Metals play a vital role in the metabolism of plants and animals and, increasingly, in medicine. The Open University has developed an undergraduate course – Metals and Life, with two books, Metals and Life and Concepts in Transition Metal Chemistry used as the main teaching material. The books and the course cover how organisms acquire metals, their transport and storage, illustrated by such diverse examples as iron in the human body, and structures such as shells and teeth. This book provides an introduction to the metals essential to life and ligands of biological importance. It considers the uptake of metals, their transport and ultimately their storage, illustrated in particular with the story of iron in the body. It also considers Na, K and Ca ion channels and biomineralisation and covers the key roles that metals and their complexes play in living systems, for example in respiration and photosynthesis. The last section (delivered online) considers metal toxicity and deficiency and the role that metals play in medicine, looking at both diagnostics and therapy and the forefront of inorganic research.

Eleanor Crabb is a Lecturer in Materials Chemistry at The Open University. She studied chemistry at the University of Reading where she continued to undertake a PhD in heterogeneous catalysis. She spent another 9 months at Reading as a postdoctoral research fellow before moving to the Ecole Normale Superieure de Chimie de Montpellier as a postdoctoral research fellow. In 1993, she joined The Open University and has worked on a number of science courses, producing both text based and multimedia materials. She has produced sequences introducing students to 3D molecular representations of molecules and proteins for a wide range of courses at different levels, and has produced a number of animated multimedia sequences on receptor binding. Her research interests remain in the field of heterogeneous catalysis and she is the author (or co-author) of around 20 papers in this area. She is currently seconded part-time to one of the Centres for Excellence in Teaching and Learning (CETL) awarded to The Open University. Rob Janes is a Staff Tutor at The Open University in Wales. He studied chemistry at the University of Leicester, where he remained to undertake PhD research in solid state chemistry of the silver halides. He spent one year as a visiting scientist at Eastman-Kodak, Rochester, New York, before moving to the University of Cambridge as a Post-Doctoral Research Associate, working on high-Tc superconductivity. He taught at the Manchester Metropolitan University, developing courses in inorganic chemistry, materials chemistry and imaging science. His publications record consists of around 50 papers in the field of solid-state/materials chemistry. His research interests centre on the synthetic routes to ceramics – more specifically nanosize ceramic oxides and composite oxides, inorganic pigments and phosphors and studies of the electronic and magnetic properties of solids. Elaine Moore is a Reader in Chemistry at The Open University. She studied chemistry at Oxford University, stayed on to complete a DPhil in theoretical chemistry and after a two year post-doctoral position at Southampton, she joined The Open University in 1975. She has produced OU teaching texts in chemistry and astronomy and her research interests are in theoretical chemistry applied to solid state systems and to NMR spectroscopy. She is author or co-author on over 40 papers in scientific journals. Lesley Smart is a Senior Lecturer in chemistry at The Open University. She studied chemistry at the University of Southampton where she stayed on to complete a PhD on Raman spectroscopy. She has written on many science courses and chaired the production of the second level chemistry course. Her research interests are in the areas of solid state chemistry and catalysis, and in particular, preparing and characterizing new materials and catalysts. Dr Rob Davies (Consultant Author) is Senior Lecturer at Imperial College, London. He graduated from the University of Bristol before going to St John’s College Cambridge to study for his PhD. He was appointed to a three year Research Fellowship at St Catharines College Cambridge and then moved to Imperial College where he was awarded a Governors’ lectureship. His research interests lie in synthetic organometallic and coordination chemistry, especially of the main group metals. Dr David Johnson (Consultant Author) is a Visiting Reader in Chemistry at The Open University. A fellow of Trinity Hall, Cambridge, he was a founding member of the Department and worked on many of the chemistry courses prior to his retirement.

