a small amount) or a few drops of each of the solutes: KMnO 4 , I 2 , sucrose, vegetable oil (one solute per test tube). Stir each mixture with a glass stirring rod. Rinse and wipe off after each use. Identify the solutes as soluble or not soluble in each of the solvents. Use the solubility of each solute to identify it as polar or nonpolar. If a solute dissolves in a polar solvent like water, it is a polar solute. If a solute dissolves in a nonpolar solvent like cyclohexane, the solute is nonpolar. Discard the solutions in the waste containers as directed by your instructor.

B. Electrolytes and Nonelectrolytes (This will be an instructor demonstration.) Materials: Conductivity apparatus, 100- or 150-mL beakers, 0.1 M NaCl, 0.1 M sucrose, 0.1 M HCl, 0.1 M HC2 H3O 2 (acetic acid), 0.1 M NaOH, 0.1 M NH 4 OH (ammonium hydroxide), 0.1 M C2 H5 OH (ethanol) 1.

2. 3.

Pour about 15–20 mL of the above solutions into small beakers. Carefully lower the electrodes into the solution and observe the lightbulb. Caution: Bare electrodes are a hazard! Skin will conduct an electric current and will cause a shock. Do not touch the electrodes when the lightbulb apparatus is plugged in. If it glows, the solution is an electrolyte; if it does not glow, it is a nonelectrolyte. If the light is very bright, it is a strong electrolyte. A conductivity meter may be substituted for the lightbulb apparatus. Record the intensity of the lightbulb as a bright glow, slight glow, or none. Select another substance and repeat the conductivity test. Identify each solute as a strong, weak, or nonelectrolyte. Indicate the kinds of particles present in each solution.

C. Electrolytes in Body Fluids Materials: Samples of IV solutions such as saline solution, Ringer’s, etc. 1. 2. 3. 4. 5. 6.

From the display of intravenous (IV) solutions provided in the laboratory, select three and record the type of solutions in each. List the positively charged electrolytes (cations) and their concentrations (mEq/L). List the negatively charged electrolytes (anions) and their concentrations (mEq/L). Calculate the total number of mEq of cations. Calculate the total number of mEq of anions. Calculate the overall sum of the positive and negative electrolytes (4 + 5).

Solutions, Electrolytes, and Concentration 159

D. Concentration of a Sodium Chloride Solution Materials: Hot plate (or Bunsen burner, iron ring, and wire screen), evaporating dish, NaCl solution, 400-mL beaker (to fit evaporating dish), 10-mL graduated cylinder (or 10-mL pipet) In this experiment, you will measure a 10.0-mL volume of a sodium chloride solution. The mass of the solution will be determined by weighing the solution in a preweighed evaporating dish. After the sample is evaporated to dryness, it is weighed again. From this data, the mass of the sodium chloride (solute) is obtained. 1. 2.

Weigh a dry evaporating dish and record. Use a small graduated cylinder or a 10.0-mL pipet to measure out a 10.0-mL sample of the NaCl solution (see Figure 14.2). Review Using a Pipet to Transfer Liquids in the preface of this lab manual. Record the volume.

◄ FIGURE 14.2 A pipet is used to measure the volume of a solution.

3.

4.

Weigh the evaporating dish and the NaCl solution. Record their mass. Fill a 400-mL beaker about half full of water. Set the beaker with water on a hot plate, or heat with a Bunsen burner using an iron ring with a wire screen (see Figure 14.3). Place the evaporating dish on top of the beaker and heat to boiling. You may need to add more water. When the NaCl appears to be dry, turn off the burner. After the evaporating dish has cooled, dry the bottom and place it directly on the hot plate or on the iron ring. Heat gently with a low flame to dry the salt completely. Turn off the burner, and allow the evaporating dish and dried NaCl sample to cool to room temperature. Weigh the evaporating dish and the dry NaCl. Record.

▲ FIGURE 14.3 A Bunsen burner or hot plate is used to evaporate an NaCl solution.

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Calculations 5. 6.

Mass of solution = mass of dish + solution (3) − mass of dish (1) Calculate the mass of the dry NaCl by subtracting the mass of the evaporating dish (1) from the total mass of the evaporating dish and the dried NaCl (4). Mass of dry NaCl = mass of dish + dry NaCl (4) − mass of dish (1)

7.

Calculate the mass percent (m/m) concentration using the mass of the dry NaCl (6) divided by the mass of the solution (5) and multiplying by 100%. Mass percent (m/m) =

8.

mass of dry NaCl (6) × 100% mass (g) of solution (5)

Calculate the mass/volume percent (m/v) concentration by dividing the mass of the dry NaCl (6) by the volume of the NaCl solution (2) and multiplying by 100%. Mass/volume percent (m/v) =

9.

mass of dry NaCl (6) × 100% volume (mL) of solution (2)

Calculate the moles of NaCl by dividing the grams of dried NaCl (6) divided by the molar mass of NaCl. The molar mass of NaCl is 58.5 g/mole. Moles of NaCl = g of NaCl (6) ×

10.

Convert the volume (mL) of the NaCl solution (2) to the corresponding volume, in liters, by dividing by 1000 mL. mL of NaCl solution (2) ×

11.

1 mole of NaCl 58.5 g of NaCl

1L = L of NaCl solution 1000 mL

Calculate the molarity of the NaCl solution by dividing the moles of NaCl (9) by the liters of NaCl solution (10). Molarity (M) of NaCl solution =

moles NaCl (9) L of solution (10)

Date

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Pre-Lab Study Questions

14

1. Why does an oil-and-vinegar salad dressing have two separate layers?

2. What is meant by the mass percent (m/m) concentration of a solution?

3. What is molarity?

4. Why are some electrolytes considered strong, whereas others are considered weak?

5. A solution is prepared with 3.26 g KCl and water to make 25.0 mL of KCl. a. What is the % (m/v) of the KCl solution?

b. What is the molarity (M) of the KCl solution?

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REPORT SHEET

Solutions, Electrolytes, and Concentration

LAB

14

A. Polarity of Solutes and Solvents Solute

1. Soluble/Not Soluble in

Water

2. Identify the Solute as Polar or Nonpolar

Cyclohexane

KMnO 4 I2 Sucrose Vegetable oil

B. Electrolytes and Nonelectrolytes Substance

1. Observations (Intensity of Lightbulb)

2. Type of 3. Type of Particles Electrolyte (Ions, Molecules, (Strong, Weak, or or Both) Nonelectrolyte)

0.1 M NaCl 0.1 M Sucrose 0.1 M HCl 0.1 M HC 2 H3O 2 , Acetic acid

0.1 M NaOH 0.1 M NH 4OH 0.1 M C 2 H 5OH, Ethanol Copyright © 2014 Pearson Education, Inc.

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Questions and Problems Q1 Why are some solutes soluble in water, but others are soluble in cyclohexane?

Q2 For the three solutes tested in B, write an equation for their dissolution in water:

HCl(aq)

___________________________________________________________

NH 4OH(aq) ___________________________________________________________ C2 H5OH(aq ) ___________________________________________________________ Q3 Classify the solutes in each of the following equations as a weak electrolyte, a strong electrolyte, or a nonelectrolyte in water: a. XY2 (s ) ⎯⎯ → X 2+ (aq) + 2Y – (aq) __________________________________________ ⎯⎯ → H + (aq) + X – (aq ) b. HX(g ) ←⎯ ⎯

__________________________________________

c. XYZ(s) ⎯⎯ → XYZ(aq )

__________________________________________

d. YOH(s ) ⎯⎯ → Y + (aq) + OH – (aq ) __________________________________________

C. Electrolytes in Body Fluids 1. Type of IV Solution

2. Cations

3. Anions

4. Total Charge of Cations (+) 5. Total Charge of Anions (−) 6. Sum of the Charges

Solutions, Electrolytes, and Concentration 165

Questions and Problems Q4 What would be the overall charge in any IV solution? Why?

D. Concentration of a Sodium Chloride Solution 1. Mass of evaporating dish

______________________________ g

2. Volume of NaCl solution

______________________________ mL

3. Mass of dish and NaCl solution ______________________________ g 4. Mass of dish and dry NaCl

______________________________ g

Calculations 5. Mass of NaCl solution

____________________________________ g

6. Mass of the dry NaCl salt

____________________________________ g

7. Mass/mass percent (Show calculations.)

____________________________________ % (m/m)

8. Mass/volume percent (Show calculations.)

____________________________________ % (m/v)

9. Moles of NaCl (Show calculations.)

____________________________________ moles

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10. Volume of sample in liters

_______________________________ L

11. Molarity of NaCl solution (Show calculations.)

_______________________________ M

Questions and Problems Q5 A 15.0-mL sample of NaCl solution has a mass of 15.78 g. After the NaCl solution is evaporated to dryness, the dry salt residue has a mass of 3.26 g. Calculate the following concentrations for the NaCl solution. a. % (m/m)

b. % (m/v)

c. molarity (M)

Q6 How many grams of KI are in 25.0 mL of a 3.0 % (m/v) KI solution?

Q7 How many milliliters of a 2.5 M MgCl2 solution contain 17.5 g MgCl2?

Soluble and Insoluble Salts 15 LABORATORY GOALS • Predict the formation of an insoluble salt. • Observe the effect of temperature on solubility. • Measure the solubility of KNO3 at various temperatures, and graph a solubility curve. • Test a variety of water samples for water hardness. • Use water treatment techniques to purify water.

LAB INFORMATION Time: Comments: Related Topics:

2½ h Tear out the lab report sheets and place them beside the matching procedures. Solubility, insoluble salts, saturated solution

CHEMICAL CONCEPTS A. Soluble and Insoluble Salts Although many ionic compounds (salts) are soluble in water, some do not dissolve. They are known as insoluble salts. In medicine, the insoluble salt BaSO 4 is used as an opaque substance to help outline the gastrointestinal tract in X-ray images. Solubility rules are shown in Table 15.1. TABLE 15.1 Solubility Rules for Ionic Compounds in Water Soluble if it contains +

+

+

Insoluble if it contains +



Li , Na , K , NH 4 , NO3 , C2 H3O 2



Cl− , Br − , I −

Ag + , Pb 2+ , Hg 22+

SO 4 2−

Ba 2+ , Pb 2+ , Ca 2+ , Sr 2 +

OH − , CO32− , S2− , PO 43−

When solutions of two ionic compounds are mixed, the formation of a solid indicates an insoluble salt. The positive ion of one substance and the negative ion combine to form the insoluble salt. For example, mixing solutions of the soluble salts NaCl and AgNO3 will produce a white solid, which is the insoluble salt AgCl(s). We can write its chemical equation as follows: AgNO3 (aq) + NaCl(aq ) Soluble salts

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⎯⎯ → AgCl(s) Insoluble salt

+ NaNO3 (aq ) Soluble salt

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▲When aqueous solutions of NaCl and AgNO3 are mixed, an insoluble precipitate of AgCl forms. From the chemical equation, we can show the reacting components by writing an ionic equation and a net ionic equation for the reaction of AgNO3 and NaCl. The ionic equation is written by showing all the ions of the reactants and products that are soluble. Any insoluble salt in the products is written as a formula unit. Chemical Equation AgNO3 (aq )

+ NaCl(aq )

⎯⎯ → AgCl(s) + NaNO3 (aq ) Insoluble

Ionic Equation

Ag + (aq) + NO3− (aq) + Na + (aq ) + Cl− (aq )

⎯⎯ → AgCl(s) + Na + (aq) + NO3− (aq) Insoluble

When all the ions that appear unchanged (spectator ions) in the equation are removed, we obtain the net ionic equation that shows only the ions involved the chemical reaction. Ag + (aq) + NO3− (aq) + Na + (aq) + Cl− (aq) ⎯⎯ → AgCl( s ) + Na + (aq) + NO3− (aq) Insoluble

Net Ionic Equation Ag + (aq ) + Cl− (aq )

⎯⎯ → AgCl( s) Insoluble

B. Solubility of KNO 3 When a solution dissolves the maximum amount of solute at a certain temperature, it is saturated. If more solute is added, the excess appears as a solid in the container. The maximum amount of solute that dissolves is called the solubility of that solute in that solvent. Solubility is usually stated as the number of grams of solute that dissolve in 100 mL (or 100 g) of water. The solubility depends upon several factors, including the nature of the solute and solvent, the temperature, and the pressure (for a gas). Generally, for a solid, solubility increases with an increase in temperature.

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▲More solute can dissolve in an unsaturated solution, but not in a saturated solution.

C. Testing the Hardness of Water Hard water contains the ions Ca 2 + , Mg 2 + , and Fe3+ . When hard water reacts with soap, the ions in the hard water and some of the soap molecules form insoluble salts called “soap scum.” The soap molecules tied up in the scum are not free to perform their cleaning function. Initially soap is used to remove the ions, and then more soap is added to produce sudsing and cleaning. The reaction of a soap solution with the ions in hard water can be used to compare hardness of water samples. Water softeners are used in homes to remove Ca 2 + and Mg 2 + by replacing them with Na + or K + ions.

D. Purification of Water In water treatment plants, chemicals—called flocculating agents—are added to water to cause the formation of insoluble substances that sink to the bottom of the settling tank (sedimentation). The clear water on top is purified and can be drawn off for use. You will compare different compounds to determine which one leads to the most rapid settling of a muddy mixture.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Soluble and Insoluble Salts Materials: Spot plate or plastic sheet, 0.1 M solutions of AgNO3 , NaCl, Na 2SO4 , Na 3PO4 , Ca(NO3 ) 2 , droppers, or dropper bottles of the 0.1 M solutions given above 1. 2.

3. 4.

Dispose of all waste chemicals as directed by your instructor. Write the ions that are in the two solutions used to form each of the mixtures listed on your report sheet. Using solubility rules, draw a circle around any pair of ions that would form an insoluble salt. Obtain a spot plate or a plastic sheet. Make mixtures of each pair of compounds listed in the report sheet by placing 2–3 drops of each solution in the same well or on the same spot on the plastic sheet. Look for a cloudy appearance, which means the formation of a solid (insoluble salt). Record your observations. Write the formula of the insoluble salt where a solid formed. For the reactions listed, for any that form an insoluble salt, complete and balance the chemical equation. Then write and balance the ionic equation and net ionic equation for those reactions. If no insoluble salt forms, write NR (no reaction). Discard the solutions and solids as directed by your instructor.

B. Solubility of KNO 3 Materials: Weighing paper or small container, spatula, stirring rod, test tube, 400-mL beaker, buret clamp, hot plate or Bunsen burner, thermometer, 5- or 10-mL graduated cylinder or a 5.0 mL pipet, KNO3 (s) To reduce the amount of KNO3 used, each group of students will be assigned an amount of KNO3 to weigh out. The results will be shared with the class. 1. Each group of students will be assigned an amount of KNO3 between 3 and 7 grams. Weigh a container and record the mass. 2. Add an amount of KNO3 that gives a mass that is close to your assigned amount. For example, if your container has a mass of 8.00 g, and you are assigned an amount of about 3 g, add KNO3 to the container until the combined mass is about 11 g, which might be 11.10 g or 11.25 g or 10.85 g. It is not necessary to add or remove KNO3 to obtain exactly 3.00 g. Record the mass of the container and KNO3 . The temperature at which the KNO3 is soluble is determined by heating and cooling the KNO3 solution. Measure 5.0 mL of water with a graduated cylinder or pipet and place the water in a test tube. Carefully add the entire sample of KNO3 to the test tube. Clamp the test tube to a ring stand and place the test tube in a beaker of water. Use a hot plate or Bunsen burner to heat the water (see Figure 15.1).

◄FIGURE 15.1 The KNO3 solution is heated in a water bath.

3.

Stir the mixture with the thermometer and continue heating until all the KNO3 dissolves. As soon as all the KNO3 dissolves, turn off the burner or hot plate. Loosen the clamp and remove the test tube from the hot water. Stir gently with a thermometer as the test tube and contents cool. When the solution becomes saturated, crystals of KNO3 will begin to form. Once the sample has crystallized, do not attempt to stir or remove the thermometer - it may break!

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Record the temperature of the solution as soon as crystals appear. Return the test tube to the hot water bath and repeat the warming and cooling of the solution until you obtain two temperature readings that agree. Dispose of waste chemicals as directed by your instructor. Calculations 4.

Calculate the mass of KNO3 by subtracting the mass of the container (1) from the combined mass of the container and the KNO3 (2).

5.

Calculate the solubility as grams of KNO3 per 100 mL of water. Solubility (g KNO3 /100 mL of H 2 O) =

6.

g KNO3 (4) × 100 5.0 mL H 2 O

From other students, obtain the solubility results for different quantities of KNO3 and their solubility temperatures. Prepare a graph of the solubility curve for KNO3 . Plot the solubility (g KNO3 /100 mL water) on the vertical axis and the temperature (°C) on the horizontal axis.

C. Testing the Hardness of Water Materials: 250-mL beaker, 50-mL (or 25-mL) graduated cylinder, two 250-mL flasks with stoppers, soap solution in dropper bottles, water samples

Place 50 mL of distilled water in a 250-mL flask. Add 1 drop of the soap solution. Stopper the flask and shake for 10 seconds. You should see a thick layer of suds. If not, add another drop of soap solution and shake again for 10 seconds. The suds that form in the distilled water sample will serve as your reference sample. Save for comparison. Shake again if necessary. 1. Add 1 drop of the soap solution to another water sample. Stopper the flask and shake for 10 seconds. Keep adding drops of the soap (up to 20 drops) solution and shake until the sample forms about the same amount of suds as the distilled water sample. Stop if no suds are formed after you have added 20 drops of soap solution. Record the number of drops required to soften each water sample. Repeat the above instructions to test an assortment of water samples available in the lab or from your home, a pool, or a well.

D. Purification of Water Materials: Test tubes, test tube rack, muddy water, 1% Al2 (SO 4 )3 , 1% Na 2SO 4 , 1% NaCl

Set up four test tubes in a test tube rack. To each, add 10 mL of muddy water. If the mud has settled, shake or stir the test tube. Add 5 mL water to the first test tube. This is your reference sample. Add 5 mL 1% NaCl to the second test tube. Add 5 mL 1% Al2 (SO 4 )3 to the third test tube. Add 5 mL 1% Na 2SO 4 to the fourth test tube. 1.

2.

Mix each test tube thoroughly by shaking and allow the test tubes to stand undisturbed for 15 min. Look for a separation of a precipitate and a clarification of the upper portion of water. Record your observations as follows; NS - no settling, still muddy SS - some settling, still cloudy MS - mostly settled, slightly cloudy SC - settled, clear After 30 min, observe the test tubes again and record.

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Pre-Lab Study Questions

15

1. What is an insoluble salt?

2. Are salts more soluble in water at higher temperature or lower temperature?

3. Indicate whether each of the following salts is soluble or insoluble in water and give a reason. Salt

Soluble or Insoluble Reason in Water?

a. BaSO 4 b. AgBr c. Pb(NO3 ) 2 d. CaCO3 e. Na 3 PO 4 4. For each of the following combinations of solutions that produces an insoluble salt, write the net ionic equation for the formation of the insoluble salt. If no reaction occurs, write NR. Solutions

Net Ionic Equation

a. NaCl(aq ) + Pb(NO3 ) 2 (aq) ⎯⎯ → b. (NH 4 )3 PO 4 (aq) + Pb(NO3 ) 2 (aq) ⎯⎯ → c. CaCl2 (aq ) + Na 2 CO3 (aq) ⎯⎯ → d. AgNO3 (aq) + K 2S(aq ) ⎯⎯ →

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REPORT SHEET

LAB

Soluble and Insoluble Salts

15

A. Soluble and Insoluble Salts Compounds in Solutions

1. Ions

2. Observations 3. Formula of Insoluble Salt (if any)

AgNO3 + NaCl AgNO3 + Na 2SO 4 AgNO3 + Na 3PO 4 Ca(NO3 ) 2 + NaCl Ca(NO3 ) 2 + Na 2SO 4 Ca(NO3 ) 2 + Na 3PO 4

→ 4. AgNO3 (aq) + NaCl(aq ) ⎯⎯ Ionic equation

_______________

+

________________

________________________________________________

Net ionic equation ________________________________________________ AgNO3 (aq) + Na 2SO 4 (aq) ⎯⎯ → _______________ Ionic equation

+

________________

________________________________________________

Net ionic equation ________________________________________________

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AgNO3 (aq) + Na 3PO 4 (aq) ⎯⎯ → Ionic equation

__________________

+ __________________

___________________________________________________________

Net ionic equation ___________________________________________________________ Ca(NO3 ) 2 (aq ) + NaCl(aq ) ⎯⎯ → Ionic equation

__________________

+ __________________

___________________________________________________________

Net ionic equation ___________________________________________________________ Ca(NO3 ) 2 (aq ) + Na 2SO 4 (aq ) ⎯⎯ → __________________ + __________________ Ionic equation

___________________________________________________________

Net ionic equation ___________________________________________________________ Ca(NO3 ) 2 (aq ) + Na 3PO 4 (aq ) ⎯⎯ → __________________ + __________________ Ionic equation

___________________________________________________________

Net ionic equation ___________________________________________________________

B. Solubility of KNO 3 2. Mass of 4. Mass of 3. Temperature 5. Solubility 1. Mass of Container + KNO 3 Container When Crystals (g KNO 3 / KNO 3 100 mL H 2O) Appear

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6. Graph for the solubility of KNO 3 vs. temperature (°C)

Q2 According to your graph, what is the effect of increasing temperature on the solubility of KNO3?

Questions and Problems Q3 According to your graph, what is the solubility of KNO3 at 65 °C?

Q4 According to your graph, at what temperature does KNO3 have a solubility of 120 g/100 mL?

Q5 The solubility of sucrose (common table sugar) at 70 °C is 320 g/100 g H 2O. a. How much sucrose can dissolve in 200 g of water at 70 °C?

b. Will 400 g of sucrose dissolve in a teapot that contains 200 g of water at 70 °C? Explain.

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C. Testing the Hardness of Water Type of Water Sample

1. Drops of Soap Required

distilled water tap water mineral water seawater

Questions and Problems Q6 Which water sample was the hardest? Why?

D. Purification of Water Agent Added

1. After 15 min

2. After 30 min

H 2O (Reference) NaCI A12 (SO 4 )3 Na 2SO 4 Questions and Problems Q5 Which chemical produced the most rapid settling? Why?

Testing for Cations and Anions 16 LABORATORY GOALS • Determine the presence of a cation or anion by a chemical reaction. • Determine the presence of the cations and anions in an unknown salt.

LAB INFORMATION Time: Comments:

Related Topics:

2½ h Use small beakers to hold the reagents. Be sure to label each. HCl and HNO3 are strong acids, and NaOH is a strong base. Handle with care! If they are spilled on the skin, rinse thoroughly with water for 10 minutes. Dispose of the solutions and solids resulting from the tests as indicated by the instructor. Tear out the report sheets and place them beside the matching procedures. Ions, chemical change, solubility rules, net ionic equations

CHEMICAL CONCEPTS Solutions such as milk, coffee, tea, and orange juice contain an assortment of ions. In chemical reactions, these ions can give a distinctive flame test, undergo color changes, or form a gas or an insoluble solid. Your observations of the results of a test are the key to identifying those same ions when you test unknown solutions. After you look at a test that produces some change when a particular ion reacts, you will look for the same change in an unknown sample. If the test result is the same, you can assume that the ion is present in the unknown. If the same test result does not occur, then you can assume that the ion is not in the unknown: the test is negative.

A. Flame Tests for K + , Ca 2+ , and Na + ions The presence of K + , Ca 2+ , or Na + is determined by the distinctive color the ions give in flame tests. A flame test can be used to detect the presence of metal ions in a solution. The color of the flame of each ion depends on the atomic spectrum of the element. When a nichrome wire loop is dipped in a solution containing a salt of a metal ion, and placed in a flame, a color may be observed. The color is produced when heat causes electrons to absorb energy and move to a higher energy level. When the electrons fall to lower energy levels, energy is emitted that is equal to the energy level difference. If the energy emitted has a wavelength in the visible range, we see a distinctive color for that metal ion.

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▲When electrons lose energy, they fall to lower energy levels and emit light of a specific energy.

B. Tests for Ammonium Ion, NH 4 + , and Iron(III) Ion, Fe 3+ When the ammonium ion (NH 4+ ) is converted to ammonia (NH3 ), a distinctive odor is emitted, and the vapor turns red litmus paper to blue. Iron (III) ion (Fe3+ ) is detected by the distinctive red color it gives when it reacts with potassium thiocyanate, KSCN.

C. Tests for Negative Ions (Anions) When AgNO3 is added to the solutions to be tested for anions, several insoluble salts may form. When nitric acid, HNO3 , is added to this second solution, any solid that is present except AgCl will dissolve. Therefore, any solid that does not dissolve is AgCl, which indicates a positive test for Cl− in the original solution. A positive test for PO 43− is its reaction with ammonium molybdate, which forms a yellow solid. A positive test for SO 42− is its reaction with barium chloride, BaCl2 , which forms a precipitate of BaSO 4 . Barium may form insoluble salts with some other anions, but the addition of HNO3 dissolves any barium salts except BaSO 4 . The insoluble BaSO 4 remains in the test tube after HNO3 is added. When HCl is added to a solution containing CO32− , bubbles of CO 2 gas form.

D. Writing the Formula of Your Unknown Salt After you have completed testing the known solutions and unknown solution for cations and anions, you can identify the ions in your unknown. You should have a positive test that indicates you have one of the cations in your unknown, and another positive test that indicates that you have one of the anions in your unknown. Once you have identified the ions in your unknown, you can write the formula and name of your unknown salt.

E. Testing Consumer Products for Some Cations and Anions Consumer products contain many of the same ions that you will test for in this experiment. Once you have gone through the procedures and identified positive test results for cations and anions, the same procedures can be applied to consumer products to identify some of the ions present. After you carry out the tests for cations and anions, you may prepare a solution of a consumer product, and carry out the same tests.

Testing for Cations and Anions

EXPERIMENTAL PROCEDURES

181

GOGGLES REQUIRED!

Tests for Positive Ions (Cations) Materials: Four small beakers, one large beaker, test tubes and test tube rack, 3 M HCl, 6 M HNO3 , 6 M NaOH, unknown solutions In this experiment, you will need 2 mL of a solution. Measure 2 mL of water in the same-sized test tube that you will be using. During the experiment, obtain 2 mL samples by comparing the volume to the measured water sample. Be sure to label each beaker and test tube. Many beakers and test tubes have a frosted section to write on. If not, use a marking pencil or a label. Use separate beakers to obtain the following solutions: 15 mL of 3 M HCl 15 mL of 6 M HNO3 15 mL of 6 M NaOH 15 mL of an unknown salt solution from your instructor. Fill a large beaker half full of distilled water for rinsing the droppers and stirring rods. 1.

Record the sample number for your unknown solution.

A. Flame Tests for K + , Ca 2+ , and Na + ions Materials: Spot plate, Bunsen burner, flame-test wire, dropper bottles of 0.1 M NaCl, 0.1 M KCl, 0.1 M CaCl2 , Obtain a spot plate. Place 5 to 8 drops of 0.1 M solution of each of the following solutions into separate indentations: KCl, CaCl2 , NaCl, unknown solution To clean the flame-test wire, dip the loop at the end of the wire in 3 M HCl and heat the wire in the tip of the inner blue flame until there is no apparent color.

▲A wire loop is placed at the tip of the inner flame.

2.

Place the clean wire loop into the KCl solution and then into the tip of the inner flame. The color of K + does not last long, so look for it immediately. Record the color produced by the K + ion. Clean the wire as described previously.

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Place the wire loop in the CaCl2 solution and then into the flame. Record the color produced by the Ca 2+ ion. Clean the wire as before.

4.

Place the wire loop in the NaCl solution and then into the flame. Record the color produced by the Na + ion. Clean the wire as before. You may need to heat longer to remove the Na + from the wire loop.

5.

Place the wire loop in the unknown solution and then into the flame. Record the color produced by the unknown. Clean the wire as before.

6.

If there is a color that matches the color of an ion you tested in the flame tests, you can conclude that you have one of the ions K + , Ca 2+ , or Na + , in your unknown. If the flame test does not produce any color, then K + , Ca 2+ , and Na + are not present in your unknown. Record the presence of K + , Ca 2+ , or Na + in your unknown solution, or write none.

B. Tests for Ammonium Ion, NH 4 + , and Iron(III) Ion, Fe 3+ Materials: 0.1 M NH 4 Cl, 0.1 M FeCl3 , 0.1 M KSCN (potassium thiocyanate), red litmus paper, warm water bath, stirring rod Two other positive ions that may be present in your unknown are the ammonium ion, NH 4 + , and iron(III) ion, Fe3+ . Use the following tests to determine the presence of these cations in your unknown. 1.

Test for ammonium ion, NH 4+ Place 2 mL of 0.1 M NH 4 Cl in one test tube and 2 mL of your unknown in another test tube. Add 15 drops of 6 M NaOH to each. Carefully fan the vapors from the test tube toward you. You should notice the odor of ammonia from the known solution. Place a strip of moistened red litmus paper across the top of each test tube and set the test tubes in a beaker of warm water. Do not let the litmus paper touch the NaOH. The NH3 (g ) given off will turn the red litmus paper blue. NH 4 + (aq) + OH − (aq) ⎯⎯ → NH3 (g )+ H 2 O(l ) Ammonia

Record your test results for the known and unknown solutions. 2.

Test for iron(III) ion, Fe3+ Place 2 mL of 0.1 M FeCl3 in one test tube and 2 mL of your unknown in another test tube. Add 5 drops of 6 M HNO3 to each test tube. Add 2–3 drops of 0.1 M KSCN to each test tube. A deep red color indicates that Fe3+ is present. A faint pink color is not a positive test for iron. Fe3+ (aq ) + SCN − (aq) ⎯⎯ → Fe(SCN) 2+ (aq) Deep red color

Record your test results for the known and unknown solutions.

3.

From your test results, determine whether your unknown contains NH 4 + or Fe3+ , or none of these.

C. Tests for Negative Ions (Anions) Materials: Test tubes, test tube rack, 0.1 M NaCl, 0.1 M AgNO3 (dropper bottle), 3 M HCl, 6 M HNO3 , stirring rod, 0.1 M Na 2SO 4 , 0.1 M BaCl2 , 0.1 M Na 3 PO 4 , (NH 4 ) 2 MoO 4 (ammonium molybdate reagent), 0.1 M Na 2 CO3 , hot water bath

Testing for Cations and Anions

1.

