Understanding Minerals & Crystals
 9781431700844, 1431700843

Citation preview

Understanding

Minerals & Crystals Bruce Cairncross Terence McCarthy

Published by Struik Nature (an imprint of Penguin Random House South Africa (Pty) Ltd) Reg. No. 1953/000441/07 The Estuaries No 4, Century Avenue (Oxbow Crescent), Century City, 7441 PO Box 1144, Cape Town, 8000 South Africa Visit www.randomstruik.co.za and join the Struik Nature Club for updates, news, events and special offers. First published in 2015 10 9 8 7 6 5 4 3 2 1 Copyright © in text, 2015: Bruce Cairncross and Terence McCarthy Copyright © in photographs, 2015: Bruce Cairncross, unless otherwise indicated Copyright © in illustrations, 2015: Terence McCarthy Copyright © in published edition, 2015: Penguin Random House South Africa (Pty) Ltd Publisher: Pippa Parker Managing editor: Helen de Villiers Editor and indexer: Tina Mössmer Project manager: Colette Alves Illustrator: Colin Bleach Design director: Janice Evans Design assistance: Neil Bester, Deirdré Geldenhuys Designer: Dominic Robson Proofreader: Thea Grobbelaar Printing and binding: Times Offset (M) Sdn Bhd, Malaysia All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the prior written permission of the copyright owner(s). ISBN: 978 1 43170 084 4 e-Pub: 978 1 77584 333 7 e-PDF: 978 1 77584 334 4 Front cover: Chalcopyrite Back cover (clockwise from top left): Quartz Citrine crystal; Fluorite; Vanadanite; Quartz crystal (see text for details) Penguin Random House independently submitted this book for critical peer review prior to its publication. This independent peer review was co-ordinated by Pippa Parker: Publisher, Struik Nature. The reviewers were: Professor Johan de Villiers Department of Materials Science and Metallurgical Engineering Hatfield Campus, University of Pretoria Pretoria, South Africa Dr Lynnette Greyling Department of Geological Sciences University of Cape Town Private Bag X3 Rondebosch 7701 Cape Town, South Africa

Note: The authors and the publishers cannot be held liable for any accident or injury resulting from the use of hydrochloric acid as an aid in mineral identification.

Sponsor’s foreword Newcomers to mineral collecting and those seeking to know what minerals are, how they form, why they differ in their physical properties and how they should go about identifying them will find this book useful. But so, too, will experienced collectors who wish to know more about the theory behind mineral formation. Students enrolled for introductory mineralogy courses at universities across the world will find the book particularly valuable. But why should it be of interest to you? The answer is that minerals form the basic building blocks of our Earth, and many are important for the health of living species, including humans, as they help in skeleton and tooth formation, blood coagulation and muscle contraction. Minerals are utilized extensively in modern societies – as a source of power and in our concrete buildings and bridges, the cabling in our telecommunication systems, cars, food and beverage cans and much more. Minerals power much of the global economy. In South Africa, the extraction of natural resources and related activities accounts for about 20% of the country’s GDP, just under 50% of the country’s merchandise exports, and 35% of the value of the JSE – and also generates foreign exchange. About 90% of South Africa’s total electricity is derived from power stations fuelled by locally mined coal. Mining directly employs half a million people and another half a million indirectly, accounting for about one (1) million of the eight (8) million people formally employed in the country. The mining industry returns billions to the economy in operating costs, salaries and wages, capital expenditure, taxes, royalties, dividends and interest. The mining industry, and the minerals on which it relies, has significant environmental and human consequences. As a mineral resource extraction company, Exxaro is especially aware of these issues, and the need to understand them, and we therefore welcome the publication of a book that explains the science behind mineralogy in non-technical language, looks attractive and has appeal to a wider audience.

Sipho Nkosi Chief executive officer Exxaro Resources Limited www.exxaro.com

contents 1 Introduction 2

3

4

Atoms: The building blocks of minerals

6 9

The nucleus

10

The distribution of electrons around the nucleus

11

The Periodic Table

13

The Periodic Table and the properties of atoms

14

Abundances of elements

18

The abundances of elements on Earth

20

Chemical bonding and the formation of minerals

22

The covalent bond

22

The ionic bond

25

Bonds intermediate between covalent and ionic

25

The metallic bond

26

Van der Waals bonds

28

Mixed bonding

29

The formation of crystalline solids

31

The structure of ionic and metallic bonded solids

31

The structure of metals

35

Structures with ionic bonding

36

Structures with covalent bonding

38

Structures with mixed bonding

39

The origin of crystal faces

42

Ionic substitution in minerals

46

Polymorphism 49 Crystal growth

51

Twinned crystals

55

Pseudomorphism 56

5

The morphology of crystals

58

Natural crystals vs crystal models

58

Elements of symmetry

59

Crystallographic systems

60

Common crystal classes

67

Describing crystals

73

Twinned crystals

73

6

The physical characteristics of minerals

78

Density and specific gravity 78 Properties that depend on light 79 Mechanical properties 86 Magnetic properties 90 Habit 91 Reaction with hydrochloric acid 97

7

Mineral classification and identification

98

Classification 98 Identification 98 Key to minerals 99

8

Native elements

110

9

Sulphides

122

10 Oxides and hydroxides

145

11 Halides

167

12 Carbonates

172

13 Sulphates

200

14 Phosphates and vanadates

209

15 Tungstates

216

16 Silicates

220

Nesosilicates 220 Sorosilicates 237 Cyclosilicates 240 Inosilicates 249 Phyllosilicates 263 Tectosilicates 276

17 Mineraloids

292

Appendix 1: Abundances of the elements 294 Appendix 2: Making your own crystal models 296 Further reading and sources 302 photographic credits 302 acknowledgements 303 glossary 303 index 308

1

Introduction

Chemical elements such as iron, aluminium, oxygen, carbon and many others are familiar to us all, but less well known is the form in which these substances occur in Nature. Chemical elements in fact occur in the form of minerals, which make up the solid Earth’s crust. A mineral can be broadly defined as a naturally occurring chemical compound of inorganic origin. Mineralogists generally add an additional qualifier, namely that a mineral should have a definite, regular internal atomic structure. This qualifier separates minerals, which are crystalline solids, from non-crystalline materials, such as amber (fossilized tree resin) for example (figure 1.1), which are termed mineraloids. Minerals can be composed of many chemical elements in combination or they may consist of only one element. Nevertheless, all have a definite, well-defined chemical composition, which is usually expressed as a chemical formula such as NaCl (sodium chloride) for the mineral halite (figure 1.2).

