Chemical Equilibria in Soils [1 ed.] 0471027049

Table of contents :
Chemical Equilibria in Soils
Contents
Chemical Equilibria in Soils (Introduction)
Methods of Handling Chemical Equilibria
Aluminum
Silica
Aluminosilicate Minerals
Carbonate Equilibria
Calcium
Magnesium
Sodium and Potassium
Iron
Manganese
Phosphates
Zinc
Copper
Chelate Equilibria
Nitrogen
Sulfur
Silver
Cadmium
Lead
Mercury
Molybdenum
Organic Transformations
Selected Standard Free Energies of Formation for Use in Soils Science
Index

Citation preview

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CHEMICAL EQUILIBRIA IN SOILS WILLARD L. LINDSAY Centennial Professor Colorado State University, Fort Collins

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A YVILEY INTERSCIENCE PUBLICATION

JOHN WILEY & SONS, New York • Chichester • Brisbane • Toronto

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Copyright © 1979 by John Wiley & Sons Inc.

All rights reserved. Published simultaneously in Canada.

Reproduction or translation of any part of this work beyond that permitted by Sections 107 or 108 of the 1976 United States Copyright Act without the permission of the copyright owner is unlawful. Requests for permission or further information should be addressed to the Permissions Department John Wiley & Sons Inc.

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Library of Congress Cataloging in

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Publication Data

Lindsay, Willard Lyman 1926-

Chemical equilibria in soils. "A

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Wiley Interscience publication.”

Includes bibliographical references. I . Soil chemistry. 2. Chemical equilibrium. I. Tide. 631.41 S592.5. L55 ISBN 0-471-02704 9

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79-12151

Printed in the United States of America 10 9 8 7 6 5 4 3 2 1

PREFACE This book is the outgrowth of approximately 20 years experience on my part in teaching soil chemistry and directing graduate research in soil science. Its objective is to help bridge the gap between soil science and chemistry and to show that most reactions taking place in soils can be understood and predicted from basic chemical relationships. Emphasis in this text is placed on minerals and solid phases in soils that dissolve and precipitate and , in doing so, control the composition of the soil solution. Solubility relationships are important because they determine the mobility of chemical elements in soils and affect their availability to plants. This book is designed for students who have had at least one year of inorganic chemistry. The text is intended for soil scientists, plant nutritionists, aquatic chemists, geochemists, sanitary and water engineers, environ mentalists, and others who are concerned with the reactions, solubility relationships, and fate of chemical substances in soils. The book is arranged sequentially ; therefore first time readers should start at the beginning and work diligently to understand each new principle as it is introduced. Problems are given at the end of each chapter to assist the reader in understanding the subject matter and in developing the skills necessary to handle solubility relationships with ease. Without such skills, most readers become discouraged before they appreciate what equilibrium relationships are all about. Redox reactions are very important in soils because they modify solubility relationships and affect many mineral transformations. In Chapter 2 the

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PREFACE

rtss fiscasvs'Ssssrarss: ns the H + ions of an overall chemical reaction into those associated the redox component of the reaction from those associated with the acid-base component. Readers are challenged to make greater use of this ft

Sith C

^^

°-

ubffity relationships used throughout this text are based on selected standard free energies of formation that are summarized and documented in the Appendix. The author was assisted in this exhaustive search and selec tion by Dr. Muhammad Sadiq . These values have been adjusted, when necessary, to be internally consistent. Hopefully this compilation of thermo dynamic data will find many uses beyond the immediate scope of this book . As new solubility data become available they can be used to expand and modify the developments given herein. No attempt has been made in this book to apply rigorous kinetic theory to predict chemical reaction rates in soils because the presence of living organisms, catalysts, and unknown constituents in soils are often the controlling factors governing reaction rates. Instead, selected solubility studies in soils have been used to develop practical guidelines for interpreting many of the thermodynamic solubility relationships developed in this book. The solubility relationships that we consider are limited to 25°C. Although temperature changes modify solubility relationships slightly, the added . complexity of including them could not be justified. Undoubtedly some readers will find that their speciality elements are not included in this work. Limits of time and effort have precluded a complete coverage of all elements that are important in soils. Readers are encouraged to apply the

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principles demonstrated herein to their speciality elements. The rewards for doing so are even more satisfying than reviewing what some one else has already done.

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WILLARD L. LINDSAY Fori Collins, Colorado May 1979 .

ACKNOWLEDGMENTS

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I am especially grateful to Dr. Muhammad Sadiq who for nearly two years untiringly assisted with the formidable task of collecting and screening the standard free energies of formation documented in the Appendix . This unifying effort has added greatly to the internal consistency of the solubility data used throughout this book. Appreciation is also expressed to the many graduate students and colleagues, who over the years, through discussion , dialogue, and research have contributed suggestions and insights that are reflected in this work. D. Special appreciation is extended to Drs. Wendell A. Norvell and Ardell of metal ions Halvorson for their contributions in measuring the activities har and Lee Abuzk . A in soils and hydroponic solutions, to Drs. Ahmed ter programs for the Sommers for their assistance in developing the compu Rai for his initial work on chelate equilibrium diagrams, to Dr. Dhanpat for his Contribution on Adamsen aluminosilicate equilibria, to Floyd J . . G. Vlek and Jimmy J . Street for organic reactions, and to Drs. Paul P. L nt. their helpful suggestions and encourageme a for typing the preliminary Nukay Special thanks goes to Lorraine K . maintaining a cheerful and helpful always drafts and final manuscript and for drawings to David W. Fanning for the line ded emen is n ciatio attitude. Appre ulous details of his worL and for taking special pride in the meuc mmed Sadiq Dr Robert D Hal Appreciation is extended to Dr. Muha , and Mr. Randal A . Gaseorfor Mr. Paul A. Schwab, Mr. Lavoir A.gBanks offerin many helpful suggestions. reviewing the manuscript and

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ACKNOWLEDCMri«

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^ Finally, heartfUl thanks goes to my wife, Lorna, and our CHIU Cm dfen cooperation and understanding, which enabled me to for their d evote so , hours to what , at times seemed an endless task. many|0 g

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W. L. L.

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CONTENTS xvii

Symbols

1

Introduction

1

1.1 Dynamic Equilibria in Soils, 1.2 Intensity and Capacity Factors, 1.3 . Elemental Composition of Soils,

2 4 6

8 9

References, Problems,

10

>( 1 Methods of Handling Chemical Equilibria 2.1 Equilibrium Constants, 2.2 Concentration Versus Activity Constants, 2.3 Ionic Strength, 2.4 Activity Coefficients, 2.5 Debye-Huckel Equations, 2.6 Activity Coefficients from Electrical Conductivities, 2.7 Transforming Equilibrium Constants, 2.8 Equilibrium Constants from Thermodynamic Data, 2.9 Redox Relationships, 2.10 pe + pH , 2.11 E° versus log K° , 2.12 pe versus Eh,

11 12 12 13 13 16 18 21 23 23 26 27 ix \

CONTENTS

x

in 2.13 Redox Measurements References,

Natural Environments

Problems,

28 30 31 34

3.1 * 32 3.3 3.4 3.5 3.6 3.7 3.9

Solubility of Aluminum Oxides and Hydroxides, Solubility of Aluminum Sulfates,

Hydrolysis of Al 3 + Fluoride Complexes of Aluminum, Other Aluminum Complexes, Estimating Al3 + Activity, Redox Relationships of Aluminum, . Exchangeable Aluminum, References, Problems,

4 Silica 4.1 Forms of Silica in Soils, 4.2 Silicate Species in Solution,

35 38 39 41 43 45 46 47

48 49 50

51 53 54 55

References, Problems, 5 Aluminumosilicate Minerals

5.1 . 5.2 5.3 5.4 5.5 5.6 5.7 5.8

56

Unsubstituted Aluminosilicates, Sodium Aluminosilicates, Potassium Aluminosilicates Calcium Aluminosilicates, Magnesium Aluminosilicates, Summary Stability Diagrams for Aluminosilicates Controls of Al 3 * Activity in Soils, General Discussion of Aluminosilicates,

57 62 65 66 68 71 73

.

75 76 77

References, Problems, 6 Carbonate Equilibria

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6.1 The COj H 20 System. 6.2 The C02 Soil System , References, Problems,

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78 ‘

79 84

84 84

CONTENTS

*1

7 Calcium

7.1 7.2 7.3 7. 4

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7.5 7.6 7.7 7.8 7.9

Calcium Silicates and Aluminosilicates. Other Calcium Minerals, Complexes of Calcium in Solution. Redox Relationships of Calcium. The CaCOj COj HjO System The Phase Rule, The C02 H 20 System. The CaO CO, H 20 System. The CaO C02 H 20 H 2S04 System, Problems,

— —

—— — — — —

.

8 Magnesium

8.1 8.2 8.3 8.4 8.5

103 105 106 112 113 114 116

Problems,

116 118

Solubility of Sodium Minerals, Solubility of Potassium Minerals, Complexes of Sodium and Potassium, Redox Relationships,

119 123 125 126

Problems,

127 128

10 Iron

10.1 10.2 10.3 10.4 10.5 10.6 10.7

87 93 95 98 98 101 102 102 102

Solubility of Magnesium Silicates, Magnesium Aluminosilicates, Oxides, Hydroxides, Carbonates, and Sulfates, Magnesium Complexes in Solution, Effect of Redox on Magnesium,

9 Sodium and Potassium

9.1 9.2 9.3 9.4

K

Solubility ofFe(III ) Oxides in Soils, Other FeflII) Minerals, Hydrolysis of Fe(III ), Fe(III) Complexes in Soils, Effect of Redox on Fe(II) Solubility, Effect of Redox on the Stability of Iron Minerals, Hydrolysis and Complexes of Fe(II),

References, Problems,

129

133 134

136 139 141 146

148 149

CONTENTS «11

y

iso

u

.

ttlul pH on Mnnguneic of Manganese

111 SoilSpecies

Solubility,

References, Problems,

v

160 162

12 Phosphates

12.1 12.2 12.3 12.4

151 157 160

Orthophosphoric Acid, Aluminum Phosphates,

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Iron Phosphates, Effect of Redox on the Stability of Iron and • Aluminum Phosphates, 12.5 Solubility of Calcium Phosphates, 12.6 Effect of Redox on the Stability of Calcium Phosphates, 12.7 Solubility of Magnesium Phosphates, 12.8 Manganese Phosphates, 12.9 Other Orthophosphates, 12.10 Reduced Forms of Phosphorus, 12.11 Stability of Polyphosphates in Soils, 12.12 Orthophosphate Complexes in Solution , 12.13 Reactions of Phosphate Fertilizer with Soils, References, Problems,

163 169 173 177 180 185 186 187 189 189 190 195 197

204 205

13 Zinc

210

13.1 Oxidation State of Zinc 13.2 Solubility of Zinc Minerals 13.3 Zinc Species in Solution , in Soils,

211 211 216

References, Problems, 14 Copper

u

l

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Solubility of Cu(II)

»,

14.4 Complexes of Cu( I),

References,

Problems.

, ;r c u < > "

Minerals in Soil ,

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219 219 221

222 228 231 234 235 236

i xiii

CONTENTS

15 Chelate Equilibria 15.1 15.2 15.3 15.4 15.5 15.6 15.7

238

Metal Chelates and Their Stability Constants, Development of Stability-pH Diagrams for Chelates, Effect of Redox on Metal Chelate Stability, Chelation in Hydroponics, Use of Chelates to Estimate Metal Jon Activities in Soils, Use of Chelating Agents as Soil Tests, Natural Chelates in Soils,

239 244 252 256 259 261 263

References, Problems,

264 265

16 Nitrogen 16.1 Oxidation States of Nitrogen , 16.2 Equilibrium Between Atmospheric N 2 and 02 , 16.3 Effect of Redox on Nitrogen Stability,

267 268 269 272 279 280

References, Problems,

281

17 Sulfur Effect of Redox on Sulfur Speciation, Dissociation of Sulfur Acids, Formation of Elemental Sulfur in Soils, Formation of Metal Sulfides, 17.5 Effect of Sulfides on Metal Solubilities,

17.1 17.2 17.3 17.4

References, Problems,

^

, Effect of Redox on the Stability of Silver Minerals Solubility of Silver Halides and Sulfides, Stability of Other Silver Minerals, Stability of Silver Halide Complexes, Hydrolysis Species and Other Silver Complexes,

References, Problems,

i

282 287 288 290 295 297 297 299

18 Silver

18.1 18 2 18.3 18.4 18.5

'

300 304 306 308 310 313

313

CONTENT*

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19 Cadmium

3ls

19.1 19.2 19.3 19.4 19.5

316 316 321 322

19.6

Oxidation States of Cadmium in Soils, Cadmium Minerals in Soils, Hydrolysis Species of Cd ( II), Halide and Ammonia Complexes of Cadmium, Other Cadmium Complexes, Need for Further Studies, References, Problems,

323 326 326 327

20 Lead

328

20.1 Solubility of Lead Minerals, 20.2 Hydrolysis Species of Lead, 20.3 Halide Complexes of Lead , 20.4 Other Complexes of Lead,

329 338 339 341

References, Problems,

341 342

21 Mercury

343

21.1 21.2 21.3 21.4 21.5 21.6

Stability of Hg( Il ) Minerals and Complexes, Stability of Hg( I) Minerals and Complexes, Stability of Elemental Mercury, Solubility of Mercury Sulfides in Soils, Summary Redox Diagram for Mercury, Organic Mercury Reactions

344 353 355 358 359 362

References, Problems,

362 362

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22 Molybdenum

22.1 Molybdenum Species in Solution , 22.2 Stability of Molybdenum Minerals in Soils, 22.3 The Effect of Redox on Molybdenum Solubility, References, Problems,

23 Organic Transformations

23.1 Oxidation States of Carbon, 23.2 Products of Glucose Metabolism,

364

365 367 369

372 372 373 374 375

CONTENTS

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23.3 Reactions of Acetic Acid , 23.4 Oxidation to CO g) and Reduction to CH 4(g), 23.5 Stability of Graphite , References, Problems,

^

380 381 382

383 383

Appendix Standard Free Energies of Formation

385

Index

423 '

SYMBOLS, CONSTANTS, AND ABBREVIATIONS

A Avog.

amorp atm

bar [] cone. (c) °C d, deg e est. £

E° EC Eh

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Constant = 0.509 at 25°C for the Debye Huckel equation Activity of species / Avogadro’s number, 6.02252 x 1023 formula units mole' 1 Amorphous solid Pressure in atmosphere, 1,013,250 dynes cm - 2 106 dynes cm 2

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Brackets indicate concentrations in moles liter " i

Concentration Crystalline solid Temperature in degrees Celsius Effective diameter of ions in the Debye-Huckel equation Degrees of temperature in Kelvin or Celsius Base of the natural logarithm Electron

Estimated Electronic charge in the Boltzman equation

Standard electrode potential (volts or millivolts) 1 Electrical conductivity (millimhos cm - ) Electrical potential relative to the standard hydrogen electrode (volts or millivolts) xvii

SYMBOI .S, CONSTANTS. AND ABBREVIATIONS

xviii

Faraday constants

F

ts = 96,487 coulombs equivalen 1

= 23,061 calories volts

"

"

1

equivalent " 1

Activity coefficient of species i Gas phase

) \ (g )

G

AG;

AG; A H? AN ; (1), ( ID, etc.

K

°K

JC

Kc

tc* ( 1)

In log

In x M MF n N ( )

.

"

"

"

‘

Kilo

k

Kf

Free energy ( kcal ) Gibbs standard free energy of formation ( kcal mole " ' ) Gibbs standard free energy of reaction (kcal mole 1 ) Standard heat of formation or enthalpy ( kcal mole ) Standard heat or enthalpy of reaction ( kcal mole 1 ) Immediately following a chemical symbol signifies the oxidation state of that element

•

Boltzman constant Equilibrium formation constant Equilibrium dissociation constant Temperature in degrees Kelvin or absolute Equilibrium constant expressed in terms of activities Equilibrium constant expressed in terms of concentrations Equilibrium constant expressed in terms of concentrations " " + except for H , OH , and e , which are expressed in terms of activities Liquid phase Natural logarithm, base e = 2.71828 . . . Common logarithm, base 10 2.302585 log * 1 Concentration in terms of moles liter Mole fraction in a reaction The number of moles of electrons participating 1 Concentration in terms of equivalents liter Placed behind a chemical compound indicates an aqueous solution species with no charge, for example, H 4Si04° eses around chemical species indicate activities (moles "

"

Parenth

liter 1 ) Negative log of base 10. activity is Negative log of electron activity where electron defined as unity for the standard hydrogen electrode Partial pressure of gas i (atm) given Quotient expressing products divided by reactants of a reaction "1 1 Universal gas constant = 1.98717 cal deg "mole I = 8.3143 joules deg 1 mole" Correlation coefficient "

P

pe

P, Q

"

R r

i

SYMBOLS, CONSTANTS, 5° A S'r

lT I1

STP Z

AND ABBREVIATIONS

xix

Standard entropy of a substance (cal deg - 1 mole - 1 ) Standard entropy of reaction (cal deg - 1 mole - * ) summation Temperature in degrees Kelvin (°K ) (25°C = 298.16*K ) Ionic strength by itself or micro ( 10 - 6) if it precedes another unit of measure Standard temperature (298.16°K ) and pressure ( 1 atmosphere) Valency Electronic charge on surface of a clay. Boltzmann equation.

