Chemistry for Colleges and Schools

The analogy in the title to this chapter applies to all science. Much of the material we will discuss here provides the basis of several branches of knowledge about ourselves and our environment, including, of course, chemistry.

Why do chemical reactions occur? The question is of fundamental importance, and many others follow from it: (a)Why do some reactions occur spontaneously, and others not?(b)Do all spontaneous reactions proceed to completion? If not, why not?(c)Can we predict which reactions will occur spontaneously?(d)Can we predict whether they will go to completion? The answers to these questions provide part of the basis of our understanding of chemistry. Attempts to answer these questions started with the idea that the heat given off in reactions was a measure of the ‘reactivity’ of the chemical reagents involved. The idea is only partly correct, as we shall see, but it provided the basis of that sub-discipline of chemistry called thermodynamics. Thermodynamics is one of the most fully and rigorously developed sections of chemistry; and, as a glance at the shelves of any chemistry library will show you, it is one of the most written about. Whole books (often longer than this one) have been devoted to the subject. We shall sacrifice some of the rigour and much of the content; and consider only those conclusions which will be of help to us in later topics.

A study of chemical thermodynamics allows us to understand why chemical reactions occur. Despite the inclusion of the term ‘dynamics’, such a study tells us nothing about the speeds of chemical reactions. The study of chemical kinetics is concerned with the measurement of the speeds (or rates) of reactions together with the factors which affect these rates. As we shall see, kinetics leads us to an understanding of how chemical reactions occur. Also, almost as a bonus, kinetic studies give access to some thermodynamic data.

Chemical equilibrium is a particular case of the more general thermodynamic equilibrium discussed in Chapter Two. We have seen that a system will cease to undergo change when it comes into a state of minimum free energy. That is when any further change of state would involve a positive change in free energy; an increase in free energy.

Many chemical reactions which reach equilibrium involve reagents in two and sometimes three different phases. The calcium carbonate reaction discussed in Chapter 4 reaches heterogeneous equilibrium in a closed system.

Our discussion of acid—base equilibria will use the Brønsted-Lowry definition of an ACID as a proton donor, and a BASE as a proton acceptor.

The class of reactions called oxidation — reduction reactions (or more commonly RedOx reactions) are similar to the acid-base reactions discussed in Chapter 6. The latter involve the transfer of proton(s) from the acid to the base, RedOx reactions involve the transfer of electrons.

For gases, the state functions describing the system are pressure, volume, and temperature. These state functions are related via an equation of state, which, for 1 mole of gas is <m:math display=&#x2019;block&#x2019;> <m:mrow> <m:mi>P</m:mi><m:mi>V</m:mi><m:mtext>&#x2009;</m:mtext><m:mo>=</m:mo><m:mtext>&#x2009;</m:mtext><m:mi>R</m:mi><m:mi>T</m:mi><m:mo>.</m:mo></m:mrow> </m:math>]]</EquationSource><EquationSource Format="TEX"><![CDATA[$$PV\; = \;RT.$$ A gas which obeys this equation is said to be an ideal gas.

As the name ‘condensed phases’ suggests, molecules in liquid and solid states of matter are much closer together than in the gaseous state. Since chemists work mostly with solids, liquids, and solutions it is of great importance to understand the nature and properties of substances in these phases. We can approach this understanding by first considering how the condensed phases are created.

Our present-day understanding of the detailed nature of atoms and molecules is based on what is known as quantum (mechanical) theory. This approach is relatively new, dating only from the early years of the century. The reasons for its development lay in the failure of so-called classical mechanics — by which we describe and explain the behaviour and interactions of relatively large bodies of matter — to describe satisfactorily the atomic system.

Before plunging into the problems associated with devising a wave description of an electron, let us consider the conceptually simpler one- and two-dimensional waves. A vibrating (guitar or violin) string is a one-dimensional wave. The sounds produced by such a string are caused by the movement of the string in the air surrounding it. There are always at least two fixed points on the string — one at each end, which correspond to the stops which hold the string in place. Thus movements which would require the string to shift at either end are forbidden. Various allowed and forbidden movements (vibrational modes) are shown in Figure 11.1. The requirements that there be zero displacement at each end of the string are known as boundary conditions, and the points of zero displacement are known as nodes. Notice how the boundary conditions severely limit the permitted modes of vibration. The various allowed modes can be described by the relation: 11.1<math display='block'> <mrow> <mi>l</mi><mo>=</mo><mi>n</mi><mrow><mo>(</mo> <mrow> <mfrac> <mi>&#x03BB;</mi> <mn>2</mn> </mfrac> </mrow> <mo>)</mo></mrow> </mrow> </math>]]</EquationSource><EquationSource Format="TEX"><![CDATA[$$l = n\left( {\frac{\lambda }{2}} \right)$$In the above ‘wave equation’ for the one-dimensional vibrator, ‘n’ is a quantum number; its value controls the wavelength of allowed vibration. If n = 2, the wavelength of allowed motion is one-half that for the mode when n = 1; and so on.