Minerals are extremely important because they are the raw materials from which we extract essential commodities, such as iron (figure 1.3), manganese, silicon or aluminium, on which our civilization depends. We also use some minerals in their natural state because of their particular chemical or physical properties, for example agricultural lime, which is very important in crop production, or talc, an essential ingredient of talcum powder. Perhaps even more importantly, minerals are the building blocks from which rocks are made and thus form the solid parts of the entire Earth (figure 1.4). We are surrounded by minerals! To really understand the Earth and how it works, knowledge of minerals is absolutely essential. Little wonder the study of geology begins with mineralogy, the study of minerals. Minerals vary widely in their chemical constitution and their physical properties, such as hardness and colour. Some form very attractive crystals (figure 1.5) and most natural history museums have a

Figure 1.1 Amber, fossilized tree resin, sometimes contains entrapped insects (Dominican Republic, 6 cm).

Figure 1.2 Halite crystals (Namibia, 5.8 cm).

6

section devoted to the display of such crystals. But well-formed crystals are rare and most mineral specimens we come across are simply irregular lumps. They are nevertheless crystalline in that they have regular internal order in their atomic structures, but this order is not immediately apparent as the samples have an irregular external shape. The ability to identify the more common minerals is an essential skill for anyone interested in geology. Moreover, mineral collecting is a popular pastime in its own right and collectors need to be able to identify mineral specimens. So how does one go about identifying a mineral? There are many different ways to do this. The more sophisticated methods involve examining thinly sliced samples under specially designed petrographic

Figure 1.4 A slab of the common rock type granite in which individual mineral grains are clearly visible (field of view 8 cm). Figure 1.3 An opencast iron ore mine (Postmasburg, South Africa).

7

microscopes, or chemically analyzing the sample or even probing its crystalline structure using X-rays. Naturally, these require very sophisticated laboratory equipment and are only used by the professional mineralogist. At the other end of the spectrum is what is usually termed ‘hand specimen’ mineralogy. The term implies identifying mineral specimens without recourse to sophisticated laboratory equipment. Hand specimen mineral identification relies on a few very simple tests, which can be carried out just about anywhere, as they require the most rudimentary equipment. Such methods are widely used by hobbyists and even by practising geologists in the course of their daily work in mines or in the field, where rapid identification of minerals is needed in the absence of laboratory equipment.

While mineral identification is a useful skill, it is also helpful to understand the principles underlying the formation of minerals, which are responsible for the differing properties and crystal forms of minerals. In this book we begin at the very beginning – by examining the nature of atoms and the way they bind together to form minerals with distinctive crystal structures. We introduce the reader to the nature and classification of these crystals. Finally, we examine the physical properties of minerals and how these can be used in identification. We hope that this book will not only help you learn how to identify minerals, but that it will equip you with a fairly good understanding of the nature of minerals at the atomic scale and why their properties differ as they do.

Figure 1.5 A selection of well-crystallized minerals illustrating a variety of habits and colours (various southern African specimens, front centre 13 cm).

8

2

Atoms: The BUILDING blocks of minerals

Our world is made of the most astonishing variety of substances. The Earth is enveloped in a gaseous atmosphere and two-thirds of its surface is submerged beneath a liquid ocean that contains an assortment of dissolved materials. The ocean floors and the continents are made of a plethora of different rock types, such as basalt, granite, limestone, sandstone and mudstone, each of which is composed of several different minerals. Most of the rocks of the continental surface are buried beneath soils of many different kinds. The Earth is home to millions of different life forms, each with its own unique characteristics. Centuries ago, alchemists, the forerunners of modern-day chemists, began to investigate natural materials; many were driven by the notion that commonplace materials could be transmuted into gold. Gold remained elusive, however, but in their quest the early chemists discovered that with few exceptions, natural materials could be broken down into simpler substances, but only up to a certain point. Beyond this, further subdivision was not possible. Progress was erratic and there were many false leads. However, chemists began to realize that these indivisible substances were the elemental materials from which all other substances were made and the notion of chemical elements was born. By the early 1800s, the idea that elements consisted of atoms with measurable relative masses began to take shape, notably through the work of John Dalton and Amedeo Avogadro. The number of known elements, their relative masses and their chemical properties grew steadily during the nineteenth century. Some sense of order was sought amongst this growing population. It was noticed that when

the elements were arranged in order of increasing relative mass, there seemed to be a periodicity in their chemical properties. In 1869, Dmitri Mendeleev hit upon the idea of arranging the 64 elements then known in order of increasing mass, but instead of a continuous sequence, he tabulated them according to the periodicity of their chemical properties, thus creating the first Periodic Table of the elements. So confident was Mendeleev that his discovery reflected the natural order of the elements, that he predicted that gaps in his tabulation would be filled by elements yet to be discovered and even went so far as to predict their chemical behaviour. He was vindicated in his own lifetime with the discovery of the elements gallium, scandium and germanium, which slotted neatly into gaps in his table. In the years following Mendeleev’s seminal contribution, further elements were discovered, reaching 92 in all. Since then, the list has grown with the synthetic manufacture of several elements in nuclear physics laboratories, but none of these is known to occur naturally. The vast array of substances that make the world around us, and ourselves too, is astonishing, but even more so with the realization that all of it is made from so few elements. In fact, only about 15 of the 92 naturally occurring elements make up most of the familiar materials around us – but more of this later. The discovery of radioactivity and X-rays in the early 1900s resulted in further advances in the understanding of atoms. Henry Moseley discovered that there is a very precise relationship between the energy of X-rays emitted by atoms and their position on the Periodic Table. He was able to number the elements in a sequence based on the energy of their