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CHEMICAL EQUILIBRIA IN SOILS

P

::::

the ocean.” In this respect

Jwettfngt

^^

Many physical, chem**

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drying, freezing, thawing, chang 0 in soils. Physical Pr ° mo< jify the surface areas of soil particles. tempera as ionic species in solution d « energy from the levels. In addition, plants capture . M croorgantsms uulnte u in the form of organic compounds pathways these products through their many biochemical be and system that take place m soil How can this maze of complex reactions atically examined ? chemistry has advanced During the past century our knowledge of . knowledge and its chemical between exists gap great a yet tremendously, . Such per application to soils where so many parameters are unknown investigations empirical to plexities have caused many soil scientists to resort of soils without first utilizing the vast knowledge of chemistry that is available. This book was designed to help bridge the gap between chemistry and soil science and to show how the principles of chemical equilibria can be used to examine many of the chemical reactions that occur in soils.

“ , UScto rt" ^Gerais * eTlSenergy Sr

; rr ^transform

.

. -

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1.1 DYNAMIC EQUILIBRIA IN SOILS

Soils comprise a multiple-phase system consisting of numerous solid phases (about 50 %), a liquid phase (about 25 %), and a gas phase (abouF %X The solids include rock consisting of many different primary and secondary minerals. Superimposed on this inorganic matrix is what Truog (1951) described as the “ living phase,” which includes bacteria, actinomycetes, fungi, algae, protozoa, nematodes, and other forms of life. These living organisms are continually breaking down organic residues and synthesizing many of the products into body tissues while others are released to the ’

rrs

constitucms r t h °

stituents may also be released back into th ions in the soil solution are buffered bv W

?

Soils contain numerous minerals, 2

‘

= “ -

plant con ?IUti°n ( ReaCti°n

MlS

^

^

antltles

L t Sh

*

^^ ^^ ^

some o

are

]

DYNAMIC

EQUILIBRIA IN

SOILS Nutrient Uptake By Plants

’1 l *

Soil Air

Exchangeable Ions Surface Adsorption

MM Organic Matter

Solid Phases

Microorganisms

•ll’ '

Minerals/

Rainfall * Evaporation Drainage Addition of Fertilizer Fig. 1.1 The dynamic equilibria that occur in soils.

amorphous. These minerals impose limits on the chemical composition of the soil solution. If the soil solution becomes supersaturated with respect to any mineral, that mineral can precipitate ( Reaction 5) until equilibrium is attained. Similarly, if the soil solution becomes undersaturated with respect to any mineral present in the soil, that mineral can dissolve until equilibrium is attained ( Reaction 6). Reactions ? and 8 depict several dynamic processes that may occur in soils. For example, rainfall adds water that dilutes the soil solution (Reaction 8). Exoess wafer may drain from the soil profile arid carry with if salts and other dissolved constituents ( Reaction 7). Fertilizers of various kinds are frequently added to soils. These may dissolve ( Reaction 8), form new other ways in soil. reaction products ( Reaction 5), or be distributed in ‘ Organic matter and microorganisms also affect the equilibrium relation ships in soils. Dvmg organisms remove constituents from the roil solution and incorporate them into their body tissues ( Reaction 9). Similarly nutrients are released during the decomposition of organic matter or upon the death of organisms ( Reaction 10). These reactions are connected with abrokerf line to indicate that true equilibrium relationships are generally relationships of not achieved but are modified by the metabolic energy microorganisms that mediate many of these reactions. ‘

,

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CH . I

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4

INTRODUCTION

, n f i |0 attain equilibrium with the soil solution. t Gases in the soil air soil a]r ( Reaction 11 ) or dissolved in reieaw be either may Gases plants and microorganisms generRca ” n aCCCptor and give off CO, from metabolic the soil solution

^ STAS f SSsss sgttnssssp . iii'i: Ei s soitsolution ^

^ ^

H

^

is affected by all o f the reactions depicted in Fig. 1.1, but THte its composition is ultimately controlled by the mineral phases of the soil . Often the rates of dissolution and precipitation of soil minerals are so slow that true equilibrium is not attained ; consequently both kinetic and thermodynamic factors must be considered. Diffusive and convective gradients are both established in soils, and these gradients must also be considered where transport processes are involved. .

.

\2 INTENSITY AND CAPACITY FACTORS

In soils two very important parameters influence the availability of an element to plants. These are ( 1 ) the intensity factor, which is the concentration of an element in the soil solution, and ( 2) the capacity factor, which is the ability of solid phases in soils to replenish that element as it is depleted from solution. As plants remove ions from solution, the concentration of immediate vicinity of roots is reduced and diffusion gradientsthose ions in the are established. The relationship between intensity and capacit y factors as they affect nutrient availability is depicted graphically in Fig. 1.2. The concen

tration

^^

b

hrthe capacity

termed factor, is represented time 0 are the three mineral phases : A B and maintaining a different level of the • same at

^^

'°S

r are eaCT* caPa le nntnV®\tm , mineral A is most soluble and hesoil solution. Initially This level causes supersaturationwill control with raoea "“ *” 1 “u “ C“, which “ slowly precipitate causing “ minera ls B and miner al A V P pleteiy dissolved, the solution dram fr dlSSolve- nce mineral A is corn-

P

' ^ ^°

'

^

'

immediate because of the This dr p of the soil that mustbuffering action nf he°adsorbed ° maynot and exchangeable attain the new equ i brium (eve! corres ponding

10ns

?'

..

^

INTENSITY AND CAPACITY FACTORS

Time

5 Q

Soil

Soil

Solution

ISSS A

h

£

b

“I

f I5

B

•

—

Capacity Factor

Time X Critical Level

[U

c Capacity Factor

ity of nutrients to plants Fig 12 influence of soil minerals on the solubility and availabil ( adapted from Lindsay. 1972 ).

.

soil solution attains equilibrium to level b in the soil solution. Even after the . The nutrient level at b is still with mineral B, chemical reactions do not cease Consequently, mineral C slowly supersaturated with respect to mineral C. ally mineral B disappears and precipitates as mineral B dissolves. Eventu only mineral C remains. t level, plants will show a At time X, when mineral C governs the nutrien n is below the critical deficiency of this nutrient because its level in solutio absorption mechanism at level needed for maximum plant growth. The Organic matter can supply some time X provides very little nutrient reserve. by microorganisms before nutrients, but generally it must be broken down ing effects of adsorbed modify the nutrients are available to plants. The important, but the mineral ions and the activity of microorganisms are solution. ts in phases ultimately control the level of nutrien

CH. 1

6

INTRODUCTION

13 ELEMENTAL COMPOSITION OF SOILS

important Soils differ in total chemical composition. These differences are elements in equilibrium considerations because they help to determine which control the solubility of other elements. in The elemental composition of the lithosphere and of soils is reported Table 1.1. These data are based in the reports of Clarice and Washington (1966), (1924), Swaine (1955), Vinogradov (1959), Jackson (1964), Bowen Mitchell (1964), and Taylor (1964). The selected average is an arbitrary reference level for soils used in this text. The last column in Table 1.1 gives the maximum concentration of each element in the soil solution if all that element at its average reference level were to dissolve in the water present at 10 % of the dry weight of the soil. This parameter is expressed as log M ( moles liter - ) and provides a limiting molar concentration for each element in the soil solution. The molecular ratio of any two elements is useful in relating the stoichio metric combinations of those elements in various minerals. For example, iron may combine with phosphorus to form FeP04 * 2 H 20(strengi.te) in soils. The molecular ratio of iron to phosphorus in an average soil is 10° -83/10 - 0.71 = IQ154 or 34.7-fold. All of the phosphorus may be present as strengite,- but all of the iron cannot, because there is not enough phosphorus to combine with it. The conclusion can be drawn that iron may control phosphorus solubility, but phosphorus can not control iron solubility. Such deductions are useful in selecting and eliminating phases that may govern the solubility of different elements in soils where many complex solubility relationships are found. The average molar concentration of elements in soils given in the last column of Table 1.1 can easily be adjusted to correspond to actual elemental compositions and moisture contents. For example, if a soil treated with sewage sludge contains 500 rather than 50 ppm of zinc, the log of the ratio (500/50) = 1.00 can be added to - 2.12 (Table 1.1) to give - 1.12 M for the maximum concentration of zinc in this soil at 10 % moisture. Furthermore, if the moisture content of this soil were 40 % instead of 10 %, the log of this ratio (10/40) = -0.60 can be added to 1.12 to give 1.72 M for the maximum concentration of zinc in this soil containing 500 ppm of zinc at 40 % moisture. If a suspension of this soil were prepared in the laboratory to contain 1 g of soil (oven-dry basis) per 100 ml of water, the maximum molar concentration of zinc is obtained as follows:

‘

-

—

—

-2.12 + log 500/50 + log 0.1/100 = - 2.12 + 1.00 + ( - 3.00)

log[Zn 2 + ] =

= - 4.12 M

TABLE 1.1 THE CONTENT OF VARIOUS ELEMENTS IN THE LITHOSPHERE AND IN SOILS Selected Average for Soils

Element Ag

A1 As

B Ba Be Br C Ca Cd Cl Co Cr

a

Cu F Fe Ga Ge He 1 K La Li Mg

Mn Mo N Na Ni O

P Pb Rb S Sc

‘

•

Atomic Weight (g)

Content in Lithosphere

Common Range for Soils

(ppm )

( ppm )

ppm

107.87 26.98 74.92 10.81 137.34 9.01 79.91 12.01 40.08 112.40

0.07 81,000 5 10 430 2.8 2.5 950 36,000 02 500 40

0.01 - 5 10,000- 300,000 1 -50 2- 100 100- 3,000 0.1 -40 1 10

0.05 71,000 5 10 430 6 5 20,000 13,700 0.06 100

35.45 58.93 52.00 132.91 63.54 19.00 55.85 . 69.72 72.59 200.59

200 32 70 625 51,000

15 7 0.1 0.3

126.90 39.10 138.91

26.000 18 65

6.94

24.33 54.94 95.94 14.01 2299 58.71 16.00 30.97

207.19 85.47 32.06 44.96

•

21,000 900 23

28.000 100 465,000 ' 1,200 16 280 600 5

-

7.000-500,000 0.01-0.70 20-900 1 -40 1 - 1,000 0.3-25 2-100 10-4,000 7,000-550,000 5-70 1 -50 0.01-0.3 0.1 -40 400-30,000 1 - 5,000

5-200 600-6,000 20-3,000 0.2-5 200-4.000 750-7,500 5 500

-

-

5,000 200 ' 2-200 50-500 30- 10,000 5-50

Molar Cone, at 10 % Moisture log M

8

100 6 30 200 38,000 14 1 0.03 5

8200 30 20 5,000 600 2

1.400

,6,300 40

490;000 600 10 10 700 7

-5.33 .42 J

- 3.18 - 2.03 - 1.50 - 2.18 - 3.20

122 0.53 527 - 1.55 2.87 1.72 3.35 233 0.98 0.83, 270 3.86 5.83 3.40 0.33 - 267 1.54 0.31 0.96

-

-

-

-

-

.

- 3.68 -0.00 0.44 217 249 0.71 3.32

- 293 -0.66 -281

( Continued) 7

.»

CM

INTRODUCTION

t TABl -E t .i

Selected Common Range for Soils Lithosphere ( ppm ) (PPm > Content in

Atomic Weight ( g)

Element

0.09

7896 28.09

Se Si Sn Sr Ti V Y Zn

276.000 40 150 6,000 150

118.69 87.62 47.90 50.94 88.91 65.37 91.22

Zr

80 220

ppm

Average for Soils

Molar Cone, at 10 % Moisture log M

0.3 0.1-2 , 000 230,000-350,000 320 10 2- 200 200 50- 1,000 4,000 , 1 000-10,000 100 20-500 50 250 2550 10-300 300 60-2,000

.

'

- 4.42

-

2.06 3.07 1.64 0.08 1.71 2 5 . 2.12 1.48

-

-^ -

Source: Based on Clarke and Washington (1924), Swaine (1955), Vinogradov (1959), Jackson ( 1964). Bowen (1966), Mitchell ( 1964), and Taylor (1964).

Thus the maximum concentration of zinc in this suspension would be lO 412 M. Data in the last column of Table 1.1 can be adjusted for soils of any elemental content or moisture content. These parameters are useful for predicting maximum solubilities of the various elements in soils and for predicting those minerals that cannot possibly persist in soils. "

REFERENCES

.

—

.

Bowen H. J . M . 1966. Trace Elements in Biochemistry. Academic Press New York' 924 l Th » « « °f earth s crust . U .S. Geo . Sur. * - '

****

*

°f SOih

.

'

F- E

“ « Ed >- Ch'm«wy of

. sou. he

Lindsay W . L. 1972. Influence of the soil matrix on the availablity or Ann. N. Y . Acad. Sci. 199: 37 45. trace elements to plants.

-

^

Comm. No. 48 .

Bear (Ed.), Chemistry of the Soil, 2nd ed.

Herald PrinTng! ' 17 °'

^

C > K>nvw:allh Bur. Soil Sci. Tech. °" "

ments m the continental

—

crust A new table.

—

PROBLEMS

9

Vinogradov. A . P. 1959 . The geochemistry of rare and dispersed chemical elements in soils. 2nd cd . Translation from Russian . Consultants Bureau. New York .

PROBLEMS 1.1 Explain how the equilibria depicted in Fig. 1.1 are affected when : a. A plant root begins to absorb K + . b. Soluble fertilizers are added to soils. c. Soil is subjected to an annual rainfall of 60 inches. d. A heavy straw residue is plowed under. / j. e. A rice field is flooded during most of the growing season. 1.2 Of what value are equilibrium relationships when it is generally recog nized that soils never attain complete equilibrium ? 1.3 How might the equilibrium depicted in Fig. 1.1 be affected by diffusive and convective gradients near plant roots ? 1.4 Why is it not possible for compounds A, B, and C in Fig. 1.2 to coexist permanently in soils ? 1.5 What changes might cause compound C in Fig. 1.2 to become unstable so that compound B could again form ? Give an example. 1.6 Explain how a mineral could maintain different levels of a nutrient in different soils. 1.7 From the average elemental contents of soils given in Table 1.1 discuss the following : a. The possibility that A1 P04 • 2 H 20(variscite) could control the solubility of aluminum in soils. solub. The possibility that PbMo04( wulfenite ) could control the bility of lead in soils. the solubility of c. The possibility that Mn 3( P04)2 could control ,

-

'

S

d

'

containing 600 ppm ThiTparLs'per'milhon of manganese(P in )a,soil can control manganese

of phosphorus below which Mn 3 04 phosphorus solubility but above which this mineral can control

'“

££

mercury in m0 e. The maximum concentration of having the indicated averag could result from shaking 5 g of soil water. mercury content (Table 1.1) with 500 ml of

A

s.

TWO

METHODS OF HANDLING CHEMICAL EQUILIBRIA

substances are mixed, they often undergo chemical \ \ Jhen . Most yV chemical reactions do not go to completion, that is, all ofchanges the reactants do not become products. Equilibrium is reached when the forward reaction

just balances the reverse reactlorrYhramounts of proSucirahTreactams present at equilibrium differ for each chemical reaction. ^

One problem in dealing with chemical systems is how to decide when equilibrium has been attained . There is a simple test. Allowjhe reaction to proceed in the forward direction until nothitigjnor.e happens and measure the composition of the system . Allow the same reaction to proceed in the reverse direction until nothing more happens. If the composition of the system is the same, regardless of the direction from which equilibrium is approached, then true equilibrium exists. Some chemical reactions in soils proceed with sufficient speed that equilibrium relationships arc immediately attained . Other reactions proceed so slowly that final equilibrium is probably never attained. Regardless of the rate at which equilibrium is attained , equilibrium relationships are useful for predicting chemical changes that can and cannot occur . Equilibrium provides a reference point for predicting which chemical reactions can take place regardless of the rate at which they occur.

_

2.1 EQUILIBRIUM CONSTANTS

Equilibrium for the reaction : AB

A+ B

can be expressed by a formation constant ( K j ) AB K B

'-

When Reaction 11 is written in the reverse order, that is A+ B AB

( 2.1 )

(12)

. --

(2.3 )

A B K =U

(14)

Hie equilibrium constant is called a dissociation constant (JC*)

‘

It is apparent that the formation and dissociation constants for any reaction are reciprocals of each other. That is, K

1

~ K,

(15)

»

CHEMICAL EQUILIBRU

METHODS OF HANDLING 12 than the example given above, • c are nre more complex ta the same manner. Consider Many chemical reactions am CH. 2

yet their equilibrium the reaction the

constants

==

K

r

„hrr«ct

to the power

53

(2.6)

* cC + dD aA + bB ? expressed as equilibrium constant is C D‘

= A‘

Bf

in “don yra g cacTtel This . of its coefficient in the

^^

C

reaction

(2.7)

. ^

ud brium expr«rion

12 CONCENTRATION VERSUS ACTIVITY CONSTANTS

So' far equilibrium constants have been defined only in general terms. Actually there are different kinds of equilibrium constants depending upon the units in which reactants and products are expressed. If they are expressed as activities, they define activity constants. If they are expressed as con centrations, they define concentration constants. Each has advantages and

-

disadvantages.