183

Test for chloride ion, Cl − Place 2 mL of 0.1 M NaCl solution in a test tube and 2 mL of your unknown in another test tube. Add 5–10 drops of 0.1 M AgNO3 to each test tube. Caution: AgNO3 stains the skin. Add 10 drops of 6 M HNO3 to each test tube. Stir with a glass stirring rod. Any white solid that remains is AgCl(s). Ag + (aq ) + Cl− (aq ) ⎯⎯ → AgCl(s )

Record the test results for the known and your unknown.

2.

Test for sulfate ion, SO 42 − Place 2 mL of 0.1 M Na 2SO 4 solution in a test tube and 2 mL of your unknown in another test tube. Add 1 mL (20 drops) of BaCl2 to each test tube. Add 10 drops of 6 M HNO3 to each test tube. Any white solid that remains is BaSO 4 . Ba 2+ (aq) + SO 4 2− (aq ) ⎯⎯ → BaSO 4 (s )

Record the test results for the known and your unknown.

3.

Test for phosphate ion, PO4 3 − Place 2 mL of 0.1 M Na 3 PO 4 solution in a test tube and 2 mL of your unknown in another test tube. Add 10 drops of 6 M HNO3 to each. Warm the test tubes in a hot water bath (60°C). To each test tube, add 15 drops of ammonium molybdate solution, (NH 4 ) 2 MoO 4 . The formation of a yellow precipitate indicates the presence of PO 43− . Record the test results of the known and the unknown.

4.

Test for carbonate ion, CO 32 − Place 2 mL of 0.1 M Na 2 CO3 solution in a test tube and 2 mL of your unknown in another test tube. To each test tube, add 10 drops of 3 M HCl. Watch for the formation of bubbles of CO 2 gas. CO32− (aq) + 2H + (aq ) ⎯⎯ → CO 2 ( g ) + H 2 O(l ) Gas bubbles

Record the test results for the known and your unknown.

D. Writing the Formula of Your Unknown Salt Your unknown solution is made from a salt that has a cation and an anion. From your test results, you can identify one of the cations (Na + , K + , Ca 2+ , NH 4 + , or Fe3+ ) and one of the anions (Cl− , SO 4 2− , PO 43− , or CO32− ). For example, if you found that in the cation tests you got the same test result as for Ca 2+ and in the anion tests you got the same result as for Cl− , then the ions in your unknown salt would be Ca 2+ and Cl− . Then you would write the formula of your salt as CaCl2 .

1. 2. 3.

From your test results, write the symbols and names of the cation and the anion that gave positive tests for your unknown solution. From the cation and anion, write the formula of the salt in your unknown. Name the salt (ionic compound) in your unknown.

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E. Testing Consumer Products for Some Cations and Anions The tests in this experiment may be used to identify ions in samples of the following consumer products. Perhaps you have other ideas for products to test. Ask your instructor. In many of the products, there will be several cations and anions that give positive tests. Describe your results on the lab report. Product

Procedure

Juices

Obtain 25 to 30 mL of a light-colored fruit juice without fiber or pulp.

Sodas

Obtain 25 to 30 mL of a soft drink or mineral water. For colas, root beers, or others with deep colors, mix the soft drink with a small amount of activated charcoal in a small beaker. Charcoal will absorb the dyes. Filter.

Milk

Obtain 30 mL of nonfat milk. Add 10 mL of 0.1 M acetic acid (HC2 H3O 2 ). Small particles of protein (curds) will form. Filter. Gently boil the filtrate to reduce the volume to 15–20 mL.

Bone meal or plant food granules

In a beaker, mix a scoop of bone meal or plant food with 15 mL distilled water and 15 mL of 6M HNO3 . Heat gently (DO NOT BOIL) until most of the material dissolves. Cool and filter.

Window cleaner

Obtain 20 to 25 mL of a window cleaner.

Date

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Pre-Lab Study Questions

16

1. How can the presence of an ion in a solution be detected?

2. What tests would you use to identify a solution of Ag3 PO 4?

3. What tests would you use to identify a solution of FeCl3?

4. A flame test of a colorless solution gives a bright yellow color. When reacted with AgNO3 a white precipitate forms that dissolves when HNO3 is added. When HCl is added to the unknown solution, bubbles form. What is the compound in the colorless solution?

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REPORT SHEET

Testing for Cations and Anions

LAB

16

1. Unknown solution number _______________________

A. Flame Tests for K + , Ca 2+ , and Na + ions Cation Tested

Color produced

2. K + 3. Ca 2+ 4. Na + 5. Unknown 6. From your test results, does your unknown contain K + , Ca 2+ , or Na + , or none of these?

Explain your choice.

B. Tests for Ammonium Ion, NH 4 + , and Iron(III) Ion, Fe 3+ Cation

Test Results (Known)

Test Results (Unknown)

Odor

Odor

Litmus

Litmus

1. NH 4 + 2. Fe3+ 3. From your test results, does your unknown contain NH 4 + or Fe3+ , or none of these? Explain your choice.

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C. Tests for Negative Ions (Anions) Anion

Observations (Known)

Observations (Unknown Solution)

1. Cl− 2. SO 4 2− 3. PO 43− 4. CO32− 5. Identification of the negative ion in the unknown solution From your test results, what negative ion (anion) is present in your unknown?

_______________

Explain your choice.

D. Writing the Formula of Your Unknown Salt 1. Cation _____________

Name __________________

Anion _____________

Name __________________

2. Formula of your unknown salt _________________________________ 3. Name of your unknown salt _________________________________

E. Testing Consumer Products for Some Cations and Anions Product tested ___________________________________ Cation tests Flame tests (Na + , K + , Ca 2+ ) NH 4 + Fe3+ Anion tests Cl− SO 4 2− PO 43− CO32 −

Observations

Ion(s) present

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Questions and Problems Q1 How do the tests on known solutions containing cations and anions make it possible for you to identify the cations or anions in an unknown solution?

Q2 You have a solution that is composed of either NaCl or CaCl2 . What tests would you run to identify the compound?

Q3 If a solution turns a deep red color with a few drops of KSCN, what cation is present?

Q4 A plant food contains (NH 4 )3PO 4 . What tests would you run to verify the presence of the NH 4 + ion and the PO 43− ion?

Q5 Write the symbol of the cation or anion that gives each of the following reactions: a. Forms a precipitate with AgNO3 that does not dissolve in HNO3 b. Forms a gas with HCl c. Gives a bright, yellow-orange flame test d. Forms a precipitate with BaCl2 that does not dissolve in HNO3

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Properties of Solutions 17 LABORATORY GOALS • • • •

Perform chemical tests for chloride, glucose, and starch. Use dialysis to distinguish between solutions and colloids. Separate colloids from suspensions. Discuss the effects of hypotonic and hypertonic solutions on red blood cells.

LAB INFORMATION Time: Comments: Related Topics:

2½ h Label the containers to keep track of different solutions. Tear out the report sheets and place them beside the matching procedures. Solutions, colloids, suspensions, osmosis, hypertonic solutions, isotonic solutions, hypotonic solutions, dialysis Dispose of all chemicals safely as directed by your lab instructor.

CHEMICAL CONCEPTS A. Identification Tests In this experiment, you will perform identification tests for Cl – , glucose, and starch. You will use the tests to determine the presence or absence of Cl – , glucose, or starch. For each test, observe and record the initial properties of the reagent and the final appearance of the solution after the reagent is added. If a test is positive, there will be a change in the original properties of the reagent, such as a color change or the formation of a precipitate. If there is no change in the appearance of the reagent, the test is negative.

B. Osmosis and Dialysis Osmosis occurs when water moves through semipermeable membranes such as the walls of red blood cells. The osmotic pressure of an isotonic solution is the same as that of red blood cells. Both 0.9% NaCl (saline) and 5% glucose solutions are considered isotonic to the cells of the body. In isotonic solutions, the flow of water in and out of the red blood cells is equal. A hypotonic solution has a lower osmotic pressure than an isotonic solution; the osmotic pressure of a hypertonic solution is greater. In either case, the flow of water in and out of the cell is no longer equal, and the cell volume changes.

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▲In osmosis, water flows into the solution with a higher solute concentration until the flow of water becomes equal in both directions. When red blood cells are placed in a concentrated salt solution, they shrink. This process, called crenation, occurs because water diffuses out of the cells into the more concentrated salt solution. If red blood cells are placed in water, they expand and may rupture. This process, called hemolysis, occurs because water diffuses into the cells where there is a higher solute concentration. In both cases, osmosis occurs because water moves through a semipermeable membrane into the more concentrated solution.

▲In an isotonic solution, a red blood cell retains its normal volume. In a hypotonic solution water flows into a red blood cell, causing it to swell and possibly burst (hemolysis). In a hypertonic solution, water leaves the red blood cell, causing it to shrink (crenation).

In dialysis, small particles and water, but not colloids, move across a semipermeable membrane from a higher concentration to a lower one. Many of the membranes in the body are dialyzing membranes. For example, the intestinal tract consists of a semipermeable membrane that allows the solution particles from digestion to pass into the blood and lymph. Larger, incompletely digested food particles that are colloidal size or larger remain within the intestine. Dialyzing membranes are also used in hemodialysis to separate waste particles, particularly urea, from the blood.

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▲(a) Suspensions settle out; (b) suspensions are separated by a filter; (c) solution particles go through a semipermeable membrane, but colloids and suspensions do not.

C. Filtration In the process of filtration, gravity separates suspension particles from the solvent. The suspension particles are larger than the pores in the filter paper, so they cannot pass through and they become trapped on it. The colloidal and the solution particles are smaller and can pass through the pores of the filter paper.

▲In dialysis, solution particles move through a dialyzing membrane, but colloidal particles do not.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Identification Tests Materials: Four small beakers (50–100 mL), six test tubes, test tube rack, stirring rod, 50-mL graduated cylinder, test tube holder, stirring rod, droppers, boiling water bath, 10% NaCl, 1% starch, 10% glucose, 0.1 M AgNO3 , iodine reagent, Benedict’s reagent. Dropper bottles with these reagents may be available in the laboratory. Place small amounts of the reagents for these identification tests in small beakers for use throughout this experiment. Label the beakers and keep them at your desk for the remainder of the experiment. Use small beakers to obtain each of the following: Caution: AgNO 3 and iodine reagent stain and iodine is toxic! 3–4 mL of 0.1 M AgNO3 4–5 mL of iodine reagent 20–25 mL of Benedict’s reagent 10 mL of distilled water Set up the boiling water bath. Place 3 mL of water in a test tube for volume comparison. Chloride (Cl – ) test Place 3 mL of 10% NaCl in a test tube. Add 2 drops of 0.1 M AgNO3 and stir. Record your observations for the positive test result for chloride (Cl− ). Starch test Place 3 mL of 1% starch solution in a test tube. Add 2–3 drops of iodine reagent and stir. Record your observations for the positive test result for starch. Glucose test Place 3 mL of 10% glucose solution in a test tube. Add 3 mL of Benedict’s reagent and stir. Place the test tube in a boiling water bath and heat for 5 min. Record your observations for the positive test result for glucose.

B. Osmosis and Dialysis Materials: Cellophane tube (15–20 cm), three test tubes, test tube rack, test tube holder, stirring rod, droppers, small graduated cylinder, funnel, 100-mL beaker, 250-mL beaker, boiling water bath, 10% NaCl, 1% starch, 10% glucose, 0.1 M AgNO3 , iodine reagent, Benedict’s reagent. In a small beaker, mix together 10 mL of 10% NaCl, 10 mL of 10% glucose, and 10 mL of 1% starch solution. Obtain a piece of cellophane tubing (dialysis bag) and tie a firm knot in one end. Place a funnel in the open end and pour in about 20 mL of the mixture prepared above. Save the rest of the mixture for testing in part C. Tie a firm knot in the open end to close the bag. Rinse the outside of the dialysis bag with distilled water. Place the dialysis bag in a 250-mL beaker and cover with distilled water (see Figure 17.1).

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◄ FIGURE 17.1 A dialysis bag is placed in distilled water.

1. Testing for Cl – , Starch, and Glucose Initial (0 min) tests Immediately pour off 15 mL of the solution from the outside of the dialysis bag into a small beaker. Divide the 15 mL into three test tubes of 5 mL each.

Cl – Test Perform the identification test for chloride ion from Part A. Record your observations. Indicate if chloride is present. Starch Test Perform the identification test for starch from Part A. Record your observations. Indicate if starch is present. Glucose Test Perform the identification test for glucose from Part A. Record your observations. Indicate if glucose is present. 30-minute tests Pour off 15 mL of the solution outside the dialysis bag into a small beaker. Divide the 15 mL into three test tubes of 5 mL each.

Cl – Test Perform the identification test for chloride ion from Part A. Record your observations. Indicate if chloride is present. Starch Test Perform the identification test for starch from Part A. Record your observations. Indicate if starch is present. Glucose Test Perform the identification test for glucose from Part A. Record your observations. Indicate if glucose is present. Testing contents of the dialysis bag Open the dialysis bag and pour off 15 mL of the solution from the inside of the dialysis bag into a small beaker. Save the remainder of the bag’s contents for use in part C. Divide the 15 mL into three test tubes of 5 mL each.

Cl – Test Perform the identification test for chloride ion from Part A. Record your observations. Indicate if chloride is present. Starch Test Perform the identification test for starch from Part A. Record your observations. Indicate if starch is present. Glucose Test Perform the identification test for glucose from Part A. Record your observations. Indicate if glucose is present.

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C. Filtration Materials: Funnel, filter paper, 50-mL graduated cylinder, three test tubes, test tube rack, test tube holder, stirring rod, droppers, two 150-mL beakers, boiling water bath, powdered charcoal, 0.1 M AgNO3 , iodine reagent, Benedict’s reagent. To the remaining mixture of glucose, NaCl, and starch, from part B add a small amount of powdered charcoal and mix. Fold a piece of filter paper in half, and then fold it in half again. Open the fold to give a cone-like shape and place the filter-paper cone in a funnel. Push it gently against the sides and moisten with distilled water. Place a small beaker under the funnel. Pour the mixture into the filter paper. Collect the liquid (filtrate) that passes through the filter (see Figure 17.2).

◄ FIGURE 17.2 A mixture is filtered by pouring it through filter paper.

1. 2. 3.

Describe the appearance of any substance on the filter paper. What substance was present on the filter paper? Divide 15 mL of the solution in the beaker (filtrate) into three test tubes of 5 mL each. Cl – Test Perform the identification test for chloride ion from Part A. Record your observations. Indicate if chloride is present.

Starch Test Perform the identification test for starch from Part A. Record your observations. Indicate if starch is present. Glucose Test Perform the identification test for glucose from Part A. Record your observations. Indicate if glucose is present. 4.

From the test results in B and C, identify chloride ion, starch, glucose, and charcoal as solutions, colloids, or suspensions, and give your reasons.

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Pre-Lab Study Questions

17

1. In making pickles, a cucumber is placed in a concentrated salt solution. Explain what happens to the cucumber when it is left in the salt solution for some time.

2. Why is it important that cell membranes are semipermeable?

3. What are the differences between solutions, colloids, and suspensions?

4. Indicate the compartment (A or B) that will increase in volume for each of the following pairs of solutions separated by a semipermeable membrane: Solution A

Solution B

a. 10% (m/v) starch

4%(m/v) starch

______________________

b. 2%(m/v) albumin

5%(m/v) albumin

______________________

c. 8%(m/v) sucrose

0.8%(m/v) sucrose

______________________

5. Indicate whether a red blood cell will undergo hemolysis, crenation, or no change in each of the following: a. 10%(m/v) NaCl

______________________

b. 2%(m/v) glucose

______________________

c. H 2 O

______________________

d. 5%(m/v) glucose

______________________

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6. Predict what is present in each of the following: a. A solution that turns orange-red when mixed with Benedict’s ______________________ reagent and heated for 5 min. b. A solution that turns blue-black when mixed with a few drops ______________________ of iodine reagent. c. A solution that forms a white precipitate when mixed with a ______________________ few drops of AgNO3 .

Date

Name

Section

Team

Instructor

REPORT SHEET

Properties of Solutions

LAB

17

A. Identification Tests Test

Observations for Positive Test

Cl – Test Starch Test Glucose Test

B. Osmosis and Dialysis 1. Testing for Cl – , starch, and glucose Time

0 min

30 min

Contents of Dialysis Bag

Cl – Test Observations Cl – present? Starch Test Observations Starch present? Glucose Test Observations Glucose present?

Questions and Problems Q1 Which substance(s) were found in the water outside the dialysis bag?

Q2 How did those substance(s) get into the water outside the dialysis bag?

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Q3 What substance(s) were retained inside the dialysis bag? Why were they retained?

C. Filtration 1. Appearance of filter paper

________________________________________________

2. Substance present on filter paper ________________________________________________ 3. Identification Test

Observations

Substance Present

Cl – Test Starch Test Glucose Test

4. Substance

Solution, Colloid, Suspension?

Reason

Cl – Starch Glucose Charcoal

Questions and Problems Q4 State whether each of the following is isotonic, hypotonic, or hypertonic: a. H 2O

_____________________________________

b. 0.9% (m/v) NaCl

_____________________________________

c. 10% (m/v) glucose

_____________________________________

d. 3% (m/v) NaCl

_____________________________________

e. 0.2% (m/v) NaCl

_____________________________________

Q5 Predict the effect on a red blood cell (crenation, hemolysis, or none) that the following solutions would have: a. 2% (m/v) NaCl

_____________________________________

b. H 2O

_____________________________________

c. 5% (m/v) glucose

_____________________________________

d. 1% (m/v) glucose

_____________________________________

e. 10% (m/v) glucose

_____________________________________

Reaction Rates and Chemical Equilibrium 18 LABORATORY GOALS • Describe how temperature, concentration, and catalysts affect the rate of a reaction. • Use the concept of reversible reactions to explain chemical equilibrium. • Use Le Châtelier’s principle to describe the changes made in equilibrium concentrations when reaction conditions change.

LAB INFORMATION Time: Comments: Related Topics:

2 to 2½ h Read the section on chemical concepts before you come to lab. Tear out the report sheets and place them next to the corresponding procedures. Carefully observe the colors and states of the components in the test tubes before and after reactions. Factors affecting rates of reactions, energy of activation, chemical equilibrium, factors affecting equilibrium, Le Châtelier’s principle

CHEMICAL CONCEPTS A. Factors That Affect the Rate of a Reaction The rate or speed at which a reaction occurs depends on the amounts of the reactants, the temperature, and the presence or absence of a catalyst. Temperature Typically a reaction goes at a faster rate as the temperature is increased. In general, the rate of a reaction doubles for a 10 °C increase in temperature. We heat food to make it cook faster. When we have a fever, our metabolic reactions increase, including our rate of breathing and our pulse. In cardiac surgery, the body temperature is lowered to slow down the metabolic reactions and the amount of oxygen required by the brain. In a chemical reaction, collisions between the reactants lead to the formation of products. The energy that is needed to change the reactants into products is called the energy of activation. At high temperatures, more collisions have the energy to react and form products. At low temperatures, fewer collisions lead to products. Concentrations of Reactants If more reactant is added, products form faster. For example, we normally breathe air that is 20% oxygen (O 2 ). However, if a person is given pure oxygen (100%), oxygenated hemoglobin (HbO 2 ) is formed faster. Hb + O 2 ⎯⎯ → (HbO 2 )

If some reactant is removed, the rate at which product forms is slowed. At higher altitudes, the ability of hemoglobin to pick up O 2 is slowed because there is a lower amount of O 2 . This results in less O 2 reaching the cells in the body and especially the brain. A lowered level of O 2 in the brain may result in mental confusion and hallucinations, which has led to disastrous results for people who have tried to climb high mountains without adequate sources of oxygen. Catalysts A catalyst increases the rate of a reaction without becoming a part of the product. Enzymes are biological catalysts that make reactions in our cells occur at rates needed for cellular survival. Industry uses catalysts to make commercial reactions go faster. For example, the reaction of hydrogen with vegetable oils to produce margarine goes faster with a platinum (Pt) catalyst. Pt

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When a catalyst is present, the reaction takes a pathway which has a lower energy of activation. As a result, more reactants have sufficient energy to change to products.

▲A catalyst lowers the activation energy of a reaction.

B. Chemical Equilibrium: Reversible Reactions When a reaction begins, reactants are converted to products. This means that the number of product atoms or molecules will increase and more collisions will occur between the products. In some of the collisions between products, reactants are reformed. Most reactions proceed in two directions—forward (reactants to products) and reverse (products back to reactants). In that case, the reaction is called a reversible reaction.

▲One sample initially contains SO 2 and O 2 . Another sample contains only SO3 . At equilibrium, both equilibrium mixtures contain small amounts of reactants SO 2 and O 2 , and a large amount of product, SO3 .

Eventually, the rate of the forward reaction becomes equal to the rate of the reverse reaction. There is no further change in the concentrations of the reactants and the products, which means that the system is at equilibrium. For example, the reaction of ammonia (NH3 ) and water produces ammonium ion (NH 4 + ) and hydroxide ion (OH − ). At equilibrium, all of the products and reactants are present and their concentrations do not change. ⎯⎯ → NH 4 + ( aq) + OH − ( aq) NH3 (aq ) + H 2 O(l ) ←⎯ ⎯

▲Equilibrium occurs when the rate of the forward reaction becomes equal to that of the reverse reaction.

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Any change in the concentration of one of the reactants or products at equilibrium will create stress. For example, adding more reactant will shift the forward reaction in the direction of the products. If more of the product is added, the reverse reaction will shift in the direction of the reactants. In a similar way, removing some of the reactant will shift the equilibrium in the direction of the reactants (reverse reaction). Removing a product will shift a reaction in the direction of the products (forward reaction).

C. Changing Equilibrium Conditions: Le Châtelier’s Principle In another reversible reaction, the yellow Fe3+ ion and the colorless SCN − (thiocyanate) ion are in equilibrium with a deep red complex ion FeSCN 2+ . ⎯⎯ → FeSCN 2 + (aq ) Fe3+ (aq ) + SCN − (aq ) ←⎯ ⎯ yellow

colorless

red

When the equilibrium mixture contains mostly reactants, the solution has the yellow color of Fe3+ . When the equilibrium system contains mostly product FeSCN 2 + , the solution is a deep red. In any system at equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction. Therefore, an equilibrium constant can be written, which represents the ratio of concentrations of the products divided by the concentrations of the reactants. Kc =

[Products] [FeSCN 2+ ] = [Reactants] [Fe3+ ][SCN − ]

Color changes allow us to see a shift in equilibrium between reactants and products. If we add more reactant Fe3+ (yellow) or SCN − (colorless), the equilibrium shifts in the direction of the red product, FeSCN 2+ . If we add a substance such as Cl− that removes Fe3+ by forming colorless FeCl4 − , the equilibrium shifts in the direction of the reactants. The decrease in the product FeSCN 2+ ion is observed by the loss of the red color. ⎯⎯ → FeCl4 − (aq) Fe3+ (aq ) + 4Cl− (aq ) ←⎯ ⎯ yellow

colorless

A heterogeneous equilibrium contains solids or liquids as well as gases. However, since the concentrations of solids and liquids remain constant, they are not included in the equilibrium constant expression. For the heterogeneous equilibrium for the decomposition of CuSO 4 , ⎯⎯ → CuSO 4 ( s ) + 5H 2 O(g ) CuSO 4 i5H 2 O(s) ←⎯ ⎯

the K c expression does not include the concentration of the solids CuSO 4 i5H 2 O(s ) or CuSO 4 (s). It is written as K c = [H 2 O]5 . The exponent indicates the coefficient of H 2 O in the equation.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Factors That Affect the Rate of a Reaction A.1 Effect of Temperature Materials: Two test tubes, two 400-mL beakers for hot and cold water baths, ice, 20 mL of 0.1 M HCl, two scoops (four spatula tips) of NaHCO3 (s ), thermometer

In this experiment, NaHCO3 is added to samples of HCl at different temperatures. NaHCO3 ( s ) + H3O + (aq ) ⎯⎯ → CO 2 ( g ) + Na + (aq ) + 2H 2 O(l )

Place 10 mL of 0.1 M HCl in each of two large test tubes. 1. 2. 3. 4.

Place test tube 1 in a 400-mL beaker half-filled with crushed ice and water. Cool to a temperature of 10 °C or lower. Place test tube 2 in a 400-mL beaker half-filled with water and heat to 50–60 °C. Remove both of the test tubes and place them in a test tube rack. Measure the temperature of each Sample and record. Immediately add one scoop or two spatula tips of NaHCO3 to each sample. Observe the fizzing (bubbles) in each test tube. Record which test tube clears first.

A.2 Effect of Changing Reactant Concentration Materials: Three test tubes, three cleaned pieces of magnesium ribbon of the same length (2–3 cm), 10-mL graduated cylinder, 10 mL each of 1.0 M HCl, 2.0 M HCl, 3.0 M HCl, watch with second hand or stop watch In this experiment, HCl solutions of different concentrations are added to equal amounts of Mg. Mg( s ) + 2HCl(aq) ⎯⎯ → MgCl( g ) + H 2 ( g )

Caution: HCl is a strong acid. Work carefully. If any HCl gets in your eyes, wash immediately at the eye fountain. Be sure to wash your hands at the end of this experiment. 1. 2. 3.

Label test tubes 1 to 3. Add one piece of Mg to each test tube. Measure 10. mL 1.0 M HCl in a graduated cylinder and add HCl to test tube 1. Immediately begin recording the time. Stop recording the time when all the Mg disappears. Record total time. Repeat with test tubes 2 and 3 using 10. mL of the 2.0 and 3.0 M HCl. Dispose of the materials as directed by your instructor.

A.3 Effect of a Catalyst Materials: Five test tubes, 10 mL of 3% H 2 O 2 , MnO 2 , spatula, piece of Zn, freshly cut potato, pieces of boiled potato,

In this experiment, different substances are added to a H 2 O 2 solution to determine their catalytic effect on the decomposition of hydrogen peroxide, H 2 O 2 . One kind of biological catalyst is the enzyme catalase, which is present in living cells, where it breaks down hydrogen peroxide, H 2 O 2 . 2H 2 O 2 (aq) ⎯⎯ → 2H 2 O (l ) + O 2 ( g )

1.

Label five test tubes 1 to 5 and place in a test tube rack.

2.

Pour 2 mL of 3% H 2 O 2 into each of the test tubes.

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205

Test tube 1 is the reference. Record its appearance. Test tube 2: Add a spatula tip of MnO 2 . Record your observations. Test tube 3: Add a small piece of zinc metal. Record your observations. Test tube 4: Add a small piece of freshly cut potato. Record your observations. Test tube 5: Add a small piece of boiled potato. Record your observations. Dispose of the materials as directed by your instructor.

B. Chemical Equilibrium: Reversible Reactions B.1 Reaction of a Hydrate Materials: Test tube containing CuSO 4 i5H 2 O, test tube holder, Bunsen burner, dropper, small beaker with water In this experiment, we will dehydrate copper(II) pentahydrate and reverse the reaction by adding water. The compound CuSO 4 i5H 2 O(s) is called a hydrate because it contains water molecules as part of its solid structure. By heating the hydrate crystals, water is lost, which leaves CuSO 4 ( s ), which is an anhydrate. The hydrate is reformed when water is added to the anhydrate. ⎯⎯ → CuSO 4 ( s ) + 5H 2 O( g ) CuSO 4 i5H 2 O(s ) ←⎯ ⎯ blue (hydrate )

1.

2.

white (anhydrate )

Record the initial appearance of the CuSO 4 i5H 2 O in the test tube. Using a test tube holder, heat the test tube containing the blue CuSO 4 i5H 2 O (hydrate) over a low flame, constantly moving the test tube in a circular pattern. Do not let the open end point toward anyone. When some of the sample has turned from blue to white, observe the upper, cooler, part of the test tube for drops of water (condensation). Record your observations. Turn off the burner and allow the white CuSO 4 ( s ) (anhydrate) to cool back to room temperature. Using a dropper, slowly add a few drops of water to the white CuSO 4 ( s ) (anhydrate). Record your observations. Return the test tube with CuSO 4 i5H 2 O(s ) to the lab supply area.

B.2 Reaction of Copper(II) Ion and Hydroxide Materials: two test tubes, test tube rack, droppers, stirring rod, 10 mL of 0.1 M CuCl2 , 5 mL of 0.1 M NaOH, 2 mL of 0.1 M HCl In this experiment, copper ion (II), Cu 2+ , is reacted with hydroxide ion, OH – , to form solid Cu(OH) 2 . When more OH – is added, the [OH – ] increases in the Cu(OH) 2 equilibrium system. ⎯⎯ → Cu(OH) 2 ( s) Cu 2+ (aq ) + 2OH − (aq) ←⎯ ⎯

When HCl is added, the H3O + it provides decreases the [OH – ] in the Cu(OH)2 equilibrium system. H3O + (aq) + OH − (aq ) ⎯⎯ → 2H 2 O(l )

1. 2.

3. 4.

Place 3 mL of 0.1 M CuCl2 in each of two test tubes and label as test tube 1 and test tube 2. Record the color of the CuCl2 (aq ) solution. Add drops of 0.1 M NaOH to test tube 1 and stir until a light blue, cloudy precipitate of Cu(OH)2 is produced. Add the same number of drops to test tube 2. Describe the initial appearance of the contents of test tube 1 and test tube 2. Using a dropper, add more drops of 0.1 M NaOH to test tube 1, which increases [OH − ]. Describe any change in the appearance in test tube 1. Using a dropper, add drops of 0.1 M HCl to test tube 2, which decreases [OH − ]. Stir. Describe any change in the appearance in test tube 2. Dispose of the materials as directed by your instructor.