9

X-ray emissions, thus defining the atomic number for the first time. Using the newly discovered radioactivity, physicists (notably Ernest Rutherford) began probing the internal structure of atoms. This research led to the notion that atoms consist of a small, positively charged nucleus surrounded by electrons. The nucleus contains positively charged particles, named protons, and accommodates virtually all of the atom’s mass. With this discovery, the significance of Moseley’s atomic number became apparent – it is in fact the number of protons in the nucleus. The final piece in the puzzle was the discovery of the neutron in the early 1930s, following which the modern concept of the nature of the atom rapidly took shape.

The nucleus At the centre of an atom lies the nucleus, which is made up of positively charged protons and uncharged neutrons, tightly clustered together (figure 2.1). Particles of the same charge repel each other in the same way that magnets repel each other when like poles are brought together. In contrast, oppositely charged particles attract each other, like opposite poles of magnets. The positively charged protons in the nucleus should, therefore, experience mutual repulsion, making the nucleus unstable. However, nuclei

proton

Figure 2.1 The atomic nucleus.

10

neutron

are generally stable because nuclear particles possess a powerful attractive force that operates exclusively at a very close range, which is known as the strong nuclear force. Notwithstanding this force, the like charges of protons still result in some net mutual repulsion, but the presence of the strong nuclear force of the neutrons overcomes this repulsion, making the nucleus stable. Neutrons are therefore essential as they help to hold the nucleus together. The number of neutrons is generally equal to or exceeds the number of protons in the nucleus of an atom. The atomic nucleus is surrounded at some distance by electrons, which carry a negative charge equal but opposite to the charge of a proton. The number of electrons equals the number of protons, making atoms electrically neutral. Protons and neutrons have almost the same mass and are about 2,000 times heavier than electrons. The mass of an atom as a whole therefore is almost entirely due to the mass of its nucleus. The number of protons in the nucleus is termed the atomic number (Z) and the number of protons plus neutrons is termed the mass number. The number of neutrons in a nucleus with a given atomic number can vary; for example, nuclei containing two protons can have either one or two neutrons. Atoms with the same number of protons (atomic number) but different numbers of neutrons are known as isotopes. The notion that elements differ in their atomic masses is more than two centuries old. The relative atomic mass or relative atomic weight is nowadays expressed by relating the mass of an atom to the mass of a carbon atom, which is assigned a mass of 12 (12 atomic mass units). Atomic mass is a bulk property, which means that it represents the average atomic mass of all of the isotopes of that element, which is the main reason why atomic masses are not integer numbers.

The distribution of electrons around the nucleus Electrons surround the nucleus. The outer electrons of an atom, termed valence electrons, may interact with outer electrons of other atoms and the manner in which the atoms interact varies with the number of electrons. These interactions define the chemical properties of an atom. Since the number of outer electrons is determined by the number of protons in the nucleus, it is the atomic number that ultimately determines the chemical properties of an element. Each proton in the nucleus is matched by an electron residing outside of the nucleus. Obviously, the higher the atomic number, the larger the number of electrons surrounding the nucleus. The study of the distribution of electrons around nuclei has revealed some remarkable surprises and gave rise to the field of quantum mechanics. This is an arcane discipline, but at its core is the discovery that the behaviour of electrons (and indeed, all things at and below the scale of atoms) is quite unlike what one might expect from everyday experience. Although the analogy is sometimes made between electrons orbiting the nucleus and planets orbiting the Sun, this notion is actually incorrect. Electrons move rapidly around a nucleus but rather than follow predictable orbits

like planets, can actually occur anywhere in relation to the nucleus (even inside it!). However, there are certain regions of space near the nucleus where they spend most of the time. These regions can be imagined as clouds created by the electrons (like the blur created by rapidly spinning helicopter blades, figure 2.2). When atoms are brought close together, their respective electron clouds repel each other, creating the impression that the atoms are rigid bodies, in spite of the fact that they consist mainly of empty space. The extent of the electron cloud around an atom is typically about 10,000 times larger than the nucleus. The shapes of the regions most frequently occupied by electrons are reasonably well defined and are called orbitals. Electrons occupy four types of orbital: s orbitals, which are essentially spherical (figure 2.3A), and p, d and f orbitals (figure 2.3B, C), which are shaped more or less like dumbbells in various combinations (the details need not concern us too much here). The p orbitals actually exist as a set of three equivalent orbitals occupied by electrons, the d as a set of five and the f as a set of seven orbitals. The shapes of orbitals are derived mathematically and cannot actually be seen. Only two electrons can occupy each orbital. As is evident from figure 2.3D, the orbitals overlap one another in complex

Ken Lovell

Figure 2.2 Rapidly moving electrons create a cloud, giving form to atoms, akin to rapidly rotating helicopter blades.