Equilibrium constants expressed in terms of activities have a special significance because they can be calculated from thermodynamic data as will be shown shortly. TheyjirejrueConstantsthat hold for solutions of all ionic strengths. These constantvhowever, have the disadvantage that many reactants and products consist of specific ionic or molecular species whose activities are difficult or impossible to measure. For this reason, concentration constants are often used. Concentration constants have the disadvantage that they change with ionic strength and must either be used in system of the same ionic strength in which they were determined, or ionic strength corrections must be applied. In this book activities are designated by round brackets ( ) and con

] whenever il is necessary ,o dis,in uish c « "

SllTL

23 IONIC STRENGTH

Ionic strength is defined as

-i , ? 2c Z

(2.8)

r DEBVE-HOCKEL EQUATIONS

1.1

,

where /i is the ionic strength, c is the concentration In moles liter" 1 of Ion i Z is the valency or that ion, and I indicates that the product of each Ion and its valency squared w ffied for all ions in solution.

,

^

EXAMPLE CALCULATIONS What is the ionic strength of a 0.01 M NaCl solution ? From Eq. 2.8 H = i[0.01 x l 2 + 0.01 x ( 1)2] \i

-

= 0.01

What is the ionic strength of a 0.01 M CaCl 2 solution ? H

= iCO.Ol

H

= 0.03

-

x 22 + 0.02 x ( 1

V

)2]

2 A ACTIYTTY COEFFICIENTS

Only in infinitely dilute solutions are activities and concentrations equal. The ratio of the activity of an ion, a{ , to its concentration, cf , is called the activity coefficient, >• :

,

* = ^Ci

(2.9)

Knowing yit we are able to convert from concentrations to activities, and vice versa. At infinitely dilute solution a{ = c,, and >•,- = 1. Generally as ionic strength increases, ions of opposite charge interact in such a way that their “ effective ” concentration or activity decreases. 2.5 DEBYE-HOCKEL EQUATIONS

on The Debye-Huckel theory of estimating activity coefficients is based ions assumes that laws of electrostatics and thermodynamics. In essence, it dielectric constant behave like point charges in a continuous medium with a activity calculating equal to that of the solvent. The resulting equation for coefficients of simple ions in aqueous solutions is 2 (2.10) log yt = AZfnu of calculated activity where A = 0.509 for water at 25°C. Comparisons values show a coefficients using Eq. 2.10 with experimentally measured

-

-

CH. 2 METHODS OF HANDLING CHEMICAL EQUILIBRIA

14

ul

aenerally unalte'

^^

tH

^

theory to account for the effective size extending h DebyeBT of hydrated ions, a more precise equation is obtained, that , HQckel

is

log Vi

=

-

AZf

H 112

(2.11)

1 + BdiU 112

where B = 0.328 x 108 for water at 25°C and d{ is the effective distance of closest approach measured in centimeters and corresponds roughly to the effective size of the hydrated ion. Values of for several selected ions as calculated by Kielland ( 1937) are reported in Table 2.1 The values of activity coefficients at various ionic strengths calculated from Eq. 2.11 are tabulated in Table 2.2 The constants A and B in this equation are temperature de pendent, but the dt values do not change appreciably with temperature. In general the extended Debye-Huckel equation holds fairly well in solutions of ionic strength up to 0.2 M . At higher concentrations ionic interactions are difficult to predict and many activity coefficients become larger than unity due to the repulsion of ions. Butler ( 1964) has summarized thiffindings of several workers on the use of various empirical modifica tions of the Debye- Huckel equation to give the best fit method of calculating activity coefficients. For estimating unknown activity coefficie nts, Guntelberg suggested a value of 3 x 10 8 be used for so that Bd in Eq 2.11 is unity, dt giving

.

.

-

-

,

"

log >’(

=

,

.

112

- AZf 1 +H n

(2.12)

112

TOs equation gives values ofy that are too small for many electrolytes. Guggenheim suggested that a linear term be included in Eq. 2.12 and that bn the 6 factor be determined by best fit of the data. After examining the values 0 « the following equation;

'^


CJI * )J (CHOHCOO)f H 2C(COO)J - (CH 21COO 1 (C002) -, H citrate

.

r. -

3'

^^

^

Organic Ions: Charge 3

5

Citrate 3

.

Source : Kielland ( 1937)

n with several adjustable experimental data can be obtained if an equatio n giving: parameters are added to the Guggenheim equatio log 7i

= ~ AZ

/it

,

‘ 1+

/2

fi 112

3 + bn + cfi2 + dp +

•••

-

(2 14)

e The coefficients, b, c, d , ... , must be determined experimentally to provid they which in the best fit approximation of activity coefficients for the system are used.

2.6 ACTIVITY COEFFICIENTS FROM CONDUCTIVITIES

ELECTRICAL

A convenient and direct method of estimating the ionic strength of a solution is to measure its electrical conductivity. Griffin and Jurinak (1973) examined 27 soil extracts and 124 river waters and obtained the following relationship: H

j j* - V ‘ L

-

0.013 EC

/ -L

r

= 0.996

(2.15)

:

ACTIVITY COEFFICIENTS FROM ELECTRICAL CONDUCTIVITIES

17

TABLE 2.2 SINGLE ION ACTIVITY COEFFICIENTS CALCULATED FROM THE EXTENDED DEBYE-HOCKEL EQ., 2.11 FOR 25°C Ionic Strength ( ji )

d ,*

0.001

0.0025

0.005

0.025

0.01

0.03

0.05

0.1

0.874 0.870 0.865 0.860

0.841

0.854 0.848 0.841 0.834 0.826 0.817 0.807

0.826 0.817 0.807 0.796 0.783 0.770 0.754

0.572 0.559 0.546 0.531 0.516

0.517 0.500 0.483 0.464 0.445

0.445 0.424 0.401 0.377 0.351

0299 0.256 0241 0.226

0242 0.178 0.161

0.178 0.128 0.111 0.095

0.135 0.089 0.080

0.098 0.054 0.047

0.063 0.026 0.020

Ionic Charge 1

9 8 7 6 5 4

3

0.967

0.950 0.950 0.949 0.948 0.947 0.946 0.946

0.966 0.966 0.966 0.965 0.965 0.965

0.934 0.933 0.931 0.930 0.928 0.927 0.925

0.914 0.911 0.909

.

0.907 0.904 0.902 . 0.899

0.881 0.877 0.873 0.868 0.863 0.858 0.852

0.854 0.848

•

Ionic Charge 2 8 7 6 5 4

9 6 5 4

0.872 ' 0.871 0.870 0.869 0.867

0.813 0.810 0.808 0.805 0303

^

0.756 0.752 0.748 0.743 0.738

0.690 0.683 0.676 0.668 0.661

0.592 0.581 0.568 0.555 0.541

^

0.737 0.731 0.728 0.726

Ionic Charge 3

0.632 0.619 0.614 0.610

0.540 0.520 0.513 0.505

0.443 0.414

0.404 0.394

0.321 0280 0266 0251

0.194

Ionic Charge 4

11 6 5

0.587 0.572 0.569

0.452 0.426 0.420

0.348 0.312 0.305

0252 0.209 0200

0.151 0.104

0.095

() * Effective ionic parameter of ion / .

rneasuringTht” mTny'dWer . ionic specte included b, .he Utter workers. " Tofp n uo rnThTu

^

0

CM.

IK




M i?

•

J j 5? t.j

:

\ >

•

I

-

TABLE 5.1 EQUILIBRIUM REACTIONS USED TO CONSTRUCT THE STABILITY DIAGRAMS FOR ALU MI NO SILICATES Reaction No.

log K °

Chemical Reaction

-

Al Si + i AI 2Sips (sillimunitc ) + 6 H + ^± 2 AIJ 2 Ali + AI 2 SiOs ( kyanite ) + 6 H + + 2 A 13 + AI 2 SiOj( andalusitc ) + 6 H AI 2 Si 2 Os ( OH )i( halloysitc ) + 6 H + ± 2 Ali + ^ AI 2Si 2 Oj( OH )4(dickite) + 6 H + ± 2 AIJ + ^ AI 2 Si 2 Os (OH )4( kuolinite ) + 6 H + ± 2 AIJ + AI 2 Si 4 O 1 0(OH )2 ( pyrophyllile ) + 6 H + + 4 H 20 ^ 2 AI 3 +

I 2 3

^ ^

4 5 6

7

^

+ H 4 SiOS + H 20 + H 4 SiOS + H 20 + H 4SiO £ + H 20 + 2 H 4 SiOS + H 20 + 2 H 4SiO$ + H 20 + 2 H 4SiOJ + H 20 + 4 H 4SiO?

15.45 15.12 14.48 8.72 5.95

5.45

- 1.92

NaAI -Si 8 9

10 tl 12 13 14 VS

^

NaAISi 04( nephcline ) + 4 H * Na + + Al 3 + + H 4Si 04 NaAISi 206 ( jadcite ) + 4 H + + 2 H 20 Na + + Al 3 + + 2 H 4Si 04 NaAISi , Oh - H ,0( analcime ) + 4 H + + H 2 O Na + + Al 3 + + 2 H 4 Si 04 NaAISijOH ( Na -glass ) + 4 H + + 4 H 20 ± Na + + Al 3 + + 3 H 4 Si 04 NaAlSi 20 B ( high albiie ) + 4 H + + 4 H 2 O ^ Na + + Ali + + 3 H 4 SiO? NaAlSij08( low albitc ) + 4 H + + 4 H 2 O Na * + Al 3 + + 3 HjSiOJ NaAliSi ^OtoCOH paragonile ) + IOH + ^ Na + + 3 AI 3 + + 3 H 4SiO? ao.ijA\ 2 JiS\ 1.b 2Ol 0(OH ) 2(bcideUite ) + 7.32 H + + 2.68 H 20 0.33 Na + -f 2.33 AlJ + + 3.67 H4 S/ 04

^

^ ^ ^

11.25

^ ^

3.67 2.74

^

7.11

8.15 10.87 17.40 6.13

I

!

-

KAI Si

10 17 18 19

20 2!

3*

K * + Al ^ ' + AI ' , 211 K 0 + , ^ ; ± K 4 H + Al ' + 20

KAISiOa ( kaliophilitc ) + 4 H '

+ H 4SiOS 3 KAISi 2 Oh ( lcucilc ) + 411 ' + 2 H 4SiO$ 4 KAISijCMK -glass ) + 411 ' + 3 H 4Si 04 KAlSijO ( high sanidine ) + 4 H 4 + 4 H 20 ?± K 4 + Al ' 4 + 3 H 4 Si 04 KAlSi 0 ( microclinc ) + 411 * + 4 H 2 O K 4 + Al 34 + 3 H 4SiOJ KAI 2 ( AISijO 10 )( OII ) 2( muscovitc ) + 10 H ' K ' + 3 AI ’ 4 + 3 H 4SiOJ

„ , „

13.05 6.72 *

•

^

^

7.87 1.40 1.00

13.44

-

CaAI Si 22 23 24 25 26 27 28

CnAI 2 SiOfl( pyroxcne ) -f 8 H 4 ± Ca 2 > + 2 AI" + H 4Si 04 + 2 H 20 ^ CaAI 2Si 208(Cn glass ) + 8 M 4 Ca 2 + + 2 Al 3 * + 2 H 4SiOJ CaAl 2 Si 20 R ( hcxagonal anorthitc) + 8 M 4 ± Ca 2 4 + 2 Al 3 * + 2 H 4 Si 04 ^ CaAl 2Si 20 H ( unorthile ) + 8 H 4 ± Ca 2 4 + 2 A 13 + + 2 H 4 Si 04 ^ CaAI 2Si 208 2 H 20( la \vsonitc ) + 8 H * ± CaI + + 2 Al 3 * + 2 H 4SiOi + 2 H 20 ^ CaAI 2 Si 40 l 2 2 U 20( wairakitc ) + 8 H 4 + 2 H 2 O CaJf + 2 Al 3 * + 4 H 4 SiOS 4 Ca 2 AI 4 SiR 024 7 H 20( lconharclite) + 16 H + H 20 2 Cn 2 + + 4 A! + + 8 H 4SiO'i

-

-

-

^

-

,

^ ^

35.25 33.91 26.10 23.33 17.54 16.05 17.29

MgAI -Si

29

30 31

5 M g 2 * + 2 A I 4 + 3 H 4SiOJ + 6 H 2 0 Mg 5 AI 2 Si ,Oj „( OH )8 (chloritc ) + I 6 H 4 4 Mg 2 AI 4Si ,OIR ( Mg -corilicritc) + I 6 H + 2 H 2 0 2 M g 2 * + 4 Al .4 * + 5 H 4 Si 04 4 4 4 KMg , AISi ,0|0 F 2 ( nuorplilogopitc ) + 8 H + 2 H 2 0 ^± K + 3 M g 2 * + A l ' + 3 M 4SiOS + 2 F

^ ^

’

60.30 45.46 7.85 { Continual )

'C

.

g TABLE

5.1 ( Continued )

Reaction No.

Chemical Reaction

log Km

-

Substituted Al Si 32

(

H )o.O 2 ^( 111 )0.46 - ^ 0.061 0.1 )Si 2 . >J I Al , . , 4Ol 0( OH ) J ( vermiculite ) + 10.36 H +

^

^15.2.

71

.

^

^

.

*

2+

+

+ 0.02 Fe ^ 2.71+ Mgi 0.1 K + 1.14 Al +

1+

+ 0.46 Fc3 * + 0.06 Ca 2 * + 2.91 H 4SiO* + 0.36 H 2 O

38.14 10.35

Ko^ Mgo.ijAl ^Sij. jO oCOH illite) + 8 H + + 2 H 2 O ^± 0.6 K + + 0.25 Mg 2 + + 2.3 AI 2 * + 3.5 H 4SiOJ Mgo. 2(Si 3. , Al . 7 lFe( III )0 22 Mg0 29 )O , 0( OH ) 2 3+ J+ 2+ ( Mg-montmorillonite) + 6.76 H \ + 3.24 l-l 20 ^ 0.49 Mg + 1.71 A 1 -1- 0.22 Fc + 3.81 H 4SiC>:

33 34

|

^

2.68

Reacting Components

35 36 37

38 39 40 41 42 43 44

_

Fe( OH ) j(soil ) + 3 H + ± FeJt + 3 H 20 FcJ + + c ~ ^ Fc 2 + a - FcOOM ( gcothitc ) + 3 H + Fc ? + + 2 H 20 + AI( OH )3(amorp) + 3 H A\i + + 3 H 20 AI ( OH )3( gibbsile ) + 3 H + Al1 + + 3 H 20 Si 02( amorp ) + 2 Hj0 tH 4Si0J ^ Si02(soil ) + 2 H 2 O H 4 SiC> Si 02(quartz) + 2 H 2 O H 4SiO$ CaF 2 ( fluoritc ) Ca 2 + + 2 F ~ Ca 2 + (soiI ) ;j± Ca 2 +

^ ^ ^ ^ ^ ^

:

‘

2.70 13.04 0.02

-

9.66 8.04 - 2.74 - 3.10 - 4.00

- 10.41

—

2.50

A

I

CH . 5 UNSUBSTITUTED ALUMINOSILICATES

61

By selecting different values for log H 4SiO£ , Eq. 5.2 can be solved for log AI 3 + + 3 pH to give :

log Al 3 +

log H 4Si 04

-2.5 -4.0 -5.5

+ 3 pH

5.23 6.73 8.23

These relationships are plotted in Fig. 5.1 and represent the solubility line for kaolinite labeled 6, which refers to Reaction 6 of Table 5.1. The other solubility lines in this figure were developed similarly. 12

11

Y

10

AI(OH ) 3 ( amorp)

9

£ r>

AI (OH )3 (gibbsite ) f

8

< O)

7 (1) ( 2)

6-

(3 )

(4 ) ( 5)

5

4

-

6

AI 2Si 205 tQHU AI 2Si 205 (0 H )4

AI2Si 2 p5(OH )4 (7) AI 2Si 4OiQ ( OH )2

(6)

V,

AI 2SI05 AI2SO5 AI2Si05

N

=s

§

s-

4

log H 4 SiQ4

8

«/>

i

S \ 55 \

*>

“ i

Fig 5.1 The solubility of various unsubstituted aluminosilicates compared to gibbsite.

.

>» CH . 5 62

_.

ALUMINOSILICATE MINERALS

are most stable and the to see which minerals easv F'om stabilities- Solubilities of the AI 2Si 20, minerals their that t ! condition andulusite. Solubilities of the der sillimanite > kyanite > ,

.

it is

level, kaolinite is more stable. line at The solubility of Al(OH )3(gibbsite) is shown by the that of Al(OH )3(amorp) lies at 9.66. A t log Al 3 + + 3 pH = 8.04, while 5,31 M gibbsite is the most stable mineral H4SiOt activities below 10 represented in Fig. 5.1. The solubilities relationships of Si02(quartz), | Si02(soil), and Si02(amorp) are included to show the leVels of H 4SiO 40 3.10 , solid phases these , by nambly , imposed 10 10 and that are 10 2 - 74 M , respectively. Since primary minerals generally have a higher Si/Al ratio than secondary minerals, silica is released during weathering and tends to maintain H 4Si04 somewhere within these limits. In soils developed under more intense weathering environments, silica is leached from soils as first kaolinite and eventually gibbsite become the stable mineral phases. Gibbsite, along with all the other aluminum oxide and hydroxide minerals discussed in Chapter 3, is unstable in soils in which H SiOS exceeds 10 5-31 M . The activity of Al 3 + in equilibrium with the various aluminosilicate minerals shown in m Fig. 5.1 can -fee readily obtained from the diagram when pH is known. For example, equilibrium with gibbsite is given as log Al 3 + + 3 pH = 8.04 which at pH 5 becomes

horizontal

"

"

"

.