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C. Changing Equilibrium Conditions: Le Châtelier’s Principle In this experiment, we will place stress on an equilibrium system of yellow Fe3+ ion, the colorless SCN − (thiocyanate) ion, and the deep red product FeSCN 2+ . ⎯⎯ → FeSCN 2+ (aq) Fe3+ (aq ) + SCN − (aq ) ←⎯ ⎯ yellow

colorless

red

When more Fe3+ or SCN − is added to the equilibrium mixture, the increase in the reactant shifts the reaction in the direction of the deep red FeSCN 2+ (aq ). When Cl− is added to the equilibrium mixture, it forms colorless FeCl4 − , which decreases [Fe3+ ] and shifts the reaction in the direction of the reactants. We will also look at the stress on equilibrium caused by raising and lowering the temperature. C.1 Effect of Changing Reactant Concentration Materials: Six test tubes, test tube rack, small 100-mL beaker, 10-mL or 25-mL graduated cylinder, 10 mL of 0.01 M Fe(NO3 )3 , 10 mL of 0.01 M KSCN, stirring rod, 1 M Fe(NO3 )3 , 1 M KSCN, 3 M HCl Caution: 3 M HCl is a strong acid. Work carefully. If any HCl gets in your eyes, wash immediately at the eye fountain. Be sure to wash your hands at the end of this experiment. 1.

Use a small graduated cylinder to measure 10 mL of 0.01 M Fe(NO3 )3 , and 10 mL of 0.01 M KSCN and pour into a small beaker. Stir

2.

Set up six test tubes in a test tube rack and label each. To each test tube, add 3 mL of the mixture. Test Tube 1 (reference) Add 10 drops of water to test tube 1. Stir. Record the color you observe in the reference test tube 1. Test tube 1 is prepared as a reference, which means that you compare it to the color in each of the other test tubes (2−6). Test Tube 2 Add 10 drops of 1 M Fe(NO3 )3 to test tube 2. Stir. Record the color you observe in test tube 2 compared to the reference. Test Tube 3 Add 10 drops of 1 M KSCN to test tube 3. Stir. Record the color you observe in test tube 3 compared to the reference. Test Tube 4 Add 10 drops of 3 M HCl to test tube 4. Stir. Record the color you observe in test tube 4 compared to the reference.

C.2 Effect of Temperature Materials: Two 400-mL beakers, ice, Bunsen burner or hot plate, thermometer 1. 2. 3. 4.

Place test tube 5 in a 400-mL beaker half-filled with crushed ice and water. Cool to a temperature of 10 °C or lower. Add 10 drops of water and stir. Place test tube 6 in a 400-mL beaker half-filled with water. Heat to 50–60 °C and then turn off the heat. Add 10 drops of water and stir. After 10 minutes, remove the test tubes from the hot and cold baths. Record the color you observe compared to the reference, test tube 1. State whether the [FeSCN 2 + ] increased or decreased in test tubes 5 and 6. Dispose of the materials as directed by your instructor.

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

18

1. What factors increase the rate of a chemical reaction?

2. When is equilibrium established in a reversible reaction?

3. How does a system at equilibrium respond to the addition of more reactant?

4. How does a system at equilibrium respond to the addition of more product?

5. In Part C, we look at the following reaction: ⎯⎯ → FeSCN 2+ (aq ) Fe3+ (aq ) + SCN − (aq ) ←⎯ ⎯

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a. Write the equilibrium constant expression for the reaction if the equilibrium constant is 78.

b. Does the equilibrium mixture contain more products or reactants?

c. If more SCN − is added to the equilibrium mixture, will the red color of the mixture intensify or lessen? Explain.

Date

Name

Section

Team

Instructor

REPORT SHEET

Reaction Rates and Chemical Equilibrium

LAB

18

A. Factors That Affect the Rate of a Reaction A.1 Effect of Temperature Test Tube Temperature (3)

Observations (4)

1.

____________

_____________________________________

2.

____________

_____________________________________

Which test tube cleared first (4)?

_____________________________________

Q1 How did an increase in temperature affect the rate of a reaction?

A.2 Effect of Changing Reactant Concentration Test Tube

Concentration of HCl

1.

1.0 M HCl

2.

2.0 M HCl

3.

3.0 M HCl

Initial time (2)

Final time (3)

Total time (3) – (2)

Rank the test tubes in order of slowest to fastest. ______ < ______ < ______ Q2 How did an increase in the concentration of HCl cause a difference in the rates of reaction of HCl with Mg?

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A.3 Effect of a Catalyst Test Tube

Observations (2)

Catalyst Present (Yes or No?)

1. Reference 2. MnO 2 3. Zn 4. fresh potato 5. boiled potato Which substance was most active as a catalyst? ____________________________ Q3 How does a catalyst change the rate of a reaction without affecting its equilibrium?

Q4 How do you explain the difference in catalytic activity of the fresh and boiled potatoes?

B. Chemical Equilibrium: Reversible Reactions B.1 Reaction of a Hydrate Appearance

Heating (1)

Addition of Water (2)

Initial Final Why did water droplets form at the mouth of the test tube when heated?

Q5 When the hydrate of CuSO 4 i5H 2O(s ) was heated, did the reaction proceed in the direction of the reactants or products?

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211

Q6 When water was added to the anhydrate of CuSO 4 (s ), did the reaction proceed in the direction of the reactants or products?

B.2 Reaction of Copper(II) Ion and Hydroxide Test tube 1

Test tube 2

Initial color of CuCl2 solution (1) Initial appearance Cu(OH) 2 ( s ) (2) Change in appearance (3) Change in appearance (4) Q7 In test tube 1, did the addition of OH − cause the reaction to go in the direction of the reactants or products? Explain.

Q8 In test tube 2, did the addition of H3O + cause the reaction to go in the direction of the reactants or products? Explain.

C. Changing Equilibrium Conditions: Le Châtelier’s Principle C.1 Effect of Changing Reactant Concentration Test tube

1. Reference 2. Fe3+ added 3. SCN − added 4. Fe3+ removed

Color

Is color deeper or [FeSCN 2+ ] lighter red increases compared to the or decreases? reference?

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Q9 In test tube 2, did the equilibrium shift in the direction of products or reactants? Explain.

Q10 In test tube 3, what is the evidence that the shift in equilibrium was in the direction of products?

Q11 In test tube 4, what is the evidence that the shift in equilibrium was in the direction of reactants?

C.2 Effect of Temperature Test tube

Color Temperature (increased or decreased)

Color is deeper or [FeSCN 2+ ] lighter red than (increased or the reference (3) decreased) (4)

1. Reference 5. Ice water 6. Warm water Q12 In test tube 5, what was the evidence that decreasing the temperature shifted the equilibrium in the direction of the products?

Q13 In test tube 6, what was the evidence that increasing the temperature shifted the equilibrium in the direction of the reactants?

Q14 Complete the following equation with heat as a reactant or a product: ⎯⎯ → FeSCN 2+ (aq ) Fe3+ (aq ) + SCN − (aq ) ←⎯ ⎯

Acids, Bases, pH, and Buffers 19 LABORATORY GOALS • Extract a naturally occurring dye from red cabbage to use as a pH indicator. • Measure the pH of several substances using the cabbage indicator and a pH meter. • Calculate pH from the [H3O + ] or the [OH − ] of a solution. • Calculate the molar concentration and percentage of acetic acid in vinegar. • Observe the changes in pH as acid or base is added to buffered and unbuffered solutions.

LAB INFORMATION Time: Comments:

Related Topics:

2½ h Students may be asked to bring a red cabbage and/or colorless household samples to class. Share test tubes with your lab neighbors to prepare the pH reference solutions. Tear out the report sheets and place them beside the matching procedures. Acids, bases, pH, buffers

CHEMICAL CONCEPTS An acid is a substance that dissolves in water and donates a hydrogen ion, or proton (H + ), to water. In the laboratory, we have been using acids such as hydrochloric acid (HCl) and nitric acid (HNO3 ).

You use acids and bases every day. For example, there are acids in oranges, lemons, and vinegar. In this experiment we will use acetic acid (HC2 H3O 2 ). Acetic acid is the acid in vinegar that gives it a sour taste. Some typical bases used in the laboratory are sodium hydroxide (NaOH) and potassium hydroxide (KOH). Most of the common bases dissolve in water and produce hydroxide ions, OH − .

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▲A common base produces a cation and an OH − anion in an aqueous solution. An important weak base found in the laboratory and in some household cleaners is ammonia. In water, ammonia accepts H + to form ammonium and hydroxide ions.

A. Reference Colors for pH Using Red Cabbage Indicator The pH of a solution tells us whether a solution is acidic, basic, or neutral. On the pH scale, values below 7 are acidic, a value equal to 7 is neutral, and values above 7 are basic. Typically, the pH scale has values between 0 and 14.

Many natural substances contain dyes that produce distinctive colors at different pH values. By extracting (removing) the dye from red cabbage leaves, a natural pH indicator can be prepared. Adding the red cabbage solution to a set of pH reference solutions will produce a series of distinctive colors. When the red cabbage solution is added to a test sample, the color produced can be matched to the colors of the pH reference set to determine the pH of the sample. A pH meter can also be used to measure pH.

B. Measuring pH The molar concentrations of H3O + and OH − are indicated by square brackets. In pure water, [H3O + ] = [OH − ] = 1.0 × 10−7 Μ at 25ºC. The ion constant product of water, K w , can thus be calculated. K w = [H3O + ][OH − ] = [1.0 × 10−7 ][1.0 × 10−7 ] = 1.0 ×10−14

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215

If the [H3O + ] or [OH − ] for an acid or a base is known, the other can be calculated. For example, if an acid has a [H3O + ] = 1.0 × 10−4 M, we can find the [OH − ] of the solution by solving the K w expression for [OH − ]. [OH − ] =

1.0 × 10−14 [H3O + ]

=

1.0 × 10−14 1.0 × 10−4

= 1.0 × 10−10 M

The pH of a solution is a measure of its [H3O + ]. It is defined as the negative log of the hydrogen ion concentration. pH = −log [H3O + ]

Therefore, a solution with [H3O + ] = 1.0 × 10−9 M has a pH of 9.00 and is basic. pH = −log [1.0 × 10−9 ] pH = − [ − 9.00] pH = 9.00 The number of significant figures to the right of the decimal point in the pH is equal to the number of significant figures in the coefficient of the [H3O + ].

C. Effect of Buffers on pH The pH of the blood is maintained between 7.35 and 7.45 by buffers in the body. If blood pH goes above or below that range, it can damage or destroy the cells in the blood. Buffers maintain the pH of a solution by reacting with and neutralizing small amounts of acids or bases. Many buffers contain a weak acid and its salt. The weak acid reacts with excess base, and the anion of the salt picks up excess H + . It is the ability of a buffer to react with excess acid or base that maintains the pH of a solution. One important buffer system in the blood is the hydrogen carbonate buffer, which is carbonic acid, H 2 CO3 (weak acid) and hydrogen carbonate anion, HCO3− (salt). When base (OH − ) is added, it reacts with the weak acid in the buffer and produces hydrogen carbonate ion and water. H 2 CO3 (aq) + OH − (aq) ⎯⎯ → HCO3− (aq) + H 2 O(l )

When acid (H + ) enters the blood, it reacts with the HCO3− anion and re-forms carbonic acid: H 2 CO3 (aq) + H 2 Ο(l ) ←⎯ ⎯ H3O + (aq) + HCO3 (aq )

In this experiment, the effect of an acid and a base on the pH of a buffer and an unbuffered system will be determined.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Reference Colors for pH Using Red Cabbage Indicator Materials: Red cabbage leaves, 400-mL beaker, distilled water, Bunsen burner or hot plate, small (100-mL or 150-mL) beaker, 13 test tubes, two test tube racks, set of buffers with pH ranging from 1 to 13 Preparing red cabbage indicator Tightly pack torn red cabbage leaves in a 150-mL beaker, then transfer leaves to a 400-mL beaker. Add about 150 mL of distilled water or enough water to cover the leaves. Heat, using a Bunsen burner or a hot plate, but do not boil. When the solution has attained a dark purple color, turn off the heat and allow the solution to cool. Preparing pH reference standards Place about 30 mL of cabbage dye indicator in a small beaker. Arrange 13 labeled test tubes in two test tube racks. You may need to combine your test tubes with those of your neighbor (your instructor may prepare a pH reference set for the entire class). Pour 2 mL of each buffer solution in separate test tubes to give a pH reference set with pH 1−13. Caution: Low pH values are strongly acidic; high pH values are strongly basic. Work with care. To each test tube, add about 2 mL of the cooled red cabbage solution. To obtain a deeper color, add more cabbage solution. Shake the test tube to mix. Record the color of the each of the solutions at different pH values. Keep this reference set for the next part of the experiment.

B. Measuring pH Materials: Shell vials or test tubes, samples to test for pH (shampoo, conditioner, mouthwash, antacids, detergents, fruit juice, vinegar, cleaners, aspirin, etc.), cabbage juice indicator from part A, pH meter, calibration buffers, wash bottle, Kimwipes. Using red cabbage indicator to determine pH 1.

2.

Place 2 mL of each liquid sample in a shell vial (or a test tube). Add 2 mL of red cabbage solution to each sample. Shake the test tube to mix. Record the color. Prepare solutions of solid/viscous samples by adding small amount to 2 mL of water and mix thoroughly. Compare the color obtained is part 1 to the colors of the pH reference set. Identify and record the pH that gives the best color match to the color of your sample. Repeat for other samples. Dispose of all chemicals as directed by your lab instructor.

Using a pH meter to determine pH Your instructor will demonstrate the use of the pH meter. 3. Using a new sample of a solution, determine the pH with the pH meter. Rinse off the pH electrode and use the pH meter to find the pH of other samples. 4. Identify each sample as acidic, basic, or neutral.

C. Effect of Buffers on pH Materials: Buffer with a high pH (9–11), buffer with a low pH (3–4), droppers, test tubes, graduated cylinder, 0.1 M NaCl, 0.1 M HCl, 0.1 M NaOH, pH meter, cabbage juice indicator from part A Use a graduated cylinder to measure each of the following solutions and place in a separate test tube: 10 mL of H 2 O

10 mL of 0.1 M NaCl

10 mL of a buffer with a high pH

10 mL of a buffer with a low pH

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C.1 Effect of adding acid 1. 2. 3. 4. 5.

Determine the pH of each by adding 2 mL of cabbage indicator to each sample and or by using a pH meter. Record. Add 5 drops of 0.1 M HCl solution to each of the four test tubes from part 1 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. Add 5 more drops of 0.1 M HCl solution to each of the four test tubes from part 2 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. Determine the pH change in each of the solutions. Identify the solutions that are buffers.

C.2 Effect of adding base Use a graduated cylinder to measure each of the following solutions and place in a separate test tube:

1. 2. 3. 4. 5.

10 mL of H 2 O

10 mL of 0.1 M NaCl

10 mL of a buffer with a high pH

10 mL of a buffer with a low pH

Determine the pH of each by adding 2 mL of cabbage indicator to each sample or by using a pH meter. Record. Add 5 drops of 0.1 M NaOH solution to each of the four test tubes from part 1 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. Add 5 more drops of 0.1 M NaOH solution to each of the four test tubes from part 2 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. Determine the pH change in each of the solutions. Identify the solutions that are buffers.

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Pre-Lab Study Questions

19

1. How is the pH of a solution related to the [H3O + ]?

2. Using the equation for K w , explain how [OH − ] changes when more H3O + is added.

3. Is a solution with a pH of 12.0 acidic or basic? 4. Is a solution with a pH of 2.0 acidic or basic? 5. What is a buffer?

6. If you add acid or base to water, how will the pH change?

7. If you add acid or base to a buffer, how will the pH change?

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Date

Name

Section

Team

Instructor

REPORT SHEET

Acids, Bases, pH, and Buffers

LAB

19

A. Reference Colors for pH Using Red Cabbage Indicator pH 1

Colors of Acidic Solutions

pH 8

2

9

3

10

4

11

5

12

6

13 pH 7

Colors of Basic Solutions

Color of Neutral Solution

B. Measuring pH Substance

1. Color with Indicator

2. pH Using Indicator

3. pH Using pH Meter

4. Acidic, Basic, or Neutral?

Household cleaners vinegar ammonia Drinks, juices lemon juice apple juice Detergents, shampoos shampoo detergent hair conditioner Health aids mouthwash antacid aspirin

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Other items

Questions and Problems Q1 Complete the following table: [H 3O + ]

[OH − ]

pH

Acidic, Basic, or Neutral?

2.0 × 10−6 9.8 3.5 ×10−3 Neutral Q2 A solution has a [OH − ] = 4.0 × 10−5 M. What are the [H3O + ] and the pH of the solution?

Q3 A sample of 0.0084 mol of HCl is dissolved in water to make a 1500-mL solution. Calculate the molarity of the HCl solution, the [H3O + ], and the pH. For a strong acid such as HCl, the [H3O + ] is the same as the molarity of the HCl solution. HCl(aq) + H 2O(l ) ⎯⎯ → H3O + (aq) + Cl− (aq)

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223

C. Effect of Buffers on pH C.1 Effect of adding acid Solution

1. Initial pH

2. pH after 5 drops of HCl

3. pH after 10 drops of HCl

4. pH change

5. Buffer yes or no?

2. pH after 5 drops of NaOH

3. pH after 10 drops of NaOH

4. pH change

5. Buffer yes or no?

H 2O 0.1 M NaCl High pH buffer Low pH buffer C.2 Effect of adding base Solution

1. Initial pH

H 2O 0.1 M NaCl High pH buffer Low pH buffer Questions and Problems Q4 Which solution(s) showed the greatest change in pH? Why?

Q5 Which solutions(s) showed little or no change in pH? Why?

Q6 Normally, the pH of the human body is fixed in a very narrow range between 7.35 and 7.45. A patient with an acidotic blood pH of 7.3 may be treated with an alkali such as sodium hydrogen carbonate. Why would this treatment raise the pH of the blood?

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Acid-Base Titration 20 LABORATORY GOALS • Set up a buret and perform a titration using appropriate technique. • Calculate the molar concentration and mass/volume percent (m/v) of acetic acid in vinegar. • Determine the acid-absorbing capacity of a commercial antacid.

LAB INFORMATION Time: Comments:

Related Topics:

2h Tear out the lab report sheets and place them next to the matching procedures. Practice observing the color change for the indicators before you do the titrations. Carefully read the markings on the buret. Students may bring their own antacid samples to test. Acid, base, neutralization, titration, percent concentration, molarity

CHEMICAL CONCEPTS In a typical neutralization reaction, the H + ions from the acid combine with OH − ions from a base to produce a salt and water (H 2 O). The salt is composed of the anion from the acid and the cation from the base. For example, the neutralization of HCl by NaOH is written as HCl(aq) + NaOH(aq) ⎯⎯ → NaCl(aq ) + H 2 O(l )

If we write the reactants and the products as ions, we see that the H + and the OH − form water. The net ionic equation is written by removing the Na + and Cl− ions, which appear on both sides of the equation as spectator ions. H + (aq ) + Cl – (aq ) + Na + (aq) + OH − (aq ) ⎯⎯ → Na + (aq ) + Cl – (aq) + H 2 O(l ) H + (aq) + OH – (aq ) ⎯⎯ → H 2 O(l )

In a complete neutralization, the amount of H + will be equal to the amount of OH − .

A. Concentration of Acetic Acid in Vinegar Vinegar is an aqueous solution of acetic acid, HC2 H3O 2 , or CH3COOH. The amount of acetic acid in a vinegar solution can be determined by neutralizing the acid with a base, in this case NaOH. As shown in the following equation, one mole of acetic acid is neutralized by one mole of NaOH. HC2 H3O 2 (aq ) + NaOH (aq) ⎯⎯ → Na C2 H3O 2 (aq ) + H 2 O(l ) Acetic acid

Base

Salt

A titration involves the addition of a specific amount of base required to neutralize the acid in a sample. When all the H + (or H3O+ ) from the acid has been neutralized, an indicator in the sample will change color. This change in the indicator color determines the endpoint, which indicates that the acid has been neutralized by the addition of base. At this point no further base is added. The volume of base used to neutralize the acid is then determined. In this experiment, the indicator phenolphthalein changes from colorless in acid to a faint but permanent pink color at the endpoint.

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Molarity (M) of Acetic Acid Using the average measured volume of the NaOH, and its molarity (on the label), the moles of NaOH used can be calculated. Moles NaOH = L NaOH ×

moles NaOH 1 L NaOH

When an acid is completely neutralized, the moles of NaOH used are equal to the moles of acetic acid (HC2 H3O 2 ) present in the sample. This occurs because there is the same number of H + and OH − ions in the reactants: Moles of HC2 H3O 2 = moles of NaOH

Using the moles of acid, the molarity of acetic acid in the 5.0-mL sample of vinegar is calculated. Molarity (M) HC2 H3O 2 =

moles HC 2 H3O 2 . 0.0050 L vinegar

Mass/Volume Percent (m/v) of Acetic Acid To calculate the percent (m/v) of HC2 H3O 2 in vinegar, we convert the moles of acetic acid to grams using the molar mass of acetic acid, 60.1 g/mole. g HC2 H3O 2 = mole HC2 H3O 2 ×

Mass/volume percent (m/v) =

60.1 g HC2 H3O 2 1 mole HC2 H3O 2

.

g HC2 H3O 2 × 100%. 5.0 mL

B. Titration of an Antacid Stomach acid is primarily hydrochloric acid (HCl), which has a concentration of about 0.1 M. When a person is under stress, excess HCl may be produced, causing discomfort. An agent called an antacid is used to neutralize some of the excess stomach acid (see Table 20.1). TABLE 20.1 Basic Compounds in Some Antacids Antacid

Base(s) in Antacids

Amphojel, Gaviscon, Mylanta Milk of magnesia Mylanta, Maalox, Di-Gel, Gelusil, Riopan, Equate, Maalox(liquid) Bisodol, Rolaids Titralac, Tums, Pepto-Bismol, Maalox (Tablet) Alka-Seltzer

Al(OH)3 Mg(OH) 2 Mg(OH) 2 , Al(OH)3 CaCO3 , Mg(OH) 2 CaCO3 NaHCO3 , KHCO3

In this experiment, the volume (mL) of 0.1 M HCl that can be absorbed by some common antacids will be determined. Milk of magnesia, Di-Gel Mg(OH)2 (aq ) + 2HCl(aq ) ⎯⎯ → MgCl2 (aq ) + 2H 2 O(l ) Tums

CaCO3 ( s ) + 2HCl(aq) ⎯⎯ → CaCl2 (aq) + CO 2 ( g ) + H 2 O(l )

Acid-Base Titration

EXPERIMENTAL PROCEDURES

227

GOGGLES REQUIRED!

A. Concentration of Acetic Acid in Vinegar Materials: Vinegar (white), two beakers (150- and 250-mL), 250-mL Erlenmeyer flask, 10-mL graduated cylinder (or a 5-mL pipet and bulb), phenolphthalein indicator, 50-mL buret (or 25-mL buret), buret clamp, small funnel to fit buret, 0.1 M NaOH (standardized), white paper or paper towel Dispose of all chemicals safely as directed by your lab instructor. 1.

Obtain about 20 mL of vinegar in a small beaker. Record the brand of vinegar and the percent acetic acid stated on the label. Using a 10-mL graduated cylinder or a 5.0-mL pipet, transfer 5.0 mL of vinegar to a 250-mL Erlenmeyer flask. (See the instructions for using a pipet.) Add about 25 mL of distilled water to increase the volume of the solution for titration. This will not affect your results. Add 2–3 drops of the phenolphthalein indicator to the solution in the flask.

Using a pipet: Place the pipet bulb on the end of the pipet and squeeze the bulb to remove air. Place the tip of the pipet in the vinegar in the beaker and allow the bulb to expand by slowly releasing your grip on it so that the change in pressure will draw the liquid into the pipet. Be careful not to allow liquid to be drawn into the bulb. When the liquid goes above the volume line of the pipet, carefully remove the bulb and quickly place your second finger (index finger) tightly over the end of the stem. By adjusting the pressure of the index finger, lower the liquid to the etched line that marks the 5.0-mL volume and stop. Lift the pipet with its 5.0 mL of vinegar out of the beaker and let it drain into an Erlenmeyer flask. Touch the tip of the pipet to the wall of the flask to remove the rest of the vinegar. A small amount that remains in the tip has been included in the calibration of the pipet, allow it to remain in the pipet. (see Figure 20.1). Caution: If you are using an automatic pipet, dispense a 5.0-mL volume. If you are using a pipet, always use a suction bulb to draw vinegar into a pipet. Do not pipet by mouth!

▲ FIGURE 20.1 A bulb is used to draw liquid into a pipet. Obtain a 50-mL or 25-mL buret and place it in a buret clamp or butterfly clamp (see Figure 20.2). Using a 250-mL beaker, obtain about 100 mL of 0.1 M NaOH solution. (If you are using a 25.0-mL buret, use a 0.2 M NaOH solution.) 2.

Record the molarity (M) of the NaOH solution. Rinse the buret with two 5-mL portions of the NaOH solution. Discard the washings.

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▲ FIGURE 20.2 A buret setup for acid–base titration provides a measured amount of NaOH solution. Observe the markings on the buret. The top is marked 0.0 mL and the bottom is marked 50.0 mL. Place a small funnel in the top of the buret and slowly pour the NaOH solution into the funnel. As the NaOH solution nears the top line, lift the funnel and allow the solution to fill above the top line (0.0 mL). Using a small beaker (waste), slowly open the stopcock and drain NaOH solution until the meniscus of the NaOH solution is at the 0.00 mL line or below. The buret tip should be full of NaOH solution, and free of bubbles. 3.

Record the initial level of NaOH solution in the buret.

Titrating the Vinegar Solution with NaOH Place the flask containing the vinegar solution on a piece of white paper under the buret tip. (Be sure you added indicator). Slowly add NaOH solution to the vinegar in the flask by opening and closing the stopcock with your left hand (if you are right-handed). Swirl the flask with your right hand to mix the acid and the base (see Figure 20.3). Initially, the pink color produced by the indicator will disappear quickly.

▲ FIGURE 20.3 During a titration, the solution in the flask is swirled as NaOH is added to the acid sample. As you near the endpoint, the pink color will be more persistent and disappear more slowly. Slow down the addition of the NaOH to drops at this time. When one drop of the NaOH solution gives a faint, permanent pink color to the sample, stop adding NaOH. You have reached the endpoint of the titration. 4.

Record the final buret reading of the NaOH solution. For trials 2 and 3, rinse the flask with distilled water and repeat the procedure with a new sample of vinegar.

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229

Calculations Molarity (M) of Acetic Acid 5.

Calculate the volume of NaOH solution used to neutralize the vinegar sample(s) for Trials 1, 2, and 3. The volume is the difference between the final and initial volume readings on the buret. Volume of NaOH = Final Volume (NaOH) (4) − initial volume (NaOH) ( 3)

6.

Calculate the average volume of the NaOH solution needed for the titration. Average volume (mL) =

7.

volume (1) + volume (2) + volume (3) 3

Convert the average volume of NaOH (6) solution, in milliliters, to its volume in liters (L). L = mL NaOH (6) ×

8.

1 L NaOH 1000 mL NaOH

Calculate the moles of NaOH using the volume (L) (7) and molarity of the NaOH (2) solution. Moles of NaOH = L NaOH (7) ×

9.

moles NaOH 1 L NaOH

The number of moles of HC2 H3O 2 in the vinegar is equal to the moles of NaOH used in the titration. Convert the moles of NaOH (8) to the moles of acid present in the vinegar. Moles of HC2 H3O 2 = moles of NaOH (8)

10.

Calculate the molarity (M) of the acetic acid (HC2 H3O 2 ) solution in the vinegar sample. Molarity (M) HC2 H3O 2 =

moles HC 2 H3O 2 (9) 0.0050 L vinegar

Mass/Volume Percent (m/v) of Acetic Acid 11.

Using the number of moles of acetic acid (9), calculate the number of grams of acetic acid in the vinegar sample using the molar mass of acetic acid, 60.1 g/mole. grams of HC2 H3O 2 = mole HC2 H3O 2 (9) ×

12.

60.1 g HC2 H3O 2 1 mole HC2 H3O 2

Calculate the percent (mass/volume) acetic acid in the vinegar using the grams of HC2 H3O 2 (11) and the initial volume of 5.0 mL of vinegar. Mass/volume percent (m/v) =

g HC2 H3O 2 (11) × 100% 5.0 mL

B. Titration of an Antacid Materials: Antacid products, two small beakers (100- and 150-mL), mortar and pestle, 250-mL Erlenmeyer flask, methyl orange, graduated cylinder, 0.10 M HCl (standardized), 0.10 M NaOH (standardized), buret, small funnel, buret clamp, white paper or paper towel 1.

Record the name of the antacid and the base(s) listed on the label. Use the mortar and pestle to crush the antacid tablet.

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2.

Weigh a 250-mL Erlenmeyer flask. Record its mass. Do not round off. Transfer some of the crushed tablet (up to 0.5 g) to the flask.

3.

Weigh the flask and the crushed antacid. Record the mass.

4.

Record the molarity of the HCl solution.

5.

Using a graduated cylinder, add 50.0 mL of HCl solution to the flask containing the antacid. Swirl to dissolve the antacid. The solution may be cloudy because of the starches in the tablet. This will not affect the titration. Add 5 drops of methyl orange indicator. The antacid solution should turn red. If the antacid solution is yellow, carefully add more HCl solution in 10.0-mL portions until the solution is red. Record the total volume of HCl solution that you add.

6.

Record the molarity of NaOH solution as stated on the reagent bottle.

7.

Fill a buret with the NaOH solution provided. The initial level should be at or below the 0.0 mark. Record the initial level of NaOH solution in the buret.

8.

Titrate the sample using NaOH solution until a yellow color forms. This is a back titration, which means we are neutralizing the excess HCl that was not neutralized by the antacid. Record the final volume of NaOH solution. Rinse the buret with distilled water.

Calculations 9.

Calculate the mass of the antacid tablet in the flask. grams of antacid = mass of flask and antacid (3) – mass of flask (2)

10.

Calculate the volume (mL) of NaOH used to back titrate the HCl. mL NaOH solution = final volume NaOH solution (8) – initial level NaOH solution (7)

11.

State the volume of excess HCl solution. This is equal to the volume of NaOH solution (10) because the molarity of the NaOH solution matched the molarity of the HCl solution.

12.

Calculate the volume of HCl solution that was neutralized by the antacid. This is the difference between the total volume of HCl solution (5) added to the antacid minus the volume of the excess HCl solution (11).