Metals and Chemical Change

By David Johnson

The Royal Society of Chemistry

Copyright © 2002 The Open University
All rights reserved.
ISBN: 978-0-85404-665-2

Contents

METALS AND CHEMICAL CHANGE David Johnson and Kiki Warr,
1 INTRODUCTION, 13,
2 REACTIONS OF METALS, 17,
3 METALS AND THEIR ORES, 27,
4 METALS AND THEIR EASE OF OXIDATION: A HYPOTHESIS, 37,
5 EQUILIBRIUM: A RESTATEMENT OF THE PROBLEM, 41,
6 THOMSENS HYPOTHESIS: TOWARDS A SOLUTION?, 42,
7 THE SECOND LAW OF THERMODYNAMICS: THE SOLUTION, 45,
8 THE FIRST LAW OF THERMODYNAMICS, 52,
9 ENTHALPIES OF REACTION: A DATABASE, 63,
10 ENTROPY CHANGES, 71,
11 THE GIBBS FUNCTION, 84,
12 METALS AND THEIR EASE OF OXIDATION, 91,
13 THERMODYNAMIC AND KINETIC STABILITY, 98,
14 REACTIVITY, 102,
15 THERMODYNAMICS AND THE OXIDATION OF METALS, 104,
16 ENTHALPY AND ENTROPY TERMS, 110,
17 METALS AMD THEIR ORES, 114,
18 THE BORN–HABER CYCLE, 122,
19 INTRODUCTION TO THE REMAINING SECTIONS, 128,
20 THE LATTICE ENERGY, 129,
21 ELECTROCHEMICAL CELLS AND REDOX POTENTIALS, 139,
22 IONIZATION ENERGIES OF ATOMS, 143,
23 THE CHEMISTRY OF GROUP I: THE ALKALI ELEMENTS, 148,
24 ALKALI METAL COMPOUNDS IN INDUSTRY, 154,
25 BINARY ALKALI METAL COMPOUNDS WITH NON-METALS, 160,
26 METAL IONS, LIGANDS AND COMPLEXES, 170,
27 ALKALI METAL COMPLEXES, 176,
28 THE GROUP II OR ALKALINE EARTH ELEMENTS, 184,
APPENDIX THERMODYNAMICS IN THIS BOOK, 203,
LEARNING OUTCOMES, 205,
QUESTIONS: ANSWERS AND COMMENTS, 208,
FURTHER READING, 234,
ACKNOWLEDGEMENTS, 234,
CASE STUDY: BATTERIES AND FUEL CELLS Ronald Dell and David Johnson,
1 INTRODUCTION, 237,
2 BATTERIES, 240,
3 BATTERY APPLICATIONS AND SIZES, 241,
4 CELL DISCHARGE AND CHARGE, 243,
5 BATTERY SPECIFICATION, 246,
6 DEGRADATION MODES IN BATTERIES, 255,
7 FUEL CELLS, 256,
ACKNOWLEDGEMENTS, 260,
INDEX, 261,

CHAPTER 1

INTRODUCTION

In the earlier books in this series, there has been an emphasis on molecular and electronic structure — that is, on the spatial arrangement of atoms within chemical substances, and on the arrangement of electrons within atoms. Very little has been said about chemical change. But here this emphasis shifts, and we ask why chemical reactions happen. There are two conditions that must be fulfilled before a chemical reaction can occur: the equilibrium constant must be sufficiently favourable, and the rate must be sufficiently fast. This Book will be concerned with both conditions, but mainly with the first. You will meet new ‘labour-saving’ properties of chemical substances; these will allow us to predict whether a chemical reaction has a favourable equilibrium constant or not; the reaction does not even have to be tried out.

The units of the properties in question are mainly those of energy, and come from a branch of science called thermodynamics. To show the relevance of this subject, we shall use it to explore an important problem about the chemical behaviour of metals. Finally, when our study of this problem is complete, thermodynamics is used again, towards the end of the Book, in a systematic study of the chemistry of the alkali and alkaline earth elements — that is, of Groups I and II of the Periodic Table. Along with thermodynamics, metals are therefore a major theme in this Book, so we begin by reminding you about them, and about the way that their properties are explained by the simplest theory of metallic bonding.