"

-

log Al 3 +

+ 3(5) = 8.04

or log Al3 +

= 8.04

—

15 =

- 6.96

which corresponds to an Al3 + activity of 10 6-96 M. Th us for any given pH, the vertical scale can be converted into Al 3 + activity. Using combined functions as axes in these plots avoids the n ecessity of separate plot for each pH. having to prepare a 5.2

"

SODIUM ALUMINOSILICATES

The solubilities of several thru 15 of Table 5.1 and sodium aluminosilicates are given by Reactions 8 are plotted in Fig. 5.2. An example of how the

63

SODIUM ALUMINOSILICATES 20 log Na*

pH 6

— — -4

7

18

%

8

%

16

-- -3

——

-2

a

14

a

ifeEGsi

/oe (, 8) a

a.

m

«

< 10

a

a

8* 8

6

/

^^

4:

55®

4

PH

6*

2

(5.4)

CH . 5

64

ALUMINOSILICATE

MlNERALs

In most well -drained soils, a realistic level of is approximately lo ~ 3 (see Chapter 9). Substituting this reference level of Na * into Eq . 5.4 giVes

Na +

log Al 3 + + 3 pH = 5.74

^

- pH - 3 log H4SiO

^

(5.5)

A plot of Eq . 5.5 at pH 7 is shown in Fig. 5.2 and is labeled low aibite (13) corresponding to Reaction 13 of Table 5.1. The other lines in this diagram were determined similarly using appropriate reactions from Table 5 l Although Fig. 5.2 is drawn for pH 7 and 10 3 M Na + , it can be used to predict solubility relationships at other pH and Na + values. The array of short lines and arrows in the. upper left of this diagram indicates that ail lines marked a shift up or down by one log unit with a one unit change in either pH or log Na + . The beidellite and paragonite lines shift less than one unit for corre sponding changes in pH and Na + . The pH changes are indicated directly on the diagram. A unit change in log Na + shifts the paragonite line by 0.33 unit and the beidellite line by only 0.14 unit These changes reflect the Na/AI ratios in these minerals. Figure 5.2 shows the relative stability of the sodium aluminosilicate minerals. The most soluble mineral is Na-glass. Of the three NaAlSi308 minerals, the solubilities decrease in the order of Na-glass > high aibite > low aibite. Sodium-glass is expected to weather most rapidly in soils. The minerals NaAlSi04( nepheline), NaAISi 206( jadeite), and NaAISi 2 Ofi H 20 (analcime) are intermediate in solubility and are expected to weather next most rapidly. The most stable sodium aluminosilicates are ( i ' "

^

NaAI 3Si 30 j 0(OH )2( paragonite) and

Nao.33Al 2.33Si3.6 70 io(OH )2( beidellite).

\

In neutral soils where H 4Si02 drops below 10 “ 3- 8 M , beidellite and paragonite become less stable than gibbsite. All of the sodium aluminosilicate minerals depicted in Fig. 5.2 are unstable in soils and will eventually disappear with weathering because kaolinite and pyrophyllite are more stable, ( Fig. .1). From tfjese relationships it |S easy to see why Na + is released in the weathering of primary minerals and ends lip m the oceans, because Na + is not strongly refixed into secondary minerals. In poorly drained soils, Na + often accumulates as the major soluble cation.

£

i

I

POTASSIUM

-

53

ALUMINOSILICATES

65

POTASSIUM ALUMINOSILICATES

The solubilities of several potassium aluminosilicate minerals are given by Reactions 16 through 21 of Table 5.1 and are plotted in Fig. 5.3. An example of how these solubility lines were developed is shown for muscovite beginning with Reaction 21 of Table 5.1:

—

KAl 2( AlSi 30 , 0 )( OH )2( muscovite) + 10 H 4

K + + 3 A134 + 3 H 4SiO; ( 5.6) log K °

= 13.44

( K 4 ) ( Al 34 )3( H 4 SiO?)3 ( H + ) ( H 4 )9

3 log Al 34 + 9 pH log Al 3 + + 3 pH

= 1013- 44

= 13.44 - log K + - pH - 3 log H 4SiOS

= 4.48 - 0.33 log K + - 0.33 pH - log H 4 SiOS

(5.7)

When K 4 is 10 “ 3 M, Eq. 5.7, becomes

log Al 3 + + 3 pH

= 5.48 - 0.33 pH - log H 4SiOS

( 5.8)

A plot of Eq. 5.8 corresponding to pH 7 is given in Fig. 5.3 and is labeled the muscovite line. Substitution of pH 6 or 8 instead of 7 into Eq. 5.8 shifts the muscovite line upward or downward by 0.33 of a unit (log Al 34 + 3 pH ). The solubility lines for the other mineral in Fig. 5.3 were obtained similarly using appropriate equations from Table 5.1. Minerals ( 16) through ( 20) in Fig. 5.3 all shift upward or downward by one unit for each unit change in pH or log K 4 as shown by the short lines and arrows in the upper left part of this diagram. Figure 5.3 shows that the solubilities of potassium aluminosilicates decrease in the order of KAlSi 308( K -glass) > KAlSi 04(kaliophilite) > KAlSi 206(leucite) > KAlSi 308(high sanidine) > KAlSi 308( microcline ) > KAl 2(AlSi30 I 0)( OH )2( muscovite ). Of these minerals K glass is expected to weather most rapidly. For a soil of pH 7, muscovite becomes more stable 48 M . The pH dependency of than gibbsite when H 4 Si02 exceeds KT to shift slightly. Muscovite activity Si muscovite causes this equilibrium H 4 04 |range expected in soils. is more stable than gibbsite over most of the H 4SiO weathering , silica has been largely when soil Only in the later stages of removed, is gibbsite stable. It will be shown later (Fig. 5.5) that muscovite is generally metastable to kaolinite.

-

^

«»»

,, _

ALUMINOSILICATE MINERALS

CH . 5 66

20

pH

6

7

18

4V 16

a"

14

8

%

log K * -4

sh -3

—*

2

%

SiJaofi,/;

£

_ *

*? 12

a'


3 . Despite this fact, it is convenient to treat all dissolved C02 as H ,CO|. For most purposes, there is no problem in doing so because the hydration status of dissolved species need not be stated in thermodynamic consideration. The equilibrium reactions relating the various carbonate species are summarized in Table 6.1 . The solubility of C02 in water at 25 °C can be represented by C02(g) + H ,Q

n2COl

log K ° = - 1.46

(6.1)

TABLE 6.1 EQUILIBRIUM REACTIONS IN THE CO, H ,0 SYSTEM AT 25°C

—

Reaction No. 1 2 3 4

5

Equilibrium Reaction C02(g ) + H 20 H ,C0$ H ,CO| H ’ + HCOj HCO 3- H + + CO5 co2(g) + H 2O ^± H + ++ HCO; ?C02(g) + H ,0 ^± 2 H + CO!

^ ^ ^

log K °

- 1.46

-6.36 - 10.33 - 7.82 - 18.15 79

CIL 6

CARBONATE EQUILIBRIA

90

giving

H 2CO°3

--

10-

, co (g) ‘ 4

(6.2)

2

(6.3) 1.46 + log C02( g) liter and C02(g) is expressed in per moles as expressed where H 2CO$ is atmospheres. to give In solution carbonic acid dissociates (6.4) log K ° = 6.36 H + + HC03H 2CO$ log H 2COJ

31

—

Thus

_

( H + XHCOJ )

H 2CO$ log

( HCOJ )

H 2CO$

-

(6.5)

JQ 6.36

(6.6)

= pH - 6.36

) is unity. This ratio | At pH 6.36, the molar ratio of ( HCOJ ) to ( H 2CO increases 10-fold for each unit increase in pH and decreases 10-fold for each unit decrease in pH. Reactions 6.1 and 6.4 can be combined to give logK °

CO 2(g) + H 2O H 2CO 3

=

;

C02(g) + H 20

‘

H 2 COS H + + HCOJ

- 1.46 - 6.36 }

H+ +

- 7.82

HCOJ

(6.7)

Thus ( HCOJ )

log HCOJ

The bicarbonate ion also

HCOj Thus

^

7 82 CQ

=l l

2 (g)

(H )

(6.8)

= -7.82 + pH + log C02(g)

(6.9)

+

dissociates to give: ± H * + C023 log K ° = - 10.33

=

;

( H + )(C02 ~ )

THCOJP = 10 ~ ‘ ( HCOj )

0.33

= PH - 10.33

(6.10)

(6.11) (6.12)

I

THE CO, H,O SYSTEM

81

At pH 10.33 the molar ratio of (COf ") to (HCOJ ) is unity. Each unit increase in pH increases this ratio by 10-fold, and each unit decrease in pH decreases it by 10-fold. Combining Reactions 6.7 and 6.10 gives log K °

co2(g) + H 2O

I

-

HCO 3

H + -f HCOJ H + + CO|~

CO2(g) + H 2O

2 H + + CO|

"

-7.82

-10.33 -18.15

(6.13)

Thus

con =

(

log COf

"

10 ~ 18 > 5CQ2(g) (H + )2

(6.14)

(6.15)

= -18.15 + 2 pH + log C02(g)

Convenient relationships that follow from Eq. 6.3, 6.9, and 6.15 are given in Table 6.2 The mole fraction distribution of the various carbonate species in solution is shown in Fig. 6.1 where activity coefficients are taken as unity. Total carbonates in solution consist of [H 2COf + HCOJ + COf ] so the mole fraction ( MF ) for HCOJ is

HCO3(6.16) H 2 COf + HCO 3- + COf Using appropriate expressions from Eq. 62, 6.8, and 6.14 for each of the terms in Eq. 6.16 gives 0 - 7.82/( H + ) _ (6.17) , , 782 ( + ) M F HOOj / H + 10 - 8 5/(H + )2 10 - 1 >6 + io

M FHCOJ

“

,

—

ATE SPECIES IN SOLUTION TABLE 62 DISTRIBUTION OF CARBON AS A FUNCTION OF C02(g) log HC03 , M ,M C02(g), atm log COjf&l atm log HjCO? 0.0003 0.003 0J01 0.1 1.0

— 3.52

- 252

- 200

-1.00 0.00

.O

-4.98 - 3.98 - 3.46 -246

- 1.46

pH - 1134 pH - 1034 pH - 9.82 pH - 8.82 pH - 7.82

log COf 2pH - 21.67 2pH - 20.67 2pH - 20.15 2pH - 19.15 2pH - 18.15

J

,,

CARBONATE EQWu

82

j'

08 U

5

0.6

o O

0.4

5

£ *»

0.2

I

Rg. 6.1 The effect of pH on the distribution of carbonate species in solution.

A plot of Eq. 6.17 is shown as the HCOJ curve in Fig. 6.1. The curves for H 2 COS and COj ~ were obtained similarly using appropriate mole fraction expressions. Since the COz (g) term is common to all carbonate species, it can be canceled from the mole fraction expressions ( Eq. 6.17). Thus the relationships shown in Fig. 6.1 are independent of the partial pressure of CO ?(g) and, therefore, tell nothing about th Lptal amount of carbonate that may be present in solution. The plot is useful in showing the important effect of pH on the distribution of carbonate species in solution. The intersection of curves at pH 6.36 and 10.33 corresponds to the pK ° values for the first and second dissociation constants of carbonic acid ( Eq. 6.4 and 6.10). Another very useful plot showing the actual activities of the various carbonate species in solution as a function of pH in equilibrium with 0.0003 atm C02(g) is given in Fig. 6.2. These relationships were plotted from Eq. 6.3 , 6.9, and 6.15. In this plot H 2 CO £ is present at 10 - 4.98 M and is independent of pH. The activity of HCOJ increases 10 fold for each unit increase in pH, whereas that of COj “ increases 100-fold. This figure shows the activities of all the carbonate species expected in soils when equilibrium is attained \»( Uh atmospheric CO,(g). The pH dependence of total carbonate solubility reacfily seen. Again the intersections of the various'lines in this figure

^

.

-

!

occur

I

—

THE COj HjO SYSTEM

83

0

-1

c

°2 (g ) = 0.0003 atm

-2 -3

*

l

o’

-

&

A

J

H 2CQ3

5

'

-6 -7

87

-8 -9

-10 4

X

X

5

6

9

8

7

10

X

J

11

12

pH

um with 0.0003 atm

s of carbonate species in equilibri Fig . 6.2 The effect of pH on the activitie of C02(g ).

the pK ° for the add dissociation reactions at pH values corresponding to iation of HCOJ to give equimolar CO| involved. For example, the dissoccorresp onding to the pK ° of 10.33 for this and HCOJ occurs at pH 10.33 reaction (Eq. 6.12). Fig. 6.2 are given for a C02(g) partial pressure The relationships depicted in be used to obtain the carbonate activities of 0.0003 atm, but the graph can (g) by 10 fold shifts all the lines in this C the 02 ing Increas . level for any COz ) s 10 fold decreases in C02(g shifts figure upward by one log unit wherea e, let us say that C02(g) increases exampl them down by one log unit. For on this graph then shift upward by from 0.0003 to 0.1 atm. The lines at pH 8.0 the HCOJ level of 10 - 3 34 M Q log (0.1/0.0003) = 2.52 log units".3Thus 34 + 2 52 or 10" 82 M . iO to shifted at pH 8 would be “

-

-

-

-

y CH . 6 CARBONATE

84

EQUILIBRIA

6.2 THE COj -SOlL SYSTEM

the atmosphere, there is continual opportunity f0r soils the partial pressure of COz (g) is expected &sorgain ofCO g). In most because C02(g) is continually being to be slightly higher than that in the ait released by the respiration of roots and other organisms in soils. In this book a level of 0.003 atm of C02(g) or approximately 10 times the atmospheric level will frequently be used as a reference level for soils. In flooded soils, the C02(g) level generally goes much higher because diffusion of gases through water is much slower than air. The C02(g) levels in flooded soils often range from 0.01 to 0.3 atm. The solubility relationships of carbonates in soils are expected to cor respond to those depicted in Fig. 6.2 with the modification that C02(g) may vary and cause the solubility lines to move up or down accordingly. The activity of CO in acid soils is quite low ( Fig. 6.2), so very few metal carbonates can form. However, in alkaline soils C02 “ reaches levels in which many metal carbonates can form. Consequently, many carbonate minerals become important in alkaline soils and impose limits on the solubilities of many metal ions. These solubility limits will be considered in the separate chapters that follow for the trace elements. The equilibrium relationships developed herein will be used to develop the solubility limits and relate them to the partial pressure of C02 (g).

^

fo

-

^

REFERENCES

, ^S° Uti°

L NcwYork pMp 74-92 ’

'

'

’

nS ’

Mincrals

Equilibria. Harper and Row .


3 Stumm, W. and J . J . Morgan, 1970. Aqun ic Chemistry. Wiley.|merscien«, New York. pp. 118- 160 .

PROBLEMS 6.1

ortonate' ta s

^

XnpSem '

^ ° Coffin ^^

*^ Vr \

co changes in the partial ?pressure off Cn < fuactlon of PH HoW W1 C 2(g) affect this mole fract on distribution ? Explain.