13.

Because stomach acid is about the same molarity as 0.1 M HCl solution, we can state the neutralizing power of the antacid as mL of stomach acid per gram of antacid. mL of 0.1 M HCl neutralized (12) grams of antacid (9)

=

mL stomach acid neutralized 1 g antacid

Record values (mL stomach acid/1 g antacid) obtained in the class for a different antacid (antacid 2). 14.

Write neutralization equations for the base(s) listed in 1 for your antacid products.

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Pre-Lab Study Questions

20

1. What is neutralization?

2. Write an equation for the neutralization of H 2SO4 by KOH.

3. What is a titration?

4. What is the function of an indicator in a titration?

5. How many milliliters of 0.258 M NaOH are required to completely neutralize 2.00 g of acetic acid HC2 H3O 2?

6. What is the volume, in milliliters, of a 0.350 M KOH needed to completely neutralize 15.0 mL of a 0.250 M H 2SO 4 solution? 2KOH(aq ) + H 2SO 4 (aq ) ⎯⎯ → K 2SO 4 (aq ) + 2H 2 O(l )

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Date

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REPORT SHEET

Acid-Base Titration

LAB

20

A. Concentration of Acetic Acid in Vinegar 1. Brand _______________ Volume 5.0 mL (% on label) _____________ % 2. Molarity (M) of NaOH _____________________________ M Trial 1

Trial 2

Trial 3

3. Initial NaOH level in buret 4. Final NaOH level in buret 5. Volume (mL) of NaOH used 6. Average volume (mL) 7. Average volume in liters (L) 8. Moles of NaOH used in titration (Show calculations.)

________________ mole NaOH

9. Moles of HC2 H3O 2 neutralized by NaOH ________________ mole HC2 H3O 2 ________________ M HC2 H3O 2 10. Molarity of HC2 H3O2 (Show calculations.)

11. Grams of HC2 H3O 2 (Show calculations.)

________________ g HC2 H3O 2

12. Percent (m/v) HC2 H3O 2 in vinegar (Show calculations.)

________________ % HC2 H3O 2

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Questions and Problems Q1 How many mL of a 0.10 M NaOH solution are needed to neutralize 15 mL of a 0.20 M H3PO 4 solution? 3NaOH(aq ) + H3PO 4 (aq ) ⎯⎯ → Na 3PO 4 (aq) + 3H 2O(l )

B. Titration of an Antacid Antacid 1

Antacid 2

1. Brand of antacid Bases(s) in antacid 2. Mass of flask 3. Mass of flask and antacid 4. Molarity of HCl solution 5. Total volume (mL) of HCl solution added 6. Molarity of NaOH solution 7. Initial volume (mL) of NaOH solution 8. Final volume (mL) of NaOH solution 9. Mass of antacid 10. Volume of NaOH solution used 11. Volume of excess HCl solution 12. Volume of HCl solution neutralized by antacid 13.

mL stomach acid 1 g antacid

14. Write the neutralization equations that take place in the stomach with the base(s) present in the antacid product. 1. ____________________________________________________________ 2. ____________________________________________________________ Q2 How many grams of Mg(OH) 2 will be needed to neutralize 25 mL of stomach acid if stomach acid is 0.10 M HCl?

Organic Compounds: Alkanes 21 LABORATORY GOALS • • • • • •

Observe chemical and physical properties of organic and inorganic compounds. Draw formulas for alkanes from their three-dimensional models. Write the names of alkanes from their structural formulas. Construct models of isomers of alkanes. Draw the condensed structural formulas for cycloalkanes and haloalkanes. Classify organic compounds according to their functional groups.

LAB INFORMATION Time: Comments: Related Topics:

2–3 h Tear out the report sheets and place them next to the matching procedures. Organic compounds, hydrocarbons, solubility, combustion, complete structural formula, alkane, cycloalkane, haloalkane, complete structural formula, condensed structural formula, constitutional isomers, naming alkanes, and functional groups

CHEMICAL CONCEPTS A. Comparison of Organic and Inorganic Compounds A.1 Physical Properties Organic compounds are made of carbon and hydrogen and sometimes oxygen, sulfur, nitrogen, or a halogen (F, Cl, Br, or I). Of all the elements, only carbon atoms bond to many more carbon atoms, a unique ability that gives rise to many more organic compounds than known inorganic compounds. The covalent bonds in organic compounds and the ionic bonds in inorganic compounds account for several of the general differences we observe in their physical and chemical properties. A.2 Solubility Typically, ionic inorganic compounds are soluble in water, a polar solvent. Many organic compounds are nonpolar and thus are not soluble in water; however, they are soluble in organic solvents, which are also nonpolar. A general rule for solubility is that “like dissolves like.”

▲ Vegetable oil, a mixture of nonpolar organic compounds, is not soluble in water.

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A.3 Combustion Many organic compounds react with oxygen, a reaction called combustion, to form carbon dioxide and water. Combustion is the reaction that occurs when gasoline burns with oxygen in the engine of a car or when natural gas, methane, burns in a heater or stove. In a combustion reaction, heat is given off; the reaction is exothermic. Equations for the combustion of methane and propane are written as follows: CH 4 (g ) + 2O 2 (g ) → CO 2 (g ) + 2H 2 O(g ) + heat Methane C3 H8 (g ) + 5O 2 (g ) → 3CO 2 (g ) + 4H 2 O(g ) + heat Propane

▲ The propane fuel in the tank undergoes combustion, which provides energy. Table 21.1 is a summary of the properties of organic and inorganic compounds. TABLE 21.1 Comparing Some Properties of Organic and Inorganic Compounds Organic Compounds

Inorganic Compounds

C and H, sometimes O, S, N, F, Cl, Br, or I Covalent bonds Soluble in nonpolar solvents, not water Low melting and boiling points Strong, distinct odors Flammable

Most metals and nonmetals Ionic or polar covalent bonds Soluble in water High melting and boiling points Usually no odor Not flammable

B. Alkanes The saturated hydrocarbons represent a group of organic compounds composed of carbon and hydrogen. Alkanes and cycloalkanes are called saturated hydrocarbons because their carbon atoms are connected by only single bonds. In each type of alkane, each carbon atom has four valence electrons and must always have four single bonds.

Organic Compounds: Alkanes

237

B.1 Structures and Names of Alkanes To learn more about the three-dimensional structure of organic compounds, it is helpful to examine or build models using a ball-and-stick model kit, in which colored wooden (or plastic) balls represent the atoms in organic compounds (see Table 21.2). TABLE 21.2 Elements and Bonds Represented in the Organic Model Kit Color

Element

Number of Bonds

Black White Red Green Orange Purple Blue

carbon hydrogen oxygen chlorine bromine iodine nitrogen

4 1 2 1 1 1 3

Bonds Sticks, springs Methane, CH 4 , is a hydrocarbon consisting of one carbon atom and four hydrogen atoms. The model of methane shows the three-dimensional shape, a tetrahedron, around a carbon atom.

CH 4

Three-dimensional model

Expanded structural formula

Condensed structural formula

To represent the three-dimensional model on paper, its shape is flattened and the carbon atom is shown attached to four hydrogen atoms. This type of formula is called an expanded structural formula. For convenience, chemists use a simplified structure called a condensed structural formula, which is written with the hydrogen atoms grouped with each carbon atom. In the formula, the number of hydrogen atoms is written as a subscript. The complete structural formula and the condensed structural formula for C2 H 6 (ethane) are shown below.

Three-dimensional model

Expanded structural formula

Condensed structural formula

Names of Alkanes The names of alkanes all end with “-ane.” The names of organic compounds are based on the names of the alkane family (see Table 21.3).

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TABLE 21.3 Names and Formulas of the First Ten Alkanes Name

Formula

Methane

CH 4

Ethane Propane Butane Pentane Hexane Heptane Octane Nonane Decane

CH3 ~ CH3 CH3 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH3 CH3 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH 2 ~ CH3

B.2 Isomers of Alkanes Isomers represent two or more different structural formulas of the same molecular formula. One isomer cannot be converted to the other without breaking and forming new bonds. Isomers also have different physical and chemical properties. This phenomenon of isomerism is one of the reasons for the vast number of organic compounds.

B.3 Cycloalkanes A cycloalkane has a cyclic or ring structure. The skeletal formula is used to depict a cycloalkane by showing only the bonds that outline the geometric shape of the compound. For example, the skeletal formula for cyclobutane is represented by a square. Examples of the various formulas that represent cyclobutane are shown below.

Expanded structural formula

Condensed structural formula

Skeletal formula

B.4 Haloalkanes In a haloalkane, a halogen atom such as chlorine (Cl) or bromine (Br) replaces one or more hydrogen atoms of an alkane or a cycloalkane.

Organic Compounds: Alkanes

Expanded Structural Formula

Condensed Structural Formula

239

Name

C. Functional Groups Although there are millions of organic compounds, they can be classified into organic families. Each family contains a characteristic structural feature called a functional group, which is a certain atom or group of atoms that give similar physical and chemical properties to each member of the organic family. Because the organic compounds in a family contain the same functional group, they undergo the same types of chemical reactions. Alkenes, Alkynes, and Aromatic Compounds Alkenes, alkynes, and aromatic compounds are hydrocarbons that consist of only carbon and hydrogen atoms. We have seen that alkanes contain only carbon-carbon single bonds (C ~ C). However, alkenes contain one or more carbon-carbon double bonds (C £ C), and alkynes contain one or more carbon-carbon triple bonds (C ¢ C). Aromatic compounds contain a six-carbon benzene ring (a hexagon) with three double bonds on alternating carbon atoms. This is symbolized by a circle in the center of the hexagon. Class

Functional Group

H 2 C £ CH 2

Alkene Alkyne

Example

~C¢C~

HC ¢ CH

Aromatic

Alcohols and Ethers Alcohols and ethers contain an oxygen atom. Alcohols have a hydroxyl group, which is the ~ OH , bonded to a carbon atom. In an ether, the oxygen atom is bonded to two carbon atoms. Class

Functional Group

Example

Alcohol

~ OH

CH3 ~ CH 2 ~ OH

Ether

~O~

CH3 ~ O ~ CH3

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Aldehydes and Ketones Aldehydes and ketones contain a carbonyl group, which is a carbon-oxygen double bond (C £ O). In an aldehyde, the carbonyl carbon bonds to at least one hydrogen atom. In a ketone, the carbon in the carbonyl group bonds to two other carbon atoms. Class

Functional Group

Example

Aldehyde Ketone Carboxylic Acids and Esters Carboxylic acids and esters contain the carboxyl functional group, which is a combination of a carbonyl and a hydroxyl group. In a carboxylic acid, the hydroxyl oxygen is bonded to a hydrogen atom, whereas in an ester this oxygen is bonded to a carbon atom and not to hydrogen. Class

Functional Group

Example

Carboxylic acid Ester Amines and Amides Amines and amides contain a nitrogen atom. In an amine, one or more carbon groups are bonded to the N atom. In an amide, the hydroxyl group of a carboxylic acid is replaced by a nitrogen group. Class Amine

Amide

Functional Group

Example CH3 ~ NH 2

Organic Compounds: Alkanes

EXPERIMENTAL PROCEDURES

241

GOGGLES REQUIRED!

A. Comparison of Organic and Inorganic Compounds A.1 Physical Properties Materials: Display of compounds in test tubes in the hood: sodium chloride, potassium iodide, toluene, and cyclohexane; chemistry handbook or access to the Internet 1. 2. 3. 4. 5.

Observe the samples in the test tubes in the hood: sodium chloride, potassium iodide, toluene, and cyclohexane. Record the molecular formula for each. Use a reference source if necessary. Record the physical state (solid, liquid, or gas). Look up the melting point of each compound using a chemistry handbook or the Internet. Record. Identify the types of bonds in each as ionic or covalent. Identify each as an organic or inorganic compound.

A.2 Solubility Materials: Display of samples in the hood containing NaCl or toluene added to cyclohexane or water. 1. 2. 3. 4. 5. 6.

Observe the sample in which NaCl is added to cyclohexane, a nonpolar solvent. Record whether NaCl is soluble (S) or insoluble (I) in cyclohexane. Observe the sample to which NaCl is added to water, a polar solvent. Record whether NaCl is soluble (S) or insoluble (I) in water. Identify NaCl as organic or inorganic. Observe the sample to which toluene is added to cyclohexane, a nonpolar solvent. Record whether toluene is soluble (S) or insoluble (I) in cyclohexane. Observe a sample to which toluene is added to water, a polar solvent. Record whether toluene is soluble (S) or insoluble (I) in water. Identify toluene as organic or inorganic.

A.3 Combustion Materials: 2 evaporating dishes, spatulas, wood splints, NaCl(s), cyclohexane (demonstration) 1. 2.

Observe your instructor ignite a splint and hold the flame to a sample of NaCl. Record whether NaCl undergoes combustion. Identify NaCl as an organic or inorganic compound. Observe your instructor ignite a splint and hold the flame to a sample of cyclohexane. Record whether cyclohexane undergoes combustion. Identify cyclohexane as an organic or inorganic compound.

B. Alkanes B.1 Structures and Names of Alkanes Materials: Display of models of methane, ethane, and propane 1. 2. 3.

Observe the model of methane. Draw its expanded and condensed structural formulas. Observe the model of ethane. Draw its expanded and condensed structural formulas. Observe the model of propane. Draw its expanded and condensed structural formulas.

B.2 Isomers of Alkanes Materials: Display of models: butane, 2-methylpropane, 2,2-dimethylpropane; chemistry handbook or access to the Internet 1.

Observe the model of butane. Draw its expanded and condensed structural formulas.

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Laboratory Manual for General, Organic, and Biological Chemistry

Observe the model of 2-methylpropane, which is an isomer of butane. Draw its expanded and condensed structural formulas. Using a chemistry handbook or the Internet, obtain the molar mass, boiling point, and density of each isomer. Draw the condensed structural formulas for the three isomers of C5 H12 : pentane, 2-methylbutane, and 2,2-dimethylpropane. Using a chemistry handbook or the Internet, obtain the molar mass, boiling point, and density of each isomer.

B.3 Cycloalkanes Materials: Display of models: cyclopentane and cyclohexane 1. 2.

Observe the model of cyclopentane. Draw its condensed structural and skeletal formulas. Observe the model of cyclohexane. Draw its condensed structural and skeletal formulas.

B.4 Haloalkanes Materials: Display of models: chloromethane, 1,2-dibromoethane, and 2-iodopropane 1. 2. 3.

Observe the model of chloromethane. Draw its expanded and condensed structural formulas. Observe the model of 1,2-dibromoethane. Draw its expanded and condensed structural formulas. Observe the model of 2-iodopropane. Draw its expanded and condensed structural formulas.

4.

Draw the condensed structural formulas of the four isomers of C3 H 6 Cl2 . Name each isomer.

C. Functional Groups Materials: Display of the models with functional groups Classify each model of an organic compound according to its functional group.

Date

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Pre-Lab Study Questions

21

1. What elements are present in alkanes?

2. Would you expect hexane to be soluble in water? Why?

3. Which is more flammable, hexane or potassium sulfate?

4. Draw the condensed structural formula of hexane.

5. Why are 1-chlorobutane and 2-chlorobutane structural isomers?

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6. Classify each of the following compounds according to its functional group. a.

______________________________

b.

______________________________

c.

______________________________

d.

______________________________

e.

______________________________

Date

Name

Section

Team

Instructor

REPORT SHEET

Organic Compounds: Alkanes

LAB

21

A. Comparison of Organic and Inorganic Compounds A.1 Physical Properties Name

1. Formula 2. Physical 3. Melting 4. Type of 5. Organic or State Point Bonds Inorganic?

Sodium chloride Potassium iodide Toluene Cyclohexane

A.2 Solubility Solute 1. Soluble in Cyclohexane?

2. Soluble in Water? 3. Organic or Inorganic?

NaCl

Solute

4. Soluble in Cyclohexane?

5. Soluble in Water? 6. Organic or Inorganic?

Toluene

A.3 Combustion Compound

Undergoes Combustion? (yes/no)

Organic or Inorganic?

1. NaCl 2. Cyclohexane

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Q1 From your observations of the chemical and physical properties of organic and inorganic compounds, complete the following table: Property

Organic Compounds

Inorganic Compounds

Elements Bonding Melting point Solubility in water Flammability

Q2 A white solid is soluble in water and is not flammable. Would you expect it to be organic or inorganic? Explain your reason.

Q3 A clear liquid forms a layer when added to water. Would you expect it to be organic or inorganic? Explain your reason.

B. Alkanes B.1 Structures and Names of Alkanes Compound Expanded Structural Formula 1. Methane

2. Ethane

3. Propane

Condensed Structural Formula

Organic Compounds: Alkanes

247

B.2 Isomers of Alkanes Isomers of C4 H10 Compound

Butane

2-Methylpropane

1. Expanded Structural Formula

2. Condensed Structural Formula

3. Molar Mass Boiling Point Density

Isomers of C5 H12 Compound

Pentane

2-Methylbutane 2,2-Dimethylpropane

4. Condensed Structural Formula

5. Molar Mass Boiling Point Density Questions and Problems Q4 What physical property is identical for the isomers of C5 H12?

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B.3 Cycloalkanes Compound

Condensed Structural Formula

Skeletal Formula

Expanded Structural Formula

Condensed Structural Formula

1. Cyclopentane

2. Cyclohexane

B.4 Haloalkanes Compound 1. Chloromethane

2. 1,2-Dibromoethane

3. 2-Iodopropane

Isomers of C3 H 6Cl 2 Condensed Structural Formula

Condensed Structural Formula

Name:

Name:

Name:

Name:

Organic Compounds: Alkanes

249

Questions and Problems Q5 Draw the condensed structural formula for each of the following: a. 1,4-dichlorocyclohexane b. 2,3-dimethylpentane

Q6 Write the correct name of the following alkanes: a.

___________________________________

b.

___________________________________

c.

___________________________________

C. Functional Groups Compound

Classification

Compound

Classification

Q7 Classify each of the following according to its functional group: a. ________________________________

b. ________________________________ c. ________________________________

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d. ________________________________

e. ________________________________

f. ________________________________

Reactions of Unsaturated Hydrocarbons

22

LABORATORY GOALS • Observe the reactions of unsaturated hydrocarbons with oxygen, bromine, and potassium permanganate. • Use chemical tests to distinguish alkanes from alkenes. • Draw the products of addition reactions of alkenes and alkynes.

LAB INFORMATION Time: Comments:

Related Topics:

2h Tear out the report sheets and place them beside the matching procedures. Caution: Hydrocarbons are flammable. Use very small amounts. Do not use any burners during these labs. Avoid touching the chemicals. Dispose of organic wastes in proper containers as instructed by your lab instructor. Saturated, unsaturated, and aromatic hydrocarbons; haloalkanes; addition reactions

CHEMICAL CONCEPTS A. Types of Unsaturated Hydrocarbons Alkanes are saturated hydrocarbons containing single bonds between carbon atoms. Alkenes and alkynes are unsaturated hydrocarbons containing double or triple bonds. The double or triple bond, which is unsaturated, is a very reactive site in an alkene or alkyne. Aromatic compounds are hydrocarbons with a benzene ring.

Naming Alkenes and Alkynes The IUPAC names for alkenes and alkynes are related to the names of alkanes with the same number of carbon atoms, but ending in ene or yne (see Table 22.1). Skeletal formulas are written for unsaturated hydrocarbons to show the position of the double or triple bond.

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TABLE 22.1 Comparison of Names for Alkanes, Alkenes, and Alkynes Alkane

Alkene

Ethane

Ethene (ethylene)

Alkyne HC ¢ CH Ethyne (acetylene)

CH3 ~ CH3

H 2 C £ CH 2

CH3 ~ CH 2 ~ CH3

CH3 ~ CH £ CH 2

CH3 ~ C ¢ CH

Propane

Propene

Propyne

B. Addition Reaction: Bromine Test When bromine (Br2 ) reacts with an alkene, the dark red color of the Br2 disappears quickly as the atoms of bromine form bonds with the carbon atoms in the double bond. If the red color disappears rapidly, we know the compound contains an unsaturated site.

The reaction of bromine alkanes is slow and requires light. In the reaction, one H in the alkane is replaced with a Br. When bromine is added to an alkane, the red color persists for several minutes before it fades. Aromatic compounds (benzene rings) are not reactive with bromine.

C. Oxidation: Potassium Permanganate Test In this test, you will discover that potassium permanganate (KMnO 4 ) oxidizes alkenes, but not alkanes or aromatic compounds. In the reaction, the purple color of KMnO 4 changes to the muddy brown of manganese dioxide (MnO 2 ). The product is a diol. H 2 C £ CH 2 Colorless

+

KMnO 4 Purple

⎯⎯⎯ → HO ~ CH 2 ~ CH 2 ~ OH + MnO 2 Colorless

Brown

D. Identification of an Unknown Use the test results to identify your unknown as one of the compounds used in this experiment.

Reactions of Unsaturated Hydrocarbons

EXPERIMENTAL PROCEDURES

253

GOGGLES REQUIRED!

A. Types of Unsaturated Hydrocarbons Materials: Display of models of unsaturated hydrocarbons: ethene (ethylene), propene, cyclobutene, cis-2-butene, and ethyne Observe the models of ethene (ethylene), propene, cyclobutene, cis-2-butene, and ethyne (acetylene). Draw each of their condensed structural formulas.

B. Addition Reaction: Bromine Test (May be an instructor demonstration) Materials: 4 test tubes, test tube rack, cyclohexane, cyclohexene, unknowns, dropper bottles of 1% bromine solution (in methylene chloride) Caution: Work in the hood. The fumes of Br2 can irritate the throat and sinuses. If bromine is spilled on the skin, flood with water for 10 minutes. 1.

2.

Cyclohexene Place 15 drops of cyclohexene in a dry test tube. Carefully add 3–4 drops of the bromine solution and mix by shaking. Record whether the red color disappears in 2–3 s or persists for a longer time. A positive test for an unsaturated compound is a fast disappearance of the red bromine color. Indicate if the test result is positive or negative. Cyclohexane Repeat steps 1 and 2 by placing 15 drops of cyclohexane in a dry test tube. Carefully add 3–4 drops of the bromine solution and mix by shaking. Record whether the red color disappears quickly or not. Indicate if the test result is positive or negative. Unknown Repeat steps 1 and 2 by placing 15 drops of an unknown in a dry test tube. Carefully add 3–4 drops of the bromine solution and mix by shaking. Record whether the red color disappears quickly or not. Indicate if the test result is positive or negative.

C. Oxidation: Potassium Permanganate Test Materials: 4 test tubes, test tube rack, 1% KMnO 4 , cyclohexane, cyclohexene, unknowns 1.

Cyclohexene Place 5 drops of cyclohexene in a dry test tube. Carefully add 15 drops of 1% KMnO 4 solution. Caution: KMnO4 stains the skin. Record your observations.

2.

A positive test for an unsaturated compound is a change in color from purple to brown in 60 seconds or less. Indicate if the test result is positive or negative. Cyclohexane Repeat steps 1 and 2 by placing 5 drops of cyclohexane in a dry test tube. Carefully add 15 drops of 1% KMnO 4 solution. Record your observations. Indicate if the test result is positive or negative. Unknown Repeat steps 1 and 2 by placing 5 drops of an unknown in a dry test tube. Carefully add 15 drops of 1% KMnO 4 solution. Record your observations. Indicate if the test result is positive or negative.

D. Identification of an Unknown From your test results, indicate if the bromine test and the KMnO 4 test were positive or negative. Identify your unknown as a saturated or unsaturated hydrocarbon. State your reasoning.

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Pre-Lab Study Questions

22

1. a. What changes in color occur when bromine reacts with an alkene?

b. What changes in color take place when KMnO 4 reacts with an alkene?

2. Why is the reaction of ethene with bromine called an addition reaction?

3. If you used a bromine test to distinguish between hexane and 2-hexene, in which would the red color of bromine persist?

4. If you used a permanganate test to distinguish between hexane and 1-hexene, in which would the purple color of KMnO 4 persist?

5. Draw the condensed structural formula of each of the following: a. 2-methyl-1-pentene

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b. 3-methylcyclopentene

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REPORT SHEET

Reactions of Unsaturated Hydrocarbons

LAB

22

A. Types of Unsaturated Hydrocarbons Models of Unsaturated Hydrocarbons Compound Condensed Structural Compound Formula Ethene cis-2-butene

Propene

Condensed Structural Formula

Ethyne (acetylene)

Cyclobutene

Questions and Problems Q1 Write the IUPAC names of the following compounds: a.

___________________________________

b.

___________________________________

c.

___________________________________

d.

___________________________________

e.

___________________________________

f.

___________________________________

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B. Addition Reaction: Bromine Test Compound

1. Disappearance of 2. Test Results Bromine Color (fast/slow) (Positive/Negative)

Cyclohexene

Cyclohexane

Unknown

C. Oxidation: Potassium Permanganate Test Compound

1. KMnO4 Observations

2. Test Results (Positive/Negative)

Cyclohexene

Cyclohexane

Unknown

Questions and Problems Q2 Complete and balance the following reactions: a. b.

D. Identification of an Unknown Results of tests with unknown Unknown #

Bromine Test

Explain your conclusion.

KMnO4 Test

Saturated or Unsaturated?

Alcohols and Phenols 23 LABORATORY GOALS • • • •

Describe the chemical and physical properties of alcohols and phenols. Classify an alcohol as primary, secondary, or tertiary. Perform a chemical test to distinguish between the classes of alcohols. Draw the condensed structural formulas of the oxidation products of alcohols.

LAB INFORMATION Time: Caution:

2h Be careful when you work with chromate solution. It contains concentrated acid. Do not use burners in lab when you work with flammable organic compounds. Tear out the lab report sheets and place them beside the matching procedures. Related Topics: Alcohols, classification of alcohols, solubility of alcohols in water, phenols, oxidation of alcohols Dispose of all chemicals as directed by your lab instructor.

CHEMICAL CONCEPTS A. Structures of Alcohols and Phenol Alcohols are organic compounds that contain the hydroxyl group (~ OH). The simplest alcohol is methanol. Ethanol is found in alcoholic beverages and preservatives and is used as a solvent. 2-Propanol, also known as rubbing alcohol or isopropyl alcohol, is found in astringents and perfumes.

Hand sanitizers are used to kill bacteria and viruses that spread colds and flu. As a gel or liquid solution, many hand sanitizers use ethanol as their active ingredient. The amount of ethanol in an alcoholcontaining sanitizer is typically 60% (v/v), but can be as high as 85% (v/v). The high volume of ethanol can make hand sanitizers a fire hazard in the home because ethanol is highly flammable. When using an ethanol-containing sanitizer, it is important to rub your hands until they are completely dry. It is also recommended that sanitizers containing ethanol be stored in areas away from heat sources in the home.

▲ Hand sanitizers that contain ethanol are used to kill bacteria on the hands. Copyright © 2014 Pearson Education, Inc.

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A benzene ring with a hydroxyl group is known as phenol. Concentrated solutions of phenol are caustic and cause burns. However, derivatives of phenol, such as thymol, are used as antiseptics and are sometimes found in cough drops.

Classification of Alcohols In a primary (1°) alcohol, the carbon atom attached to the ~ OH group is bonded to one other carbon atom. In a secondary (2°) alcohol, the carbon with the ~ OH is attached to two carbon atoms, and in a tertiary (3°) alcohol it is attached to three carbon atoms.

B. Properties of Alcohols and Phenol The polarity of the hydroxyl group (~ OH) makes alcohols with one to three carbon atoms completely soluble with water because they can form many hydrogen bonds. An alcohol with four carbon atoms is somewhat soluble, whereas the large hydrocarbon portion in longer chain alcohols makes them insoluble in water.

▲ Hydrogen bonds form between the hydroxyl group and H and O atoms in water.

Alcohols and Phenols

261

Acidity of Phenol In water, phenol acts as a weak acid because the hydroxyl group ionizes slightly. Although phenol has six carbon atoms, the polarity of the hydroxyl group makes it soluble in water.

C. Oxidation of Alcohols Primary and secondary alcohols are easily oxidized. An oxidation consists of removing an H from the ~ OH group and another H from the C atom attached to the ~ OH group. Tertiary alcohols do not undergo oxidation because there are no H atoms on that C atom attached to the ~ OH. Primary and secondary alcohols can be distinguished from tertiary alcohols using a solution with chromate, CrO 4 2− . When an oxidation has occurred, the orange color of the chromate solution turns green.

D. Iron(III) Chloride Test Phenols react with the Fe3+ ion in iron(III) chloride (FeCl3 ) solution to give complex ions with strong colors from red to purple. Phenol Colorless

+

Fe3+ Yellow

⎯⎯ →

Fe3+ (phenol complex) Purple

E. Identification of an Unknown The group of tests for alcohols and phenols described in this experiment will be used to identify the functional group and family of an unknown.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Structures of Alcohols and Phenol Materials: Display of models: ethanol, 2-propanol, and 2-methyl-2-propanol

Draw the condensed structural formula of each model in the display and for phenol. Classify each as a primary, secondary, or tertiary alcohol.

B. Properties of Alcohols and Phenol Materials: 6 test tubes, pH paper, stirring rod, ethanol, 2-propanol, 2-methyl-2-propanol (t-butyl alcohol), cyclohexanol, 20% phenol, and unknown, Caution: Avoid contact with phenol. pH 1.

Place 10 drops of ethanol, 2-propanol, t-butyl alcohol (2-methyl-2-propanol), cyclohexanol, 20% phenol, and the unknown into six separate test tubes. Obtain some pH paper. Use a stirring rod to place a drop of each on a piece of pH paper. Clean the stirring rod between applications. Compare the color of the pH paper with the color chart on the container to determine the pH. Record your observations. Save the test tubes and alcohols for part 2.

Solubility in water 2.

Add 2 mL of water (40 drops) to each test tube. Shake and determine the solubility in water of each alcohol. If the compound is soluble in water, you will see a clear solution with no separate layers. If it is insoluble, a cloudy mixture or separate layer will form. Record your observations. DISPOSE OF ORGANIC COMPOUNDS IN DESIGNATED WASTE CONTAINERS!

C. Oxidation of Alcohols Materials: 6 test tubes, ethanol, 2-propanol, 2-methyl-2-propanol (t-butyl alcohol), cyclohexanol, 20% phenol, unknown, 2% chromate solution 1.