11

The single s orbital

The three p orbitals

(A)

(B)

y x

x

y

z z z

The five d orbitals (D)

(C)

y x

x

y

z Figure 2.3 The shapes and arrangement of the (A) s, (B) p and (C) d orbitals. The seven f orbitals are extremely complex and are not shown. Orbitals interpenetrate each other, as illustrated in (D) by the s and p orbitals. ways and the spatial distribution of all the electrons around a nucleus is therefore impossible to visualize except for the simplest of atoms. How electrons know to which orbital they belong is one of the many mysteries of quantum mechanics. Orbitals are grouped together into layers or shells, referred to by their so-called principal quantum number (n), from 1 to 6. The shells are also often referred to alphabetically as the K, L, M, N, and so on. Quantum number n=1 (K shell) corresponds to the lowest energy shell and lies closest to the nucleus. It contains only a single s orbital and can therefore host a maximum of two electrons. The next shell, with n=2 (L shell), hosts a single s orbital and three p orbitals and can therefore hold a maximum of eight electrons (2s+6p electrons). This

12

is followed by n=3 (M shell), which contains one s, three p and five d orbitals and can hold a maximum of 18 electrons (2s+6p+10d electrons); n=4 (N shell), which contains one s, three p, five d and seven f orbitals and can hold a maximum of 32 electrons (2s+6p+10d+14f electrons); n=5 and n=6, which have the same arrangement of orbitals as the N shell (maximum of 32 electrons ), and finally n=7, where only the s orbital is known to exist. Electrons occupying successively higher shells lie progressively further from the nucleus and thus the effective size of atoms increases with increasing shell occupancy (and atomic number). Electrons will always occupy the lowest available orbital (unless they are given additional energy that could elevate them to a higher orbital). Hydrogen, for example, has only one electron, which resides in

the 1s orbital. If supplied with energy, however, the electron will move to a higher energy orbital. A peculiar property of electrons is that the move from a lower to a higher energy orbital is not a progressive, cumulative move, like ascending a steep hill on a bicycle, but rather occurs in a single jump. To accomplish this move, the electron has to be supplied with a certain minimum quantity of energy, known as a quantum. The energy levels are therefore said to be quantized. The sequence of increasing energy of the orbitals and hence the order in which they are occupied as atomic number increases is not quite regular as one might expect. The sequence commences with the 1s orbital, followed by 2s, 2p, 3s, 3p, then 4s, 3d, and so on, as schematically illustrated in figure 2.4.

The Periodic Table This discussion of orbitals and their electrons may seem rather complicated and academic. However, the topic is very important because it underpins the fundamental classification of the elements in the Periodic Table (figure 2.5; see also Appendix 1). In the Table, elements are listed in order of increasing atomic number but are also grouped according to which orbitals are filled, following the sequence shown in figure 2.4. Thus, the first row or period represents the filling of the first shell (n=1, K shell), which can host just two electrons – the first position corresponding to the element hydrogen (atomic number = 1) and the second to the element helium (atomic number = 2). The next row corresponds to the filling of the second shell (n=2, L shell), which is made up of an s orbital (maximum of two electrons) and three p orbitals (maximum of six electrons) and thus has two elements on the left (s orbital filling) and six on the right (p orbital filling). Row three represents the filling of the next set of orbitals, and so forth, as shown in figure 2.4.

As a consequence of this form of representation, the elements fall into four blocks on the Periodic Table (figure 2.5): on the left is a block consisting of two columns, which represents the filling of the s orbitals (known as the s block); on the right is a block consisting of six columns, which represents the filling of the three p orbitals (the p block); in the centre is a block composed of ten columns, representing the filling of the five d orbitals (the d block, or more commonly referred to as the transition elements); and finally there is a block at the bottom of the table with fourteen columns representing the filling of the seven f orbitals (the f block; elements in the first row of this block are also referred to as the lanthanides and the second as the actinides, after lanthanum and actinium). The anomalous position of helium in the table (i.e. in the p rather than the s block) is discussed under the next heading.

A B

1s

C 2p

2s

D E

3s

3p

3d

4s

4p

4d

5s

6s

5p

6p

F G

5d

H

4f 5f

6d

7s Figure 2.4 A schematic diagram indicating the sequence in which orbitals are filled.

13

Groups Periods 1

1s

1

18

1 H

2 He

2 4 Be

s block p block d block f block

13

14

15

16

17

5 B

6 C

7 N

8 O

9 F

10 Ne

13 Al

14 Si

15 P

16 S

17 Cl

18 Ar

2

2s2p

3 Li

3

3s3p

11 Na

12 Mg

3

4

5

6

7

8

9

10

11

12

19 K

20 Ca

21 Sc

22 Ti

23 V

24 Cr

25 Mn

26 Fe

27 Co

28 Ni

29 Cu

30 Zn

31 Ga

32 Ge

33 As

34 Se

35 Br

36 Kr

37 Rb

38 Sr

39 Y

40 Zr

41 Nb

42 Mo

43 Tc

44 Ru

45 Rh

46 Pd

47 Ag

48 Cd

49 In

50 Sn

51 Sb

52 Te

53 I

54 Xe

55 Cs

56 Ba

57 La

72 Hf

73 Ta

74 W

75 Re

76 Os

77 Ir

78 Pt

79 Au

80 Hg

81 Tl

82 Pb

83 Bi

84 Po

85 At

86 Rn

87 Fr

88 Ra

89 Ac

4

4s3d 4p

5

5s4d 5p

6

6s4f 5d6p

7

7s5f 6d

4f

5f

58 Ce

59 Pr

60 Nd

61 Pm

62 Sm

63 Eu

64 Gd

65 Tb

66 Dy

67 Ho

90 Th

91 Pa

92 U

93 Np

94 Pu

95 Am

96 Cm

97 Bk

98 Cf

99 Es

68 Er

69 Tm

70 Yb

71 Lu

100 101 102 103 Fm Md No Lr

Figure 2.5 The Periodic Table showing the orbitals filled in each row, the atomic numbers and element symbols. The names and atomic masses of the elements are listed in Appendix 1. Perhaps the most significant feature of the Periodic Table is that all the elements in the same column (referred to as a group and numbered as shown in figure 2.5) have the same number and arrangement of outermost electrons. It will be recalled that the chemical properties of an atom are determined by the manner in which the outer electrons interact with electrons of other atoms. Atoms with the same arrangement of outer electrons might therefore be expected to show similar chemical properties, which is indeed the case. It was these similarities in chemical properties, combined with knowledge of the relative atomic masses of the elements known at the time, that enabled Dmitri Mendeleev to formulate the original Periodic Table. The significance of his tabulation of the elements only became evident decades later with the discovery of the nucleus, protons, electrons, neutrons and the nature of orbitals.

