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Modern inorganic chemistry AN INTERMEDIATE TEXT

Senior Chemistry Master, Bolton School

Professor of Inorganic Chemistry, The University of Liverpool

ENGLAND Butterworth & Co (Publishers) Ltd London: 8 Kingsway, WC2B 6AB

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First published 1975

© Butterworth & Co (Publishers) Ltd 1975

Printed and bound in Great Britain by R. .).

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1 The periodic table 1 2 Structure and bonding 25 3 Energetics 62

4 Acids and bases: oxidation and reduction 84 5 Hydrogen 1

6 Groups I and I 119 7 The elements of Group I 138 8 Group IV 160 9 Group V 206 10 Group VI 257 1 Group VII: the halogens 310 12 The noble gases 353

13 The transition elements 359 14 The elements of Groups IB and IIB 425

15 The lanthanides and actinides 440 Index 447


The welcome changes in GCE Advanced level syllabuses during the last few years have prompted the writing of this new Inorganic Chemistry which is intended to replace the book by Wood and Holliday. This new book, like its predecessor, should also be of value in first-year tertiary level chemistry courses. The new syllabuses have made it possible to go much further in systematising and explaining the facts of inorganic chemistry, and in this book the first four chapters—-the periodic table; structure and bonding; energetics: and

acids and bases with oxidation and reduction—provide the necessary grounding for the later chapters on the main groups, the first transition series and the lanthanides and actinides. Although a similar overall treatment has been adopted in all these later chapters, each particular group or series has been treated distinctively, where appropriate, to emphasise special characteristics or trends.

A major difficulty in an inorganic text is to strike a balance between a short readable book and a longer, more detailed text which can be used for reference purposes. In reaching what we hope is a reasonable compromise between these two extremes, we acknowledge that both the historical background and industrial processes have been treated very concisely. We must also say that we have not hesitated to simplify complicated reactions or other phenomena—thus, for example, the treatment of amphoterism as a pH-dependent sequence between a simple aquo-cation and a simple hydroxo-anion neglects the presence of more complicated species but enables the phenomena to be adequately understood at this level.

We are grateful to the following examination boards for permission to reproduce questions (or parts of questions) set in recent years in

Advanced level (A), Special or Scholarship (S), and Nuffield (N) papers: Joint Matriculation Board (JMB). Oxford Local Examinations (O). University of London (L) and Cambridge Local Examina- tion Syndicate (C). We also thank the University of Liverpool for permission to use questions from various first-year examination papers. Where appropriate, data in the questions have been converted to SI units, and minor changes of nomenclature have been carried out; we are indebted to the various Examination Boards and to the University of Liverpool for permission for such changes.

1 The periodic table

We now know of the existence of over one hundred elements. A century ago, more than sixty of these were already known, and naturally attempts were made to relate the properties of all these elements in some way. One obvious method was to classify them as metals and non-metals; but this clearly did not go far enough.

Among the metals, for example, sodium and potassium are similar to each other and form similar compounds. Copper and iron are also metals having similar chemical properties but these metals are clearly different from sodium and potassium—the latter being soft metals forming mainly colourless compounds, whilst copper and iron are hard metals and form mainly coloured compounds.

Among the non-metals, nitrogen and chlorine, for example, are gases, but phosphorus, which resembles nitrogen chemically, is a solid, as is iodine which chemically resembles chlorine. Clearly we have to consider the physical and chemical properties of the elements and their compounds if we are to establish a meaningful classification.

By 1850. values of atomic weights (now called relative atomic masses) had been ascertained for many elements, and a knowledge of these enabled Newlands in 1864 to postulate a law of octaves. When the elements were arranged in order ot increasing atomic weight, each

2 THE PERIODICTABLE successive eighth element was 4a kind of repetition of the first'. A few years later, Lothar Meyer and Mendeleef, independently, suggested that the properties of elements are periodic functions of their atomic weights. Lothar Meyer based his suggestion on the physical properties of the elements. He plotted 'atomic volume'—the volume (cm3) of the

Atomic weight

Figure Ll. Atomic volume curve (Lothar Meyer] atomic weight (g) of the solid element- against atomic weight. He obtained the graph shown in Figure L We shall see later that many other physical and chemical properties show periodicity (p. 15).

Mendeleef drew up a table of elements considering the chemical properties, notably the valencies, of the elements as exhibited in their oxides and hydrides. A part of Mendeleef s table is shown in Figure

1.2 -note that he divided the elements into vertical columns called groups and into horizontal rows called periods or series. Most of the groups were further divided into sub-groups, for example Groups


IA, IB as shown. The element at the top of each group was called the "head' element. Group VIII contained no head element, but was made up of a group of three elements of closely similar properties, called "transitional triads'. Many of these terms, for example group, period and head element, are still used, although in a slightly different way from that of Mendeleef.

Group I

Li No

A sub- < group fK Cui

Rb B

Ag \ sub- Cs group

H EZ ¥ in MEITTTf —

Fe Co Ni Ru Rh Pd

Os Ir Pt

* Francium. unknown to Mendeleef, has been added Figure 1.2. Arrangement oj some elements according to Mendeleef

The periodic table of Mendeleef, and the physical periodicity typified by Lothar Meyer's atomic volume curve, were of immense value to the development of chemistry from the mid-nineteenth to early in the present century, despite the fact that the quantity chosen to show periodicity, the atomic weight, was not ideal. Indeed, Mendeleef had to deliberately transpose certain elements from their correct order of atomic weight to make them Hf into what were the obviously correct places in his table; argon and potassium, atomic weights 39.9 and 39.1 respectively, were reversed, as were iodine and tellurium, atomic weights 126.9 and 127.5. This rearrangement was later fully justified by the discovery of isotopes. Mendeleef s table gave a means of recognising relationships between the elements but gave no fundamental reasons for these relationships.

In 1913 the English physicist Moseley examined the spectrum produced when X-rays were directed at a metal target. He found that the frequencies v of the observed lines obeyed the relationship v = a(Z ~ b)2 where a and b are constants. Z was a number, different for each metal, found to depend upon the position of the metal in the periodic table.


It increased by one unit from one element to the next, for example magnesium 12, aluminium 13. This is clearly seen in Figure 1.3. Z was called the atomic number; it was found to correspond to the charge on the nucleus of the atom (made up essentially of protons and neutrons), a charge equal and opposite to the number of ext ra nuclear

Figure 1.3. Variation of (frequency]' with Z

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