2. 3. 4.

Place 8 drops of ethanol, 2-propanol, 2-methyl-2-propanol (t-butyl alcohol), cyclohexanol, 20% phenol, and the unknown into six separate test tubes. Carefully add 2 drops of chromate solution to each. Stir carefully to allow the alcohol to react. Caution: Chromate solution contains concentrated H2 SO4 , which is corrosive. Look for a color change in the chromate solution. If a test tube becomes hot, place it in a beaker of ice-cold water. Record your observation of the color after 2 minutes. Draw the condensed structural formula of each alcohol. Classify each alcohol as primary (1°), secondary (2°), or tertiary (3°). If the orange color turned green within 2 min, oxidation of the alcohol has taken place. If the color remained orange, no reaction has occurred. If a reaction occurred, draw the condensed structural formulas of the product. When there is no change in color, no oxidation took place. In this case, write “no reaction” (NR). DISPOSE OF ORGANIC COMPOUNDS IN DESIGNATED WASTE CONTAINERS!

D. Iron(III) Chloride Test Materials: 6 test tubes, ethanol, 2-propanol, 2-methyl-2-propanol (t-butyl alcohol), cyclohexanol, 20% phenol, unknown, 1% FeCl3 solution

Place 5 drops of the alcohols and unknown in separate test tubes. Add 5 drops of 1% FeCl3 solution to each. Stir and record observations. DISPOSE OF ORGANIC COMPOUNDS IN DESIGNATED WASTE CONTAINERS!

E. Identification of an Unknown Use the test results to identify your unknown as one of the five compounds used in this experiment. Draw the condensed structural formula of the unknown.

Date

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Pre-Lab Study Questions

23

1. What is the functional group of an alcohol and a phenol?

2. Why are some alcohols soluble in water?

3. Classify each of the following alcohols as primary, secondary or tertiary. a. 3-pentanol

_______________________

b. 2-methyl-2-butanol _______________________ c. 1-propanol

_______________________

4. If you add chromate, an oxidizing agent, to each of the following, would a green Cr 3+ solution be formed? a. 3-pentanol

_______________________

b. 2-methyl-2-butanol _______________________ c. 1-propanol

_______________________

5. If an alcohol solution has a pH of 5, would it be a primary alcohol, a secondary alcohol, a tertiary alcohol, or a phenol?

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REPORT SHEET

Alcohols and Phenols

LAB

23

A. Structures of Alcohols and Phenols Ethanol

2-Propanol

Classification: 2-methyl-2-propanol (t-butyl alcohol)

Phenol

Classification: Questions and Problems Q1 Draw the condensed structural formula and give the classification for each of the following alcohols: 1-Pentanol

3-Pentanol

Classification: Cyclopentanol

1-Methylcyclopentanol

Classification:

B. Properties of Alcohols and Phenols Alcohol

1. pH

Unknown # _______________ 2. Soluble in Water?

Ethanol 2-Propanol 2-Methyl-2-propanol Cyclohexanol Phenol Unknown Copyright © 2014 Pearson Education, Inc.

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C. Oxidation of Alcohols Alcohol

1. Color After 2 min

2. Condensed Structural Formula

3. Classifica- 4. Condensed tion Structural Formula of Oxidation Product

Ethanol 2-Propanol 2-Methyl2-propanol Cyclohexanol Phenol Unknown Q2 Draw the condensed structural formula of the product of the following reactions (if no reaction, write NR): a.

b.

c.

D. Iron(III) Chloride Test Alcohol Ethanol 2-Propanol t-Butyl alcohol Cyclohexanol Phenol Unknown

FeCl 3 Test (Color)

Alcohols and Phenols

E. Identification of an Unknown Unknown # _______________ Summary of Testing

Test Results

Conclusions

B.1 pH B.2 Soluble in water? C.1 CrO 4 2− D

FeCl3

Name of Unknown

Condensed Structural Formula

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Aldehydes and Ketones 24 LABORATORY GOALS • Draw the condensed structural formulas of aldehydes and ketones. • Determine chemical and physical properties of aldehydes and ketones. • Perform chemical tests to distinguish between aldehydes and ketones.

LAB INFORMATION Time: Comments:

2h Flammable compounds are used in this experiment. Do not use burners. In tests with color changes, carefully observe the color of the reactants before they are mixed. Tear out the lab report sheets and place them beside the matching procedures. Related Topics: Aldehydes, ketones, oxidation of aldehydes Dispose of all chemicals as directed by your lab instructor.

CHEMICAL CONCEPTS A. Structures of Some Aldehydes and Ketones Aldehydes and ketones both contain the carbonyl group. In an aldehyde, the carbonyl group has a hydrogen atom attached; the aldehyde functional group occurs at the end of the carbon chain. In a ketone, the carbonyl group is located between two of the carbon atoms within the chain.

B. Properties of Aldehydes and Ketones Many aldehydes and ketones have sharp odors. If you have taken a biology class, you may have noticed the odor of Formalin™, which is a solution of formaldehyde. Aromatic aldehydes have a variety of odors. Benzaldehyde, the simplest aromatic aldehyde, has an aroma of almonds. Vanillin, an aromatic aldehyde, is found in the seed pods of the vanilla plant. When you use fingernail polish remover, you may notice the strong odor of acetone, the simplest ketone, which is used as the solvent.

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▲The vanilla bean is the dried fruit of the vanilla plants. Vanilla extract is prepared by soaking vanilla beans in ethanol and water.

▲Acetone (propanone) is used as a solvent for paint and nail polish.

C. Solubility, Iodoform Test, and Benedict’s Test Ketones containing a methyl group attached to the carbonyl give a reaction with iodine (I 2 ) in a solution of NaOH. The reaction produces solid, yellow iodoform, CHI3 . Iodoform, which has a strong medicinal odor, is used as an antiseptic.

Oxidation of Aldehydes and Ketones Aldehydes are oxidized using Benedict’s solution, which contains cupric ion, Cu 2+ . Because ketones cannot oxidize, this test can distinguish aldehydes from ketones. In the oxidation reaction, the bluegreen Cu 2+ is reduced to cuprous ion (Cu + ), which forms a reddish-orange precipitate of Cu 2 O.

Identification of an Unknown Using the results of the tests, an unknown substance can be identified as an aldehyde or a ketone.

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GOGGLES REQUIRED!

A. Structures of Some Aldehydes and Ketones Materials: Display of models: formaldehyde, acetaldehyde, propionaldehyde, acetone, butanone, and cyclohexanone 1. 2.

Observe models of formaldehyde, acetaldehyde, propionaldehyde, acetone, butanone, and cyclohexanone. Draw each of their condensed structural formulas. Write the IUPAC or common names (if any) for each.

B. Properties of Aldehydes and Ketones Materials: Test tubes, droppers, 5- or 10-mL graduated cylinder, acetone, benzaldehyde, camphor, vanillin, cinnamaldehyde, propionaldehyde, and cyclohexanone, chemistry handbook or Internet access 1.

2. 3.

Carefully detect the odor of samples of acetone, benzaldehyde, camphor, vanillin, and cinnamaldehyde. Your instructor will demonstrate how to fan the fumes above the test tube sample. Record your observations. Draw each of the condensed structural formulas. You may need a chemistry handbook or the Internet. Identify each as a ketone or aldehyde.

C. Solubility, Iodoform Test, and Benedict’s Test Solubility in Water Materials: Test tubes (5), dropper, 5 or 10 mL graduated cylinder 10% NaOH, propanal, benzaldehyde, acetone, cyclohexanone, and an unknown 1.

Place 2 mL of water in each of five test tubes. Add 5 drops propanal, benzaldehyde, acetone, cyclohexanone, and an unknown, to separate test tubes. Mix. If two separate layers form or the mixture turns cloudy, the aldehyde or ketone is not soluble. Save test tubes for Part C.2.

Iodoform Test for Methyl Ketones Materials: Test tubes from Part C.1, dropper, 10% NaOH, warm water bath, and iodine test reagent 2. Using the test tube samples from Part C.1, add 10 drops of 10% NaOH to each. Warm the tubes in a warm water bath to 50–60°C. Add 20 drops of iodine test reagent. Mix. Look for the formation of a yellow solid. Record your observations. 3. State if a methyl ketone is present or not. Benedict’s Test Materials: Test tubes, propionaldehyde (propanal), benzaldehyde, acetone, cyclohexanone, unknoenedict’s reagent, droppers, boiling water bath 4. Place 10 drops of propionaldehyde (propanal), benzaldehyde, acetone, cyclohexanone, and unknown in separate test tubes. Label. Add 2 mL of Benedict’s reagent to each test tube. Mix. Place the test tubes in a boiling water bath for 5 min. After 5 min, record the color of the samples. The appearance of the red-orange color of Cu 2 O indicates that oxidation has occurred. Moderate amounts of Cu 2 O will blend with the blue Cu 2+ solution to form green or rust color. 5. 6.

Identify the compounds that were oxidized. If you were given an unknown compound, you can now compare the results of the tests for the unknown with the tests you performed with known aldehydes and ketones. Identify your unknown.

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Pre-Lab Study Questions

24

1. Draw the condensed structural formula, if any, of the product from each of the following:

a.

b.

c.

d.

2. Draw the condensed structural formula of each of the following: a. 2-methylbutanal b. 3-pentanone

c. 3-bromopropanal

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REPORT SHEET

Aldehydes and Ketones

LAB

24

A. Structures of Some Aldehydes and Ketones Formaldehyde

Acetaldehyde

IUPAC Name __________________________ Propionaldehyde

IUPAC Name _______________________ Acetone

IUPAC Name ___________________________ Butanone

IUPAC Name _______________________ Cyclohexanone

Common Name __________________________

B. Properties of Aldehydes and Ketones 1. Odor

2. Condensed Structural Formula

3. Aldehyde or Ketone?

Acetone

Benzaldehyde

Camphor

Vanillin

Cinnamaldehyde

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C. Solubility, Iodoform Test, and Benedict’s Test 1. Soluble? 2. Iodoform 3. Methyl 4. Color Test Ketone?

5. Oxidation? (yes/no)

Propanol

Benzaldehyde

Acetone

Cyclohexanone

Unknown

6. Identify the unknown as propanal, benzaldehyde, acetone, or cyclohexanone. Give your reasoning.

Questions and Problems Q1 Indicate the test results for each of the following compounds in the iodoform test and in the Benedict’s test:

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Q2 Two compounds, A and B, have the formula of C3H 6O. Determine their condensed structural formulas and names using the following test results. a. Compound A forms a red-orange precipitate with Benedict’s reagent but does not react with iodoform.

b. Compound B forms a yellow solid in the iodoform test but does not react with Benedict’s reagent.

Q3 What chemical tests could you use to distinguish between 2-pentanone and 3-pentanone?

Q4 What aldehyde or ketone might be present in the following everyday products? Almond-flavored cookies

_______________________

Candies with cinnamon flavor

_______________________

Nail polish remover

_______________________

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Carboxylic Acids and Esters 25 LABORATORY GOALS • • • •

Write the condensed structural formulas of carboxylic acids and esters. Determine the solubility and acidity of carboxylic acids and their salts. Write equations for neutralization and esterification of acids. Prepare esters and identify their characteristic odors.

LAB INFORMATION Time: Comments:

Related Topics:

2h When noting odors, hold your breath, fan across the top of a test tube, then inhale slightly to detect the odor. The formation of esters requires concentrated acid. Use carefully. Tear out the report sheets and place them beside the matching procedure. Carboxylic acid functional group, ester functional group, ionization of carboxylic acids, neutralization, esterification, hydrolysis, saponification Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS A. Carboxylic Acids and Their Salts A salad dressing made of oil and vinegar tastes sour due to the vinegar, which contains acetic acid (ethanoic acid). The sour taste of fruits such as lemons is due to acids such as citric acid. Some face creams contain alpha hydroxy acids such as glycolic acid. All these acids just described are carboxylic acids, which contain the carboxyl group: a carbonyl group attached to a hydroxyl group. A dicarboxylic acid, such as malonic acid, has two carboxylic acid functional groups. The carboxylic acid of benzene is called benzoic acid.

Ionization of Carboxylic Acids in Water Carboxylic acids are weak acids because the carboxylic acid group ionizes slightly in water to give a hydrogen ion (H + ) and a carboxylate ion as shown in the equation below. The polarity of the carboxylic acid group makes acids with one to four carbon atoms soluble in water.

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Neutralization of Carboxylic Acids An important feature of carboxylic acids is their neutralization by bases such as sodium hydroxide to form carboxylate salts and water. Neutralization is the reaction of an acid with a base to give a salt and water. Even insoluble carboxylic acids with five or more carbon atoms can be neutralized to give salts that are usually soluble in water. For this reason, acids used in food products or medications are in their soluble carboxylate salt form rather than the acid.

B. Esters Carboxylic acids may have tart or unpleasant odors, but many esters have pleasant flavors and fragrant odors. Octyl acetate gives oranges their characteristic aroma and flavor; pear flavor is due to pentyl acetate. The flavor and smell of raspberries come from isobutyl formate.

▲The flavor of raspberries is due to isobutyl formate. An ester of salicylic acid is methyl salicylate, which gives the flavor and odor of oil of wintergreen used in candies and ointments for sore muscles. When salicylic acid reacts with acetic anhydride, acetylsalicylic acid is formed. Known as aspirin, it is used to reduce fever and inflammation.

▲Ointments containing methyl salicylate are used to soothe sore muscles.

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Esterification and Acid Hydrolysis In a reaction called esterification, the carboxylic acid group combines with the hydroxyl group of an alcohol. The reaction, which takes place in the presence of an inorganic acid catalyst, produces an ester and water.

The reverse reaction, hydrolysis, occurs in acid or base. When an acid catalyst is used, the ester decomposes to yield the carboxylic acid and alcohol. The ester product is favored when an excess of acid or alcohol is used; hydrolysis is favored when more water is present.

C. Saponification When an ester is hydrolyzed in the presence of a base, the reaction is called saponification. The products are the salt of the carboxylic acid and the alcohol. Although the ester is usually insoluble in water, the salt and alcohol (if short-chain) are soluble.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Carboxylic Acids and Their Salts Materials: Test tubes, glacial acetic acid, benzoic acid, dropper, spatula, pH paper, red and blue litmus paper, stirring rod, 400-mL beaker, hot plate or Bunsen burner, 10% NaOH, 10% HCl 1. 2.

3.

4.

5. 6.

Write the condensed structural formulas for acetic acid and benzoic acid. Place about 2 mL of cold water in two test tubes. Add 5 drops of acetic acid to one test tube and a small amount of benzoic solid (enough to cover the tip of a spatula) to the other. Tap the sides of the test tubes to mix or stir with a stirring rod. Record your observations. Save the test tubes for Part 3. Test the pH of each carboxylic acid by dipping a stirring rod into the solution, then touching it to a piece of pH paper. Compare the color on the paper with the color chart on the container and record the pH. Save the test tubes for Part 4. Place the test tube of benzoic acid (solid should be present) in a hot water bath and heat for 5 min. Describe the effect of heating on the solubility of acid. Allow the test tube to cool. Record your observations. Save the test tubes for Part 5. Add drops of 10% NaOH to each test tube until the solution turns red litmus paper blue. Record your observations. Save the test tubes for Part 6. Add drops of 10% HCl to each test tube until the solution turns blue litmus paper red. Record your observations.

B. Esters Materials: Models of acetic acid, methanol, and methyl acetate. Dropper bottles containing methanol, 1-pentanol, 1-octanol, benzyl alcohol, 1-propanol, glacial acetic acid, 85% H3 PO 4 , salicylic acid (s), test tubes, test tube rack hot plate or Bunsen burner, 400-mL beaker, stirring rod, spatula, small beaker Caution: Use care in dispensing glacial acetic acid. It can cause burns and blisters on the skin. 1. 2.

Observe models of acetic acid, methanol, and methyl acetate. Write the equation for the formation of methyl acetate. As assigned, prepare one of the mixtures listed by placing 3 mL of the alcohol in a test tube, and label it with the mixture letter. Mixture A B C D E

Alcohol Methanol 1-Pentanol 1-Octanol Benzyl alcohol 1-Propanol

Carboxylic Acid Salicylic acid Glacial acetic acid Glacial acetic acid Glacial acetic acid Glacial acetic acid

3.

Add 2 mL (liquid) or the amount of solid that covers the tip of a spatula of the assigned carboxylic acid to the alchohol. With the test tube pointed away from you and the other students, cautiously add 15 drops of the concentrated phosphoric acid, H3 PO 4 . Stir gently. Place the test tube in a hot water bath for 15 min. Remove the test tube and cautiously fan the vapors toward you as previously demonstrated by your instructor. Record the odors that you detect. To obtain a stronger odor, place 15 mL of hot water in a small beaker and pour the ester mixture into the beaker. Record the odors of esters produced by other students for the other combinations.

4.

Write the condensed structural formulas and names of the esters produced. Dispose of the ester products as instructed.

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C. Saponification Materials: Test tube, test tube holder, droppers, stirring rod, hot plate or Bunsen burner, 250- or 400-mL beaker, methyl salicylate in a dropper bottle, 10% NaOH, 10% HCl, blue litmus paper 1. 2. 3.

4. 5. 6.

Draw the condensed structural formula of methyl salicylate. Place 3 mL of water in a test tube. Add 5 drops of methyl salicylate. Record the appearance and odor of the ester. Add 1 mL (20 drops) of 10% NaOH to the solution. There should be two layers in the test tube. Place the test tube in a hot water bath for 30 min or until the top layer of the ester disappears. Record any changes in the odor of the ester. Remove the test tube and cool in cold water. Write the equation for the saponification reaction. After the solution is cool, add about 20 drops (1 mL) of 10% HCl or until a drop of the solution turns blue litmus paper red. Record your observations. Draw the condensed structural formula of the solid that forms.

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Pre-Lab Study Questions

25

1. Give the IUPAC name for each of the following:

a.

__________________________________

b.

__________________________________

c.

__________________________________

2. Write the condensed structural formula of the organic product for propanoic acid reacting with each of the following: a. NaOH

b. CH3 ~ OH

c. H 2 O

3. Write the condensed structural formula of the organic products for ethyl ethanoate when it reacts with each of the following: a. NaOH

b. H 2 O and HCl

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REPORT SHEET

Carboxylic Acids and Esters

LAB

25

A. Carboxylic Acids and Their Salts Acetic Acid

Benzoic Acid

1. Condensed structural formulas

2. Solubility in cold water 3. pH 4. Solubility in hot water 5. NaOH 6. HCl

Q1 a. Write the balanced equation using condensed structural formulas for the reaction of acetic acid with NaOH.

b. Write the balanced equation using condensed structural formulas for the reaction of benzoic acid with NaOH.

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B. Esters 1. Use condensed structural formulas to write the equation for the formation of methyl acetate.

2. Alcohol and Carboxylic Acid A. Methanol and salicylic acid

B. 1-Pentanol and acetic acid

C. 1-Octanol and acetic acid

D. Benzyl alcohol and acetic acid

E. 1-Propanol and acetic acid

3. Odor of Ester

4. Condensed Structural Formula and Name of Ester

Carboxylic Acids and Esters

C. Saponification 1. Condensed structural formula

2. Appearance and odor

3. Changes in appearance and odor

4. Equation for saponification

5. Observations

6. Condensed structural formula of the compound formed by adding HCl

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Aspirin and Other Analgesics 26 LEARNING GOALS • • • •

Use an esterification reaction to synthesize aspirin. Test the purity of prepared aspirin and commercial aspirin products. Determine the physical and chemical properties of aspirin. Use thin-layer chromatography to separate and identify components in analgesics.

LAB INFORMATION Time: 2½ h Comments: Tear out the report sheets and place them beside the matching procedures. Related Topics: Esters, carboxylic acids Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS Aspirin and Willow Bark For many centuries, relief from pain and fever was obtained by chewing on the leaves or bark from the willow tree. By the 1800s, chemists had discovered that salicin was the agent responsible for the relief of pain. However, the body converts salicin to salicylic acid, which irritates the stomach lining. In 1899, the Bayer chemical company in Germany produced an ester of salicylic acid and acetic acid, called acetylsalicylic acid (known as aspirin), which is less irritating. In some aspirin preparations, a buffer is added to neutralize the carboxylic acid group, which reduces irritation and ulcers. Today, aspirin is used as an analgesic (pain reliever), antipyretic (fever reducer), and anti-inflammatory agent.

▲The discovery of salicin in the leaves and bark of the willow tree led to the development of aspirin.

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In commercial aspirin products, a small amount of acetylsalicylic acid (300 mg to 400 mg) is bound together with starch, and sometimes caffeine and buffers, to make an aspirin tablet. The basic conditions in the small intestine break down the acetylsalicylic acid to yield salicylic acid, which is absorbed into the bloodstream. The addition of a buffer reduces the irritation caused by the carboxylic acid group of the aspirin molecule.

A. Preparation of Aspirin Aspirin (acetylsalicylic acid) can be prepared from acetic acid and the hydroxyl group on salicylic acid. However, this is a slow reaction. The ester more when acetic anhydride is used to provide the acetyl group in the presence of acid acting as a catalyst. Caution: the aspirin you will prepare in this experiment is impure and must not be taken internally!

Using the following equation, the maximum theoretical yield of aspirin that is possible from 2.00 g of salicylic acid can be calculated. 2.00 g salicylic acid ×

1 mole salicylic acid 138.1 g salicylic acid

×

1 mole aspirin 1 mole salicylic acid

×

180.2 aspirin 1 mole aspirin

= 2.61 g aspirin (maximum theoretical yield)

Suppose the total amount of aspirin you obtain has a mass of 2.25 g. A percent yield can be calculated as follows: Percent yield (%) =

2.25 g aspirin obtained × 100% = 86.2% yield of aspirin product 2.61 g aspirin calculated

B. Testing Aspirin Products The purity of the crude sample and the recrystallized aspirin product can be tested with iron(III) chloride, FeCl3 . The Fe3+ ion reacts with the phenol group on salicylic acid and gives a purple color. This test can also be used to determine the purity of commercially prepared aspirin. Sometimes old aspirin breaks down to give salicylic acid and acetic acid. Then the aspirin in the bottle smells like vinegar and should be discarded.

C. Analysis of Analgesics Aspirin is one of several analgesics that are used to relieve pain. Other analgesics include acetaminophen, ibuprofen, and naproxen. These products, including aspirin, are used to reduce fever, which means they are also antipyretics. However, aspirin also has anti-inflammatory properties and may reduce the risk of a heart attack.

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Thin-Layer Chromatography (TLC) Thin-layer chromatography (TLC) is a technique used to separate substances and identify components in a mixture. A TLC plate is typically a sheet of plastic, coated with a thin layer of a solid adsorbent such as silica gel. Small amounts of known and unknown substances are placed as small spots near one end of the TLC plate. Then the end of the plate with the spots is placed in a solvent contained in a developing chamber. The solid silica layer on the TLC plate is called the stationary phase. The solvent moves up the silica layer on the TLC plate, carrying the substances in the spot with it. The more soluble a substance is in the solvent, the higher the solvent will carry it up the plate. A substance that adheres strongly to the stationary silica gel moves only a short distance with the solvent. Thus, differences in the chemical and physical properties of the substances determine the distances they travel up the plate. As the solvent front nears the top of the TLC plate, the plate is removed, marked, and dried. Then the substances are visualized. If they have colors, they can be seen directly. In this experiment, they are colorless. Because the silica material on the plate contains a fluorescent compound, ultraviolet light (254nm) from a UV lamp can be used to visualize the substances, which appear as dark spots on the plate. Calculating R f Values A value called the R f value can be calculated for each substance on a plate. The R f is the distance that a substance moves on the plate divided by the distance the solvent moves. An unknown substance is identified if its R f value matches the R f value of one of the known substances used on the plate (see Figure 26.1.)

▲ FIGURE 26.1 Distances moved by a substance and solvent on a TLC plate. In this experiment, you will use TLC to determine the R f values for several known analgesics. You will also determine the R f values and identify the types of analgesics in a variety of over-the-counter drugs used to relieve pain.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Preparation of Aspirin Materials: 125-mL Erlenmeyer flask, 400-mL beaker, hot plate or Bunsen burner, ice, salicylic acid, acetic anhydride, 5- or 10-mL graduated cylinder, stirring rod, pan or large beaker, dropper, 85% H3 PO 4 in a dropper bottle, Büchner filtration apparatus, filter paper, spatula, watch glass 1. 2.

Weigh a 125-mL Erlenmeyer flask. Record its mass. Add about 2 g of salicylic acid. Reweigh the flask with salicylic acid and record new mass. Working in the hood, carefully add 5 mL of acetic anhydride to the flask. Caution: Acetic anhydride is irritating to the nose and sinuses. Handle carefully. Slowly add 10 drops of 85% phosphoric acid, H3 PO 4 . Stir the mixture with a stirring rod. Place the flask and its contents in a boiling water bath and stir until all the solid dissolves. Remove the flask from the hot water and let it cool to room temperature. Working in the hood, cautiously add 20 drops of water to the cooled mixture. Keep your face away from the top of the flask: acetic acid vapors are irritating When the reaction is complete (when no further change is observed), add 50 mL of cold water. Cool the mixture by placing the flask in an ice bath for 10 min. Stir gently. Crystals of aspirin should form. If no crystals appear, gently scratch the inside of the flask with a stirring rod.

Collecting the Aspirin Crystals Add a piece of filter paper to a Büchner filtration apparatus. Place the funnel in the filter flask, making sure that the neck fits snugly in a rubber washer. Moisten the filter paper. Turn on the vacuum or water aspirator and pour the aspirin product slowly onto the center of the filter paper (see Figure 26.2). Push down gently on the funnel to create the suction needed to pull the water off the aspirin product. The aspirin crystals will collect on the filter paper. Use a spatula to transfer any crystals left in the flask. Rinse the inside of the flask with a 10-mL portion of ice-cold water and transfer all the crystals to the funnel. Wash the aspirin crystals on the filter paper with two 10-mL portions of cold water. Spread the aspirin crystals out on the filter paper and draw air through the funnel. This helps dry the crystals. Remove the Büchner funnel and turn off the water. Use a spatula to lift and transfer the filter paper and aspirin to a paper towel. Don’t touch it; it may still contain acid. Allow the crystals to air dry.

▲ FIGURE 26.2 Apparatus for suction filtration with a Büchner funnel.

Aspirin and Other Analgesics

3. 4.

295

Weigh a clean, dry watch glass. Record its mass. Transfer the aspirin (acetylsalicylic acid) crystals to the watch glass. Weigh the watch glass and the aspirin. Record the mass.

Calculations: 5. 6.

Calculate the mass of salicylic acid. (2 − 1) Calculate the maximum theoretical yield of aspirin possible from the salicylic acid (see Part A).

7. 8.

Calculate the mass of the aspirin you collected. (4 − 3) Calculate the percent yield of aspirin. (7 ÷ 6) × 100%

9.

Use a melting point apparatus to determine the melting point of your aspirin product. Pure aspirin has a melting point of 135 °C. Pure salicylic acid melts at 157 – 159 °C. Compare the melting point of your aspirin with the known melting points of aspirin and salicylic acid. Your instructor will demonstrate the use of the melting point apparatus.

B. Testing Aspirin Products Materials: Test tubes, spatula, aspirin from Part A, commercial aspirin tablets, buffered aspirin, acetylsalicylic acid that you made, 0.15% (m/v) salicylic acid, pH indicator paper, stirring rod, 1% FeCl3 Preparation of aspirin product samples Place 3 mL of 0.15% salicylic acid in test tube 1. In test tubes 2 to 4, place a few crystals (the amount on the tip of a spatula) of the following substances and add 3 mL of water to each: test tube 1 0.15% salicylic acid (reference sample) test tube 2 commercial aspirin (crushed) test tube 3 buffered aspirin (crushed) test tube 4 aspirin product from Part A pH Test Stir each mixture and touch the stirring rod to a piece of pH indicator paper. Compare the color of the paper to the chart on the container. 1.

Record the pH of each solution in test tubes 1 to 4. Save these test tubes and samples for Part B.

Testing Aspirin Purity To each of the samples from Part B.1, add 5 drops of 1% FeCl3 solution. Any free salicylic acid (unreacted during synthesis or resulting from hydrolysis in the breakdown of aspirin) reacts with the FeCl3 to give a purple color.

2.

The maximum salicylic acid allowed in commercially prepared aspirin products is 0.15%. A higher amount of salicylic acid in the sample will form a deeper color, which indicates that the product is impure and not safe for ingestion. If the sample test has a lighter color than a 0.15% standard, the sample would be considered pure by USP standards. However, no matter what the results of the test, your laboratory-prepared aspirin must not be ingested. Record the colors of each aspirin product with FeCl3 .

3.

Compare the purity of the tested products to the reference sample of 0.15% salicylic acid.

C. Analysis of Analgesics Materials: 400-mL beaker (developing chamber), plastic wrap, rubber band to fit beaker, solvent (75% ethyl acetate and 25% hexane), TLC plate coated with silica gel, UV lamp (short wave 254 nm), micropipettes, spot plates, dropper bottle containing 1% solutions in ethanol of aspirin, ibuprofen, acetaminophen, naproxen, caffeine, over-the-counter drugs (unknown), ruler

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▲ FIGURE 26.3 Spots of analgesics on a TLC plate. Preparing the TLC Developing Chamber Obtain a 400-mL beaker, a piece of plastic wrap that can cover the mouth of the beaker, and a rubber band. Carefully pour a small amount of solvent into the beaker to a level of 0.5 to 0.6 cm. It is important that the solvent level is below the spots you place on the TLC plate. Cover the beaker with plastic wrap and secure with a rubber band. Spotting the TLC Plate Obtain a TLC plate that is 6 cm × 10 cm. To avoid transferring substances from your fingers to the plate, be sure to handle it at the edges only. Draw a light line with a pencil about 1 cm above the end of the TLC plate. This is your starting line, or origin. On the line, mark 6 equally spaced dots and label them 1 through 6. Place a few drops of each of the 1% solutions in a spot plate. Number the wells as follows. 1. aspirin 2. ibuprofen 3. acetaminophen 4. naproxen 5. caffeine 6. over-the-counter drug (unknown) Using clean capillary pipettes (micropipettes), one for each substance, spot a tiny amount of each substance on a separate dot. Lightly touch the micropipette to deliver a small amount. When dry you can apply the same substance again. The spots must be kept small (see Figure 26.3.) Placing TLC Plate in Developing Chamber Carefully set the plate in the beaker containing the solvent that you prepared as the developing chamber. The level of the solvent must be lower than the line drawn on the plate (see Figure 26.4). Cover the beaker with plastic wrap and leave undisturbed as the solvent moves up the TLC plate. When the solvent has risen almost to the top of the plate, open the chamber and remove the plate. Before the solvent dries, draw a pencil line along the solvent front. Place the plate in the hood and allow the solvent to evaporate. Place used solvent in the organic solvent container as directed by your instructor. Measurements and Calculations 1. 2. 3. 4. 5.