14

The Periodic Table is an extremely powerful tool for bringing together diverse information about the different chemical elements in a way that allows for easy synthesis and hence better understanding of their chemical properties.

The Periodic Table and the properties of atoms The most striking example of the similar chemical behaviour of elements in the same group is afforded by the so-called noble gases (Group 18). All of these elements have completely filled s and p orbitals and this configuration of electrons is extremely stable. The noble gases hold their electrons very tightly and do not react chemically – hence the term ‘noble’. Helium is somewhat of an anomaly. It has a filled s orbital but no p orbital electrons, yet it is also extremely unreactive and hence is placed in Group 18. In contrast to helium, other elements with filled s

1

2

3

1s

2s2p

3s3p

4

4s3d 4p

5

5s4d 5p

6

6s4f 5d6p

7

7s5f 6d

1

18

1 H

2 He

32

2

13

14

15

16

17

93

3 Li

4 Be

5 B

6 C

7 N

8 O

9 F

10 Ne

128 90

82

77

75

73 72

71

12 Mg

13 Al

14 Si

15 P

16 S

17 Cl

18 Ar

98

11 Na

154 136 19 K

20 Ca

3

4

5

6

7

8

9

10

11

12 118 111 106 102 99

21 Sc

22 Ti

23 V

24 Cr

25 Mn

26 Fe

27 Co

28 Ni

29 Cu

30 Zn

31 Ga

32 Ge

33 As

34 Se

35 Br

36 Kr

203 174 144 132 122 118 117 117 116 115 117 125 126 122 120 116 114 112 37 Rb

38 Sr

39 Y

41 Nb

40 Zr

42 Mo

44 Ru

43 Tc

47 Ag

46 Pd

45 Rh

48 Cd

50 Sn

49 In

51 Sb

52 Te

54 Xe

53 I

216 191 162 145 134 130 127 125 125 128 134 148 144 141 140 136 133 131 55 Cs

56 Ba

57 La

73 Ta

72 Hf

75 Re

74 W

76 Os

77 Ir

80 Hg

79 Au

78 Pt

82 Pb

81 Tl

84 Po

83 Bi

85 At

86 Rn

235 198 169 144 134 130 128 126 127 130 134 149 148 147 146 146 145 87 Fr

88 Ra

89 Ac

-

-

4f

5f

58 Ce

59 Pr

60 Nd

61 Pm

62 Sm

63 Eu

64 Gd

65 Tb

66 Dy

67 Ho

68 Er

69 Tm

70 Yb

-

71 Lu

165 165 164 163 162 185 161 159 159 158 157 158 174 156 90 Th

165

91 Pa

-

92 U

138

93 Np

-

94 Pu

-

95 Am

-

96 Cm

-

97 Bk

-

98 Cf

-

99 Es

-

100 Fm

-

101 Md

-

102 No

-

103 Lr

-

Figure 2.6 Atomic radii of the elements (in pm=10-12m). orbitals, such as magnesium and calcium, are quite different and very reactive, losing electrons very readily. The unreactive nature of helium is ascribed to the fact that its filled s orbital also represents a filled K shell. Although elements in the same group have the same arrangement of outer electrons, they differ in that atoms with higher atomic numbers in the group are larger because of the filled orbits inside the atoms. Consequently, the size of the atoms (see How big is an atom? p20) increases down a group (figure 2.6). In general, the larger the atom, the further the outer electrons are from the nucleus and consequently the less tightly they are held by the positively charged nucleus. The number and arrangement of outer electrons is important in influencing chemical properties, but so too is the strength with which the outer electrons are held by the nucleus. Therefore, whereas chemical properties of elements in the same group are generally similar, they

change systematically down the group as the size of the atoms increases. The electron configuration of the noble gases in which both the s and p orbitals are filled is particularly stable. This enhanced stability is also echoed in the elements neighbouring the noble gases. Elements that have a few electrons less than the noble gas arrangement will tend to take up extra electrons to fill the vacancies in the p orbitals to attain the noble gas electron arrangement, thereby becoming negatively charged due to the extra electrons. Such negatively charged atoms are referred to as anions. On the other hand, elements that have a few more electrons than the noble gas configuration will tend to shed the extra electrons to attain a stable electron arrangement, thereby becoming positively charged. Such positively charged atoms are called cations. Elements in Group 1 of the Periodic Table, which have one more electron than the noble gas configuration, readily lose that electron and always form