6.2

Calcul te the H 2COl and increasing C02(g) from 0.0003 o 0

,

£

PH

, “

*

,

at which the P

“

*

I

PROBLEMS

*

ft

6.3 Calculate the pH of distilled water in equilibrium with a. Air containing 0.0003 atm C02(g). b. Pure C02(g) at I atm. 6.4 A liter of distilled water has been equilibrated with air containing 0.0003 atm C02(g). This solution is placed in a closed system where no exchange of C02(g) is permitted and slowly titrated with NaOH. Calculate the milliequivalents of base required to reach a. pH 6.36 b. pH 8.00 c. pH 10.33. 6.5 A liter of distilled water is equilibrated with 0.003 atm of C02(gX Assuming activity coefficients are unity, calculate a hypothetical titration curve: a. For a closed system where no additional C02(g) is admitted. b. For an open system where 0.003 atm of C02(g) is maintained and the titration is continued to pH 11.0.

t. I

%

SEVEN

*

CALCIUM

-j

r comprises approximately 3.6 % of the lithosphere while the content of soils is near 1.37 % ( Table l . l ). Calcium is somewhat

* ^.suable m soils and its content is largely influenced by vei ape

,

\

parent material and

lAtnfall Soils that develop from calcareous parent materials often have calcium carbonate somewhere in the profile. With advanced weathering and high rainfall, calcium carbonate, and most other calcium minerals disappear from soils. In this chapter the stabilities of various calcium minerals and solution complexes arc examined . Reference solubilities of Ca 2 + for both calcareous and noncalcarcous soils arc proposed . The solubility relationships of calcium phosphates, sulfides, and molybdates are examined in Chapters 12, 17, and 22, respectively. Since the anions of these minerals are generally less abundant 2+ in soils than calcium, these minerals are not expected to control Ca solubility. 7.1 CALCIUM SILICATES AND ALUMINOSILICATES

The solubilities of several calcium silicates and aluminosilicates are given by Reactions 1 through 12 of Table 7.1 and are plotted in Fig. 7.1. The 3 reference levels of Al * and H 4 SiO£ used to develop this diagram are the equilibrium levels with AI 2Si 205(OH )4( kaolinite) and Si02(quartz). te can also coexist with kaolinite and Furthermore, Mg - montmorilloni + 2 is near 10 3 M , and Fe 3 + is controlled by quartz if pH is near 7.8, Mg soil- Fe ( Fig. 5.6). For these reasons the kaolinite quartz equilibrium was selected as the logical weathering environment for examining the stability of calcium minerals in soils. An example of how the solubility lines in Fig 7.1 were developed is given for wairakitc, using the following equilibrium reactions: log K ° CaAI 2Si4012 2 H 20( wair) + 16.05 Ca 2 + + 2 A13 + + 4 H 4SiOS + 8 H + 2 H 20 * Al 2 Si 205(OH )4 ( kaol ) + 6 H - 5.45 2 AI 3 + + 2 H 4 Si 04 + HjO 2(4.00) 2 Si02(quartz) + 4 H 20 2 H 4SiOS ; “

-

-

Ca 2 + + AI 2 Si ,Os ( OH )4( kaol ) 18.60 + 2 Si02(q ) + H ,0 ( 7.1 )

CaAUSi 40 , 2 2 H 20( wair )+ + 2H

-

( Ca 2 + ) (H + )

log Ca 2 +

b

-

IQlS.b O

18.60

- 2 pH

(7.2) 87

S

TABLE 7.1 EQUILIBRIUM , REACTIONS OF CALCIUM AT 25°C i

Reaction No.

log X °

Equilibrium Reuction

Silicates 1

2 3 4

/J-CaSiOj( wolIastonite) + 2 H + + H 20 =^ CaJ + + H 4SiOJ CaSiOj( pseudowolIastonite ) + 2 H + + H 2 O Ca 2 + + H 4 SiOJ ^ /?-Ca 2 Si 04(larnitc) + 4 H + ^ 2 Ca 2 ++ + H 4SiOJ y-Cu 2Si04(Cu olivine) + 4 H * 2 CaJ + H 4SiOJ ^

13.27 14.23 39.62 37.82

Aluminosilicates 5 6

7

/

I

10 11

12

CaAI 2 Si06(pyroxene) + 8 H + CaI + + 2 AI 3 + + H 4Si02 + 2 H 20 CaAi 2 Si 208(Ca glass) + 8 H + ± CaJ + + 2 AIJ + + 2 H 4SiOS ^ CaAI 2Si 208( hexagonal anorthitc) + 8 H + iCa 2 + + 2 A13 + + 2 H 4Si 04 ^ CaAl 2Si 20 (anorthite) + 8 H + ± CaJ + + 2 AI 3 + + 2 H 4SiOJ CaAI 2 Si 208 2 HjO( lawsonite) + 8 H + ^ Ca 2 + H 2 A13 + + 2 H 4SiOJ + 2 H 20 CaAi 2Si40 2 2 H 20( wairakite) + 8 H + + 2 H 2 O CaJ + + 2 A13 + + 4 H 4SiOJ Ca 2 AI 4Si 024 7 H 20(leonhurdite) + 16 H + + H 20 ± 2 Ca 2 + + 4 Al 3 + + 8 H Si 4 04 CaMg(Si03)2 (diopside) + 4 H + + 2 H 2 O ^ Ca 2 + + MgJ + 2 H SiO + 4 £

^

-

-

,-

„

^ ^ ^

„ -

-

_

3 5 25 33.91 26.10

23.33 17.54 16.05

1729 21.16

Carbonates 13 14 15 16

CaCOj(calcite ) + CaC03(aragonitc) + CaC02 6 H 20( ikaite) +

2 H + Ca 2 + + 2 H + Cu 3 + + 2 H* M * + CaMg(COj) 2(d0|omite) + 4 H + Cu 2 * +

-

^

^ ' ^

C02(g) + H ,0 CO,(g) + H ,0 C02(g) + 7 H ,0 Mgi + + 2C02( g ) + 2 H,0

9.74 9.97 11.78 18.46

1

1

17 18 19 20

Soil, Oxides, Hydroxides, Ferrites Soil Ca Ca 24 / CaO( limc) + 2 H 4 ;± Ca 24 + H 20 Ca(OH ) j( portlunditc) + 2 H 4 Ca 24 + 2 H 20 CaFe 204(c) + 8 H + CaJ + + 2 Fe + + 4 H 20

- ^ ^ ^

V

,

Sulfates 22 23 24

CaS04( insolubIe) ;± Ca 24 + SOj a -CaS04( soluble ) Ca 2 4 + SOi /? CaS04(solubIe ) : Ca 2 4+ + SO} CaS04 • 2 H jO(gypsum ) Ca 2 + SO? + 2 HzO

25

Fluorides CaF2( nuorile) Ca 2 + + 2 F’

21

"

35 36 37

3

'

"

Solution Complexes Co24 + Cr CaCl + Cn 2 + + 2 CI CaCl ? + C02( g) 4- HjO CaHCOj + H 4 + C02(g) + H 20 CaCO$ + 2 H 4 Ca 2 + + NOJ CaNOj4 Ca 24 + 2 NO; CafNO )? Ca 24 + H 2O CaOH * + H 4 Ca 24 + 2 H 20 ^± Ca (0 H )5 + 2 H 4 Ca 24 + HjPO; CaH 2 P044

^ ^ ^ =^ ^ , ^ ^ ^ + H PO; CaHPO; + II + HjPO; ^ ^ CaPOi + 2 H + SOi s± CaSO: '

Ca 2 +

Ca 2 +

Ca 24

Ca 24

Ca 24

4

2

’

3195

2180

21.42

-4.41 -145 - 1.75 -4.64 - 10.41

^

26 27 28 29 30 31 32 33 34

^^

-

"

-150*

4

- 1.00 0.00

-6.70 - 15.01 - 4.80 - 4.50 - 1170 - 27.99 1.40 -4.46

- 13.09 131

( Continued )

17 18 19 20

21 22 23 24

Soil , Oxides , Hydroxides, Ferrites Soil -Ca CaJ 4 ( Tr . CaO( limc ) + 2 H 4 Ca 24 + H 20 Ca (OH ) 2( portlanditc) + 2 H + Ca 2 + + 2 H 20 , CaFe 204(c) + 8 H + CaJ + + 2 Fe + + 4 H 20

^ ^ ^

Sulfates CaS04( insoIuble ) ;± Ca 24 a - CaS04( soluble ) ?± Ca 24 /?-CaS04(soluble ) ;± Cn 244 CaS04 • 2 H 20(gypsum ) Ca 2

3

"

"

+ SOi + SOj + 2 H 20 "

"

Fluorides CaF2( lluorite ) Ca 24 + 2 F ~

25

26 27 28 29 30 31 32 33 34 35 36 37

+ SOi + SOi

Solution Complexes Ca 2 + + cr cncr Cn 2 + + 2 CP ^± CaC15 + C02(g ) + H 20 ^± CaHC0 j+ + H + C02( g) + H 20 CaC05 + 2 H 4 Ca 2 + + NO; CaNOj Ca 24 + 2 NO; Ca( N 02)$ Ca 24 + H 2 O CaOH * + H 4 Cn 24 + 2 H 2 O CB ( OH )$ + 2 H 4 Cn 24 + H 2 PO; 5± CaH 2 PO; Ca 24 + H 2 PO; CaHPO; + II 4 Cn 24 + HjPO; CaP04 + 2 H 4 Ca 24 + SOi 5* CnSOJ

^

4

'

^ ^ ^ ^ ^ ^ ^

‘

3195 2180 21.42

-4.41 - 145 - 1.75

-4.64

- 10.41

^

Ca 24 Cn 24

-150*

- 1.00 0.00 - 6.70 - 15.01 - 4.80 - 4.50 - 1170

- 27.99 1.40 -4.46

- 13.09 131

{ Continued )

s

TABLE 7.1 ( Continued ) Reaction No.

Equilibrium Reaction

-

logiC

Redox and Other Reactions 38 39 40 41

Cai + + 2 e

(c) ^ Ca HPOr + H * ^ HPOi POr + PT H po;^ ± por + 2 H *

H 2 P04

’

2

* Arbitrary

_

^

reference level for CaJ * used for noncalcarcous soils .

97.16

-7 ^0 -1235 -1935

CALCIUM SILICATES AND

ALUMINOSILICATES

91

10

8

\c

6

\

0\ P , S

4

+\x '

\

2

SnU>

\

g5 o

"?0

r2

N

Soil - Ca

yS - CaSi 03 ( 2) CaSi 03 ( 3 ) 3 - Ca 2 Si04 / (1 )

-4 -

7- Ca 2 Si04 15 ) CaAI2 Si06

(4 )

-6

-8

-

_

(6 ), (7),(8) CaAI2 Si2Oa

20 -

( 9) CaAI2 Si fc 2,H2O (10) CaAI2Si40|2 2H20

8024 - 7H20

( 11 ) Ca 2 AI4 Si (12 ) CaMgSi 206

1 4

5

6

7 pH

'

8

9

10

.

Fig.7.1 The solubility of calcium minerals in equilibrium with kaolinile ( K ). quartz (Q) and /or gibbsite (G ) as indicated .

Equation 7.2 is plotted as the wairakite solubility line in Fig. 7.1. The other lines in this diagram were developed similarly. The number (10) accompanying this mineral identifies its chemical formula in Fig. 7.1 and refers to Reaction 10 in Table 7.1 where the equilibrium constant for the indicated reaction is given. The equilibrium constants in this table were calculated from the AG°f values documented in the Appendix. The KQ associated with wairakite, indicates that kaolinite and quartz were used to establish the

CH. 7 92

CALCIUM

order:

pyroxene) > /?-Ca 2Si04( lamite) CaAl Si 2 Oa (Ca -gla5s) > CaAI 2 SiO„( (hexagonal anorthite) > y-Ca 2 Si04(Ca-olivine) > CaAi 2Si 2 Oa > CaAl 2Si 40 I 2 - 2 H 20(wairakite) > CaSi 03(pseudowollastonite) > CaAl 2Si 208(anorthite ) > /?-CaSi03(wollastonite) > CaMgSi 206(diopside ) > CaC03(caIcite) > CaAl 2Si 208 2 H 20( lawsonite ) > Ca 2Al4Si8024 • 7 H 20(leonhardite). , All minerals included in Fig. 7.1 are unstable in acid soils and can be expected to dissolve eventually. In alkaline soils calcite, leonhardite, and lawsonite appear as possible stable minerals. The soil-Ca line drawn at Ca 2 + = 10 2 - 5 M is given as a reference activity of Ca 2 + in acid and nearneutral soils. In such soils soluble Ca 2 + is largely buffered by exchangeable calcium. As exchangeable bases are depletea from soils, the pH drops causing H + and Al 3 + to enter the exchange sites displacing Ca 2 + and other cations.._ This displacement helps to maintain fairly, constant levels of Ca 2 + “

overa wide range of pH. It is a common agricultural practice to lime soils that become highly acidic. The twomaior benefits of liming are that ( 1) it neutralizes soil acidity and (2) it replenishes exchangeable Ca 2 + . Liming reactions can be rep resented as follows:

-

CaC03(c) + 2 H + -exch

>

CaC03(c) + 0.66 Al 3 + -exch

+ 2 H 20

Ca 2 + -exch + C02(g) + HaO (7.3) Ca 2 + -exch + 0.66 Al (OH) (solid )

^

^

+ C02(g) + HzO (7.4) The calcium silicates and aluminosil icates included in Fie 7 1 can all

CaAl 2Si06(pyroxene ) + 2 H + ^ exch

+ 2 H 20 — Ca 2 + -exch + Si (solid ) 02 + 2 Al (OH)3(solid ) (7.5)

Thus 1 mole of pyroxene is equivalent to only 1 mole ofCaC . Based on the solubility 03 informati on used to develop Fig. 7.1, leonhardite and lawsonite would appear to be more stable than calcite.. However,

OTHER CALCIUM MINERALS

93

experimental evidence supporting this conclusion was not found . To the contrary, calcium carbonate generally accumulates in alkaline soils. Possibly the thermodynamic measurements used to calculate the solubilities of lawsonite and leonhardite may not be sufficiently accurate to permit this close an interpretation. Increasing C02(g) above that of the atmosphere suppresses the calcite solubility lines in Fig. 7.1. Also the solubility of leonhardite increases as silica is weathered from soils and the activity of H 4Si04 drops below 10 * M . For these reasons, calcite at a designated partial pressure of C02(g) is used in this text as the solid phase most likely to control the Ca 2 + solubility in alkaline soils. “

12 OTHER CALCIUM MINERALS

The solubilities of several other calcium minerals are given by Reactions 13 through 25 of Table 7.1 and several are plotted in Fig. 1.2.Calcite is the least soluble calcium carbonate mineral, and CaC03(aragonite) is only slightly more soluble*. Because of their similar solubilities, both minerals may occur together unefer certain conditions. The hydrated mineral CaC03 6 H 20 (ikaite) is considerably more soluble and is not expected in soils. and The mineral CaO(lime) is much too soluble to persist in soils form to is even too soluble to appear in Fig 7.2 It readily hydrates soluble than calcite Ca( OH )2 ( portlandite), which is considerably more ite can coexist is portland and (Fig 7.2). The C02(g) level at which calcite calculated as follows: log Kc

-

Ca 2 + + 2 H 20 + CaC03(calcite) + 2 H CaC03(calcite) +

from which

HzO

;

+ Ca( OH) 2( portlandite) + 2 H Ca 2 + + C02(g) + H 20

± Ca (OH)2(portlandite)

+ C02(g)

- 2280 9.74

-13.06 (7.6)

13.06 C02(g) = io -

(7.7)

" 13 06 atm, both calcite and portlandite can Thus when COa(g) drops to 10 level, calcite becomes metastable to coexist. If C02(g) drops below this portlandite (unlikely in soils ). in Table 7.1 ( Reactions 21 to 24), Of the four sulfate minerals included . The solubility line for gypsum in CaS04 • 2 H 20(gypsum ) is most stable

CALCIUM

94

5

PH Fig. 7.2 The solubility of various calcium

minerals in soils.

Fig. 7.2 is drawn for a sulfate activity of 10 3 M showing shifts upward and downward. Gypsum is too soluble to persist in soils unless SOl approaches 10 2 M . Generally gypsum is found only in arid soils were very little leaching occurs. It also occurs temporarily in some soils rich in sulfides that oxidize to release large amounts of SO , for example, some of the polder soils of the Netherlands. * Application of gypsum to sodium to 10 raises soluble Ca 2 + above that affected soils in the pH range of 8 5 of calcite (Fig. 7.2) and leads to the precipitation of calcite with the release of protons: Ca 2 + + CO,(g) + HzO CaC03(calcite) + 2 H + (7-8) "

"

"

-

,

COMPLEXES OF CALCIUM IN SOLUTION

95

The pH then drops into the range of 7.5 to 8.0 where gypsum and calcite can coexist. In this way soluble Ca 2 + is restored to approximately 10 “ 2 5 M, keeping the soil colloids flocculated and predominantely calcium saturated. Displaced Na 4 can then be leached from the soil as drainage is supplied. The solubility of CaF2 (fluorite) ( Reaction 25 of Table 7.1) in equilibrium with soil-Ca can be represented as

-

log K °

CaF2(fluorite) Ca 24 CaF 2( fluorite )

Ca 2 + + 2 F Soil -Ca

"

=

2 F " + Soil-Ca

( F )2

=

-

IO

"

- 10.41 2.50

- 7.91

(7.9)

-

7 91

which gives log F

"

=

- 3.96

( 7- 10)

~4 The presence of fluorite in soils limits F activity to approximately J 0 M . Further consideration will be given to KMg 3 AlSi 3O 10 F2 ( fluorphlogopite) in Chapter 8 and to Ca 5( P04)3F(fluorapatite) in Chapter 12 as other possible fluoride-controlling minerals in soils. "

73 COMPLEXES OF CALCIUM IN SOLUTION

of an element in the soil It is sometimes important to know the ionic forms 2+ given in Table 7.1 are Ca of es complex solution. Several inorganic by these equations depicted ships relation ( Reactions 26 through 37). The activity ranges used to develop are plotted in Figs. 7.3 and 7.4. The anion . soils in these plots are those commonly found and nitrate complexes of calcium Figure 7.3 shows that the chloride calcium. The activities of CP and N03 contribute very little to total soluble es both in the presence of soil -Ca and affect the stability of these complex extra lines and arrows were omitted to calcite. In the case of calcite, the . The CaHCOj ion is insignificant in avoid further cluttering of the diagram io submerged soils where pH normally acid soils, but it becomes important pressure of C02(g) increases. Such approaches neutral and the partial Ca 2 + . changes increase CaHC03 and decrease (g ) by 10 fold depresses log Ca 2 + , increasing C02 In the presence of calcite + and Ca ( NQ 3 )l by one log unit ( Fig. 7.3). The Ca CP, CaCl 2 , CaNQ 3 , *

1

.

CH 7

96

CALCIUM

0

-

b » 2.52

Ca * *

-5 *

log

cr

A-

2

cacr

£ S

J

w

eft

•

-

-10

—

CaN03 log

—

NOf 2

-

r.

a-

- 15 4

\

V 'A

CaCI2* o,

o>

2

^

log NO

2 !