Observe the TLC plate under UV light. Circle each spot. Draw and label a diagram with the samples and the position of their spots as they appear on your TLC plate. Measure the distance, in centimeters, from the origin (bottom line) to the solvent front. Measure the distance, in centimeters, from the origin (bottom line) to the center of each spot. Calculate the R f value for each spot of analgesic. (3 ÷ 2) Use the R f values of the known analgesics to identify the analgesic(s) in the over-the-counter drug (unknown).

▲ FIGURE 26.4 A developing chamber containing TLC plate.

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Pre-Lab Study Questions

26

1. Circle and label the functional group of aspirin that causes it to irritate the stomach.

Aspirin

2. Why are buffers added to some aspirin products?

3. What quantity of aspirin is contained in most over-the-counter aspirin products?

4. A student measures the solvent height on a TLC plate as 11.5 cm. Spot A was measured at a height of 4.7 cm, spot B at 2.3 cm, and spot C at 8.2. What are the R f values for spots A, B, and C?

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REPORT SHEET

Aspirin and Other Analgesics

LAB

26

A. Preparation of Aspirin 1. Mass of flask 2. Mass of flask and salicylic acid 3. Mass of watch glass 4. Mass of watch glass and aspirin product Calculations 5. Mass of salicylic acid 6. Maximum theoretical yield of aspirin (Show calculations.)

7. Mass of aspirin 8. Percent yield of aspirin (Show calculations.)

9. Melting point (°C) of aspirin product

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Questions and Problems Q1 Circle and label the ester group and the carboxylic acid group on the aspirin structure drawn below.

Q2 If a typical aspirin tablet contains 325 mg aspirin (the rest is starch binder), how many tablets could you prepare from the aspirin you made in lab?

B. Testing Aspirin Products Samples Tested

1. pH

2. Color with FeCl 3

3. Purity compared to reference sample

1. 0.15% salicylic acid (reference sample) 2. commercial aspirin brand:______________ 3. buffered aspirin brand:______________ 4. aspirin from Part A Questions and Problems Q3 Give an explanation for any differences in the pH values in the samples from Part B.

Q4 How does the pH of buffered aspirin product compare to the pH of the nonbuffered aspirin product?

Aspirin and Other Analgesics

C. Analysis of Analgesics 1. Diagram of Spots on TLC Plate

1

2

3

4

5

6

2. Distance moved by solvent ___________________________ Spot #

Analgesic substance

3. Distance spot moved

4. R f value

1 2 3 4 5

6

5. What substance(s) are present in the over-the-counter drug (unknown)?

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Amines and Amides 27 LABORATORY GOALS • • • • •

Draw the condensed structural formulas and give the names of amines. Classify amines as primary, secondary, or tertiary. Observe some physical properties of amines and amides. Write an equation for the formation of an amine salt. Write an equation for the formation of an amide and its hydrolysis in acid and base.

LAB INFORMATION Time: Comments:

2½ –3 h Some amines have an odor that irritates the nose and sinuses. Work in the hood. Tear out the report sheets and place them next to the matching procedures. Related Topics: Amines, solubility and pH of amines, amidation, amides, hydrolysis of amides Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS A. Structure, Classification, and Solubility of Amines Amines are considered derivatives of ammonia in which one or more hydrogen atoms are replaced with alkyl or aromatic groups. The number of alkyl groups attached to the nitrogen atom determines the classification of primary, secondary, or tertiary amines. The aromatic amines used the name aniline.

Amines are often found as part of compounds that are physiologically active or used in medications.

Physiological Effects of Some Amines Histamine Histamine is synthesized in the nerve cells of the hypothalamus from the amino acid histidine when a carboxylate group is converted to CO 2 . Histamine is produced by the immune system in response to pathogens and invaders or injury. When histamine combines with histamine receptors, it causes allergic

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reactions, which may include inflammation, watery eyes, itchy skin, and hay fever. Histamine can also cause smooth muscle constriction, such as the closing of the trachea in persons allergic to shellfish. Histamine is also stored and released in the cells of the stomach, where it stimulates acid production. Antihistamines, such as Benadryl, Zantac, and Tagamet, block the histamine receptors and stop the allergic reactions.

Serotonin Serotonin (5-hydroxytryptamine) helps us to relax, sleep deeply and peacefully, and think rationally, and it gives a feeling of well-being and calmness. Serotonin is synthesized from the amino acid tryptophan, which can cross the blood-brain barrier. A diet that contains foods such as eggs, fish, cheese, turkey, chicken, and beef, which have high levels of tryptophan, will increase serotonin levels. Foods with a low level of tryptophan, such as whole wheat, will lower serotonin levels. Psychedelic drugs such as LSD and mescaline stimulate the action of serotonin at its receptors.

Low levels of serotonin in the brain may be associated with depression, anxiety disorders, obsessivecompulsive disorder, and eating disorders. Many antidepressant drugs, such as fluoxetine (Prozac) and paroxetine (Paxil), are selective serotonin reuptake inhibitors (SSRIs). When the reuptake of serotonin is slowed, it remains longer at the receptors, where it continues its action; the net effect is as if additional quantities of serotonin were taken. Solubility of Amines In water, ammonia and amines with one to four carbon atoms act as weak bases because the unshared pair of electrons on the nitrogen atom attracts protons. The products are an ammonium ion or alkyl ammonium ion and a hydroxide ion.

B. Neutralization of Amines with Acid Because amines are basic, they react with acids to form the amine salt. These amine salts are much more soluble in water than the corresponding amines.

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305

C. Amides When a carboxylic acid reacts with ammonia or an amine, the product is an amide. The functional group is called the amide group.

In a reaction called amidation, an amide forms when a carboxylic acid is heated with ammonia or an alkyl or aromatic amine.

D. Hydrolysis of an Amide When an amide is hydrolyzed, the amide bond is broken and the carboxylic acid and the amine are separated. Hydrolysis takes place in either an acid or a base. Acid hydrolysis produces the carboxylic acid and the ammonium salt. In a base, the hydrolysis reaction produces the salt of the carboxylic acid and ammonia. The odor of ammonia and the reaction of ammonia with litmus paper are used to detect the hydrolysis reaction.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Structure, Classification, and Solubility of Amines Materials: Aniline, N-methylaniline, triethylamine, test tubes, test tube rack, stirring rod, pH paper

1. 2. 3.

4. 5.

WORK IN THE HOOD. THE VAPORS OF AMINES ARE IRRITATING TO THE NOSE AND SINUSES. Draw the condensed structural formulas of aniline, N-methylaniline, and triethylamine. State the classification (1°, 2°, 3°) of each. To three separate test tubes, add 5 drops of aniline, N-methylaniline, and triethylamine. Cautiously note the odor of each. Remember to hold a fresh breath of air while you fan the vapor toward you. The instructor will demonstrate this technique. Record the odor. Add 2 mL of water to each test tube and stir. Clean and dry the stirring rod between test tubes. Describe the solubility of each amine in water. Determine the pH of each solution by dipping a stirring rod in each solution and then touching it to pH paper. Record the pH. Save these test tubes and samples for Part B.

B. Neutralization of Amines with Acid Materials: Test tubes from Part A, blue litmus paper, 10% HCl 1. 2. 3.

Using the amines from Part A, add drops of 10% HCl to each amine solution until blue litmus paper turns red. Record any changes in solubility of each amine. Record any changes in odor. Draw the condensed structural formulas for the reactants and products for the neutralization equations of aniline, N-methylaniline, and triethylamine with HCl.

C. Amides Materials: Acetamide, benzamide, test tubes, spatula 1. 2. 3.

Draw the condensed structural formulas for the reactants and products of the equation for the amidation of (a) acetic acid and ammonia, (b) aniline and ammonia. Place small amounts (tip of a spatula) of acetamide and benzamide in separate test tubes. Using caution, note the odor of each. Record the odors. Add 2 mL of water to each test tube. Record the solubility of each amide in water Save these test tubes for Part D.

D. Hydrolysis of an Amide Materials: Acetamide, benzamide, test tubes, spatula, 10% HCl, 10% NaOH, 250-mL beaker for water bath, hot plate, red litmus paper, small graduated cylinder 1. 2.

Draw the condensed structural formulas for the reactants and products of the equation for the hydrolysis of (a) acetamide and (b) benzamide using HCl. Using the test tubes from Part C.3, add 2 mL of 10% HCl to each. Place the test tubes in a boiling water bath and heat gently for 5 min. Using caution, note any odor coming from each mixture. Record your observations.

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3.

4. 5.

307

Place small amounts (tip of a spatula) of acetamide and benzamide in separate test tubes. Add 2 mL of 10% NaOH to each. Place the test tubes in a boiling water bath. Wet a piece of red litmus paper and place it over the mouth of each test tube. Heat the test tubes gently for 5 min. Record any change in the color of each of the red litmus papers. Using caution, note any odor coming from each mixture. Record your observations. Draw the condensed structural formulas for the reactants and products of the equation for the hydrolysis of (a) acetamide and (b) benzamide, using NaOH.

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27

1. What is the functional group in amines? In amides?

2. What products are formed when amides are hydrolyzed?

3. Draw the condensed structural formulas for each of the following amines and amides: a. 1-propanamine

b. 2-methyl-3-hexamine

c. N-methylpentamide

d. N,N-dimethylbenzamide

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REPORT SHEET

Amines and Amides

LAB

27

A. Structure, Classification, and Solubility of Amines Aniline

N-Methylaniline Triethylamine

1. Condensed structural formula 2. Classification (1°, 2°, 3°) 3. Odor 4. Solubility in water 5. pH Questions and Problems Q1 In the discussion, the structures are given for Neo-Synephrine and methamphetamine. Give the amine classification of each of the compounds.

Q2 What type of compound accounts for the “fishy” odor of fish?

Q3 Explain why amines are basic.

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B. Neutralization of Amines with Acid Aniline

N-Methylaniline

1. Solubility after adding HCl 2. Change in odor after adding HCl 3. Equations for the neutralization of amines with HCl Aniline + HCl

N-methylaniline + HCl

Triethylamine + HCl

Questions and Problems Q4 How does lemon juice remove the odor of fish?

Q5 Write an equation for the reaction of butylamine with HCl.

C. Amides 1. Amidation a. Acetic acid + ammonia

b. Aniline + ammonia

Triethylamine

Amines and Amides

Acetamide

Benzamide

2. Odor 3. Solubility

D. Hydrolysis of an Amide 1. Acid Hydrolysis

a. Acetamide

b. Benzamide

Acetamide

Benzamide

2. Odor

Acetamide 3. Change in color of red litmus paper 4. Odor after adding NaOH

Benzamide

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5. Base Hydrolysis

a. Acetamide

b. Benzamide

Questions and Problems Q6 You have unknowns that are a carboxylic acid, an ester, and an amine. Describe how you would distinguish among them.

Synthesis of Acetaminophen 28 LABORATORY GOALS • Write equations for amidation. • Prepare the common analgesic acetaminophen. • Use solubility to purify a crude sample of acetanilide.

LAB INFORMATION Time: Comments:

2½–3 h When recrystallizing product, use ice water to rinse crystals. Tear out the report sheets and place them beside the matching procedures. Related Topics: Amides, amidation Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS Synthesis of Acetaminophen Compounds used to relieve pain are called analgesics, and compounds used to reduce a fever are called antipyretics. Aspirin is both an analgesic and an antipyretic and so is acetaminophen,

People who are sensitive to aspirin may use products such as Tylenol (acetaminophen) or Motrin, Advil, and Nuprin, which contain ibuprofen. Ibuprofen is a carboxylic acid, not an amide.

Aspirin and acetaminophen are used in several common analgesic preparations, which may also contain caffeine and buffers (see Table 28.1). TABLE 28.1 Some Products with Aspirin and/or Acetaminophen Product

Aspirin

Aspirin

325 mg

Anacin

400 mg

32 mg

Bufferin

324 mg

32 mg

Excedrin

250 mg

250 mg

65 mg

Stomach upset; long bleeding time; possible toxic levels

325 mg Possible liver damage in high dosages for long-term users

Increase in pulse and heart rate

Tylenol Side effects

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Acetaminophen

Caffeine

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Marketed as Tylenol, acetaminophen is an amide that we will prepare in the laboratory from p-aminophenol and a 2-carbon group obtained from acetic anhydride.

Percent Yield The maximum theoretical yield of acetaminophen (C8 H9 NO 2 ) from p-aminophenol (C6 H 7 NO) is calculated using the molar masses of the reactant p-aminophenol (109.1 g/mole) and the product acetaminophen (151.2 g/mole). grams of p -aminophenol ×

1 mole p -aminophenol 109.1 g p-aminophenol

×

1 mole acetaminophen 1 mole p -aminophenol

×

151.2 g acetaminophen 1 mole acetaminophen

= grams of acetaminophen (theoretical yield)

The percent yield (%) is calculated from the actual yield of acetaminophen obtained in the reaction divided by the theoretical yield. Percent yield (%) =

actual yield (g) of acetaminophen × 100% theoretical yield (g) of acetaminophen

Synthesis of Acetaminophen

EXPERIMENTAL PROCEDURES

317

GOGGLES REQUIRED!

Synthesis of Acetaminophen Materials: 125-mL Erlenmeyer flask, p-aminophenol, 85% H3 PO 4 , acetic anhydride, 150-mL beaker, hot plate, stirring rod, ice bath, Büchner filtration apparatus, filter paper, melting point apparatus, watch glass 1. 2.

Weigh a 125-mL Erlenmeyer flask. Record the mass. Add about 1.5 g of p-aminophenol to the flask. Avoid contact with skin. You may wish to wear gloves. Weigh the Erlenmeyer flask and p-aminophenol. Record the mass.

Preparation of Acetaminophen To the Erlenmeyer flask containing p-aminophenol, add 25 mL of water and 20 drops of 85% H3 PO 4 . Place the flask and contents on a hot plate and heat to boiling until the p-aminophenol has dissolved. When all the p-aminophenol has dissolved, remove the flask from the hot plate. Working in the hood, carefully add 2 mL of acetic anhydride to the flask. Stir. Place the flask in an ice bath. Continue to stir gently to crystallize the acetaminophen. You may need to scratch the walls of the flask with a glass stirring rod to start the crystallization. If no crystals appear, add a small crystal of acetaminophen to start the formation of solid acetaminophen. Leave the flask in the ice-water bath for 30 min. Collecting the Crystals Using a Büchner Funnel Vacuum Add a piece of filter paper to a Büchner filtration apparatus. Place the funnel in the filter flask, making sure that the neck fits snugly in a rubber washer. Moisten the filter paper. Turn on the vacuum or water aspirator and slowly pour the acetaminophen onto the center of the filter paper in the Büchner funnel (see Figure 28.1) Push down gently on the funnel to create the suction needed to pull the water off the aspirin product. The acetaminophen crystals will collect on the filter paper. Use a spatula to transfer any crystals left in the flask. Rinse the inside of the flask with a 10-mL portion of ice cold water and transfer all the crystals to the funnel. Wash the crystals on the filter paper with two 10-mL portions of cold water. Spread the crystals out on the filter paper and draw air through the funnel. This helps dry the crystals. Remove the Büchner funnel and turn off the water. Use a spatula to lift and transfer the filter paper and crystals to a paper towel. Don’t touch it; it may still contain acid.

▲ FIGURE 28.1 Apparatus for suction filtration with a Büchner funnel.

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Recrystallization of Crude Acetaminophen

3. 4. 5.

Place the crude acetaminophen in a 125-mL flask and add 20 mL of water and heat on a hot plate until all of the solid dissolves. If the solution reaches boiling and crystals remain, add more water, a few mL at a time. Remove the flask and allow the solution to cool. When crystals begin to appear, place the flask in an ice bath for 20 min. If no crystals appear, scratch the inside walls with a glass stirring rod. Collect the crystals using the Büchner filtration apparatus. Wash with 10 mL of cold water. Transfer the filter paper and crystals to a paper towel and let dry until the next laboratory class. Weigh a dry watch glass. Record the mass. Transfer the pure acetaminophen to the watch glass. Weigh the watch glass and product. Record the combined mass. Using a melting point apparatus, determine the melting point of your dry, pure acetaminophen product.

Calculations 6.

Calculate the mass of starting material p-aminophenol (2 − 1).

7. 8.

Calculate the actual yield of pure acetaminophen (4 − 3). Calculate the theoretical yield of acetaminophen product.

grams of p -aminophenol (6) ×

1 mol p-aminophenol 109.1 g p-aminophenol

×

1 mol acetaminophen 1 mol p -aminophenol

= grams of acetaminophen (theoretical yield)

9.

Calculate the percentage yield of the pure acetaminophen. Percent yield (%) =

actual yield (7) × 100% theoretical yield (8)

×

151.2 g acetaminophen 1 mol acetaminophen

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Pre-Lab Study Questions

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1. Draw the condensed structural formula of acetaminophen.

2. What method is used to remove the impurities in a product?

3. A student measures 1.62 g of p-aminophenol, which undergoes reaction to produce acetaminophen. After recrystallization, 1.57 g of pure acetaminophen is obtained. a. What is the theoretical yield, in grams, of acetaminophen from the reaction?

b. What is the percent yield of acetaminophen for the reaction?

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Synthesis of Acetaminophen

LAB

28

Synthesis of Acetaminophen 1. Mass of flask 2. Mass of flask and p-aminophenol 3. Mass of watch glass 4. Mass of watch glass + pure acetaminophen 5. Melting point of the pure acetaminophen 6. Mass of p-aminophenol 7. Mass of pure acetaminophen 8. Theoretical yield (g) of acetaminophen (Show calculations)

9. Percent yield (Show calculations).

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Questions and Problems Q1 Acetaminophen does not act as a base in water, but p-aminophenol does. Explain.

Q2 How does the melting point of the pure acetaminophen (5) compare to the known melting point of acetaminophen, which is 169–171 °C?

Q3 Phenacetin is another over-the-counter medication for reducing fever and relieving pain.

a. Circle and label the functional groups in phenacetin.

b. How might you prepare phenacetin?

Types of Carbohydrates 29 LABORATORY GOALS • Identify the characteristic functional groups of carbohydrates. • Describe common carbohydrates and their sources. • Distinguish between monosaccharides, disaccharides, and polysaccharides.

LAB INFORMATION Time: Comments: Related Topics:

2h Tear out the report sheets and place them next to the matching procedures. In the study of carbohydrates, it is helpful to review stereoisomers and the formation of hemiacetals. Carbohydrates, monosaccharides, disaccharides, polysaccharides, hemiacetals, stereoisomers, aldohexoses, ketohexoses, chiral compounds, Fischer projections, Haworth structures

CHEMICAL CONCEPTS Carbohydrates in our diet are our major source of energy. Foods high in carbohydrates include potatoes, bread, pasta, and rice. If we take in more carbohydrates than we need for energy, the excess is converted to fat, which can lead to a weight gain. The carbohydrate family can be organized into three classes: monosaccharides, disaccharides, and polysaccharides.

▲Carbohydrates contained in foods such as pasta and bread provide energy for the body.

A. Monosaccharides Monosaccharides contain C, H, and O in units of (CH 2 O)n . Most common monosaccharides have 6 carbon atoms (hexoses) with a general formula of C6 H12 O6 . They contain many hydroxyl groups (~OH) along with a carbonyl group. The aldoses are monosaccharides with an aldehyde group, and ketoses contain a ketone group.

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Monosaccharides

Sources

Glucose

C6 H12 O6

Fruit juices, honey, corn syrup

Galactose

C6 H12 O6

Lactose hydrolysis

Fructose

C6 H12 O6

Fruit juices, honey, sucrose hydrolysis

The letters D and L refer to the orientation of the hydroxyl (~OH) group on the carbon that is one above the bottom carbon. In the isomers of glyceraldehyde, the position of the ~OH is on the right in the D-isomer and on the left in the L-isomer.

The hexoses glucose, galactose, and fructose are the most important monosaccharides. The most common hexose, D-glucose, C6 H12 O6 , also known as dextrose, is found naturally in fruits, vegetables, corn syrup, and honey; it is also the form of glucose in circulating blood, known as blood sugar. D-glucose is a building block of the disaccharides sucrose, lactose, and maltose, and polysaccharides such as amylose, cellulose, and glycogen. Galactose, C6 H12 O6 , is an aldohexose that is obtained from the disaccharide lactose, which is found in milk and milk products. Galactose is important in the cellular membranes of the brain and nervous system. The only difference in D-glucose and D-galactose is the arrangement of the ~OH group on carbon 4. In contrast to glucose and galactose, fructose, C6 H12 O6 , is a ketohexose. The structure of fructose differs from glucose at carbons 1 and 2 by the location of the carbonyl group. Fructose is the sweetest of the carbohydrates; it is almost twice as sweet as sucrose (table sugar).

Haworth Structures Most of the time glucose exists in a ring structure, which forms when the ~OH on carbon 5 bonds to carbon 1 in the carbonyl group. In the Haworth structure, the new hydroxyl group may be drawn below carbon 1 (the α form or anomer) or above carbon 1 (the β form or anomer).

Types of Carbohydrates 325

B. Disaccharides A disaccharide contains two monosaccharides bonded together. Some common disaccharides include maltose, sucrose (table sugar), and lactose (milk sugar). Disaccharides

Sources

Monosaccharides

Maltose

Germinating grains, starch hydrolysis

Glucose + glucose

Lactose

Milk, yogurt, ice cream

Glucose + galactose

Sucrose

Sugar cane, sugar beets

Glucose + fructose

In a disaccharide, two monosaccharides form a glycosidic bond with the loss of water. For example, in maltose, two glucose units are linked by an α-1,4-glycosidic bond.

Lactose, milk sugar, is a disaccharide found in milk and milk products. The bond between the monosaccharides in lactose is a β-1,4-glycosidic bond because it is the ~OH group of a β form of galactose that forms a bond with the ~OH group on carbon 4 of glucose. Lactose is a reducing sugar because the bond at carbon 1 can open to give an aldehyde that can reduce other substances.

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▲α-Lactose, a disaccharide found in milk and milk products, contains β-D-galactose and α-D-glucose.

C. Polysaccharides Polysaccharides are long-chain polymers that contain many thousands of monosaccharides (usually glucose units) joined together by glycosidic bonds. Three important polysaccharides are starch, cellulose, and glycogen. They all contain glucose units, but differ in the type of glycosidic bonds and the amount of branching in the polymer. Polysaccharides

Found in

Monosaccharides

Starch (amylose, amylopectin)

Rice, wheat, grains, cereals

Glucose

Glycogen

Muscle, liver

Glucose

Cellulose

Wood, plants, paper, cotton

Glucose

Starch is an insoluble storage form of glucose found in rice, wheat, potatoes, beans, and cereals. Starch is composed of two kinds of polysaccharides, amylose and amylopectin. Amylose, which makes up about 20% of starch, consists of α-D-glucose molecules connected by α-1,4-glycosidic bonds in a continuous chain. A typical polymer of amylose may contain from 250 to 4000 glucose units. Amylopectin is a branched-chain polysaccharide that makes up as much as 80% of starch. In amylopectin, α-1,4-glycosidic bonds connect most of the glucose molecules. However, at about every 25 glucose units, there are branches of glucose molecules attached by α-1,6-glycosidic bonds between carbon 1 of the branch and carbon 6 in the main chain. Glycogen is similar to amylopectin but it is even more highly branched, with α-1,6-glycosidic bonds about every 10 to 15 glucose units.

Types of Carbohydrates 327

Cellulose is the major structural material of wood and plants. Cotton is almost pure cellulose. In cellulose, glucose molecules form a long unbranched chain similar to amylose except that β-1,4-glycosidic bonds connect the glucose molecules. The β isomers are aligned in parallel rows that are held in place by hydrogen bonds between the rows. This gives a rigid structure for cell walls in wood and fiber and makes cellulose more resistant to hydrolysis.

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Types of Carbohydrates 329

EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Monosaccharides Materials: Organic model kits or prepared models 1. 2.

Make or observe models of L-glyceraldehyde and D-glyceraldehyde. Draw their Fischer projections. Draw the Fischer projection for D-glucose, D-galactose, and D-fructose.

3.

Draw the Haworth (cyclic) structures for the α and β forms (anomers) for D-glucose.

4.

Draw the Haworth (cyclic) structures for the α forms (anomers) of D-galactose and D-fructose.

B. Disaccharides 1. 2.

Using Haworth structures, draw the structure for α-maltose. Using Haworth structures, write an equation for the hydrolysis of α-maltose by adding H 2 O to the glycosidic bond.

3.

Using Haworth structures, write an equation for the formation of α-lactose from β-D-galactose and α-D-glucose. Draw the Haworth structure of sucrose and circle the glycosidic bond.

4.

C. Polysaccharides 1.

Using Haworth structures, draw a portion of amylose consisting of four units of α-D-glucose. Indicate the glycosidic bonds.

2.

Using Haworth structures, draw a portion of cellulose consisting of four units of β-D-glucose. Indicate the glycosidic bonds.

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Pre-Lab Study Questions

29

1. What are some sources of carbohydrates in your diet?

2. What does the D in D-glucose mean?

3. What is the bond that links monosaccharides in di- and polysaccharides?

4. Draw the Haworth structure of each of the following: a. α-D-galactose.

b. β-lactose

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REPORT SHEET

Types of Carbohydrates

LAB

29

A. Monosaccharides 1.

L-Glyceraldehyde

2.

D-Glucose

3. α-D-Glucose

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D-Glyceraldehyde

D-Galactose

D-Fructose

β-D-Glucose

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4. α-D-Galactose

α-D-Fructose

Questions and Problems Q1 Describe how the structure of D-glucose compares to the structure of D-galactose.

B. Disaccharides 1. α-maltose

2. Equation for the hydrolysis of α-maltose

3. Equation for the formation of α-lactose

4. Sucrose

Types of Carbohydrates 335

Questions and Problems Q2 What is the type of glycosidic bond in maltose?

Q3 Why does maltose have both α and β forms (anomers)? Explain.

C. Polysaccharides 1. Amylose

What is the difference in the structure of amylopectin and amylose?

2. Cellulose

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Questions and Problems Q4 What is the monosaccharide that results from the complete hydrolysis of amylose?

Q5 What is the difference in the structure of amylose and cellulose?

Tests for Carbohydrates 30 LABORATORY GOALS • Observe physical and chemical properties of some common carbohydrates. • Use physical and chemical tests to distinguish between monosaccharides, disaccharides, and polysaccharides. • Identify an unknown carbohydrate.

LAB INFORMATION Time: 3h Comments: Tear out the report sheets and place them next to the matching procedures. Related Topics: Carbohydrates, aldohexoses, ketohexoses, reducing sugars, fermentation Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS A. Benedict’s Test for Reducing Sugars All of the monosaccharides and most of the disaccharides can be oxidized to carboxylic acids. When the cyclic structure opens, the aldehyde group is available for oxidation. Benedict’s reagent contains Cu 2+ ion that is reduced. Therefore, all the sugars that react with Benedict’s reagent are called reducing sugars. Ketoses also act as reducing sugars because the ketone group on carbon 2 isomerizes to give an aldehyde group on carbon 1.

When oxidation of a sugar occurs, the Cu 2+ is reduced to Cu + , which forms a red precipitate of copper(I) oxide, Cu 2 O(s ). The color of the precipitate varies from green to gold to red depending on the concentration of the reducing sugar. If the concentration of the reducing sugar is low, the sample will have a gold or green color; if the concentration is high, the sample will have a red color. Sucrose is not a reducing sugar because it cannot revert to the open-chain form that would provide the aldehyde group needed to reduce the copper(II) ion.

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B. Seliwanoff’s Test for Ketoses Seliwanoff’s test is used to distinguish between hexoses with a ketone group (ketoses) and hexoses that are aldehydes (aldoses). With ketoses, a deep red color is formed rapidly. Aldoses give a light pink color that takes a longer time to develop. The test is most sensitive for fructose, which is a ketose.

C. Fermentation Test Most monosaccharides and disaccharides undergo fermentation in the presence of yeast. The products of fermentation are ethyl alcohol (CH3 ~ CH 2 ~ OH) and carbon dioxide (CO 2 ). The formation of bubbles of carbon dioxide is used to confirm the fermentation process. yeast

C6 H12 O6 ⎯⎯⎯→ 2C 2 H5 OH + 2CO 2 ( g ) Glucose

Ethanol

Although enzymes are present in yeast for the hydrolysis of most disaccharides, they are not available for lactose or galactose. Therefore, lactose and galactose give negative results with the fermentation test.

D. Iodine Test for Polysaccharides When iodine (I 2 ) is added to amylose, the helical shape of the unbranched polysaccharide traps iodine molecules, producing a deep blue-black complex. Amylopectin, and cellulose, react with iodine to give red-to-brown colors. Glycogen produces a reddish-purple color. Monosaccharides and disaccharides are too small to trap iodine molecules and do not form dark colors with iodine.