15

1

2

3

4

5

6

7

1s

2s2p

3s3p 4s3d 4p 5s4d 5p 6s4f 5d6p 7s5f 6d

1

18

1 H

2 He

2.1

2

3 Li

4 Be

Metals Semi-metals Non-metals Noble gases

1.0 1.5 11 Na

12 Mg

0.9 1.2 19 K

20 Ca

13

14

15

16

17

-

5 B

6 C

7 N

8 O

9 F

10 Ne

-

2.0 2.5 3.0 3.5 4.0 14 Si

13 Al

3

4

5

6

7

8

9

10

11

12

21 Sc

22 Ti

23 V

24 Cr

25 Mn

26 Fe

27 Co

28 Ni

29 Cu

30 Zn

16 S

15 P

17 Cl

18 Ar

-

1.5 1.8 2.1 2.5 3.0 31 Ga

32 Ge

33 As

34 Se

35 Br

36 Kr

-

0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 37 Rb

38 Sr

39 Y

41 Nb

40 Zr

42 Mo

44 Ru

43 Tc

47 Ag

46 Pd

45 Rh

48 Cd

50 Sn

49 In

51 Sb

52 Te

54 Xe

53 I

-

0.8 1.0 1.3 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 55 Cs

56 Ba

57 La

73 Ta

72 Hf

75 Re

74 W

76 Os

77 Ir

80 Hg

79 Au

78 Pt

82 Pb

81 Tl

84 Po

83 Bi

85 At

86 Rn

-

0.7 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 87 Fr

88 Ra

89 Ac

0.7 0.9 1.1 4f 5f

58 Ce

59 Pr

60 Nd

61 Pm

62 Sm

63 Eu

64 Gd

65 Tb

66 Dy

67 Ho

68 Er

69 Tm

70 Yb

71 Lu

1.1 1.1 1.1 1.1 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 90 Th

91 Pa

92 U

1.1 1.3 1.5

93 Np

94 Pu

95 Am

96 Cm

97 Bk

98 Cf

99 Es

100 Fm

101 Md

102 No

103 Lr

-

-

-

-

-

-

-

-

-

-

-

Figure 2.7 Pauling’s electronegativities of the elements. singly charged cations; those in Group 2 have two more electrons than a noble gas arrangement and form cations with a charge of +2; those in Group 3 form cations with a charge of +3; while those in Group 4 form cations with a charge of +4. Cationic charges more positive than +4 are uncommon. Elements in Group 17 (known as the halogens) are one electron short of the noble gas configuration and accordingly take on an extra electron, resulting in an anion with a charge of -1. Elements in Group 16 are two electrons short and form anions with a charge of -2. The situation with elements that are three or more electrons short of the noble gas configuration is complicated by other factors, as is the behaviour of heavier elements in the p block, and will be discussed later. What makes some atoms lose electrons to form cations whereas others gain electrons to form anions is ultimately determined by the strength with which

16

the nuclei attract and hold their electrons. Many methods have been devised to express this attraction quantitatively. One of these, known as electronegativity, will be used here. The Nobel Prizewinning chemist Linus Pauling defined electronegativity as the ability of an atom to attract electrons to itself. The electronegativities of the elements are shown on the Periodic Table in figure 2.7. There are many ways of calculating electronegativity and each method has its shortcomings. Therefore, the values listed in figure 2.7 are not fully quantitative and should be seen only as a general guide. Electronegativities have not been calculated for the noble gases (Group 18) because their outer electrons are tightly held and they neither lose nor attract electrons. The highest electronegativity is that of the element fluorine (Z=9), with a value of 4.0. The lowest is that of caesium (Z=55), with a value of 0.7. Fluorine is an extremely powerful attractor of electrons, making it a very dangerous substance.

1 1

1s

18 2 He

1 H 1-

2

13

14

15

16

17

3 Li

4 Be

5 B

6 C

7 N

8 O

9 F

3+ 23

4+ 16

3+ 16

13 Al

14 Si

15 P

2

2s2p

3

3s3p

4

4s3d 4p

1+ 133 2+ 99

5

5s4d 5p

1+ 147 2+ 112 3+ 92

6

6s4f 5d6p

1+ 167 2+ 134 3+ 114 4+ 78

7

7s5f 6d

1+ 68

11 Na

2+

35

12 Mg

1+ 97

2+ 66

19 K

20 Ca

37 Rb

38 Sr

55 Cs 87 Fr

56 Ba 88 Ra

1+ 180 2+ 143

3 21 Sc

3+ 81

39 Y 57 La

5

4 22 Ti

3+ 76 4+ 68

6

7

24 Cr

23 V

2+ 88 3+ 74

25 Mn

3+ 63 6+ 52

2+ 80 4+ 60

43 Tc

40 Zr

41 Nb

42 Mo

4+ 79

4+ 74

4+ 70

72 Hf

8

-

9

26 Fe

10

27 Co

2+ 74 3+ 64

11 29 Cu

28 Ni

2+ 69

2+ 72

3+ 51

12

31 Ga

30 Zn

2+ 74

2+ 72

4+ 39

33 As

44 Ru

45 Rh

46 Pd

47 Ag

48 Cd

49 In

50 Sn

51 Sb

4+ 67

3+ 68

2+ 80

+ 126

2+ 97

3+ 81

4+ 71

3+ 76

78 Pt

79 Au

80 Hg

81 Tl

82 Pb

83 Bi

73 Ta

74 W

75 Re

76 Os

77 Ir

5+ 68

4+ 70

4+ 72

6+ 69

4+ 68

-

16 S

17 Cl

18 Ar

34 Se

35 Br

36 Kr

52 Te

53 I

54 Xe

2- 174 1- 181 4+ 30

3+ 58 5+ 46

4+ 53

10 Ne

2- 140 1- 133

5+ 35

32 Ge

3+ 62

1-178

2- 193 1- 196

-

-

-

2- 211 1- 220

84 Po

2+ 80 1+ 137 2+ 110 1+ 147 2+ 120 3+ 96

85 At

86 Rn

-

6+ 67

-

89 Ac 3+

119

58 Ce

4f 5f

59 Pr

60 Nd

61 Pm

3+ 107 3+ 106 3+ 104 4+ 94

-

91 Pa

93 Np

90 Th

92 U

4+ 102 3+ 113 4+ 97 6+ 80

-

62 Sm

63 Eu

64 Gd

3+ 100 2+ 112 3+ 97 3+ 98

94 Pu -

95 Am -

96 Cm -

65 Tb

66 Dy

67 Ho

68 Er

69 Tm

70 Yb

71 Lu

3+ 93 4+ 81

3+ 92

3+ 91

3+ 89

3+ 87

3+ 86

3+ 85

97 Bk

98 Cf

99 Es

100 Fm

101 Md

102 No

103 Lr

-

-

-

-

-

-

-

Figure 2.8 Ionic radii of the elements (only selected values shown; radii in pm=10-12m). In contrast, caesium’s hold on its outer electron is so weak that the electron can be dislodged by exposure to a flash of light. Electronegativities vary systematically through the Periodic Table. Firstly, they increase from left to right across each row or period. The lowest value in each period occurs in the first column of the s block (Group 1). Atoms in this group have one more electron than the noble gas configuration. This electron is loosely held and is easily lost. The highest values occur in the second-last column (Group 17) in which the elements have one less electron than the noble gas configuration and readily accept an electron to attain this stable electron arrangement. The increase in electronegativity across the periods occurs because the charge on the nucleus increases steadily across each row as atomic number increases, but because the p, d and f orbitals are dumbbell-shaped, the electrons are not fully shielded from the increasing positive charge on the nucleus. Hence the electrons