Of course the hydrogen atom, discussed at length in Chapter Eleven, is a one-electron system, but it is not the only one. The ions He+ and Li2+ for example also possess one electron. It is appropriate to consider all such systems together, because their properties are similar. In many-electron atoms the picture, regrettably, is more complex, as electron-electron repulsion terms must be included in any expression for electron energy. It is only for one-electron systems that the energy-degeneracy of orbitals within a particular shell is retained (Figure 12.1). Not only are orbitals of (say) the 3-shell of higher energy than those of the 2-shell, but the ‘size’ of orbitals of the same type increases with energy. We have used the expression 11.8<math display='block'> <mrow> <mi>E</mi><mo>=</mo><mfrac> <mrow> <mo>&#x2212;</mo><msub> <mi>E</mi> <mn>0</mn> </msub> </mrow> <mrow> <msup> <mi>n</mi> <mn>2</mn> </msup> </mrow> </mfrac> </mrow> </math>]]</EquationSource><EquationSource Format="TEX"><![CDATA[$$E = \frac{{ - {E_0}}}{{{n^2}}}$$

In this brief chapter we will begin to examine that form in which most matter exists — molecules. We can develop an understanding of molecules by considering how they resemble their constituent particles — atoms — and how and why they differ. We will discover that many of the principles discussed in the last few chapters about atoms apply equally well to molecules. But first it is necessary to ask: what are molecules and why do they form?

We have seen that, in principle, molecules can be treated in much the same way as atoms. The electrons in each molecule can be described by Schrödinger’s equation — and by solving the equation we should be able to discover the permitted eigen pairs of energy and wave function. Unfortunately, it is not possible to solve exactly Schrödinger’s equation for atoms more complex than hydrogen: and, of course, molecules are even more complex. The fact that an exact solution cannot be obtained does not mean Schrödinger’s equation is invalid — but it does mean that it is so complicated that solutions can only be approximated to. A common technique for finding best possible solutions is to apply what is called the variation principle. A wave function is chosen and used in Schrödinger’s equation to calculate the energy of the system. Then the wave function is changed slightly (‘varied’) and the energy computed again. The procedure is repeated until energy is minimised — at which point it is assumed that the wave function is a ‘good’ function.

The use of electron dot diagrams provides a simple pictorial representation of the valence bonds in a particular molecule. The technique involves using dots to represent electrons from the valence shells of the atoms in the molecule: the atoms are then combined in such a way as to pair the electrons (dots) in (single) valence bonds. For example: <m:math display=&#x2019;block&#x2019;> <m:mtable columnalign=&#x2019;left&#x2019;> <m:mtr> <m:mtd> <m:msub> <m:mi>H</m:mi> <m:mn>2</m:mn> </m:msub> <m:mtext>&#x2009;</m:mtext><m:mi>H</m:mi><m:mo>&#x00B7;</m:mo><m:mtext>&#x2009;</m:mtext><m:mo>+</m:mo><m:mtext>&#x2009;</m:mtext><m:mi>H</m:mi><m:mo>&#x00B7;</m:mo><m:mtext>&#x2009;</m:mtext><m:mo>&#x2192;</m:mo><m:mtext>&#x2009;</m:mtext><m:mi>H</m:mi><m:mo>:</m:mo><m:mi>H</m:mi> </m:mtd> </m:mtr> <m:mtr> <m:mtd> <m:mi>H</m:mi><m:mi>F</m:mi><m:mtext>&#x2009;</m:mtext><m:mi>H</m:mi><m:mo>&#x00B7;</m:mo><m:mtext>&#x2009;</m:mtext><m:mo>+</m:mo><m:mtext>&#x2009;</m:mtext><m:mo>&#x00B7;</m:mo><m:munderover> <m:mi>F</m:mi> <m:mrow> <m:mo>&#x00B7;</m:mo><m:mo>&#x00B7;</m:mo></m:mrow> <m:mrow> <m:mo>&#x00B7;</m:mo><m:mo>&#x00B7;</m:mo></m:mrow> </m:munderover> <m:mo>:</m:mo><m:mtext>&#x2009;</m:mtext><m:mo>&#x2192;</m:mo><m:mtext>&#x2009;</m:mtext><m:mi>H</m:mi><m:mo>:</m:mo><m:mtext>&#x2009;</m:mtext><m:munderover> <m:mrow> <m:mi>F</m:mi><m:mo>:</m:mo></m:mrow> <m:mrow> <m:mo>&#x00B7;</m:mo><m:mo>&#x00B7;</m:mo></m:mrow> <m:mrow> <m:mo>&#x00B7;</m:mo><m:mo>&#x00B7;</m:mo></m:mrow> </m:munderover> </m:mtd> </m:mtr> </m:mtable> </m:math> ]]</EquationSource><EquationSource Format="TEX"><![CDATA[$$\begin{gathered} {H_2}\;H\; + \;H\; \to \;H:H \hfill \\ HF\;H\; + \;\mathop F\limits_{}^{} :\; \to \;H:\;\mathop {F:}\limits_{}^{} \hfill \\ \end{gathered}$$ Electrons from only the valence shells are shown because it is these which are used in bond formation. Electron dot diagrams give immediate visual confirmation of there being any remaining unpaired electrons or residual bonding capacity in a proposed molecule.