1.1 Metals and their physical properties

Figure 1.1 shows a full Periodic Table, colour-coded to reveal the periodic distribution of metals, semi-metals and non-metals. Of the 114 known elements, 90 are, or are likely to be, metallic. This, and the other books in the series, concentrate on the 46 typical elements. Here, metals are not so predominant, but, even so, they still outnumber each of the other two categories.

Some metallic elements, such as bismuth and manganese, are brittle, but most, when pure, are malleable and ductile. Malleable materials are those that can be reshaped by hammering; ductile materials can be drawn out under tension into wires. Figure 1.2 shows a piece of early British gold jewellery dated 1600 BC. Such things were made by hammering out gold into sheets. This is possible because the metal is malleable, and at the same time strong.

Malleability and ductility are especially associated with those metals that possess one of the two close-packed structures discussed in The Third Dimension: Crystals.

* What are the names of these two types?

* Hexagonal close-packed and cubic close-packed.

Both close-packed structures consist of layers of atoms of the metallic element. It follows that if we can explain why such layers can slip over one another fairly easily, we can account for both malleability and ductility. Now a simple model of a metal consists of a regular array of positively charged ions in a ‘pool’ of freely moving electrons. The interaction between the positive ions and the negatively charged electrons, which surround the ions, holds the metal together. Let us contrast the situations in a metal and in an ionic solid, such as NaCl, when the layers are displaced relative to one another. Look first at Figure 1.3b.

* Why should such a displacement be unfavourable in an ionic solid?

* The displacement brings like charges in adjacent layers into close proximity. Repulsion between the charges will then push the layers apart.

This explains why fracture and not deformation is usually the result of beating an ionic solid. The contrast with the situation in a metal (Figure 1.3a) is obvious: all the ions are of like charge, the situation after displacement is similar to what it was before, and the freely moving electrons can adjust to the change without further disruption. Consequently, in a pure metallic crystal, layers can usually slip easily over one another, thus accounting for the properties of ductility and malleability.

You will be familiar with other properties of metals from your everyday contact with iron, aluminium, copper, silver and tin, for example. They often have a lustrous appearance, and are good conductors of heat. However, the most characteristic property of metals is their high electrical conductivity. This is explained by the free electrons that roam at random through the metallic structure. When a voltage difference is applied across two points on a piece of metal, the motion of the electrons becomes less random, there is an overall movement of electrons between the two points, and an electric current flows.

The unit of electrical conductivity is siemens per metre (Sm-1). Those elements classified as metals in Figure 1.1 have an electrical conductivity at or below room temperature of at least 3 x 105 S m-1 along any direction in a single crystal of any known form of the element. Although it is a good electrical conductor, carbon in the form of graphite does not meet this criterion. The structure of graphite is shown in Figure 1.4. It consists of sheets of carbon atoms, and each atom is bonded to three others in the same sheet. Carbon has four outer electrons in the shell structure (2, 4); in graphite, each carbon atom shares three of these with three other carbon atoms in the same sheet by forming three covalent C — C bonds. The fourth electron is mobile, just as the bonding electrons in a metal are mobile; it binds carbon atoms within its sheet more strongly together by contributing to a pool of electrons concentrated around the plane of the sheet. Consequently, there is high conductivity parallel to the sheets, but very low conductivity at right-angles to them.

Because of this property, graphite is sometimes called a two-dimensional metal. However, our decision to treat as metals only those elements with high electrical conductivity in all three dimensions rules graphite out, and carbon is not classified as a metal in Figure 1.1.

1.2 Summary of Section 1

1 Most of the chemical elements are metals, and even among the typical elements, metals predominate.

2 Most, but not all, metallic elements are malleable and ductile. Elements classified as metals have high electrical conductivities in all directions in a single crystal of the element. Using this criterion, carbon is a non-metal.

3 These properties can be explained by a model in which metals are regarded as positive ions immersed in a pool of free electrons.