J

Ca( N03 ) 2

J

/

5

6

7 pH

8

9

10

Fig. 7 J Chloride, nitralc. and bicarbonate complexes of Ca ’ t in equilibrium with »oil Ca or calcite as affected by COi (g ), anion activities, und pH .

-

activity of CaC03 is constant in the presence of calcite. The exact value can be obtained by combining Reactions 13 and ?9 of Table 7.1 to give log K °

CaC03(calcite) + 2 H + Ca 2 + + CO ,(g ) + HjO CaC03(calcite)

Ca 2 + + C02(g) + H 20 CaCO 0 + 2 H +

,

-

9.74 15.01

-5.27

CaC03

(7.11) Thus calcite supports the neutral Ion pair, CaCOj , at 10’3 , 27 M which is independent of C02( g ) and pH. t

COMPLEXES OF CALCIUM IN SOLUTION

n

Or log S04"

_Ca* *

Li

\

CaSO/

'og H2PO4

'

-5

-r

ggH 2 POj

*

1 - «, d

&X

j

-10

fc#log CP2 ( g )

--

a - 3.32 b > 2.32

Ca(OH )2

-15 6

5

A

7 pH

8

9

10

-

Fig . 7 A Sulfate , phosphate, and hydroxyl complexes of calcium in equilibrium with *oil Ca or calcitc .

calcium Figure 7.4 shows that CaSOJ contributes significantly to total 2+ ~4 complexes ofCa phosphate include in solution ifSO is > 10 M . The none of these soils complexes acid ; . In CaPO CaHjPO , CaHPOS , and ^ contribute significantly to total soluble calcium . In neutral and calcareous soils, CaHPC> 4 and CaP04 are significant, especially when the activity complexes depicted of H 2 PO is > 10 5 M . In the presence of calcite all * in C02(g), but their in Fig. 7.4 decrease one log unit for each 10-fold increase 2 Changes in SO and H 2 P04 affect relative positions remain unchanged . of both calcite and soil -Ca. the stability of calcium complexes in the presence of calcite to simplify The extra lines and arrows were deleted in the case

^

"

“

"

plotting.

CALCUJV,

CH. 7

98

The hydrolysis species CaOH + and Ca(OH )2 are unimportant in th pH range of soils. Speculation that CaOH + acts as an exchangeable ion th 3 enables more than 100 % calcium saturation of soils is highly unlikely sin * the ratio CaOH + /Ca 2 + never exceeds 1 (T 4 ( Fig. 7.4).

/

7.4 REDOX RELATIONSHIP OF CALCIUM

The only important oxidation state of calcium is + 2 The redox reaction relating Ca 2 + to Ca(c) is

Ca 2 + + 2e ~

Ca(c)

log K ° =

-97.16

from which pe

= -48.58 + 0.5 log Ca 2 +

(7.12)

Thus an extremely low pe is required for Ca (c) to be stable. The only way that redox can affect calcium solubility relationships in soils is indirectly as it may affect other cations or anions with which Ca 2 + combines. 73 THE CaC03

—C02 — H ,0 SYSTEM

When calcium carbonate is present in soils, it has a dominating influence on many soil properties. Most calcareous soils fall in the pH range of 7.3 to 8.5, and only in the case of sodium -affected soils does the pH normally rise above 8.5. Because of the importance of calcium carbonate, further consideration will be given to its equilibrium relationships. The solubility of CaC03(calcite) can be expressed as follows : fn terms of C02(g) log K°

, CaC03(caIcite ) + 2 H + log Ca 2 +

In terms of H 2 C03 CaC03(caIcite ) 4- 2 H + C02(g) + H 2 O + CaC03(calcite ) + 2 H

Ca 2 + + CO,(g) + H 20

9.74 (7.13) (7.14)

+ 2 pH = 9.74 - log CO,(g) Ca 2 + + C02(g) + H 20 H 2 COI Ca 2 + + H 2COS

-

9.74 1.46

8.28 (7.15)

I *«* THE CaC03 C02 — HjO SYSTEM

99

log Ca 2 + + 2 pH = 8.28

In terms of HC03 CaC03(calcite ) + 2 H 4 C02( g) + H 2 O

Ca 24 + C02(g) + H 20 H 4 + HC03

CaC03(calcite) + H 4

( 7.16)

- log H ,CO$

-

9.74 7.82

1.92

Ca 24 + HC03

"

(7.17) (7.18)

1.92 - log HC03

log Ca 2 4 + pH

In terms of CO 2 CaC03(calcite) + 2 H 4 C02( g) + H 20

;

CaC03( calcite)

Ca 24 + CO,(g ) + H 20 2 H 4 + COf "

Ca 24

+ CO1

“

-

9.74 18.15

- 8.41 ( 7.19 )

log Ca 24

= - 8.41 - log CO 2 -

( 7.20)

Equations 7.14, 7.16, 7.18, and 7.20 are all equivalent expressions for the solubility of calcite. Rearranging Eq . 7.14 gives ( 7.21 ) pH - 0.5 pCa = 4.87 - 0.5 log C02(g ) which is referred to as the lime potential relationship. In this text the expres sion log Ca 24 + 2 pH is used rather than pH - 0.5 pCa to express the pH dependent solubility of calcium minerals. Convenient solubility relationships of calcite that follow from Eq. 7.13 through 7.21 are summarized below : Ca 24 , log C02 , CO 2 (g), pH O.SpCa log Ca 2 + + 2 pH mol/1 atm atm

-

-

0.0003 0D03 0.01 0.1 1.0

- 3.52

- 2.52 - 2.00 - 1.00 0.00

+ 2

10 ,I 3.26( H ) 4 2 10 2'26( H ) 101174( H 44 )22 10 ' °-74( H 4 ) 2 109‘74( H )

13.26 12.26 11.74 10.74 9.74

6.63 6.13 5.87 5.37 4.87

um relationships of calcite A useful plot of Eq. 7.14 showing the equilibri 7.5. This graph shows the relation at various CO,(g) levels is given in Fig.24 (g). The solubility line ships among the variables : pH, log Ca , and C02

-

l:

Of . 7

CALCIUK

m 0

-1

yOOl

sai *

2

0

-3 -4 U

o»

-5

- ft -7 -8

-9 - 10

1 5

4

J. 6

1

1

J.

X

7

8 PH

9

10

J

1 11

12

. 7.3 The solubility of CaCO, (calcue ) to term* of pH . CO ,( j). and Ca1 * * connecting line and points show the composition of pure CaQ Fi

-^ . CO

-

indicated levels of CO (g )

—

H .O

activity. The at lb*

systems

for calciie moves upward as CO:(g ) decreases and downward as it increases. As shown earlier, the calcite line moves up to coincide with that of Ca(OH )2( portlandite) when C02(g ) drops to 10 ~ 13 06 atm. The curve that cuts across the calcite lines corresponds to the solution compositions of pure CaO C02 H systems as C (g) changes. Let 20 02 us derive the composition of the . pure calcite water system cor responding to C02(g) at 0.0003 atm. The electroneutrality equation applicable for this system is 2[Ca2 *] + [CaHCOH + [H *] « [ HCOj ] + 2CCOf -] + [OH ~ ] (7.22)

—

—

-

t^ 101

THE PHASE RULE

and in terms of activities 2(Ca 2 * ) (CaHCQ 3 ) t y&s* VCaHCOj

[

(H * )

_ (HCQ3“ )

>V

yucoi

2(COaa “ ) ycoj -

H

( OHT yoH

-

(7< 23)

The activities of Ca 2 \ CaHCO , HC03", CO ] ", and OH " can be expressed in terms of pH and the formation constants for calcite, CaHC03 , carbonic acid, and water. The activity coefficients can be calculated from the Debye Huckel Eq. 2.11 and ionic strength from Eq . 2.8 using as a first approximation the concentrations of Ca 2 * and HCOJ estimated from Fig. 7.5 and 6.2. Substituting these parameters into Eq . 7.23 gives :

^

,-

10

3 69

( H + )4

+ 10° O 2( H + )3 = 10 - 11.31( H * ) + 10 -

-

2 i . au

(7.24)

Several iterations can be used to obtain precise estimates of ionic strength. By substituting various values for ( H * ) into Eq . 7.24, a pH value of 8.34 is found to give the correct solution to this equation. Thus a pure CaO C02 H 20 system in equilibrium with 0.0003 atm C02(g) comes were to pH 8.34. The other points along the transecting curve - in Fig. 7.5 obtained similarly using the indicated C02(g) levels.

— —

7.6 THE PHASE RULE

The chemical phase rule states that

F

= C- P + 2

(7.25)

of freedom or the smallest number of where F is the number of degrees completely define a system (it can be thought independent variables needed to that may be changed without causing of as the maximum number of variables a phase, thereby destroying the system) ; the appearance or disappearance of is, the least number of chemically C is the number of components, that the composition of every phase in independent species required to describe present. A phase is a region of uni the system ; and P is the number of phaseswithin a system. The last term in form chemical and physical propertiesand pressure as variables. If temper Eq. 7.25, “ 2,” accounts for temperature degrees of freedom are removed. ature and pressure are fixed, these two fields are involved, the last term In certain cases where electric or magnetic , etc. must be increased from 2 to 3 to 4

-

CH. 7

102

7.7 THE COi

—

HjO SYSTEM

CALCIUM

——

When temperature and pressure are fixed, F C P. The components in this system can be selected as C02(g) and H 20. The phases consist of C02(g) and a solution. Applying the phase rule gives: F=C P= 2 2=0

-

-

This means that when the partial pressure of C02(g) is specified, the solution composition is also fixed. The composition of the solution phase in this system can be calculated for any given C02(g) level.

— —

7 JS THE CaO C02

H 20 SYSTEM

If temperature and pressure are fixed for the CaO

phase rule states:

F

— 02—

— —

P

P=3

— 3=0

C

C

H 20 system, the

There are now' three components: CaO, C02 , and H 20. At equilibrium three phases are possible, for example: C02(g), CaC03(calcite), and solution could exist at equilibrium giving:

F=C

—

This means that CaC03(caJcite) placed in pure water at a fixed C02(g) has a unique solution composition. The equilibrium constants and activity coefficients of the species involved make it possible to calculate the com position of the equilibrium solution as demonstrated in Section 7.5.

-

— — —

7.9 THE CaO C02

H 20 H 2SO* SYSTEM

Again if temperature and pressure are fixed

F=C

—

P

The components could include CaO, C02 , H 20, and H 2S04 giving four components. At equilibrium four possible phases may exist such as C02(g), CaC03(calcite), CaS04 2 H 20(gypsum ), and the solution phase. Again the phase rule indicates

-

F=C

—=— P

4

4=0

Thus at a specified C02( g ) if either CaC03( calcite ) or CaS04 - 2 H 20 ( gypsum ) is present but not both, there is 1 degree of freedom and the i

I

^

PROBLEMS

103

addition of a component such as H 2S04 can alter the composition of the solution phase. If both CaC03(calcite) and CaS04 - 2 H 20(gypsum ) are present, the equilibrium solution will have a unique composition, and the

addition of a component such as H 2S04 merely dissolves calcite and precipitates gypsum according to the reaction

H 2S04 + CaC03(calcite) + H 20

.

(7.26) CaS04 •2 H 20(gypsum ) + C02(g) and the composition of the solution phase remains fixed. It is possible to calculate the composition of the solution phase that will have a fixed com position so long as C02( g) is fixed and both solid phases are present. Using the electroneutrality equation appropriate for this system gives a unique solution of pH 7.8 when C02(g) is 0.0003 atm . Since soils contain many different chemical components in addition to CaO, C02 , H 20, and H 2S04 , the calcite-gypsum equilibria can occur in soils at pH values slightly displaced from 7.8, but generally the displacement will not be very far. Addition of one or more of the four components to soils, however, will not affect the pH of a soil as demonstrated by Eq. 7.26. For example, addition of H 2 S04 , CaO, CaC03, CaS04 2 H 20, or other combinations of these four components, to a soil containing calcite and gypsum will not affect the pH of that soil. Increasing C02(g), however, will lower the pH at which the two minerals can coexist.

-

-

PROBLEMS 7.1

7.2

Develop the equations and plot the solubility line for leonhardite shown in Fig. 7.1. Also develop the equations and plot this line when 3+ gibbsite and kaolinite control Al and H4Si 04 activities. Write the balanced chemical reactions showing the weathering of

wairakite: a. In acid soils. b. In calcareous soils. c. Discuss how the pH of the two soils is affected. Develop the necessary relationships to show that CaHCO in a calcareous soil is independent of C02(g). Calculate the contribution that CaHC03 makes to total soluble calcium in a submerged soil of pH 7.3 when C02(g) is 0.05 atm and 3 5 S04 activity is 10 - M (assume activity coefficients are unity ). From Fig. 7.4 estimate the activity of CaHPOJ in a soil of pH 8.2 having 10 4 M HPOi and 0.003 atm C02(g).

-

73 7.4

"

“

7.5

^

"

"

. CALCIUM

« 04 7.6

7.7

7.8

7.9

CH 7

Develop the necessary equations and plot the log Ca 2 * versus pH f0 r calcite in equilibrium with 0.003 atm of C02(g). Calculate the pH of a CaO C02 H 20 system that attains equilib rium with 0.003 atm C02(g ). Hint : Begin with the electroneutralit equation and use the Debye- HQckel equation to estimate ion activity coefficients through successive approximations of ionic strength Also calculate the activities of other constituents in the equilibrium solution and plot this solution composition on the diagram developed in Problem 7.6. Calculate the pH of the CaO C02 H 20 H 2S04 system in equilibrium with 0.003 atm C02(g), CaC03(calcite) and CaS04 • 2 H 20(gypsum ) using the procedure suggested for Problem 7.7 above. Also calculate the activities oT the other parameters in the equilibrium solution. . Discuss the effect of adding the following constituents to a CaO C02 H 20 system initially in equilibrium with CaC03

— —

—

-

—

—

— —

(calcite).

a. CaCl 2 . b. HCI. c. NaOH. d. NaCl. e. H ,S04 . f. Soil. g. Increasing C02(g). to the system 7.10 Discuss the effect adding the following cohstituents gypsum are CaO COz H 20 H 2S04 system where calcite and initially present :

— — —

a. CaCI 2. b. HCI. c. NaOH. d. e. f. g.

CaO.

H 2S04. Soil.

Increasing C02(g). 8.5 M solution of NaHC03 at pH e 0.5 a to happen will what 7.11 Calculate is for available phosphorus) that (Olsen’s bicarbonate extract open to the atmosphere.

I

I

!

EIGHT MAGNESIUM

i I

! i

; I

I 1

i

HThc magnesium content of the lithosphere is estimated at 2.1 %, while 1 the average content of soils is only 0.5 % (Table 1.1). These levels reflect

the removal of magnesium from soils during weathering. The solubility relationships of several magnesium minerals are examined in this chapter to determine which minerals are stable in soils and which ones are most likely to control the activity of Mg 2 * and its complexes in the soil solution. ' The solubility relationships of magnesium phosphates, sulfides, and mo lybdates are examined in Chapters 12, 17, and 22, respectively. 8.1 SOLUBILITY OF MAGNESIUM SILICATES

The solubilities of several magnesium silicates are given by Reactions 1 through 9 of Table 8.1 and are plotted in terms of log Mg 2 + versus pH in Fig. 8.1. In this development H 4SiC> 2 was fixed by SiO,(soil ) representing the early stages of weathering 'where most primary magnesium silicates are weathering. To plot Mg ,’ 6 Fe( II )0 4Si04(olivine), Fe2 * was fixed by Fe(OH)3 (soil ). Redox levels are represented at pe + pH of 10 and 17. The magnesium silicates included in Fig. 8.1 decrease in solubility in the order: Mg , .6 Fe( II )0.4SiO4(olivine ) > Mg2Si04(foresterite ) > MgSi03(clinoenstatite) > Mg3Si 4O , 0( OH ) 2 • 2 H 20( vermiculite) > MgCa (Si03) 2 (diopside) in equilibrium with calcite > Mg 3Si 2 Os( OH )4(chrysotolite ) > Mg 2Si 3 Os ( OH )4(sepoIite ) > Mg6Si 4O10(OH )8 > (serpentine) > Mg 3Si4O10(OH )2( talc). Dolomite, MgCa(C03)2 , in equilibrium with CaC03(calcite) included as a reference mineral falls between serpentine and talc when C02(g) is.10 - 3.52 atm. Soils generally have higher C02 . Increasing C02 by 10-fold causes dolomite to become the most stable mineral shown in Fig. 8.1. The silicate minerals in Fig. 8.1 shift with changes in H 4Si04 activity. The talc line was established using the following equilibria : log K°

Mg 3Si 4O 10(OH ) 2 ( taIc) + 4 H,0 + 6 H * 4 H 4Si02

Mg3Si 4O|0( OH )2 ( talc ) + 6 H *

3Mg 2 + + 4 H 4Si04 4 Si 02(soil ) + 8 H 20

= = 3 Mg + 4 Si02 soiI ±

2+

. m

. 22.26 4(3.10)

2* 3 )

( Mg ( H * )6

=

(

-

lo34 66

) + 4 HzO

34.66 (8.1)

TABLE 8.1

EQUILIBRIUM REACTIONS OF MAGNESIUM AT 25 C i

•

Reaction

No.