A pure covalent bond exists only between two like atoms; e.g. in homonuclear diatomic molecules. As the constituent atoms become larger the equilibrium separation of the nuclei becomes greater and the extent of overlap between valence orbitals on the adjacent atoms becomes less. Thus we expect to see an inverse correlation between atomic size and bond dissociation energy. The most marked effect will be seen down a group in the periodic table — the small expected effect across the table (increased bond energy with decreasing size) is masked to a certain extent by variations in electronegativity and atomic orbital occupancy (Table 16.1).

In times past the chemist wishing to determine the structure of an unknown molecule was forced to undertake a quite complex series of tests, including degradation analyses, functional group analyses, and so on. Often a certain amount of inspired guesswork was also involved. Nowadays we are fortunate to be able to bring a formidable array of instrumental weapons to the task. Over the last few decades the techniques of: infra-red (IR), ultra-violet/visible (UV-Vis), atomic absorption (AA), Raman, nuclear magnetic resonance (nmr) and mass spectrometry have been developed or improved by enormous technological advances. In addition, sophisticated diffraction and electron microscopy techniques are also now available to the chemist.

Inorganic chemistry concerns the chemistry of all known elements, with the single exception of most of the chemistry of carbon, which is conventionally covered by organic chemistry. This gives a formidable total of 105 elements. If we exclude the man-made radioactive elements, mostly actinides, and the lanthanide rare earth elements, and the three rows of the transition metals, we are still confronted with about 40 elements. Clearly it is impossible to have the fullness and detail of knowledge about each of these which we acquire about carbon in organic chemistry. Nevertheless we have to learn facts about each element and its compounds, and this entails a certain amount of memory work. Sometimes people complain about learning facts in chemistry, forgetting that every field of study requires a basis of facts to be learnt. Facts are a little like the vocabulary of a foreign language. Without learning the vocabulary, you cannot communicate; it becomes hazardous or impossible to travel in that country; you cannot appreciate the literature and culture; perhaps you cannot even survive there.

A chemist handles chemicals in all three phases, and it is important that he knows something about what structure the atoms adopt in each phase and what effect this has on the properties of a substance. This leads to an understanding of why some solids are soluble and others not, what determines the choice of solvent, and why some materials are volatile and others not. A more advanced study (not attempted here) leads into the very important field of metallurgy where detailed theories of metals and their alloys are studied. The relevance of these to the economy of a country and to the development of materials better suited to particular needs is self-evident.

Nearly every element in the periodic table forms at least one compound both with hydrogen and with oxygen. The resulting hydrides and oxides are often important compounds in their own right. They have properties which fcllow neat trends in the periodic table, helping us to correlate the chemistry of many other elements, and they introduce us to several important concepts. It is thus useful to examine the chemistry of these two elements first.