E. Hydrolysis of Disaccharides and Polysaccharides Disaccharides hydrolyze in the presence of an acid to give the individual monosaccharides. H+

Sucrose + H 2 O ⎯⎯→ glucose + fructose

In the laboratory, we use water and acid to hydrolyze starches, which produce smaller saccharides such as maltose. Eventually, the hydrolysis reaction converts maltose to glucose molecules. In the body, enzymes in our saliva and from the pancreas carry out the hydrolysis. Complete hydrolysis produces glucose, which provides about 50% of our nutritional calories. H + or amylase

H + or amylase

H + or maltase

Amylose, amylopectin ⎯⎯⎯⎯ → dextrins ⎯⎯⎯⎯ → maltose ⎯⎯⎯⎯ → many D-glucose units

F. Testing Foods for Carbohydrates Several of the tests, such as the iodine test, can be carried out with food products such as cereals, bread, crackers, and pasta. Some of the carbohydrates we have discussed can be identified.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Benedict’s Test for Reducing Sugars Materials: Test tubes, 400-mL beaker, droppers, hot plate or Bunsen burner, 5- or 10-mL graduated cylinder, Benedict’s reagent, 2% carbohydrate solutions: glucose, fructose, lactose, sucrose, starch, and an unknown Place 10 drops of solutions of water (reference), glucose, fructose, lactose, sucrose, starch, and an unknown in separate test tubes. Label each test tube. Add 2 mL of Benedict’s reagent to each sample and mix by shaking. Place the test tubes in a boiling water bath for 3–4 min. The formation of a greenish to reddish-orange color indicates the presence of a reducing sugar. If the solution is the same color as the Benedict’s reagent, no oxidation reaction occurred. 1. Record your observations. 2. Classify each as a reducing or nonreducing sugar.

B. Seliwanoff’s Test for Ketoses Materials: Test tubes, 400-mL beaker, droppers, hot plate or Bunsen burner, 5- or 10-mL graduated cylinder, Seliwanoff’s reagent, 2% carbohydrate solutions: glucose, fructose, lactose, sucrose, starch, and an unknown Place 10 drops of solutions of glucose, fructose, lactose, sucrose, starch, water, and unknown in separate test tubes. Label each test tube. Add 2 mL of Seliwanoff’s reagent to each. Mix by shaking. The reagent contains concentrated HCl. Use carefully. Place the test tubes in a boiling hot-water bath and note the time. 1. After 1 min, observe the colors in the test tubes. A rapid formation of a deep red color indicates the presence of a ketose. 2. Classify each as a ketose or aldose.

C. Fermentation Test Materials: Fermentation tubes (or small and large test tubes), baker’s yeast, 2% carbohydrate solutions: glucose, fructose, lactose, sucrose, starch, and an unknown Fill the fermentation tubes with solutions of glucose, fructose, lactose, sucrose, starch, water, and the unknown. Add 0.2 g of yeast to each fermentation tube and mix well by shaking (see Figure 30.1). If fermentation tubes are not available, use small test tubes placed upside down in larger test tubes. Cover the mouth of the large test tube with filter paper or cardboard. Place your hand firmly over the paper cover and invert. When the small test tube inside has completely filled with the mixture, return the larger test tube to an upright position (see Figure 30.2). Set the tubes aside.

▲FIGURE 30.1 Fermentation tube filled with a carbohydrate solution.

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▲FIGURE 30.2 Test tubes used as fermentation tubes.

1. 2.

At the end of the laboratory period, and again at the next laboratory period, look for gas bubbles in the fermentation tubes or inside the small tubes. Record your observations (see Figure 30.3). State whether fermentation occurred or not.

▲FIGURE 30.3 Fermentation tubes with CO2 bubbles.

D. Iodine Test for Polysaccharides Materials: Spot plate or test tubes, droppers, iodine reagent, 2% carbohydrate solutions in dropper bottles: glucose, fructose, lactose, sucrose, starch, and an unknown Using a spot plate, place 5 drops of each solution of glucose, fructose, lactose, sucrose, starch, water, and the unknown in the wells. (If you do not have a spot plate, use small test tubes.) Add 1 drop of iodine solution to each sample. A dark blue-black color is a positive test for amylose in starch. 1. Record your observations. 2. Indicate whether amylose is present.

E. Hydrolysis of Disaccharides and Polysaccharides Materials: Test tubes, 10-mL graduated cylinder, 400-mL beaker (boiling water bath), hot plate or Bunsen burner, spot plate or watch glass, 10% HCl, 10% NaOH, red litmus paper, iodine reagent, Benedict’s reagent, 2% starch and sucrose solutions in dropper bottles Place 3 mL of 2% starch in 2 test tubes. Place 3 mL of 2% sucrose in 2 more test tubes. To one sample each of sucrose and starch, add 20 drops of 10% HCl. Mix by shaking. To the other samples of sucrose and starch, add 20 drops of H 2 O. Mix by shaking.

Label the test tubes and heat in a boiling water bath for 10 min. Remove the test tubes from the water bath and let them cool. To the samples containing HCl, add 10% NaOH (about 20 drops) until 1 drop of the mixture turns litmus paper blue, indicating the HCl has been neutralized.

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Test separate samples for hydrolysis using the iodine test and Benedict’s test as described below. Iodine Test

1.

Place 5 drops of each solution on a spot plate or watch glass. Add 1 drop of iodine reagent to each. Record your observations.

2. 3.

Benedict’s Test Add 2 mL of Benedict’s reagent to each of the samples. Heat each sample in a boiling water bath for 3–4 min. Record your observations. Determine if hydrolysis has occurred in each.

F. Testing Foods for Carbohydrates Materials: Samples of carbohydrates (refined, brown, “natural,” powdered), honey, syrups (corn, maple, fruit), foods with starches: cereals, pasta, bread, crackers, potato, Benedict’s reagent, Seliwanoff’s reagent, iodine reagent Obtain 2 carbohydrate samples to test. Perform the Benedict’s test on each. Perform the Seliwanoff’s test on each. Perform the iodine test on each. 1. Describe the kinds of carbohydrates you identify in each sample.

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1. What happens to glucose or galactose when the Cu 2+ in Benedict’s reagent is reduced?

2. Would you expect fructose or glucose to form a red color rapidly with Seliwanoff’s reagent? Explain your answer.

3. Why don’t all the disaccharides undergo fermentation with yeast?

4. How can the iodine test be used to distinguish between amylose and glycogen?

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REPORT SHEET

Tests for Carbohydrates Test

LAB

30

Water Glucose Fructose Lactose Sucrose Starch Unknown #_______

A. Benedict’s Test for Reducing Sugars 1. Observations 2. Reducing/ Nonreducing

B. Seliwanoff’s Test for Ketoses 1. Colors after 1 min 2. Ketose/Aldose

C. Fermentation Test 1. Observations 2. Fermentation (Yes/No)

D. Iodine Test for Polysaccharides 1. Observations 2. Amylose (Yes/No) Questions and Problems Q1 From the results in Part A, list the carbohydrates that are reducing sugars.

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Q2 From the results in Part B, which carbohydrates are ketoses?

Q3 From the results in Part C, which carbohydrates give a positive fermentation test?

Q4 From the results in Part D, which carbohydrates give a blue-black color in the iodine test?

Identifying an Unknown Carbohydrate Unknown No.__________ Results with Unknown

Carbohydrates Present

Benedict’s (A) Seliwanoff’s (B) Fermentation (C) Iodine (D) What carbohydrate(s) is/are in your unknown?

Questions and Problems Q5 What carbohydrate(s) would give the following test results? a. Produces a reddish-orange solid with Benedict’s reagent and a red color with Seliwanoff’s reagent in 1 minute

b. Produces a color change with Benedict’s reagent, a light orange color with Seliwanoff’s reagent after 5 min, and produces no bubbles during fermentation

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c. Produces no color change with Benedict’s or Seliwanoff’s reagent, but turns a blue-black color with iodine reagent

E. Hydrolysis of Disaccharides and Polysaccharides Sucrose + H 2O Sucrose + HCl Starch + H 2O Sucrose + HCl

Results

1. Iodine test 2. Benedict’s test 3. Hydrolysis (Yes/No) Questions and Problems Q6 How do the results of the Benedict’s test indicate that hydrolysis of sucrose and starch occurred?

Q7 How do the results of the iodine test indicate that hydrolysis of starch occurred?

Q8 Indicate whether the following carbohydrates will give a positive (+) or a negative (–) result in each type of test listed below: Benedict’s Test Seliwanoff’s Test Fermentation Test Iodine Test Glucose Fructose Galactose Sucrose Lactose Maltose Amylose

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F. Testing Foods for Carbohydrates Food Item Tested Benedict’s test

Seliwanoff’s test

Iodine test

Carbohydrates present

Lipids 31 LABORATORY GOALS • • • •

Observe the physical and chemical properties of some common lipids. Draw the structure of a typical triacylglycerol. Distinguish between saturated and unsaturated fats. Prepare a hand lotion and determine the function of its components.

LAB INFORMATION Time: Comments:

3h Bromine can cause severe chemical burns. Use carefully. Tear out the report sheets and place them beside the matching procedures. Related Topics: Fatty acids, saturated and unsaturated fatty acids, lipids, triglycerides Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS Lipids Lipids are a family of biomolecules that have the common property of being soluble in organic solvents but not in water. The word “lipid” comes from the Greek word lipos, meaning “fat” or “lard.” Within the lipid family, there are specific structures that distinguish the different types of lipids. Lipids such as waxes, fats, oils, and triacylphospholipids are esters that can be hydrolyzed to give fatty acids along with the alcohol glycerol. Steroids, which have a completely different structure, do not contain fatty acids and cannot be hydrolyzed. Steroids are characterized by the steroid nucleus of four fused carbon rings.

▲Lipids are naturally occurring compounds that are soluble in organic solvents and not in water.

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A. Physical Properties of Lipids and Fatty Acids Lipids are a family of compounds that are grouped by similarities in solubility rather than structure. As a group, lipids are more soluble in nonpolar solvents such as ether, chloroform, or benzene than in water. Important types of lipids include fats and oils, glycerophospholipids, and steroids. Compounds classified as lipids include fat-soluble vitamins A, D, E, and K; cholesterol; hormones; portions of cell membranes; and vegetable oils. Table 31.1 lists the classes of lipids. TABLE 31.1 Classes of Lipid Molecules Lipids

Components

Waxes

Fatty acid and long-chain alcohol

Fats and oils (triacylglycerols)

Fatty acids and glycerol

Glycerophospholipids

Fatty acids, glycerol, phosphate, amino alcohol

Sphingolipids

Fatty acids, sphingosine, phosphate, amino alcohol

Steroids

A fused structure of three cyclohexanes and a cyclopentane

The structural formulas of three typical lipids are shown here:

B. Triacylglycerols Fatty acids are long-chain carboxylic acids, usually 12 to 18 carbons in length.

▲The structural formula for lauric acid, a fatty acid, can be drawn in several ways.

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The triacylglycerols are esters of glycerol and fatty acids. When a fatty acid contains one or more carbon-carbon double bonds, the triacylglycerol is referred to as an unsaturated fat. When the fatty acid consists of only carbon-carbon single bonds, the triacylglycerol is a saturated fat. Table 31.2 gives the formulas of the common fatty acids and their melting points. At room temperature, saturated fats are usually solid and unsaturated fats are usually liquid. Fats that contain mostly saturated fatty acids have a higher melting point than fats with more unsaturated fatty acids.

▲When glycerol reacts with stearic acid, the triacylglycerol that forms is glyceryl tristearate (tristearin). TABLE 31.2 Structures and Melting Points of Common Fatty Acids Name

Carbon Source Atoms

Melting Structures Point (°C)

Saturated Fatty Acids Lauric acid

12

Coconut

44

Myristic acid

14

Nutmeg

55

Palmitic acid

16

Palm

63

Stearic acid

18

Animal fat

69

Monounsaturated Fatty Acids Palmitoleic acid

16

Butter

Oleic acid

18

Olive, pecan, grapeseed

0 14

Polyunsaturated Fatty Acids Linoleic acid

18

Soybeans, sunflowers

−5

Linolenic acid

18

Corn

−11

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C. Bromine Test for Unsaturation The presence of unsaturation in a fatty acid or a triacylglycerol can be detected by the bromine test, which you used in an earlier experiment to detect double bonds in alkenes. If the orange color of the bromine solution fades quickly, an addition reaction has occurred and the oil or fat is unsaturated.

D. Preparation of Hand Lotion

▲Triacylglycerols are used to thicken creams and lotions. Lotions are applied to the skin to soften, smooth, and hydrate the skin. The lotion contains components soluble in water and oils, which are not soluble in water. A lotion is prepared by mixing the water and oils to form an emulsion. Stearic acid and triethanolamine form a salt that helps emulsify the components of the hand lotion. Glycerin is used to improve the texture of the lotion or cream. Lanolin is a waxy substance that helps the skin absorb water.

Lipids

Typical Composition of a Hand Lotion Nonpolar Components

Polar Components

stearic acid cetyl alcohol

CH3 ~(CH 2 )15 ~OH

lanolin

complex mixture of waxes

triethanolamine glycerol

water

H2O

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Physical Properties of Lipids and Fatty Acids Materials: Test tubes and stoppers, dropper bottles or solids: stearic acid, oleic acid, olive oil, safflower oil, lecithin, cholesterol, spatulas, CH 2 Cl2 1.

2. 3.

To 6 separate test tubes, add 5 drops or the amount of solid lipid held on the tip of a spatula of stearic acid, oleic acid, olive oil, safflower oil, lecithin, cholesterol. Classify each as a triacylglycerol, fatty acid, steroid, or glycerophospholipid. Describe the appearance of each. Describe the odor of each.

May be done as a demonstration 4. Add about 2 mL of water to each of the test tubes. Stopper and shake each test tube. Record the solubility of the lipids in water (a polar solvent). May be done as a demonstration 5. In the hood: To 6 separate test tubes, add 5 drops or the amount of solid lipid held on the tip of a spatula of stearic acid, oleic acid, olive oil, safflower oil, lecithin, cholesterol. Add 1 mL (20 drops) methylene chloride, CH 2 Cl2 , to each sample. Record the solubility of the lipids in a nonpolar solvent. Save the test tubes and samples of stearic acid, oleic acid, olive oil, and safflower oil for Part C.

B. Triacylglycerols Materials: Organic model kits or models 1. 2. 3.

Use an organic model kit or look at prepared models of a molecule of glycerol and three molecules of ethanoic acid. What are the functional groups on each? Write the equation for the formation of the glyceryl triethanoate by drawing the condensed structural formulas of the reactants and the products. Write the equation for the hydrolysis of the glyceryl triethanoate by drawing the condensed structural formulas of the reactants and the products.

C. Bromine Test for Unsaturation Materials: Samples from Part A.5, 1% Br2 in methylene chloride 1.

2.

In the hood: To the samples of stearic acid, oleic acid, safflower oil, and olive oil from Part A.5, add 1% bromine solution, drop by drop, until a permanent red-orange color is obtained or until 20 drops have been added. Caution: Avoid contact with bromine solution; it can cause painful burns. Do not breathe the fumes. Determine if the red-orange color fades rapidly or persists. Record your observations. Indicate whether each sample contains a saturated or unsaturated lipid.

D. Preparation of Hand Lotion Materials: Stearic acid, cetyl alcohol, lanolin (anhydrous), triethanolamine, glycerol, ethanol, distilled water, commercial hand lotions, 100-mL graduated cylinder, 50-mL beaker, 100-mL beaker, 250-mL beaker for water bath thermometer, Bunsen burner, iron ring, wire screen or hot plate, stirring rods, pH paper

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Hand Lotion Place the substances listed in the table below, into beaker 1 and beaker 2. For beaker 1, use a laboratory balance to weigh out the solid substances. For beaker 2, measure 50 mL water in the 100 mL graduated cylinder. Then add glycerol to the cylinder for a total of 52 mL liquid. Stir and transfer into beaker 2. Hand lotion

Hand lotion without triethanolamine

Hand lotion without stearic acid

Beaker 1 Stearic acid

3g

Stearic acid

3g

(50 mL)

Cetyl alcohol

1g

Cetyl alcohol

1g

Cetyl alcohol

1g

Lanolin (anhydrous)

2g

Lanolin (anhydrous)

2g

Lanolin (anhydrous)

2g

Triethanolamine

1 mL

Triethanolamine

1 mL

Glycerol

2 mL

Beaker 2 Glycerol (100 mL) Water

2 mL Glycerol 50 mL Water

2 mL 50 mL

Water

50 mL

Procedure: Fill a 250-mL beaker about 2/3 full of water. Place the beaker on an iron ring covered with a wire screen. Heat the water using a Bunsen burner. (Alternatively, a hot plate may be used.) Place beaker 1 in the water bath and heat until all the compounds have melted. Remove beaker 1 from the water bath. Place beaker 2 in the water bath and heat until all the compounds have melted and reached 80 °C. Remove beaker 2 from the water bath. While they are still warm, slowly pour the contents from beaker 1 into beaker 2 as you stir. Add 5.0 mL of ethanol and stir until a smooth, creamy lotion is obtained. If the resulting product is too thick, add more warm water. 1. 2. 3.

Describe the appearance of the hand lotion. Use pH paper to determine the pH of the hand lotion you prepared. Use pH paper to determine the pH of any commercial hand lotions.

If there is time or if your instructor suggests, repeat the procedure for preparing a hand lotion but without adding triethanolamine. If there is time or if your instructor suggests, repeat the procedure for preparing a hand lotion but without adding stearic acid.

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31

1. What is the functional group in a triacylglycerol?

2. Draw the skeletal formula of linolenic acid. Why is it an unsaturated fatty acid?

3. What type of solvent is needed to remove an oil spot? Why?

4. Write the equation for the esterification of glycerol and three palmitic acids.

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Lipids

LAB

31

A. Physical Properties of Lipids and Fatty Acids Lipid

1. Type

2. Appearance 3. Odor

4. Soluble 5. Soluble in Water in CH 2Cl 2 (yes/no) (yes/no)

Stearic acid Oleic acid Olive oil Safflower oil Lecithin Cholesterol Questions and Problems Q1 Why are the compounds in Part A classified as lipids?

B. Triacylglycerols 1. Functional group of glycerol Functional group of ethanoic acid 2. Equation for the esterification of glycerol and three ethanoic acids.

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3. Equation for the hydrolysis of glyceryl triethanoate.

C. Bromine Test for Unsaturation 1. Color fades rapidly/persists

2. Saturated/Unsaturated

Fatty acids Stearic acid Oleic acid Triacylglycerols Safflower oil Olive oil Questions and Problems Q3 a. Draw the condensed structural formulas of stearic acid and oleic acid. Stearic acid

Oleic acid

b. From the results of Part C, which is more unsaturated: oleic acid or stearic acid? Explain your reason.

c. From the results of Part C, were safflower oil and olive oil saturated or unsaturated? Explain your reason.

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D. Preparation of Hand Lotion 1. Appearance Hand Lotion

(without triethanolamine)

(without stearic acid)

2. pH

3. pH of commercial hand lotions Brand ________________

Brand ________________

Brand ________________

pH ___________________ pH ___________________ pH ___________________ Questions and Problems Q4 How does omitting triethanolamine affect the properties and appearance of the hand lotion?

Q5 How does omitting stearic acid affect the properties and appearance of the hand lotion?

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Saponification and Soaps 32 LABORATORY GOALS • Prepare soap by the saponification of a fat or oil. • Observe the reactions of soap with oil, CaCl2 , MgCl2 , and FeCl3 .

LAB INFORMATION Time: Comments:

2h You will be working with hot oil and NaOH. Be sure to wear your goggles. Tear out the report sheets and place them beside the matching procedures. Related Topics: Esters, saponification, soaps, hydrophobic, hydrophilic Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS A. Saponification: Preparation of Soap For centuries, soaps have been made from animal fats and lye (NaOH), which was obtained by pouring water through wood ashes. The hydrolysis of a fat or oil by a base such as NaOH is called saponification, and the salts of the fatty acids obtained are called soaps. The other product of hydrolysis is glycerol, which is soluble in water. Saponification occurs when a fat is heated with a strong base such as sodium hydroxide to give glycerol and the sodium salts of the fatty acids, which is soap. When NaOH is used, a solid soap is produced that can be molded into a desired shape; KOH produces a softer, liquid soap. Polyunsaturated oils produce softer soaps. Names like “coconut” or “avocado shampoo” tell you the sources of the oil used in the reaction.

The fats that are most commonly used to make soap are lard and tallow from animal fat and coconut, palm, and olive oils from vegetables. Castile soap is made from olive oil. Soaps that float have air pockets. Soft soaps are made with KOH instead of NaOH, to give potassium salts.

B. Properties of Soaps and Detergents A soap molecule has a dual nature. The nonpolar carbon chain is hydrophobic and attracted to nonpolar substances such as grease. The polar head of the carboxylate salt is hydrophilic and attracted to water.

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Nonpolar tail (hydrophobic)

Polar head (hydrophilic)

▲The dual polarity of a soap (salt of a fatty acid). Preparation of Soaps Soaps are typically prepared from fats such as coconut oil. Essential or fragrance oils are added to give a pleasant-smelling product. A soap is the salt of a long-chain fatty acid. It has a long hydrocarbon-like chain with a carboxylate negative end, combined with a positive sodium or potassium ion. Soap has dual properties because parts of the soap molecule have different solubilities. The sodium or potassium carboxylate end is ionic and very soluble in water (“hydrophilic”), but it is not soluble in oils or grease. However, the long hydrocarbon end is not soluble in water (“hydrophobic”), but it is soluble in nonpolar substances such as oil or grease. When soap is used, the hydrocarbon tails of the soap molecules dissolve in the nonpolar fats and oils that accompany dirt. The soap molecules coat the oil or grease, forming clusters called micelles, in which the water-loving salt ends of the soap molecules extend outside, where they can dissolve in water. As a result, small globules of oil and fat coated with soap molecules are pulled into the water and rinsed away. In hard water (high concentration of minerals), the carboxylate ends of soap react with Ca 2+ , Fe3+ , or Mg 2 + ions and form an insoluble substance, which we see as a gray line in the bathtub or sink. This does not occur with detergents.

▲The hydrocarbon tails of soap molecules dissolve in grease and oil to form micelles, which are pulled by their ionic heads into the rinse water.

Detergents Detergents are synthetic cleaning agents because they are not derived from naturally occurring fats or oils. They are popular because they do not form insoluble salts with ions, which means they work in hard water as well as in soft water. A typical detergent is sodium lauryl sulfate.

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Sodium lauryl sulfate, a nonbiodegradable detergent

As detergents replaced soaps for cleaning, it was found that they were not degraded in sewage treatment plants. Large amounts of foam appeared in streams and lakes that became polluted with detergents. Biodegradable detergents such as an alkylbenzenesulfonate detergent eventually replaced the nonbiodegradable detergents.

Laurylbenzenesulfonate salt, a biodegradable detergent

In addition to the sulfonate salts, commercial detergents contain phosphate compounds along with brighteners and fragrances. However, phosphates accelerate the growth of algae in lakes, which depletes the dissolved oxygen in the water. As a result, fish and other living organisms die. The use of phosphates in detergents is now decreasing and nonphosphate replacements have been developed.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Saponification: Preparation of Soap Materials: 150-mL beaker, hot plate, graduated cylinder, stirring rod or stirring hot plate with stirring bar, large watch glass, 400-mL beaker, Büchner filter system, filter paper, plastic gloves, fat (lard, solid shortening, coconut oil, olive or other vegetable oil), ethanol, 20% NaOH, saturated NaCl solution Weigh a 150-mL beaker. Add about 5 g of fat or oil to the beaker. Add 15 mL ethanol (solvent) and 15 mL of 20% NaOH. Use care when pouring NaOH. Wear disposable gloves, if available. Place the beaker on a hot plate, cover with a watch glass, and heat to a gentle boil. A magnetic stirring bar may be used with a magnetic stirrer. Heat the mixture in the beaker for 30 min or until saponification is complete and the solution becomes clear with no separation of layers. Be careful of splattering; the mixture contains a strong base. Do not let the mixture overheat or char. Add 5-mL portions of an ethanol–water (1:1) mixture to maintain volume. If foaming is excessive, reduce the heat. Caution: Oil and ethanol will be hot, and may splatter or catch fire. Keep a watch glass nearby to smother any flames. NaOH is caustic and can cause permanent eye damage. Wear goggles at all times. Obtain 50 mL of a saturated NaCl solution in a 400-mL beaker. (A saturated NaCl solution is prepared by mixing 30 g of NaCl with 100 mL of water.) Pour the soap solution into this salt solution and stir. This process, known as “salting out,” causes the soap to separate out and float on the surface. Collect the solid soap using a Büchner funnel and filter paper (see Figure 32.1). Wash the soap with two 10-mL portions of cold water. Pull air through the product to dry it further. Place the soap curds on a watch glass or in a small beaker. Use disposable, plastic gloves to handle the soap. Handle with care: The soap may still contain NaOH, which can irritate the skin. Describe the soap you prepared. Save the soap for Part B.

▲ FIGURE 32.1 Apparatus for suction filtration with Büchner funnel.

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B. Properties of Soaps and Detergents Materials: Test tubes, stoppers to fit, droppers, small beakers, 50- or 100-mL graduated cylinder, stirring rod, laboratory-prepared soap (from Part A), commercial soap product, detergent, pH paper, oil; 1% CaCl2 , 1% MgCl2 , and 1% FeCl3 solutions Using 3 small beakers, prepare solutions of the soap you made in Part A, a commercial soap, and a detergent by dissolving about 1 g of each in 50 mL of distilled water. If the commercial soap or detergent is a liquid, use 20 drops. 1.

pH test Place 10 mL of each solution in separate test tubes. Label. Dip a stirring rod into one of the solutions, touch the rod to the pH paper, and record the pH. Rinse the stirring rod before it is placed in the next solution. Save the test tubes for Part B.2.

2.

Foam test Stopper each of the tubes from Part B.1. Shake for 10 s. The soap should form a layer of suds or foam. Record your observations. Save the test tubes for Part B.3.

3.

Reaction with oil Add 5 drops of oil to each test tube from Part B.2. Stopper and shake each one for 10 s. Record your observations. Compare the sudsy layer in each test tube to the sudsy layers in Part B.2.

4.

Hard water test Place 5 mL of each solution in three separate test tubes. Add 20 drops of 1% CaCl2 solution to the first sample. Add 20 drops of 1% MgCl2 solution to the second tube. Add 20 drops of 1% FeCl3 solution to the third tube. Stopper each test tube and shake it for 10 s. Compare the foamy layer in each of the test tubes to the sudsy layer obtained in Part B.2. Record your observations for each.

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1. What happens when a fatty acid is reacted with NaOH?

2. Why is ethanol added to the reaction mixture of fat and base in the making of soap?

3. Why is the product of saponification a salt?

4. Draw the condensed structural formulas in the equation for the saponification of tristearin with KOH.

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LAB

32

A. Saponification: Preparation of Soap Describe the appearance of the soap you prepared.

B. Properties of Soaps and Detergents Tests

Lab Soap

Commercial Soap

Detergent

1. pH 2. Foam 3. Oil 4. 1% CaCl2 1% MgCl2 1% FeCl3 Questions and Problems Q1 Which of the solutions in Part B were basic?

Q2 How do soaps made from vegetable oils differ from soaps made from animal fat?

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Q3 How does soap remove an oil spot?

Q4 Draw the condensed structural formulas in the equation for the saponification of trimyristin with KOH.

Amino Acids 33 LABORATORY GOALS • Use R groups to determine if an amino acid will be nonpolar (hydrophobic), polar (hydrophilic), acidic, or basic. • Use paper chromatography to separate and identify amino acids. • Calculate R f values for amino acids.

LAB INFORMATION Time: Comments:

2 ½ –3 h Ninhydrin spray causes stains. Use it carefully. Tear out the report sheets and place them next to the matching procedures. Related Topics: Amino acids Dispose of all chemicals as directed by your lab instructor.

CHEMICAL CONCEPTS A. Amino Acids In our bodies, amino acids are used to build tissues, enzymes, skin, and hair. Essential amino acids, about half of the naturally occurring amino acids, must be obtained from the proteins in the diet because the body cannot synthesize them. Amino acids are similar in structure because each has an ionized amino group (–NH3+ ) and an ionized carboxylate group ( ~ COO − ). Individual amino acids have different organic groups (R groups) attached to the alpha (α) carbon atom. Variations in the R groups determine whether an amino acid is nonpolar (hydrophobic), polar (hydrophilic), acidic, or basic.

Some R groups contain carbon and hydrogen atoms only, which makes the amino acids nonpolar and hydrophobic (“water-fearing”). Other R groups contain ~ OH or ~ SH atoms and provide a polar area that makes the amino acids soluble in water; they are hydrophilic (“water-loving”). Other hydrophilic amino acids contain R groups that are carboxylic acids (acidic) or amino groups (basic). The R groups of some amino acids used in this experiment are given in Table 33.1.

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TABLE 33.1 Amino Acids Found in Nature Amino Acid Glycine

Symbol Gly. G

Polarity Nonpolar

Reaction to Water Hydrophobic

Alanine

Ala, A

Nonpolar

Hydrophobic

Phenylalanine

Phe, F

Nonpolar

Hydrophobic

Serine

Ser, S

Polar

Hydrophilic

Aspartic acid

Asp, D

Acidic

Hydrophilic

Lysine

Lys, K

Basic

Hydrophilic

Ionization of Amino Acids In acidic solutions (low pH), the ionized amino acid accepts a proton (H + ) to form an ion with a positive charge. When placed in a basic solution (high pH), the ionized amino acid donates a proton (H + ) to form an ion with a negative charge (see Table 33.2). This is illustrated using alanine.

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TABLE 33.2 Ionized Forms of Nonpolar and Polar Neutral Amino Acids

B. Chromatography of Amino Acids Chromatography is used to separate and identify the amino acids in a mixture. Small amounts of amino acids and unknowns are placed along one edge of Whatman #1 paper, which makes the chromatogram. The chromatogram is then placed in a container with solvent. With the paper acting like a wick, the solvent flows up the chromatogram, carrying amino acids with it. Amino acids that are more soluble in the solvent will move higher on the paper. Those amino acids that are more attracted to the paper will remain closer to the origin line. After removing and drying the chromatogram, the amino acids can be visualized by spraying the dried chromatogram with ninhydrin. The distance each amino acid travels from the starting line is measured and the R f values calculated (see Figure 33.1). Rf =

distance traveled by an amino acid distance traveled by the solvent

To identify an unknown amino acid, its R f value and color with ninhydrin are compared to the R f values and colors of known amino acids. In this way, the amino acids present in an unknown mixture of amino acids can be separated and identified.

▲ FIGURE 33.1 A developed chromatogram (R f values calculated for A, B, and C).

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Amino Acids Materials: Organic model kits or prepared models 1. 2. 3. 4.