Equilibrium Reaction

log

Silicates 1 2 3 4 5 6 7 8 9

MgSiOj(clinoenslatile ) 4- 2 I I 4 4- HjO Mg 2 ' 4- H 4 Si 04 MgCa ( SiO ,) 2 ( cliopsidc ) + 411 * + 2 ll 20 ;± Mg 2 * 4- Ca 2 * 4- 2 H 4 SiOS MgjSiO.drorsteritc ) + 4 I I 4 ± 2 Mg 2 ' + H 4SiO; ^ 1.6 Mp 21 0.4 Fe 3 ‘ -t ll ., Si Mgi . f, l e( ll )0 4 Si 04( olivinc ) + 411 * + - 04 4 ” . ) ) Mg 6 ( chrysololitc H Si 3 ( 2 ll i OM Si + Mg .< 2 Os 4 04 + U 20 4 2 * M g 2 Si , (.) ,,( (.) H )4( scpolitc ) + 21120 + 4111 2 Mg + 3 H 4 SiOJ Mg .,Si 4 OI 0( OH ) 2 ( lalc ) + 4 H ,0 + 611 ' 3 Mg 2 * + 4 H 4 Si 04 3 Mg 2 4 + 4 H 4SiOJ Mg , Si 401„( 01- 1 ) 2 • 21120( vcrmiculitc ) + 2 U 20 + 6 U ‘ 4 6 Mg 2 * +, 4 U 4SiQ4 + 2 H 20 ' Mg,,Si 40|o( OM )8(scrpentinc ) + I 2 H

^

'

^ ^ ^ ^ ^ ^

11.42 21.16 28.87 26.18 32.87 13.89 22.26 30.39 61.75

Aluminosilicates

10 II 12

13 12 14

,

4 Mg, AI 2 Si 0 I 1( OM ) H ( chloritc ) + 161! * ^± 5 Mg 2 * 4- 2 AI ' 4- 3 H 4 SiOS + 6 M:0 * 2 Mg 2 ' 9- 4 A 1 x ’ 4- 5 H 4 SiOS Mg 2 Al 4 Si , OIM ( corclicritc ) -l- 2 H 2 ( ) -I- 16 II 4 4 24 Ko . hMgo . t jAlj . jSi.,.*Olt,( OII ) ,( illitc ) + 811 ' -I- 2 M 2 O 0.6 K 4- 0.25 Mg + 2.3 AI ' (+) 3.5 H 4SiOS ' - t- 3.24112 Mgo . 2 ( Si HI Al|. 71 Fe( III ),, 22 Mg „ 2 , )0 M,( OII ) 2( Mg montmorillonitc ) 4- 6.7611 4 f 2 ^ 0.49 Mg 4- 1.71 Al ' 4- 0.22 lV * 4- 3.81 II 4SiOS 4 Mg 2. Fc( ll )0.02 Fed II )0 4 (.Ca „ . iw, K „ , ( Si 2 . » I Al i . i J ) Olo( OII ) 2 ( vermiculile ) 4- 10.36 H } < 2 * 2.71 Mg 2 • 4 - 0.02 Fc 2 • I- 0.46 Fc 4- 0.06 Ca • 4 O. I K + 1.14 Al ' ‘ 4- 2.91 H 4 Si 04 4- 0.3611 ,0 (

.

^ - ^

^

*

„

^

60.12 45.46 10.34 2.67

38.05 ( Continued )

1 S

I

TABLE 8.1 ( Continued )

Reaction No.

log

Equilibrium Reaction

**

1

1 1

Oxides, H droxidcs, and Carbonates

*

15 16

17 18 19 20

MgCKpericlase ) + 2 H + Mg(OH ) ,( brucilc) + 2 H * MgC03( mugnesitc) + 2 H * MgCOjOH O( nesquehonitc) + 2 H * MgCOj 5 H Oflunsfordite ) + 2 H + ^ MgCatCOjWdolomite) + 4 H +

-

.

21.74

MgJ * + HjO ^ MgJ 2 H ,0 ^ Mg * ++ CXh (g) + H ,0 1+

MgJ *

+

CO ,( g)

16,$4

11X69 13.49 13.62 16.46

+ 4 H.O

MgJ + + C03(g) + 6 H,0

^ ^ Mgi

+

+ Ca1* + 2 C03(g) + 2 H 30

Sulfates

21

MgS04(c) ?iMgIf +

SOi -

8.18

Solution Species 22 23 24 25 26 27 28 29 30

-

Soil Mg Mgi + 2 d MgJ + + CO,(g ) + H 2 O MgJ + + CO,( g ) + H 1O MgJ + + 2 NO,MgJ + + HjO MgJ + + 2 HjO +

Mg1 *

^ MgHCO; + H * ^ MgCOS + 2 H ^ Mg ( N03 )S +

^ MgOH + H * ^ Mg(OH )J + 2 H * +

^ MgHPOJ + H MBU + soj^- ^ Mgso;

Mg1 + + HjPO

- 3.00* - ao3

MgCIJ

——

- 6 76 ,

14.92 - aoi 11.45

- 27.99 4.29

-

2.23

i

Redox Rcuclion

MgJt + 2c

3!

"

^ Mg(c)

- 79.92

Ollier Reactions 32 33 34 35 36

Fc( OH ) j(soil ) + 3 H + Fe 3 + + 3 H 20 Fc( OH ) ,(soil ) + 3 H + + c ± FcJ + + 3 H 20 ^ , AI( OM ) j(gibbsilc ) + 3 Hf Al ‘ + + 3 H 20 AljSi Oj( OH )4 ( kaolinilc ) + 6 H * 2 AI 3 f + 2 H 4 Si 04° + H 20 CuCOj(calcilc ) + 2 M * Ca 2 ' + C02(g ) + H 20 '

,

.

* From Eq 8.6 in text.

s

^ ^

^ ^

2.70 15.74 8.04 5.45 9.74

SOLUBILITY

110

OF

MAGNESIUM SILICATES

2

1

o

-1

log

-

co2 (g )

a - 3.52 b« - 2.52

cn

Q - Quartz S - Soil Si

2 -2 o’

-

Soil Mq (1 )

-4

MgSiQ 3

^

(2 ) MgCaCSiO (3) MgSiQ4 (4 ) Mgi 6 Fenr )Q4Si04 (OH )4 ( 5) Mg 3Si

.

'

O

205

Mg2Si 306(OH )4 Mg 3SI 4 Oi0 (OH )2 (8) Mg 3Si 4Oio(0H ) 2 ' 2 H2 (9) Mg 6 Si 4O1p( OH )0

(6) (7 ) *

5

-6

.

4

°

5

6

1 7

-v

a

9

10

pH

-

-

Fig 8.1 The solubility of several magnesium silicates in equilibrium with soil Si and soil Fe with indicated changes for COjIg) quartz, and redox.

.

from which log Mg 2 *

=

11.55 - 2 pH

-

log Mg 2 +

=

12.75

- 2pH

(8.2)

(8.3)

HSE SS iss unTsTLIe . byole of the mineral. For example, decreasing log H SiO;

SOLUBILITY OF MAGNESIUM SILICATES

taJc line up by f

^^^

Several u , . 0

= 1.33|og units of Mg 2 + whjJe lhat of serpentine moves up

meiore Sobblf

states bc

ill

H Si

5

‘

°S deClinK a"d

Suesium silicates may form in alkaline soils that are high in ese include talc, serpentine, sepolite, and chrysotolite. 4‘ . As H 4 biU 4 declines or C02(g) increases, dolomite calcite becomes the stable magnesium phase. The dolomite-calcite combination is used in this text as a limit of Mg solubility in calcareous soils. The equilibrium reactions describing this limit are as follows :

^

»

^

1

log K c

MgCa(C03)2(dolomite) + 4 H * Mg2 + -f Ca 2 + + 2 C02(g) + H ,O Ca 2 + + C02(g) + H 20 CaC03(calcite) + 2 H + MgCa (C03),(dolomite) + 2 H + Mg2 * + C02(g) + H 20 + CaC03(calcite)

18.46

-9.74 8.72 (8- 4)

from which log Mg2 +

= 8.72 - log CO2(g) - 2pH

-

(8 5)

Equation 8.5 is plotted in Fig. 8.1 and shows the pH -dependent solubility of Mg 2 * in equilibrium with dolomite calcite. . Below pH 7.5 most magnesium minerals are too soluble to persist in soils 2* 2+ by exchangeable buffered Mg soils acid in Mg of A reference solubility is estimated at 10 - 3 M. This reference level can be expressed as follows:

-

Mg2 + (soil )

MS2 +

logK ° =

- 3.0

(8.6)

In soils Mg2 + may vary slightly from this value depending on the rate of Mg' increases tn soil solution weathering compared to rate of leaching. If by leaching watem. This rapidly much above 10 3 M, it is removed more ightly below that of Ca reference level for Mg was selected to be s Ca 2 is generally higher exchangeable (10- 2 -5 AO to reflect the fact that magnesium may have than exchangeable Mg 2 * Occasionally soils rich in cases the activity of Mgexchangeable MgB 2 * than Ca” , and in such stability of magnesium more cAcndug f the great ^ 2 * depletion from acid soils can be 'under amd Mg m

-

rr ^

expected.

'

”

indhS

3' CH . 8

112

MAGNESIUM

8.2 MAGNESIUM ALUMINOSILICATES

The solubilities of several magnesium aluminosilicates are given by Reactions 10 through 14 of Table 8.1 and are plotted in Fig. 8.2. To develop this diagram, Al3 + and H SiOJ were fixed by kaolinite and quartz, Fe3 + by ^ 3 soil Fe, Ca 2 + at 10 “ 2 5 M or calcite, K + at 10 “ M , and pe + pH at 17. Under these conditions the solubilities of illite, cordierite, and vermiculite are much higher than Mg-montmorillonite, chlorite, and dolomite-calcite. The more soluble minerals will disappear, whereas the more stable secondary minerals will regain until they, too, are removed in later stages of weathering as soils become more acid. The solubilities of aluminosilicate minerals depicted in Fig. 8.2 are affected

-

”cn

2

o’ -2

Soil - Mg

-4

tog CO2 (g) a * - 3.52 b * - 2.52 Q - Quartz

-6

- 84

S - Soil Si G - Gibbsite (10) Mg 5 AI Si O10(OH )8 2 3 (11) Mg 2 AI4Si 50i 3

5

6

1 7

1 8

J. 9

10

pH

Fig. 8.2 The solubility of . magnesium aluminosilicates in equilibrium with kaolinite, quartz. CaJ + at \0 ~ 2 i M , K * at 10 3 M , and soil - Fe. '

OXIDES, HYDROXIDES, CARBONATES, AND SULFATES

m

by changes in H 4Si 04 . The

kaolinite-quartz equilibria was chosen as a chemical environment typical of the intermediate stages of weathering where secondary magnesium aluminosilicates are most likely to be weathering. Shifts in the equilibrium lines of Fig. 8.2 corresponding to changes in H 4Si04 can be readily obtained from the atomic ratios of silicon, aluminum, and magnesium in a mineral from the equation : A log Mg 2 +

=

^—

A log H 4SiO

~

MgS- j

(8-7)

With Mg 5 Al 2 Si 3OI 0(OH )8(chlorite) as an example, increasing log H 4Si02 by 0.9 units, as would result in going from quartz to soil Si. shifts the chlorite line by 0.9( 2 3)/5 = 0.18 log units of Mg 2 + . In the case of Mgmontmorillonite the shift would be 0.9(1.71 3.81)/0.49 = 4.76 log units. These shifts are indicated in Fig 8.2 by the short lines and arrows going from Q(quartz) to S(soil-Si). Only illite and Mg montmorillonite are affected appreciably by changes in H 4Si 04 . As H 4Si04 approaches that of soil-Si, Mg-montmorillonite becomes highly stable, even more so than dolomite calcite. In the advanced stages of weathering H 4Si 04 is expected to drop to 10 5 31 M where kaolinite and gibbsite can coexist (Fig 5.7). Under these conditions, the Mg-montmorillonite line in Fig 82 shifts upward by

—

—

-

-

-

"

(4.00

- 5.31)

1.71 - 3.81 0.49

(

5.61 log units

reflecting the great instability of Mg-montmorillonite at low H 4SiO 4. Other smectic minerals containing various substitutions of ions can be included in these solubility relationships once reliable free energy values are available.

83 OXIDES,

HYDROXIDES, CARBONATES, AND SULFATES

The solubilities of several oxides, hydroxides, carbonates, and sulfates of magnesium are given by Reactions 15 through 21 of Table 8.1, and are plotted in Fig. 8.3. Periclase ( MgO ) is much too soluble to persist in soils. It can hydrolyze to Mg(OH )2 ( brucite) or it can precipitate as one of the other less soluble minerals. The carbonates of magnesium are also too soluble to persist in soils. They decrease in solubility in the order : MgC03 - 5 H 20(lansfordite) > MgC03 3 H 20( nesquehonite) > MgC03(magnesite). Increasing C02(g) by a factor of 10 depresses Mg 2 + activity in equilibrium with each of the

-

CH. 8

114

MAGNESIUM

8

6

4

2

2 O

81 -2

Soil - Mq -4

( 15 )

MgO

( 16 ) Mg (OH )2 ( 17 ) MgCQ 3

-8

MgC03 * 3H20

(18)

-6

4

( 19 )

MgCOySHjjO

( 20)

MgCa (C03)2

1 5

6

log CO2 ( g ) a - - 3 - 52

1 7

t> « - 2. 52

1

8

9

10

pH

Fig. 83 The solubilities of oxides, hydroxides, and carbonates of magnesium.

magnesium carbonate minerals by one log unit. Of the magnesium minerals included in Fig. 8.3 only dolomite-calcite is stable. The reported solubility for MgS04(c) with a dissociation log K ° of 8.15 reflects a very high solubility for this mineral. It is much too soluble form in well-drained soils.

8.4 MAGNESIUM COMPLEXES IN SOLUTION

22 The stabilities of several magnesium complexes are given by Reactions through 30 of Table 8.1 and are plotted in Fig. 8.4

.

i

MAGNESIUM

COMPLEXES IN SOLUTION

115

0

042 -

log S

Mg 2 *

-3 -5

MgSQ4* 4

*°

w

£ 3

a

A

o\

,6

O)

MgtNC )'

MgCI2

-10

^

-4

cO

log NO3' or Cl " .

3 species is insignificant in acid soils but becomes important in slightly alkaline soils and in submerged soils where C02(g) is generally "

”

-

“

*

Hgh.

CH.

116

• MAGNESIUM

,

, ,

and CT found in soils. Also th are insignificant at the normal level of NO hydrolysis species MgOH * and Mg(OH )3 are insignificant and eontribm very little to total soluble magnesium.

gj

EFFECT OF REDOX ON MAGNESIUM

For magnesium only the + 2 oxidation state is important in soils. From Eq. 31 of Table 8.1 comes the relationship pe = -39.96 + log Mg 2 + (8.8) which reflects the great instability of Mg(c) in aqueous environments and soils. The only way changes in redox can affect magnesium in soils is in directly by affecting other cations or anions with which magnesium is associated.

-

PROBLEMS

8.1 Using reactions from Table 8.1, derive the necessary equations and plot the solubility of Mgt 6 Fe(II)0 4Si04(olivine) as shown in Fig. 8.1. Include shifts in this line for changes in pe + pH and for equilibrium with Si02(quartz). 8.2 Prepare a plot to show how the solubility line for serpentine in Fig. 8.1 changes when H 4Si04 is in equilibrium with a. Quartz. b. Kaolinite and gibbsite. 8.3 How much does the Mg-montmorillonite line in Fig. 8.2 shift when equilibrium is attained with tridymite rather than quartz ? 8.4 Develop an appropriate expression to calculate shifts in the solubility lines of Fig. 8.2 if equilibrium is attained with dickite rather than kaolinite. From this expression calculate the shift for Mg montmoril Ionite as well. 8.5 Justify the selection of dolomite as a solid phase limiting the activity of Mg in alkaline soils when it is obvious from Figs. 8.1 and 8.2 that there are minerals that may be less soluble (see also Fig. 5.6). 8.6 Examine the equilibrium conditions in a system prepared by adding

-

-

*

PROBLEMS 117

10 g each of dolomite, calcite, and magnesite to 50 ml of water open to the atmosphere. Indicate the following: a. The solids present. b. The ratio of Ca ^ Mg2 / ’ in solution. 8.7 Calculate the contribution of MgSO£, MgHPOS, and MgHCOj to total soluble magnesium in a soil having the following activities: 10 7 2 M (H +), 10 ~ 3 5 M ( SOI ), 10 4-6 ( 4 ) when the electrical conductivity of this soil is 0.77 millimho cmHPO 2 and C02(g) is 0.02 atm.

--

-

"

"

"

NINE SODIUM AND POTASSIUM

r " lithosphere is 2.8 %, whereas the average s 15 es r alcd at 0.63 % For potassium the lithosphere * contams approximately 2.6 % with an average orO.83 for soils (Table l . l ). % esc con en s re cct the removal of sodium and potassium during the weal enng o soi s. he removal of sodium slightly exceeds that of potassium. Both elements are essential to animals, and potassium is one of the three major fertilizer nutrients required by plants. Sodium and potassium are present as exchangeable cations in soils. In poorly drained soils, sodium salts generally accumulate as the major contributor of salinity. For these and other reasons, it is important to understand the chemical reactions of these two elements in soils. In this chapter the solubility relationships of sodium and potassium minerals are examined under the chemical weathering matrix of soils to see which minerals are stable and to determine if the complexes of these cations are important in soils.