The elements of group I are called the alkali metals. They each have one s electron in the outer shell and we shall see that there are very strong family resemblances between them. The elements of group II are called the alkaline earth metals. They have two s electrons in the outer shell and they too show obvious family resemblances. We thus fmd that the chemistry of the elements in these two groups can be briefly summarised by a number of simple generalisations.

The elements of groups IIIb, IVb, and Vb show a wide variety of properties, and there are fewer useful family relationships within groups than in the s-block. By the time we reach the other end of the periods, with groups VIb (the chalcogens), VIIb (the halogens), and VIIIb (the noble gases) family relationships become more marked again.

In the previous two chapters the chemistry of the forty-three main group elements has been outlined. The underlying principle to their chemistry is the filling of s and p subshells with consequent changes in ionisation energy and many other related properties. Between the alkaline earth elements of group IIa and the elements of group IIIb lie ten other groups of elements, corresponding to the filling of d subshells. These thirty elements are called transition metals, and although we shall only make a brief survey of their chemistry, some of them, such as iron, copper, and zinc, are of fundamental importance to our civilisation. The elements fall into groups based on their nuclear charge, etc., but family relationships within a group in the d block are often less marked than the general similarities between an element and its horizontal neighbours, and the general similarities of all the transition elements. This is least true for the first and last groups, IIIa and IIb, which show strong family similarities. Group VIIIa has three members from each period: this anomaly has arisen because the periodic table was not constructed on the basis of orbital filling, but on the basis of numerical theories such as Newland’s sevens (see section 18.2). Because this chapter makes only a brief survey of the transition metals, rather than summarising the chemistry of each group, the main trends will be picked out, and illustrated, after which more details will be given of some of the more important members.

In the last six chapters we have approached the huge sub-discipline of descriptive inorganic chemistry from the points of view of unifying trends across each row of the periodic table, and of the similarity in properties and reactions of the elements comprising each column. These unifying trends in group and period provide a useful, even essential, aid to the mastery of the plethora of chemical knowledge which the professional chemist must have at his fingertips. We must be aware also of the more significant exceptions to these trends and rules. Those elements and reactions which are exceptions are often of equal importance to those which conform.

Organic chemistry is the study of the compounds of the element carbon, compounds which also contain hydrogen and often other non-metal elements like oxygen, nitrogen, the halogens, sulphur, and phosphorus. You might think these ought to be covered by inorganic chemistry just as we learn in inorganic chemistry about compounds formed from sulphur and hydrogen, etc. But there are two reasons for not doing this. The first is historical: in the years before about 1830 it was thought that organic compounds could be made only in plants or animals, that is, in living organisms, and that they contained a mysterious ‘vital force’ in addition to carbon and other chemical elements. Therefore much knowledge was amassed about carbon compounds during that time, but kept separate from knowledge about other compounds. The second reason is practical: carbon forms so many different compounds, often of such complexity, and we now know so much about how they react and what they produce, that we have to study carbon compounds separately. The very obvious danger is that people are tempted never to compare the chemistry of carbon compounds with that of other compounds, or even never to consider there may be a link. But of course there are links, many useful or interesting comparisions can be made, and much exciting work is being done in the ‘inter-disciplinary’ field of study called organometallic chemistry.

The sensible place to begin is with compounds containing only carbon and hydrogen: the ‘hydrocarbons’. Hydrocarbons are divided into three types of compound depending on the type of bonding between carbon atoms: alkanes have only carbon—carbon single bonds, alkanes contain a carbon—carbon double bond, and alkynes a triple bond. The reactions of each type of hydrocarbon are rather different.

We have seen in section 26.5 how one or more hydrogen atoms in an alkane can be replaced by chlorine atoms. The products are chloroalkanes: the general term for such a compound would be haloalkane. In general, regardless of the identity of the halogen, all their reactions follow the same pattern. More space will be devoted to alkane derivatives containing just one halogen atom, because these are most common, but the reactions of those containing two or more halogens are similar.

Two compounds are said to be isomers when they have the same molecular formula but different structures. We have met several ways in which this can occur already, and we can classify them under the following headings, although in practice two isomers may differ according to two or more headings at the same time.

It is instructive to make repeated comparisons with water in this chapter. Alcohols are compounds which have one of the hydrogens in water replaced by an alkyl group, ROH, and ethers have both hydrogens replaced by alkyl groups, ROR′.