become progressively more strongly held across each row. Towards the end of each row there is the added effect of the stability afforded by the noble gas electron configuration, which causes the electronegativity to rise sharply, reaching a peak in Group 17. Electronegativity generally decreases down each group (column) in the table due to the steady increase in size of the atoms with increasing atomic number. There are some important deviations from these general trends that should be noted. Certain d block or transition elements, particularly those in groups 8–12, show marked deviation. In this region of the Periodic Table, electronegativity increases down each group and also shows a tendency to increase from left to right in each period, reaching a peak in Group 11. The electronegativity of gold is particularly high and approaches those of the halogens of Group 17. The reason for this is the poor shielding of the nuclear charge provided by electrons in the d and especially the

17

f orbitals in the case of gold (and its neighbouring elements in the period). This poor shielding restricts the increase in the size of atoms. For example silver (Ag, Z=47) has a radius of 134 pm, the same as that of gold (Au, Z=79). The outer electrons in gold are therefore more strongly held than in silver, so that the electronegativity of gold is 2.4, whereas that of silver is 1.9. The relatively inert character of the so-called noble metals (gold, silver, platinum, palladium, osmium, rhodium, ruthenium and iridium) is largely due to their high electronegativity. Cations are smaller than their neutral atom, whereas anions are larger than their uncharged equivalent (figure 2.8 on p17). The higher the ionic charge of a cation, the smaller the ion. For example, Fe2+ has a radius of 74 pm, whereas the radius of Fe3+ is 64 pm. This occurs because when an electron is lost, the remaining electrons are more strongly attracted by the nucleus – the same nuclear charge is spread over fewer electrons, so each feels a stronger attraction. Having discussed the nature of the atom and the distribution of electrons around the atom, we are now in a position to examine the manner in which atoms combine to form chemical compounds, and in our case, minerals. However, before we do, we need to look briefly at the natural abundances of elements.

Abundances of elements The natural abundance of elements varies quite dramatically. Some are common, others rare. Obviously, minerals composed of rare elements will also be rare. So to begin with, we need to know something about the relative abundances of the elements. The Sun and all of the planets formed about 4,600 million years ago from a cloud of dust and gas, the so-called solar nebula. Studies of meteorites that date back to this time and spectroscopic analysis of the light emitted by the Sun have enabled scientists to determine the abundances of the elements in the solar nebula fairly

18

precisely. These are often referred to as cosmic abundances, although this is something of a misnomer, as other nebulae in the universe have different compositions, so solar abundance is a more appropriate term. The relative abundances of non-radioactive elements in the Sun (and probably the solar nebula) are shown diagrammatically in figure 2.9. Along the horizontal axis is the atomic number. The vertical axis is the abundance. By convention, this axis is shown as the number of atoms of each element relative to one million atoms of silicon. Note that the scale is logarithmic: each division on the vertical axis represents a ten-fold change in relative abundance. Thus, for every million silicon atoms in the Sun there are about 2x1010 atoms of hydrogen (H), which means that there are 20,000 atoms of hydrogen for each atom of silicon. The rarest non-radioactive element, tantalum (Ta), is about 100 million times less abundant than silicon. There are several interesting features of the solar abundance curve: • Abundance decreases exponentially with increasing atomic number; • Elements with even atomic numbers are more abundant than neighbours with odd atomic numbers, which results in the saw-tooth appearance of the curve; • Lithium (Li), beryllium (Be) and boron (B) have unusually low abundances compared to other light elements; • Iron and nickel have slightly elevated abundances, forming a local peak in the curve; and • Lead is abnormally abundant for a heavy element. The relative abundances of the elements are a consequence of the processes that created them. The solar abundance curve has provided great insight into our understanding of the origin of the elements. The Universe was formed in the

H

10

10

He

abundance relative to one million (106) silicon atoms

109

108 O

10

C

7

Mg N Ne Si

106

S Ar

105

Fe

Ca

Al

Ni

Na

10

Cr

4

P

103

F

Mn

Ti Cl K

Co Zn

2

10

Li

Be

10

Cu

V B

Kr Ge Se

Sc

Ga Br

Sr

Rb

Y

As

1

Nb

10-1

10-2

0

5

10

15

20

25

30

35

Pb

Zr

40

Te Sn Cd Ru Pd

Mo

Xe

Ba

Pt Nd Dy Os Gd Er Hg Yb I La Sm Ir Cs Ti Bi Hf W Ag Sb Pr Ho Rh Au Eu In Tb Tm Re Lu Ta

45

50

Ce

55

60

65

70

75

80

85

atomic number Figure 2.9 Diagram illustrating the relative abundances of elements in the Sun (the solar abundance curve). By convention, abundance is expressed as the number of atoms of an element relative to 1 million atoms of the element silicon. Big Bang, which occurred about 13,700 million years ago. During this event hydrogen and helium and a smattering of slightly heavier elements formed, making up the embryonic Universe. Virtually all of the elements heavier than helium have been forged since the Big Bang inside stars by processes of nuclear fusion and the addition of neutrons (nucleosynthesis). The formation of different chemical

elements is closely related to the character of stars and their evolution. Bright blue stars, like our Sun, are powered by the conversion of hydrogen atoms to form helium. As the hydrogen fuel is exhausted, the star swells and reddens, becoming a Red Giant. In such stars, helium is being fused into heavier elements, a process that culminates in the formation of iron and its neighbouring elements. Stars that

19

reach this stage may become unstable and explode as supernovae, scattering their debris into space. The intense flux of neutrons associated with these events is largely responsible for the formation of elements with atomic numbers greater than about 30. Heavier elements are made from lighter elements, which is why abundance decreases with increasing atomic number. Elements with even atomic numbers are more stable than those with odd numbers, which is why the former are more abundant. The relatively higher abundances of iron and its neighbours reflects the pre-supernova culmination in stellar evolution.