°,, " '

.

9.1 SOLUBILITY OF SODIUM MINERALS

Sodium silicates are too soluble to form in soils, but several sodium aluminosilicate minerals are important. The solubility relationships of several aluminosilicates found in soils are given by Reactions 1 through 7 of Table 9.1 and are plotted in Fig. 9.1. These minerals are plotted in equilibrium with kaolinite. For those minerals having a ratio of Si/Al > 1 equilibrium with soil-Si is also used. These reference states were selected because they represent the silica-rich environments in which most sodium minerals generally

WAnhexample of how the solubility lines in Fig. 9.1 were developed is given for low albite :

log K‘

+ , NaAlSi308(low albite) + 4 H + 4 H+ 0 ^3 Na + Al + 3 H 4SiOg , * Al 3 + + H 4SiOS + 0.5 H O * ) + 3H + ( kaol ) 0.5 AI 2 Si 2O 5 (OH 4

=

=

* 2 H 4Si04 2 SiO,(soil ) + 4 H ,0

2.74

—

0.5( 5.45)

5

. .

+ NaAlSi308(]ow albite) + H + 0.2 H 20 ) (kaol ) + 2 S 02 (so l) Na + -t- 0.5 Al 2 Si 205( OH 4 + log Na

= 6.22 - pH

2( 3.10)

6.22 (9.1)

(9.2) 119

K> G

TABLE 9.1 EQUILIBRIUM REACTIONS OF SODIUM AND POTASSIUM MINERALS AND SOLUTION SPECIES Reaction

No.

log K °

Equilibrium Reaction

Na Aluminosilicates NaAISi 04 ( ncpheline) + 4 H + Na + + AIJ + + H 4SiOS NaAISijO,,( jadcite) + 4 H + + 2 H 2 O Na + + Al 3 + + 2 H 4 Si 02 NaAISi 206 H 20(analcimc) + 4 H + + H 20 ; Nu + '+ Al 3 + + 2 H 4 SiOJ NaAISi 30 B ( Na glass ) + 4 H + + 4 H 2 O Nn + + Al 3 + + 3 H 4 SiC> 2 NaAlSi308( liigl > albitc) + 4 H + + 4 H 2 0 ± N a + + Al 3 f + 3 H 4 SiOJ ^ NaAISijOu ( low albitc) + 4 H + + 4 HaO Na + + Al 3 * + 3 H 4 SiOJ Na * + 3 AI 3 + + 3 H 4SiOJ NaAI 2( AlSijO|0 )(OH ) 2 ( parugonite) + lOH Na0.33AI 2.33Si 3.07OI 0(OH )2( bcidelIitc) + 7.32 H + + 2.68 H 20 0.33 Na + + 2.33 AI 3 + + 3.67 H 4SiOJ

1 2 3 4

^ ^

^ ^ ^ ^=^

*

5 6

7 8

11.25 7.11 8.15 10,87

3.67 174 17.40 6.13

iJ
0

-2

-4 —

( 9) KAISi 04 (10) KAISi 206 (11), (12) (13 ) KAISi Og 3 (14- - KA ( AIS« 12 301o) (OH)2

.

>

K0.6 M90.25 AI 2.3si

-8 4

3.5010(OH)2

1

i

5

1

6

7 pH

.

8

9

J 10

Fig 9.2 The stability of potassium minerals in equilibrium with kaolinite, and when indicated, soil -Si (S ) with shifts to quartz ( QJ.

maintained with kaolinite and soil-Si, these minerals decrease in the order: KAlSi 308( K -glass) > KAlSi 04(kaliophilite) > Ko . 6 Mgo. 25 Al 2.3Si 3.5Olo(OH)2(illite) = KAlSi 206(leucite) > KAI 2 (AlSi 3O10) (OH )2( muscovite) > KAlSi 3Oa (sanidine) > KAlSi308(microcline).

These minerals are generally more soluble than the reference level of soil K, which is arbitrarily set at 10 3 M K. +. In alkaline soils, microcline, high sanidine, and muscovite are sufficiently stable to keep soluble K + from rising above the reference level. Again as H 4Si04 drops to quartz or lower, the lines of minerals having Si/Al > 1 rise, whereas only muscovite and kalio philite remain fixed. "

-

-

COMPLEXES OF SODIUM AND POTASSIUM

125

Potassium is known to be refixed in secondary clay minerals such as illite, yet the solubility value for illite used here (see the appendix) appears to be too soluble to permit its formation in soils (Fig. 9.2). Further details on the refixation of potassium by clay minerals must await more accurate determination of the free energy values of these minerals and the exchange reactions they undergo in soils. Such data is vitally needed to develop mean ingful equilibrium relationships of these important minerals in soils. Comparing Figs. 9.1 and 9.2 shows that potassium minerals in general are less soluble than sodium minerals. The rates of weathering of these minerals also depends upon particle size since chemical reactions are largely limited to surface reactions. [ 93 COMPLEXES OF SODIUM AND POTASSIUM

Several complexes of sodium and potassium in solution are given by Reactions 16 through 25 of Table 9.1 and are plotted in Figs. 9.3 and 9.4. 0

r

-2 Na" log SO

-4

'

log

1 -3

*

/

cr

NaSO/

- 3 NaCI *

6

'

CO

'

8 f -

-10 -12

-14

4

5

7 B 9 10 pH Fig. 9.3 Sodium complexes in equilibrium with 10's WNa * 6

.

.

CH 9 SODIUM AND

126

POTASSIUNJ

0

-2

K* log SO4 -2

-4

"

KSO4~

-3 6

£ Iu

f

'

8

’

.

-10

-12 -14

5

4

6

7

8

9

10

pH

Fig . 9.4

Potassium complexes in equilibrium with 10 - J M K ~ .

These plots were prepared on the basis of equilibrium with the reference levels of 10 3 M Na + or K + . All lines on these diagrams except Na C0$ 2 and K 2COJ change by one log unit with corresponding changes of K + or Na + . Since two Na + or K + ions are included in the latter complexes, they change by two log units. In general the complexes of Na and K + in well-drained soils are of only minor importance and can be fairly well ignored . These include sodium and potassium complexes ofSO “ , Cr , HC03“ , OH - , and CO . ”

^

9.4 REDOX

^

RELATIONSHIPS

Ind" 2

^=

Na + + e

"

°f

and

;

Na(c)

'

8iVe

"

log K°

= -45.89

M (9.4)

PROBLEMS

pe =

and

K+

+e

‘

=

±

pe

-

127

45.89 + log Na +

K(c)

—

(9.5)

log A' 0

—

= 49.49

= 49.49 + log Na +

1

(9.6) (9.7)

These redox relationships show that only the oxidation state of + 1 is important for these elements in aqueous or soil environments. Changes in redox can not affect sodium and potassium directly, only indirectly as it affects other constituents with which they react.

PROBLEMS 9.1 Develop equations and plot the solubility lines for high albite and paragonite in Fig. 9.1. How do these lines shift if equilibrium is main tained with Si 02(amorp) rather than soil -Si ? 9.2 Develop equations and plot the solubility line for microcline in terms of log K versus pH for equilibrium with a. Kaolinite and soil-Si. b. Kaolinite and quartz. c. Kaolinite and amorp-Si. d. Mg-montmorillonite, soil -Si, soil -Fe, and 10 3 M Mg2 + . Discuss these relationships in terms of the weathering of microcline in soils at different stages of weathering. 93 Under what conditions, if any, can beidellite be expected to form in soils. 9.4 Develop the relationships expressed in Eq. 9.3.

-

"

‘i

'

V

3

TEN

IRON

t

f Z ) ri •

e

y

I

ron is a major constituent of the lithosphere, comprising approximately

5.1 % ; the average content of soils is estimated at 3.8 % (Table 1.1 ). In primary minerals iron occurs largely as ferromagnesium minerals. During weathering these minerals dissolve (see olivine Fig. 8.1), and the released iron precipitates as ferric oxides and hydroxides. In this chapter it will be shown that the solubility of iron in soils is largely governed by Fe( III ) oxides while hydrolysis, complexation, and redox are important modifying factors. The solubility relationships of iron oxides are examined in great detail, whereas those of phosphates, iron chelates, and iron sulfides will be given in Chapters 12, 15, and 17, respectively.

log A

15.98 13.04 35.69 3-42 ..

-

43.75 12.42

10.1 SOLUBILITY OF Fe(Ill ) OXIDES IN SOILS

The solubilities of Fe(III) oxides commonly found in soils are given by Reactions 1 through 6 of Table 10.1. These equilibria are written as acid dissociation reactions, but they can easily be converted to basic dissociation reactions. Taking Fe(OH)3(soil ), for example : log K °

Fe(OH )3(soil) + 3 H + 3 H 20

Fe3 + + 3 H 20 3 H + + 3 QH -

Fe(OH)3(soil )

Fe3 + + 30 H ~

-

2.70

- 39.30

3( 14.00) ( 10.1)

Solubilities of the Fe(III) oxides are plotted in Fig. 10.1. These oxides decrease in solubility in the order : Fe(OH )3(amorp) > Fe(OH)3(soil) > y-Fe,03( maghemite) > y-FeOOH(lepidocrocite) > a-Fe203(hematite) > a-FeOOH(geothite).

- 6.74

- 16.04 - 31.99

- 46.38 -45.39

The activity of Fe3 + maintained by each of these oxides decreases 1000-fold for each unit increase in pH. The solubility of Fe 3 + maintained by Fe( OH )3 (amorp) is 103- 54 10 o 02 103 56 or 3631 times greater than that maintained / = by a-FeOOH (goethite). When soluble Fe(III ) salts are added to soils, Fe3 + readily precipitates. Within a few hours its solubility approaches that of by Fe(OH)3(amorp). To determine where in this wide range of solubilities Fe 3 + in soils might be, Norvell and Lindsay (1980) used the chelation method to determine the activity of Fe3 + in soils (see Chapter 15). Their findings indicate that soils generally maintain anrFS'3 ^activity slightly below that ofFe(OH )3

_

'

:

13.48 12.90 9.00 7.92 2.65 2.46 14.79 19.76

129

0.00

-0.07

-

0.03 2.70 3.60 2.20

31

IRON MINERALSAND TABLE 10.1 EQUILIBRIUM REACTIONS OF AT 25°C

Reaction No.

COMPLEX

Equilibrium Reaction

^

log

—

Fe( III ) Oxides and Hydroxides I 2 3 4 5 6

--

^

3.54 3.54 2.70 2.7C 1.59 1.39 0.09 -0.02

-

^ ^

Other Fe( III) Minerals ""1 7 8 9

\

Fe(0 H ) j(amorp) 4- 3 H + FeJ + 4 3 H 20 Fe( OH )3(soil) 4- 3 H * ;± Fc 3 * 4 3 H 20 F/- Fe203(maghemite) 4- 3 H * ^± Fe3 + 4 f H20 y- FeOOH( lepidocrooite ) 4- 3 H * ± Fe 3 + 4 2 H 20 . . -Fe 203( hem4tite) + 3 Pi + ^ Fe3 * +\ H,0 + i- FeOOHfgoethit?) + 3 H. Fe3 + 4- 2 H 20

^

——

’

*

FeCl 3( molysite ) ?± Fe3 + 4- 3 CT Fe 2(SOJ 3(c) 2 Fe3 + + 3 SOi KFe3(S04)2(OH )6( jarosite ) + 6 H + K * + 3 Fe3 + + 2 SOi

13.25 13.25 18? 2.89

"

'

+ 6 H 20

- 1151

Fe( III ) Hydrolysis 10 II 12 13 14

Fe3 + 4- H ,O FeOH 2 + + H + Fe3 + + 2 H ,0 tFe(0H )2+ + 2 H + ^ ) + 3H + Fe3 + + 3 H 20 ^± Fe(0 H| Fe3 + + 4 H 2 O Fe(OH ) + 4 Hr 2 Fe 3 + + 2 H 2 O Fe 2(OH ) + + 2 H +

^

r t

^

^

- 119

5.69 -13.09

-21.59 190 -

Fe( III ) Complexes

15 16 17

18 19 20 21 22 23 24 25 26 27

Fe

3*

+ Cl Fe + 2 Cr Fe3 * + 3Cr Fe3 + + Br

"

3+

"

3*

S

\

1.48

2+

^ FeCl FeCU ^ FeCI | +

^ FeBr * ^ FeBr5 ^ FeF ^ FeF * 2

Fe 4 3 Br 2 Fe3 * + F ; ~ 3* Fe + 2 F 2 Fe 3* + 3 F “ FeF | Fe 3 * 4- N03 FeN0|+ Fe 3 + + SOi FeSO; Fe3 + + 2 S0 Fe(S04)2 3+ Fe + H 2 P04 :?iFeH 2 P0i + Fe3 + + H 2 PO; ± FeHPO; + H + ^ “

"

_^

^ r- ^^ '

'

113 0.77 0.60 -0.04 6.00

9.20 11.70 1.00 4.15 5.38

* -

4

i

, J

jj

5.43

3.7J

130

A

r

V

TABLE

( Continued )

10.1

Reaction No .

Equilibrium Reaction

log A’

Redox Reactions 28

29 30. 31 .

32 33

Fc(c ) 3’

^ Fe Fe

24

+ 2e

“

Fe 4 e Fe 304( magnetite) + 8 H 4 3 Fc 2 + + 4 H 20 + 2e ' Fe 3Q4( magnetite ) + 8 H 4 4- 04- ^ Fc 4 SO; ^ FcSOi 2+ 2+

24

"

,

0.00

- 0.07 0.03 2.70 - 3.60 2.20

131

i

CH. 10

132

IRON

-6 Fe(OH )3 ( amorp)

-8 Fe (OH )3 ( Soil - Fe )

-10 y - Ft? 2

) 3 ( maghemite

°

-12

y- FeOOH (lepidocrocite)

A &

—

-14

-16

- 18

a - Fe2C>3 (hematite)

- 20 - 22

a - FeOOH (goethite)

- 24

4

5

6

7

a. 8

9

pH

.

Fig 10.1 The activity of

-

Fei maintained by Fe( III) oxides and soil Fe. +

(amorp) as represented by Fe(OH)3(soil ) ( Reaction 2, Table 10.1). When they added soluble FeCl 3 to their soils, the activity of Fe3 + temporarily

increased, but over a period of several weeks, it slowly approached the solubility of Fe(OH)3(soil). Throughout this text Fe(OH)3(soil) or soil'*; is used as a reference solid phase controlling the solubility for Fe 3 + in 501 •’ Rather than think of soil-Fe as a discrete crystalline phase, it can be thoug 8 of as an amorphous phase having a greater degree of structural order th ^ freshly precipitated Fe(OH)3(amorp), hence it has a slightly more negatlV^ AGy (see the appendix).

OTHER Fe( lll > MINERAL

^ crystalline Since the

tn

-

iron oxides in Fig, 10.1 fall below nofl Fc, they are more stable and cun be expected to crystallize slowly. The activity of Vt* * maintained by these iron minerals is very low ; consequently at high pH transformations from one mineral to another and attainment of final

equilibrium is an extremely slow process, Sfihwcrtrnann and Taylor (1977) have discussed some of the conditions believed to affect the formation of various forms or Fe(lll ) oxides in sbils. Maghemite ( y Fe 203) is a magnetic mineral and can be separated from soils with a magnet. The percent composition of varioiik iron oxides in a soil may have little bearing on the Fe 3 ,t activity maintained by that soil. Soils generally contain some of several different iron oxides. The solubility of Fe 3 + is usually controlled by the most soluble oxide present. For this reason soil - Fe generally controls the activity of Fc 3 * in most soils. Only in well -drained , highly weathered soils not subject to frequent reductions are hematite and gocthite expected to lower the solubility of Fe 3 * toward their equilibrium levels. The solubilities of hematite and gocthite arc nearly identical, but goethite is generally considered the ultimate weathering product

-

.

of iron in soils From Fig. 10.1 it is easy to see why iron deficiencies arc more prevalent on alkaline soils than on acid soils, and why highly weathered soils are lower in available iron than less weathered soils. “ Limonite,” once considered a common form of iron oxide in soils, is now recognized as finely divided goethite with adsorbed water.

10.2 OTHER Fe( III ) MINERALS

The solubilities of other Fe(lII ) minerals arc given by Reactions 7 through 9 ofTable 10.1. The solubilities of FeCI 3( molysitc) and Fe 3(S04)2(c) are much too high to permit their formation in soils. The mineral KFe 3(S04 )2 (OH )6 ( jarosite) is often found in acid sulfate soils ( Breemen, 1976). The con . Only in ditions necessary for its precipitation are depicted in "Fig. 10.2 + ; SO on and depending K activities, soils below pH 4,0,. or slightly higher can jarosite form. Such conditions are found in acid sulfate soils containing ' FeS2 ( pyrite), which oxidizes to give sulfuric acid and precipitated Fe(lII ) oxides. Figure 10.2 indicates that in acid pH range of soils jarosite can form gpothite. Release of from OHUlmorp') and Fe(O.H)3(sQjl ) byt not from + Fe 2 + from pyfite and its subsequent oxidation to Fcs favor jarosite forma tion. Potassium ions are also needed to precipitate jarosite, but they are by Na , relatively minor constituents that can be replaced to some extent H 30 + , or other cations ( Breemen, 1976)

-

'

^

-

.

CH. 10 IRQ

*

134

-© -8

-10 -12

% ,