An oxygen doubly bonded to a carbon atom is the functional group called a carbonyl group. The carbon atom must be part of the carbon skeleton of a molecule: if it is at one end of the carbon chain, the carbonyl group will be also attached to a hydrogen. Such compounds are called aldehydes, <m:math display='block'> <m:semantics> <m:mrow> <m:mi>R</m:mi><m:mo>&#x2212;</m:mo><m:mover> <m:mi>C</m:mi> <m:mo>&#x2225;</m:mo> </m:mover> <m:mo>&#x2212;</m:mo><m:mi>H</m:mi><m:mo>.</m:mo> </m:mrow> </m:semantics> </m:math>]]</EquationSource><EquationSource Format="TEX"><![CDATA[$$\[R - \mathop C\limits^\parallel - H.\]$$ If the <m:math display='block'> <m:semantics> <m:mrow> <m:mo>&#x227B;</m:mo><m:mi>C</m:mi><m:mo>=</m:mo><m:mi>O</m:mi> </m:mrow> </m:math>]]</EquationSource><EquationSource Format="TEX"><![CDATA[$$\[ \succ C = O\]$$ group is part way along a carbon chain, being attached to two alkyl groups, the compound is called a ketone, <m:semantics> <m:mrow> <m:mi>R</m:mi><m:mo>&#x2212;</m:mo><m:mover> <m:mi>C</m:mi> <m:mo>&#x2225;</m:mo> </m:mover> <m:mo>&#x2212;</m:mo><m:msup> <m:mi>R</m:mi> <m:mo>&#x2032;</m:mo> </m:msup> <m:mo>.</m:mo> </m:mrow> </m:math>]]</EquationSource><EquationSource Format="TEX"><![CDATA[$\[R - \mathop C\limits^\parallel - R'.\]$$ In spite of the different names there are only slight differences in their chemistry, so the two types of compound are considered together. We shall see that nearly all the reactions occur at the carbonyl group, regardless of the nature of the alkyl groups.

This chapter examines the properties of compounds in which the carbonyl group is directly bonded to other functional groups such as hydroxyl or halogen. We shall see that the properties of the individual groups are modified somewhat, and that sometimes new properties are found.

Amines are derivatives of ammonia with one or more hydrogens replaced by alkyl groups. They are divided into three classes: primary amines have one alkyl group attached to the nitrogen;secondary amines have two alkyl groups attached to the nitrogen;tertiary amines have three alkyl groups attached to the nitrogen. The nature of the alkyl group is unimportant.

In the early days of organic chemistry, scientists found that a fair number of compounds having pleasant aromatic smells were all derivatives of a six-carbon molecule called benzene, C₆H₆. The benzene part of the molecules seemed surprisingly stable, especially considering how many multiple bonds there would have to be to satisfy the valencies; and for some years there was no satisfactory explanation of the bonding. But during this period the term ‘aromatic’ was applied to many other benzene derivatives, even when they had no smell at all. By the time a suitable description of the bonding in benzene and its derivatives had been produced, the misleading term ‘aromatic’ had become too widely used to be dropped. It is now used to refer to compounds derived from benzene, or benzene-like compounds, with absolutely no reference to their smell. Aromatic compounds may of course have an aromatic smell, but their smell could just as well be foetid, sweet, pungent, or ethereal, or they may have no smell at all.

The electrophilic substitution reactions of benzene have been described in the previous chapter. Two additional reactions of benzene are covered in this chapter. The rest of the chapter is devoted to the reactions of benzene derivatives, having electron-withdrawing substituents on the ring. Since most of these have been covered in earlier chapters and the reactions of the ring have been covered in Chapter 34, this is virtually a chapter of review and revision: very few new reactions are met.

Since an electron-donating substituent will usually also increase the π electron density all over the aromatic ring, such rings react with electrophiles faster than benzene; sometimes so much so that milder conditions have to be used. Thus a hydroxybenzene is as reactive to electrophiles as an alkene, but still giving substitution reactions rather than addition.

Diazonium ions are prepared by treating aromatic primary amines with nitrous acid in solutions of pH ~ 3 and at temperatures below about 5°C. Although they are not very stable, diazonium ions are immensely important for synthesis in the laboratory and in industry.

You will already have solved some synthetic problems, where it has been necessary to convert one compound into another using one or more steps. Although these have used simple molecules, and the products in fact could probably be bought cheaply, the practice you have had in interconverting compounds is valuable. The reactions for changing, for example, an ester to an amine are the same, in principle, regardless of the complexity of the molecule.