The abundances of elements on Earth The debris from innumerable supernovae is spread throughout the Universe and appears as vast clouds of gas and dust, which have been beautifully illustrated by numerous images captured by the Hubble Space Telescope. The chemical elements forming these vast clouds have chemically combined to create an assortment of substances including methane (CH4), ammonia (NH3), water (H2O), carbon dioxide (CO2), carbon monoxide (CO), various hydrocarbon compounds like oils and tars, compounds involving metals, oxygen and silicon, and many more. In the frigid conditions of deep space most of these compounds occur as tiny solid, dustlike grains. Our solar nebula formed as one such cloud began to collapse under the influence of gravity 4,600 million years ago, an event most likely triggered by a nearby supernova. As the cloud collapsed, collisions between particles led to heating and in the inner portions of the cloud the temperature became so high that most of the elements were vaporized. The Earth and other inner planets formed in this hot, inner part of the solar nebula as it slowly cooled and began to condense. This hot birth, as well as Earth’s weak gravity,

20

How big is an atom? Atoms are the smallest subdivision of matter. They are very small – one million atoms of carbon strung together would form a chain about the thickness of a human hair. Atoms don’t have solid surfaces, but rather have an apparent volume created by the rapidly moving electrons that surround the atomic nucleus. In reality, they are probably more like balls of cotton wool than pool balls. Atoms are more or less spherical and the most convenient measure of the size of a sphere is its radius. Atomic radii cannot be measured directly but are generally inferred from the distances between layers of atoms in regular solid structures which are determined by means of X-rays. However, in such structures, the various atoms are bonded together and may overlap, so the radius measured may not be that of a free atom. Moreover, it has been found that the radius of a particular atom varies depending on the other constituents in the solid and the types of linkages between them. Some sources therefore list separate atomic radii for different types of bonds. As a result, atomic radii differ quite widely, depending on the bonding and also the method of determination. The unit of measurement of atomic radii generally accepted today is the picometre (pm), which is 10-12 m, or one million millionth of a metre. Another unit still in fairly common use is the angstrom (Å) which is 10-10 m, or one ten thousand millionths of a metre. One angstrom is therefore equal to 100 pm.

has meant that the relative abundances of the elements on Earth differ substantially from solar abundances. Elements with low boiling points and those that existed as gaseous compounds in the nebula are under-abundant in the Earth, whereas elements that have high boiling points or that formed compounds with high

boiling points are over-abundant in the Earth compared to solar abundance. It is for this reason that the Earth and the other inner planets are composed mainly of rocky material and have large, dense cores. Planets that formed further from the Sun in cooler parts of the nebula are composed mainly of gaseous substances – the so-called ‘gas giants’, such as Jupiter and Neptune. The Earth has also experienced chemical differentiation since its formation, separating into a core consisting mainly of iron and nickel metal, a mantle and a crust. The crust itself is of two types, one forming the floors of the oceans (oceanic crust) and the other the continents (continental crust), which differ in chemical composition. The estimated compositions of the Earth as a whole, the mantle and core are shown in table 2.1. (Only the more abundant constituents are listed in this table; a more complete list appears in Appendix 1.) By convention, the compositions are expressed as percentages

Element

Bulk Earth Mantle

Core

by mass. Solar abundances, expressed as percentages by mass (both including and excluding hydrogen and helium) are also listed in table 2.1. The main differences between solar abundances and the abundances in the Earth as a whole are the lower abundances of hydrogen, helium, carbon, oxygen and sulphur in the Earth. Most of the minerals we commonly encounter were formed in the continental crust, so it is useful to know something about its average composition (table 2.1 and Appendix 1). Eleven elements (oxygen, silicon, aluminium, iron, calcium, magnesium, sodium, potassium, titanium, manganese and phosphorus) constitute 99.7% of the total, two of which, manganese and phosphorus, are very minor. The remaining 63 elements together represent only 0.29% of the total. Given these abundances, it is not surprising that minerals composed of silicon, oxygen, iron, calcium, aluminium, magnesium and sodium make up most of the continental crust.

Continental crust

Sun

Sun minus H and He

Fe

32.0

6.26

85.5

5.2

0.14

10.9

O

29.7

44

0

46.5

0.65

51.4

Si

16.1

21

6

28.3

0.083

6.54

Mg

15.4

22.8

0

2.8

0.071

5.61

Ni

1.82

0.20

5.2

0.006

0.0084

0.66

Ca

1.71

2.53

0

4.6

0.0075

0.54

Al

1.59

2.35

0

8.4

0.0068

0.53

S

0.64

0.03

1.9

0.04

0.042

3.29

Cr

0.47

0.26

0.9

0.01

0.002

0.16

Na

0.18

0.27

0

2.3

0.004

0.31

P

0.07

0.009

0.20

0.06

0.00047

0.04

Mn

0.08

0.10

0.03

0.06

0.0015

0.12

C

0.07

0.01

0.20

uncertain

0.253

19.86

H

0.03

0.01

0.06