CHAPTER 2

REACTIONS OF METALS

Between now and the end of this Book, we shall investigate an important problem about the chemical behaviour of metals. Certain features of the problem are familiar to everybody. Gold jewellery (Figure 2.1) can survive essentially unchanged for thousands of years. Many bronze busts and statues (Figure 2.2) are much more recent, but the green stains on their stone bases show that significant corrosion of their copper content has already taken place. Again, the uranium metal intended for Nazi Germany’s first nuclear reactor went up in flames when a physicist took a shovel to it; rubidium reacts violently with water, and inflames in air, without the assistance of a shovel (Figure 2.3). Few chemists would quarrel with you if you said that copper was more reactive than gold, and that rubidium was more reactive than uranium, but what exactly do we mean by the word ‘reactive’? Again, is our statement about reactivities true only in moist air, or does it remain correct in the presence of other chemicals? It is these and other questions that we shall examine for the remainder of this Book.

To begin with, you must learn a little more about the reactions of metals. The next three Sections give you the opportunity to do this, partly by reading, and partly by watching experiments on the CD-ROM application, Reactions of metals.

ACTIVITY 2.1 Introduction: reactions with dilute acid

This activity introduces the problem of metal reactivity, and then asks you to record what happens when the metals copper, iron, magnesium, tin and zinc are added to dilute acid. It is the first sequence of Reactions of metals on the CD-ROM associated with this Book.

As noted in the latter part of Activity 2.1, in those cases where a reaction happened, the metal dissolved in the dilute acid to form a dipositive aqueous ion. For example, in the case of iron, this ion is Fe2+(aq). Moreover, when a reaction occurred, bubbles of gas were evolved. This happened for iron, magnesium, tin and zinc. You may remember that magnesium reacts with acids to yield hydrogen gas. Thus, for iron, a likely equation for the reaction is:

Fe(s) + 2H+(aq) = Fe2+(aq) + H2(g) (2.1)

Notice that this equation has been pared down to just the chemical species that change during the reaction; species like the chloride ion, Cl-(aq), which are present but do not change, have been left out.

* Try to write equations for the other three reactions that you observed.

* Following the example of iron in Equation 2.1, you should get:

Mg(s) + 2H+(aq) = Mg2(aq) + H2(g) (2.2)

Sn(s) + 2H+(aq) = Sn2(aq) + H2(g) (2.3)

Zn(s) + 2H+(aq) = Zn2(aq) + H2(g) (2.4)

In Reactions 2.1–2.4, you can see one of the most prominent chemical characteristics of metals.

* What is it?

* They form positive rather than negative ions in aqueous solution.

Metals also tend to form positive ions in solid compounds. For example, we say that solid sodium chloride contains Na ions, and this contrasts with the behaviour of non-metals, which do not form positive ions, and indeed often form monatomic negative ions such as Cl- or O2. There are important terms for describing changes of this type, to which we now turn.

2.1 Oxidation and reduction

The words oxidation and reduction tend to take on broader meaning as one learns more chemistry. You are now ready to take a step down this road. You know that a substance is said to be oxidized when it reacts with oxygen. Thus, magnesium is oxidized when it burns in oxygen gas (Figure 2.4):

Mg(s)+ 1/2 O2(g) = MgO(s) (2.5)

Let us consider how electrons are redistributed during this reaction by using an ionic bonding model.

* How are magnesium and oxygen held together in MgO?

* Each magnesium atom loses two electrons and forms the ion Mg2+, which has the electronic structure of neon; each oxygen atom gains two electrons and forms the ion O2-, which also has the electronic structure of neon. The forces between the oppositely charged ions in Mg2+O2- hold the compound together.

Thus, when the magnesium atoms on the left of Equation 2.5 are oxidized, they lose electrons and become Mg2+ ions in MgO. Chemists fasten on to this electronic change, and use it in a broader definition of oxidation:

The loss of electrons from an atom, compound or ion is described as oxidation.

Another feature of Equation 2.5 is that the oxygen atoms, bound together on the left-hand side as diatomic molecules, take on electrons and form O2- ions. Thus, whereas the magnesium atoms lose electrons, the oxygen atoms gain them. The taking-up of electrons, a process that is the reverse of oxidation, is called reduction:

The gain of electrons by an atom, compound or ion is described as reductions.