Use an organic model kit or observe models of glycine, alanine, and serine. Draw their condensed structural formulas in the ionized form. Indicate whether each of the amino acids is polar or nonpolar. Draw the condensed structural formula of glycine, alanine, and serine in an acidic solution. Draw the condensed structural formula of glycine, alanine, and serine in a basic solution.

B. Chromatography of Amino Acids Materials: 600-mL beaker; plastic wrap; plastic gloves; Whatman chromatography paper #1 (12 cm × 24 cm); toothpicks or capillary tubing; drying oven (80 °C) or hair dryer; metric ruler; stapler; amino acids (1% solutions): phenylalanine, alanine, glycine, serine, lysine, aspartic acid, and unknown amino acid; chromatography solvent: isopropyl alcohol, 0.5 M NH 4 OH; 0.2% ninhydrin spray Keep your fingers off the chromatography paper because amino acids can be transferred from the skin. Do not use pen, the ink will dissolve in the solvent and ruin your results. Using forceps, plastic gloves, or a paper towel, pick up a piece of Whatman #1 chromatography paper that has been cut to a size of 12 cm × 24 cm. Draw a pencil line about 2 cm from the long edge of the paper or plate. This will be the starting line (origin). With a pencil, mark off 7 equally spaced points about 3 cm apart along the line (see Figure 33.2). Label each with the abbreviation of one of the amino acids. Place your name or initials in the upper corner with the pencil. Use the toothpick applicators or capillary tubes provided in each 1% amino acid solution to make a small spot (the size of the letter o) by lightly touching the tip to the paper or plate. Always return the applicator to the same amino acid solution. Dry the spot either with a hair dryer or by allowing it to air dry. After the spot dries, repeat the application of the amino acids twice more, for a total of three applications.

▲ FIGURE 33.2 Preparation of a chromatogram.

Amino Acids 377

Prepare the solvent by mixing 10 mL of 0.5 M NH 4 OH and 20 mL of isopropyl alcohol. Pour the solvent into a 600-mL beaker to a depth of about 1 cm but not over 1.5 cm. The height of the solvent must not exceed the height of the starting line on your chromatography paper. Cover the beaker tightly with plastic wrap. This is your chromatography tank. Label the beaker with your name and leave it in the hood. Roll the paper with the amino acids into a cylinder and staple the edges without overlapping. The edges should not touch. Slowly lower the cylinder into the solvent of the chromatography tank with the row of amino acids at the bottom. Make sure that the paper does not touch the sides of the beaker (see Figure 33.3).

▲ FIGURE 33.3 Chromatogram in a solvent tank. Cover the beaker with the plastic wrap and leave the tank and paper undisturbed. Let the solvent flow up the paper until the solvent front is 2–3 cm from the top edge of the paper. It may take 45–60 min. Do not let the solvent run over the top of the paper. Carefully remove the paper from the tank. Take out the staples and spread the chromatogram out on a paper towel. Immediately, use a ruler and pencil to mark the solvent line. Allow the chromatogram to dry completely. A hair dryer or an oven at about 80 °C may be used to speed up the drying process. Dispose of the solvent as directed by your instructor. Working in the hood, spray the paper lightly, but evenly, with a ninhydrin solution. Distinct, colored spots will appear as the ninhydrin reacts with the amino acids. Caution: Do not breathe the fumes or get spray on your skin. Calculations: 1. 2. 3. 4. 5.

Draw the chromatogram on the report sheet, or staple the original to the report sheet. Outline each spot with a pencil. Place a dot at the center of each spot. Record the color of each spot on the drawing or original. Measure the distance, in centimeters, from the starting line to the solvent line to obtain the distance traveled by the solvent. Record. Measure the distance, in centimeters, from the starting line to the center dot of each spot. Record. Calculate and record the R f values for the known amino acid samples and the unknown. Rf =

6.

distance traveled by an amino acid distance traveled by the solvent

Compare the color and R f values produced by the unknown amino acids to those of the known samples. Identical amino acids will give similar R f values and form the same color with ninhydrin. Identify the amino acid(s) in the unknown.

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Pre-Lab Study Questions

33

1. Which two functional groups are in all amino acids?

2. How does an R group determine if an amino acid is acidic, basic, or nonpolar?

3. In a chromatography experiment, a student calculated an R f value for alanine of 0.70 and 0.91 for leucine. Which amino acid traveled higher on the chromatography paper? Explain your reasoning.

4. In a chromatography experiment, a student calculated that the solvent front was 6.8 cm above the starting line. Arginine traveled a distance of 4.9 cm and glycine traveled 3.4 cm. What R f values would the student calculate for arginine and glycine?

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REPORT SHEET

Amino Acids

LAB

33

A. Amino Acids 1. Condensed structural formulas (ionized) of amino acids Glycine

Alanine

Serine

2. Polar or nonpolar?

3. Condensed structural formulas of amino acids in an acidic solution

4. Condensed structural formulas of amino acids in a basic solution

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Questions and Problems Q1 Write the structure of the ionized form of phenylalanine.

B. Chromatography of Amino Acids 1. Draw your chromatogram including the colors of the spots or attach the original paper here.

Calculation of R f Values Amino acid

2. Color

3. Distance (cm) solvent traveled

4. Distance (cm) amino acid traveled

5. R f value

Phenylalanine Alanine Glycine Serine Lysine Aspartic acid Unknown 6. Unknown # ______ Identification of amino acid(s) __________________________

Peptides and Proteins 34 LABORATORY GOALS • • • •

Identify the structural patterns of proteins. Observe the denaturation of proteins. Use the isoelectric point of casein in milk to isolate the protein. Use chemical tests to identify proteins and amino acids.

LAB INFORMATION Time: Comments:

2–3 h In protein tests, be sure to note the color of the reagent before you add it to the samples. Concentrated HNO3 , 10% nitric acid, and 10% NaOH are extremely corrosive and damage skin and eyes. 10% acetic acid and 10% silver nitrate can stain the skin and clothing. Tear out the report sheets and place them beside the matching procedures. Related Topics: Amino acids, peptide bonds, structural levels of proteins, denaturation of proteins Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS A. Peptides A dipeptide forms when two amino acids bond together. A peptide (or amide) bond forms between the carboxylic acid of one amino acid and the amino group of the next amino acid, with the loss of H 2 O. In the name of a peptide, each amino acid beginning from the N-terminal end has the ine (or ic acid) replaced by yl. The last amino acid at the C-terminal end of the peptide has its full name.

In the reverse reaction, hydrolysis, water adds to the peptide bond to yield the individual amino acids. For example, the tripeptide alanylglycylserine hydrolyzes to give the amino acids alanine, glycine, and serine.

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Proteins When several amino acids are joined by peptide bonds they form a polypeptide. If more than 50 amino acids are in the peptide chain that has biological activity, it is considered to be a protein. Proteins make up many important features in the body including skin, muscle, cartilage, hair, fingernails, enzymes, and hormones. Proteins have specific structures, which are determined by the sequence of the amino acids. The peptide bonds that join one amino acid to the next are the primary level of protein structure. Secondary structures include the alpha (α ) helix formed by the coiling of the peptide chain and a β-pleated sheet structure formed between protein strands. The secondary structures are held in place by many hydrogen bonds between the oxygen atoms of the carbonyl groups and the hydrogen atoms of the amide group. At the tertiary level, interactions between the side groups such as ionic bonds or salt bridges, disulfide bonds, and hydrophilic interactions give the protein a compact shape. Such tertiary structures are evident in the spherical shape of globular proteins. Similar interactions between two or more tertiary units produce the quaternary structure of many active proteins.

▲Proteins consist of (a) primary, (b) secondary, (c) tertiary, and sometimes (d) quaternary structural levels.

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B. Denaturation of Proteins Denaturation of a protein occurs when there is a disruption of the interactions between the R groups that stabilize the secondary, tertiary, or quaternary structure. However, the covalent amide bonds of the primary structure are not affected. The loss of secondary and tertiary structures occurs when conditions change, such as increasing the temperature or making the pH very acidic or basic. If the pH changes, the basic and acidic R groups lose their ionic charges and cannot form salt bridges, which causes a change in the shape of the protein. Denaturation can also occur by adding certain organic compounds or heavy metal ions or through mechanical agitation. When there is a disruption of the interactions between the R groups of a globular protein, it unfolds. With the loss of its overall shape (tertiary structure), the protein is no longer biologically active.

▲Denaturation of a protein occurs when the interactions between R groups that stabilize secondary, tertiary, or quaternary structures are disrupted and biological activity is lost.

C. Isolation of Casein (Milk Protein) A typical source of protein is milk, which contains the protein casein. When the pH of a sample of nonfat milk is decreased, it reaches its isoelectric point (the same number of positive as negative charges on the molecule), and the protein separates out of the solution. The change in pH disrupts the bonds that hold the tertiary structure together. Adjusting the pH of the mixture causes the casein to solidify so it can be removed. A similar process is used in making yogurt, cheeses, and cottage cheese. An enzyme provides the acid for the denaturation of the protein for those products. In this experiment, the mass of a quantity of milk and the mass of isolated casein will be determined. From this data, the percent casein in milk will be calculated.

D. Color Tests for Proteins Certain tests give color products with amino acids, peptides, and/or proteins. The results of the test can be used to detect certain groups or type of bonds within proteins or amino acids. Biuret test The biuret test is positive for a peptide or protein with two or more peptide bonds. In the biuret test, the blue color of a basic solution of Cu 2+ turns to a violet color when a tripeptide or larger peptide is present. Individual amino acids and dipeptides do not react with the reagent, and the solution will remain blue (negative). Ninhydrin test The ninhydrin test is used to detect amino acids and most proteins. In the test, most amino acids produce a blue-violet color. Proline and hydroxyproline give a yellow color. Xanthoproteic test This test is specific for amino acids that contain an aromatic ring. Concentrated nitric acid reacts with the side chains of tyrosine and tryptophan to give nitro-substituted benzene rings that appear as yellow-colored products.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Peptides Materials: Organic model set 1. 2. 3. 4. 5.

Use the model set or observe models of the ionized forms of the amino acids glycine and serine. Draw their ionized, condensed structural formulas. By removing the components of water (H~OH) from the amino and carboxylate groups, draw the condensed structural formulas of the dipeptides glycylserine and serylglycine. Draw the condensed structural formulas of the reactants and products for the hydrolysis of the dipeptide serylglycine by breaking the peptide bond and adding the components of H 2 O. Draw the condensed structural formula for the tripeptide serylglycylalanine. Write the 1- and 3-letter abbreviations for this tripeptide.

B. Denaturation of Proteins Materials: Test tubes, test tube holder, 10-mL graduated cylinder, 1% egg albumin (or an egg, cheesecloth, and beaker), 10% HNO3 , 10% NaOH, 95% ethyl alcohol, 1% AgNO3 (dropper bottle) If making the egg albumin solution, mix the white from one egg with 200 mL of water. Filter the mixture through cheesecloth into a beaker. Place 2–3 mL of egg albumin solution in each of 5 test tubes and label 1 to 5. For each test, record your observations and give a reason for the changes in the egg albumin protein. 1. 2.

Heat Using a test tube holder, heat the egg albumin solution over a low flame. Acid Add 2 mL of 10% HNO3 .

3. 4. 5.

Base Add 2 mL of 10% NaOH. Alcohol Add 4 mL of a 95% ethyl alcohol solution. Heavy metal ions Add 10 drops of 1% AgNO3 . Caution, AgNO3 stains the skin.

C. Isolation of Casein (Milk Protein) Materials: 150-mL beaker, hot plate or Bunsen burner, thermometer, funnel or Büchner filtration apparatus, filter paper, watch glass, nonfat milk, 10% acetic acid, dropper, pH paper, stirring rod 1. 2. 3. 4.

5.

6. 7. 8.

Weigh a 150-mL beaker. Record its mass. Add about 20 mL of nonfat milk to the beaker. Weigh the beaker and nonfat milk. Record the mass. Calculate the mass of the nonfat milk sample (2–1). Using pH paper, dip a stirring rod into the milk sample, and determine its pH. Record. Warm the sample on a hot plate or a Bunsen burner until the temperature reaches about 50 °C. Remove the beaker and milk from the heat. Add 10% acetic acid, drop by drop. You may need 2 to 3 mL. Using a stirring rod, stir continuously. When no further precipitation occurs, stop adding acid. If the liquid layer is not clear, heat the mixture gently for a few more minutes. Using pH paper, determine the pH at which the casein becomes insoluble in solution. This is the pH of the isoelectric point of casein. Record the pH. Collect the solid protein using a funnel and filter paper or the Büchner filtration apparatus. Wash the protein with two 10-mL portions of water. Weigh a dry watch glass. Record its mass. Transfer the protein to the watch glass and let the protein dry. You may use a drying oven or wait until the next laboratory session. Weigh the watch glass and dry protein and record the mass. Calculate the mass of casein (protein) (7–6). Save the casein for Part D.

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387

Calculate the percentage of casein in the nonfat milk. % Casein =

mass (g) of casein (8) × 100% mass (g) of milk (3)

D. Color Tests for Proteins Materials: Test tubes, test tube rack, 10-mL graduated cylinder, boiling water bath, cold water bath, pH paper, spatula, dropper bottles of 1% amino acid solution (glycine, tyrosine), dropper bottles of 1% solutions of proteins (gelatin, egg albumin), casein from Part C, 0.2% ninhydrin solution, concentrated HNO3 (dropper bottle), 10% NaOH, 5% CuSO 4 (biuret), red litmus paper 1.

2.

3.

Biuret test: Place 2 mL of glycine solution in a test tube. Add 2 mL of 10% NaOH. Stir. Add 5 drops of biuret reagent (5% CuSO 4 ). Stir. Record the color. Positive biuret test: The formation of a pink-violet color indicates the presence of a protein with two or more peptide bonds. If such a protein is not present, the blue color of the copper(II) sulfate will remain (negative result). Repeat the procedure for the biuret test using the amino acid tyrosine, and proteins gelatin, egg albumin, and a small amount (tip of a spatula) of solid casein (from Part C). Record the colors obtained for these remaining samples. Ninhydrin test: Place 2 mL of glycine solution in a test tube. Add 1 mL of 0.2% ninhydrin solution. Place the test tube in a boiling water bath for 4–5 min. Record your observations. Positive ninhydrin test: Look for the formation of a blue-violet color. Repeat the procedure for the ninhydrin test using the amino acid tyrosine, proteins gelatin and egg albumin, and a small amount (tip of a spatula) of solid casein (from Part C), heating all the samples in a boiling water bath for 4–5 min. Record the colors obtained for these remaining samples. Xanthoproteic test (optional): Place 2 mL of glycine solution in a test tube. Cautiously add 10 drops of concentrated HNO3 to the sample. Place the test tube in a boiling water bath and heat for 3–4 min. Remove the test tube, place in cold water, and let it cool. Carefully add 10% NaOH, drop by drop, until the solution is just basic (turns red litmus blue). This may require 2–3 mL of NaOH. Caution: heat will be evolved. Positive xanthoproteic test: Look for the formation of a yellow-orange color, which may vary in intensity. Repeat the procedure for the xanthoproteic test using the amino acid tyrosine, proteins gelatin and egg albumin, and a small amount (tip of a spatula) of solid casein (from Part C), heating all the samples in a boiling water bath for 3–4 min. Record the colors obtained for these remaining samples.

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Pre-Lab Study Questions

34

1. What is a peptide bond?

2. How does the primary structure of proteins differ from the secondary structure?

3. Write the 1- and 3-letter abbreviations of the 6 tripeptides possible from Ala, Pro, Ser.

4. Draw the condensed structural formula of the dipeptide Phe-Val (FV).

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REPORT SHEET

Peptides and Proteins

LAB

34

A. Peptides 1. Condensed structural formulas of glycine and serine Glycine

Serine

2. Condensed structural formulas of dipeptides Glycylserine

Serylglycine

3. Condensed structural formulas of the reactants and products for the hydrolysis of serylglycine

4. Condensed structural formula for the tripeptide

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5. Abbreviations for this tripeptide 3-letter abbreviation

1-letter abbreviation

B. Denaturation of Proteins Treatment

Observations

Reason for Changes

1. Heat

2. Acid

3. Base

4. Alcohol

5. Heavy metal ions

Questions and Problems Q1 Why are heat and alcohol used to disinfect medical equipment?

Q2 Why is milk given to someone who accidentally ingests a heavy metal ion such as silver or mercury?

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C. Isolation of Casein (Milk Protein) 1. Mass of beaker

__________________________________

2. Mass of beaker and milk

__________________________________

3. Mass of milk

__________________________________

4. pH of milk

__________________________________

5. pH when casein precipitates

__________________________________

6. Mass of watch glass

__________________________________

7. Mass of watch glass and casein

__________________________________

8. Mass of casein

__________________________________

9. Percent casein (Show calculations.)

__________________________________

Questions and Problems Q3 Compare the pH of the milk sample and the pH at which the casein solid forms.

Q4 How does a change in pH affect the structural levels of a protein?

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D. Color Tests for Proteins Observations of Color Tests Sample

1. Biuret

2. Ninhydrin

3. Xanthoproteic

Glycine

Tyrosine

Gelatin

Egg albumin

Casein (milk protein)

Questions and Problems Q5 After working with HNO3 , a student noticed that she had a yellow spot on her hand. What might be the reason?

Q6 Which samples give a negative biuret test? Why?

Q7 What functional group gives a positive test in the xanthoproteic test?

Q8 What tests could you use to determine whether an unlabeled test tube contained an amino acid or a protein?

Enzymes 35 LABORATORY GOALS • • • •

Prepare a solution of the enzyme amylase. Describe the role of an enzyme as a catalyst in biological systems. Set up chemical tests that measure the rate of an enzyme-catalyzed reaction. Observe the effects of enzyme concentration, temperature, pH, and inhibitors upon enzyme activity.

LAB INFORMATION Time: Comments: Related Topics:

2½ h Tear out the report sheets and place them beside the matching procedures. Enzymes, active site, lock-and-key theory, enzyme activity, factors affecting enzyme activity, inhibition, denaturation Dispose of all chemicals as directed by your laboratory instructor.

CHEMICAL CONCEPTS In biological systems, reactions are catalyzed by enzymes, which speed up reactions when operating at mild conditions of temperature and pH. As catalysts, enzymes lower the activation energy for chemical reactions (see Figure 35.1). Less energy is required to convert reactant molecules to products, which increases the rate of a biochemical reaction compared to the rate of the uncatalyzed reaction. The rates of enzyme-catalyzed reactions are much faster than the rates of the uncatalyzed reactions. Some enzymes can increase the rate of a biological reaction by a factor of a billion, a trillion, or even a hundred million trillion compared to the rate of the uncatalyzed reaction. For example, an enzyme in blood called carbonic anhydrase converts carbon dioxide and water to carbonic acid. The enzyme catalyzes the reaction of about 35 million molecules of CO 2 every minute.

▲The enzyme carbonic anhydrase lowers the activation energy for the reaction that converts CO2 and H 2O to bicarbonate and H + .

Testing Enzyme Activity In this experiment you will use amylase, an enzyme that catalyzes the hydrolysis of amylose. In the presence of amylase, a sample of starch will undergo hydrolysis to give smaller polysaccharides, dextrins, maltose, and glucose. Amylase

Amylase

Starch (amylose) ⎯⎯⎯⎯→ Smaller polysaccharides, dextrins, maltose ⎯⎯⎯⎯→ Glucose

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Visual Color Reference As you proceed with each experiment, you will check enzyme activity by adding iodine to the starch mixture. When enzyme activity is high, the time required for the starch to be hydrolyzed will be very short. When the enzyme is slowed down or inactive, the blue-black color will be seen for a longer time. By observing the rate of disappearance of starch, you can assess the relative amount of enzyme activity using Table 35.1. TABLE 35.1 Iodine test color and enzyme activity comparison table. Iodine Test for Starch

Amount of Starch Remaining

Enzyme Activity

Dark blue-black Blue

All

0

Most

1

Some

2

None

3

Light brown Gold

A. Effect of Enzyme Concentration During catalysis, an enzyme combines with the reactant or substrate of a reaction to give an enzyme−substrate complex. To form this complex, the substrate fits into the active site, where reaction takes place. The products are released and the enzyme is ready to catalyze another reaction. E Enzyme

+

S

⎯⎯ → ←⎯ ⎯

Substrate

ES

⎯⎯ →

Enzyme − Substrate complex

EP

⎯⎯ →

Enzyme − Product complex

E Enzyme

+

P Product

If the enzyme concentration is increased while substrate concentration is constant, the rate of the reaction will increase. With more enzyme present, more substrate molecules can react.

▲Increasing the enzyme concentration increases the rate of reaction.

B. Effect of Temperature The optimum temperature is the temperature at which an enzyme operates at maximum efficiency. At low temperatures, the rate of reaction is slowed. At high temperatures, the enzyme protein is denatured.

▲An enzyme attains maximum activity at its optimum temperature.

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C. Effect of pH An enzyme is most active at its optimum pH. At pH values above and below optimum, the protein structure of the enzyme is altered, which can severely reduce the enzyme’s activity.

▲Enzymes are most active at their optimum pH.

D. Inhibition of Enzyme Activity Substances that limit or stop the catalyzing activity of an enzyme are called inhibitors. A competitive inhibitor competes for the active site of an enzyme, while a noncompetitive inhibitor binds to the surface of the enzyme at another site and disrupts the structure of the active site. An irreversible inhibitor forms bonds with side-chains of the amino acids in the active site, which makes the enzyme inactive, while a reversible inhibitor can disassociate from the enzyme and restore enzyme activity.

▲A competitive inhibitor, which has a structure similar to that of the substrate, fits the active site and competes with the substrate.

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EXPERIMENTAL PROCEDURES

GOGGLES REQUIRED!

A. Effect of Enzyme Concentration Materials: Test tubes, test tube rack, thermometer, 37 °C water bath, 250- or 400-mL beaker about half full of water, droppers, spot plate (or plastic sheet), 0.05% amylase preparation, 1% starch (buffered pH 7.0), iodine reagent, 5- or 10-mL graduated cylinder Preparation of warm water bath: Fill a 250-beaker about 2/3 full of tap water. Warm to about 37 °C and try to maintain the temperature. Your lab may have a commercial water bath set at 37 °C. Testing for Enzyme Activity Add 5 mL of 1% starch in four separate test tubes labeled 1−4. Tube 1 without enzyme is the reference. Place the test tubes with 1% starch in a 37 °C water bath. After 5 min, add the following number of drops of amylase to the test tubes and mix.

1.

2. 3.

Test tube 1

Test tube 2

Test tube 3

Test tube 4

starch

5 mL

5 mL

5 mL

5 mL

amylase

None

3 drops

6 drops

9 drops

Immediately (0 min), using four separate droppers, transfer 5 drops of each reaction mixture in test tube 1−4 to a spot plate (or plastic sheet). Add 1 drop of iodine reagent to each solution. Record your observations. Use the visual color reference (Table 35.1) to assess the enzyme activity. Place all of the test tubes back into the 37 °C water bath. Rinse the spot plate or plastic sheet. After 5 min, test the samples for starch again and record your results. After 10 min, test the samples for starch again and record your results.

B. Effect of Temperature Materials: Test tubes, test tube rack, test tube holder, droppers, 5- or 10-mL graduated cylinder, beakers for water baths: 250- or 400-mL beaker about half full of ice-water (0 °C − 5 °C), warm (37 °C), and boiling (100 °C) 250- or 400-mL beaker about half full of water heated to boiling, amylase preparation, spot plate (or plastic sheet), 1% starch, iodine reagent Preparing Starch Solution Pour 5 mL of 1% starch solution into each of three test tubes. Place one tube in a boiling water bath, one tube in the 37 °C water bath, and one tube in the ice bath. Allow the test tubes to remain in the water baths for 5 min to allow the solutions to reach the bath temperature. While these test tubes are reaching the experimental temperatures, prepare the amylase solutions. Preparing Amylase Solution Pour 1 mL of amylase solution into each of three other test tubes. Place one tube in a boiling water bath, one tube in the 37 °C water bath, and one tube in the ice bath. Allow the test tubes to remain in the water baths for about 5 min to allow the solutions to reach the bath temperature. Mix Starch and Amylase Solutions Remove the test tubes from the hot water bath and pour the 5 mL of starch into the test tube containing the amylase. Mix by shaking. Return the test tube to the hot water bath. Wait 5 min.

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Remove the test tubes from the warm water bath and pour the 5 mL of starch into the test tube containing the amylase. Mix by shaking. Return the test tube to the warm water bath. Wait 5 min. Remove the test tubes from the cold water bath and pour the 5 mL of starch into the test tube containing the amylase. Mix by shaking. Return the test tube to the cold water bath. Wait 5 min Test tube 1

Test tube 2

Test tube 3

1% starch

5 mL

5 mL

5 mL

temperature

0 °C

37 °C

100 °C

amylase

1 mL

1 mL

1 mL

Testing for Enzyme Activity 1. 2. 3.

Transfer 4 drops of the mixture from the boiling water bath to a spot plate (or plastic sheet). Add 1 drop of iodine. Record the color and assess the enzyme activity. Transfer 4 drops of the mixture from the warm water bath to a spot plate (or plastic sheet). Add 1 drop of iodine. Record the color and assess the enzyme activity. Transfer 4 drops of the mixture from the cold water bath to a spot plate (or plastic sheet). Add 1 drop of iodine. Record the color and assess the enzyme activity.

C. Effect of pH Materials: Test tubes, test tube rack, amylase preparation, 37 °C water bath, buffers (pH 2, 4, 7, 10), spot plate (or plastic sheet), 1% starch, iodine reagent Pour 5 mL of 1% starch solution into each of four test tubes. Place the test tubes in a 37 °C water bath. Wait 5 min. Pour 4 mL of buffer solutions of pH 2, 4, 7, and 10 into four other test tubes. Add 1 mL of amylase solution to each. Place the test tubes into a 37 °C water bath. Wait 5 min. After 5 min, pour each of the 1% starch solutions into each of the pH buffer–amylase tubes. Mix by shaking and return the mixtures to the 37 °C water bath. Wait 5 min.

1% starch Buffer solution amylase

Test tube 1

Test tube 2

Test tube 3

Test tube 4

5 mL

5 mL

5 mL

5 mL

4 mL of pH 2

4 mL of pH 4

4 mL of pH 7

4 mL of pH 10

1 mL

1mL

1 mL

1 mL

Testing for Enzyme Activity Remove 5 drops from each of the pH 2, 4, 7, and 10 test tubes and place in a spot plate or on the plastic sheet. Add 1 drop of iodine reagent. Record your observations and assess the enzyme activity.

D. Inhibition of Enzyme Activity Materials: Test tubes (6), test tube rack, 5- or 10-mL graduated cylinder, amylase preparation, 37 °C water bath, solutions in dropper bottles: 1% AgNO3 , 95% ethanol, 1% starch, iodine reagent, spot plate (or plastic sheet) Place 5 mL of 1% starch into three test tubes. Place 1 mL of amylase into three other test tubes. To test tube 1, add 10 drops of water (reference). To test tube 2, add 10 drops of 1% AgNO3 . To test tube 3, add 10 drops of 95% ethanol. Place all six tubes in a 37 °C water bath. Wait 5 min.

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Laboratory Manual for General, Organic, and Biological Chemistry

Mix Starch, Inhibitors, and Amylase After 5 min, pour each of the 1% starch solutions into each of the inhibitor–amylase tubes. Mix by shaking and return the mixtures to the 37 °C water bath. Wait 5 min. Test tube 1

Test tube 2

Test tube 3

1% starch

5 mL

5 mL

5 mL

inhibitor

10 drops water

10 drops 1% AgNO3

10 drops 95% ethanol

amylase

1 mL

1mL

1 mL

Testing for Enzyme Activity Remove 5 drops from each of the test tubes and place in a spot plate or on a plastic sheet. Add 1 drop of iodine reagent. Record your observations and assess the enzyme activity.

Date

Name

Section

Team

Instructor

Pre-Lab Study Questions

35

1. What is the substrate of the enzyme amylase?

2. What are the products of amylase action?

3. What happens to enzymes at high temperatures?

4. What happens to an enzyme when a competitive inhibitor binds to it?

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Date

Name

Section

Team

Instructor

REPORT SHEET

Enzymes

LAB

35

A. Effect of Enzyme Concentration Time

1. 0 min

Test tube 1

Test tube 2

Test tube 3

Test tube 4

No amylase

3 drops amylase 6 drops amylase 9 drops amylase

color activity

2. 5 min

color activity

3. 10 min color activity Questions and Problems Q1 a. At which enzyme concentration was starch hydrolyzed the fastest? Explain.

b. At which enzyme concentration was starch hydrolyzed the slowest? Explain.

Q2 Describe the effect of enzyme concentration on enzyme activity.

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Laboratory Manual for General, Organic, and Biological Chemistry

B. Effect of Temperature 1. 0 °C color

2. 37 °C

activity

color

3. 100 °C

activity

color

activity

Questions and Problems Q3 What was the optimal temperature for amylase?

Q4 a. Why would the enzyme activity differ in the sample at 0 °C from the sample at 37 °C?

b. Why would the enzyme activity differ in the sample at 37 °C from the sample at 100 °C?

C. Effect of pH pH 2 color

activity

pH 4 color

activity

pH 7 color

activity

pH 10 color

activity

Enzymes

405

Questions and Problems Q5 a. How is amylase activity affected by a low pH? Explain

b. How is amylase activity affected by a high pH? Explain.

Q6 What was the optimum pH for amylase?

Q7 a. During digestion, the pH in the stomach is 2. What does this indicate about the optimum pH of pepsin, an enzyme that hydrolyzes protein in the stomach?

Q7 b. What happens to the activity of pepsin when it enters the small intestine where the pH is 8?

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Laboratory Manual for General, Organic, and Biological Chemistry

D. Inhibition of Enzyme Activity 2. AgNO 3

1. Water color

activity

color

activity

3. Ethanol color

activity

Questions and Problems Q8 In which reaction mixture(s) did hydrolysis of starch occur?

Q9 What substances added to the mixture were inhibitors?

Q10 What are some differences and/or similarities in the type of inhibition caused by heat, acid or base, heavy metal ions, and ethanol on enzyme activity?