A useful mnemonic that will help you to remember this is OILRIG: Oxidation Is Loss; Reduction Is Gain.

In Equation 2.5, notice that oxidation and reduction occur together; they are complementary processes: the magnesium is oxidized and the oxygen is reduced. Apart from certain exceptional cases, a reaction that includes oxidation, also includes reduction, and vice versa. To mark this fact, such reactions are often called redox reactions, and Reactions 2.1–2.4 are typical examples.

* In Equations 2.1–2.4, is the metal oxidized or reduced?

* The metal loses electrons and forms positive ions; it has been oxidized.

* In Equations 2.1–2.4, what happens to the hydrogen ions?

They are reduced; they gain electrons and form neutral hydrogen atoms, which are combined in diatomic molecules.

QUESTION 2.1

Which of the following are redox reactions? In each redox reaction, identify the element that is oxidized, and the element that is reduced:

(i) [MATHEMATICAL EXPRESSION NOT REPRODUCIBLE IN ASCII]

(ii) [MATHEMATICAL EXPRESSION NOT REPRODUCIBLE IN ASCII]

(iii)[MATHEMATICAL EXPRESSION NOT REPRODUCIBLE IN ASCII]

(iv) [MATHEMATICAL EXPRESSION NOT REPRODUCIBLE IN ASCII]

(v) [MATHEMATICAL EXPRESSION NOT REPRODUCIBLE IN ASCII]

2.2 Oxidation of metals by aqueous hydrogen ions

We have now established that Reactions 2.1–2.4 are all redox reactions. Furthermore, in all of them, a metal atom reacts with two hydrogen ions to give an ion with two positive charges and a molecule of hydrogen gas. However, Activity 2.1 revealed a significant difference in the vigour of the reactions. In particular, the magnesium reaction was quite violent; by contrast, the iron, tin and zinc reactions were slow. Indeed, the very slow tin reaction only became perceptible on heating. Similar reactions, which are as violent or more violent than the magnesium reaction, are observed when the alkali metals (Group I of the Periodic Table) react with acids. For example, steady evolution of hydrogen occurs from the surface of lithium metal, but with sodium the gas evolution is extremely vigorous; for potassium, rubidium and caesium, the reaction is explosive, and the hydrogen gas catches fire. The reactions with plain water are very similar.

* Write equations for the reactions of lithium and caesium with acids, in which hydrogen gas is produced.

* Like magnesium, the alkali metals form ions with noble gas structures, losing one electron to form singly charged cations:

Li(s) + H+(aq) = Li+(aq) + 1/2 H2(g) (2.6)

Cs(s) + H+(aq) = Cs+(aq) + 1/2 H2(g) (2.7)

So experiments show that magnesium and the alkali metals react much more violently with aqueous hydrogen ions than do zinc, iron and tin. However, in Activity 2.1 there was one element that did not react at all with dilute hydrochloric acid.

* Which element was this?

* It was copper; when copper is dropped into dilute hydrochloric acid, no hydrogen gas is evolved, and the dipositive aqueous ion, Cu2+(aq), is not formed.

Now hydrated forms of copper sulfate, CuSO4, are used by gardeners as a fungicide. The solution of this compound in water is blue because of the blue ion Cu2+(aq) (Figure 2.5):

CuSO4(s) = Cu2+(aq) + SO42-(aq) (2.8)

Thus, the ion Cu2+(aq) exists.

* Write an equation for a reaction of metallic copper with acid to produce hydrogen gas.

* Cu(s) + 2H+(aq) = Cu2+(aq) + H2(g) (2.9)

As Activity 2.1 has demonstrated, this reaction does not happen. This is true of the analogous reaction for certain other metals, in particular silver and gold. Silver forms a well-defined colourless ion Ag+(aq) and, for our purposes, the most convenient description of aqueous oxidized gold is the ion Au3+(aq). Thus, we can write two further equations for reactions that do not occur.

(Continues…)Excerpted from Metals and Chemical Change by David Johnson. Copyright © 2002 The Open University. Excerpted by permission of The Royal Society of Chemistry.
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