Publisher Summary

This chapter presents that solvent molecules can become attracted to ions with varying degrees of firmness, depending on the characteristics of both, the ion and the solvent. The introduction of ions into a solvent can also have a marked effect on its properties. That is particularly true when, as in the case of water, the solvent has a pronounced structure of its own. Indeed, the commonest and most important solvent, water, is one of the most interesting in that respect, since in aqueous solutions of salts interactions between ions and solvent molecules profoundly affect interactions between the solvent molecules themselves. The chapter discusses mainly aqueous solutions, as these have received the most attention, owing both to the importance of water and to its interest as a solvent. However, the study of non-aqueous solvents has developed greatly in some areas and the chapter will discuss such solutions where appropriate. It should not be thought that all possible metal ions able to exist in aqueous solution have already been discovered and characterized. In the last few years, aqua-ions of palladium(II), platinum(II), and molybdenum(III) have been properly established.

1.1 DISSOLUTION OF SALTS

One of the best ways to appreciate the importance of ion–solvent interactions in electrolyte solutions is through the cycle shown as the top half of Fig. 1.1. This relates the enthalpy of solution to ion solvation enthalpies and to lattice enthalpy. Invariably the enthalpy of solution of a salt is the small difference between the large enthalpy needed to separate the ions from each other in the crystal lattice and the enthalpy gained when these ions are introduced into the solvent. The lower half of Fig. 1.1 shows that for the specific case of sodium chloride, the enthalpy of solution is only about 0.5% of the lattice or ion solvation enthalpies. Fig. 1.1 should give some idea of the strength of ion–solvent, particularly ion–water, interactions. In the following pages various aspects of the chemistry of solvated ions — their natures, properties, and reactions — are introduced.

Solvent molecules can become attracted to ions with varying degrees of firmness, depending of course on the characteristics of both the ion and the solvent. The introduction of ions into a solvent can also have a marked effect on its properties. This is particularly true when, as in the case of water, the solvent has a pronounced structure of its own. Indeed the commonest and most important solvent, water, is one of the most interesting in this respect, since in aqueous solutions of salts interactions between ions and solvent molecules profoundly affect interactions between the solvent molecules themselves. Many sections will deal mainly with aqueous solutions, as these have received most attention, owing both to the importance of water and to its interest as a solvent. However, the study of non-aqueous solvents has developed greatly in some areas, and such solutions will be discussed where appropriate.

Much of this book, especially the earlier chapters, will be concerned with the nature and properties of metal ions, though much of the discussion is equally relevant to anions and to complex ions. The main reason for this imbalance is simply that some of the more fundamental aspects of the chemistry of ions in solution are better documented and understood for solvated metal ions than for other solute species. However, an aqua-metal ion is in reality only a special case of a complex, with water acting as ligand (see section 1.4).

1.2 METAL IONS AROUND THE PERIODIC TABLE

Before embarking on the various aspects of the chemistry of solvated ions, it may prove helpful to summarise the distribution of aqua-metal ions in relation to the Periodic Table. In Fig. 1.2, elements giving one or more well-established species of this type are shown with a tinted background. There are also several much less well-established aqua-cations, some of which are listed in Table 1.1. Many of these have been postulated in connection with the measurement of physical properties such as redox potentials or stability constants. Simple representations such as Au+, Bi3+, or Zr4+ are really shorthand for species which are more complicated and indeed are very difficult to characterise properly. Difficulties arise from strong tendencies to hydrolyse and polymerise, to form complexes, to disproportionate, or to oxidise or reduce solvating water.

Salts dissolve to a greater or lesser extent in a range of polar solvents, both hydroxylic solvents such as the alcohols and dipolar aprotic solvents such as acetonitrile or dimethyl sulphoxide. In all these cases the cations will be solvated, as in aqueous solution. Methanol, ethanol, and acetone are rather less effective in solvating metal ions than water, but dimethyl sulphoxide or pyridine solvate some cations considerably more effectively than water. In general metal ions which give hydrated cations in water can usually give analogous solvated cations in polar organic solvents. Simple inorganic anions, on the other hand, such as halides or oxoanions,tend to be weakly solvated in organic solvents. It is this feeble anion solvation which so often makes simple inorganic salts very sparingly soluble in organic solvents. To increase the chances of high solubility in organic media the choice of anion should fall on an essentially hydrophobic anion such as tetraphenylboronate.

1.3 NEW AQUA-METAL IONS

It should not be thought that all possible metal ions able to exist in aqueous solution have already been discovered and characterised. In the last few years aqua-ions of palladium(II), platinum(II), and molybdenum(III) have, rather belatedly, been properly established. There is no reason to believe that aqua-ions of, for example, rhenium(IV) and technetium(III) and (IV) may not soon be characterised, as suggested at the foot of Table 1.1. The types of approach which have proved successful are illustrated in Fig. 1.3, which shows the methods of preparation used for generating aqua-ions of molybdenum(III), platinum(II), and iridium(III).† The preparative methods shown in Fig. 1.3 illustrate several important points, especially in relation to complex formation. If it is necessary to add acid to control pH, the anion of the acid added must not form a complex with the potential aqua-cation. Chloride forms quite stable complexes with many metal ions, and thus hydrochloric acid is to be avoided. From this point of view perchloric acid, p-toluenesulphonic acid, and trifluoromethane sulphonic acid are to be preferred, but are still not ideal. It has been demonstrated recently that there are significant interactions between p-toluenesulphonate and lanthanide cations, while the coordination of trifluoromethylsulphonate to such metals as tin, iron, and palladium is well-established through X-ray structural studies. Indeed there are now extensive kinetic results on solvolysis and base hydrolysis of [M(O3SCF3)(NH3)5]2+ cations, with M = e.g. Co, Rh, and Ru. In like vein, there is much evidence for complex formation involving perchlorate; titanium(IV) perchlorate, Ti(ClO4)4, is in fact an uncharged complex with four bidentate perchlorate ligands firmly bonded to the titanium. Tetraphenylboronate, [BPh4]−, is a non-coordinating anion, but has the disadvantage of forming rather a large number of salts which are practically insoluble in water. It is also not particularly stable in acidic solution. However, it is useful for avoiding complex formation and minimising ion-pairing in solutions containing hydrophobic cations in organic solvents, where its hydrophobic nature enhances solubility (see section 1.2). Even better is the [B11CH12]− anion, in 1986 awarded the title of ‘least coordinating anion’. In organic solvents of low dielectric constant and low anion solvating power, perchlorate is quite liable to form complexes, as are ions such as 6−,BF4−, or 3SO3−. These are all more reluctant to form complexes in aqueous media. The other requirement for the anion in the present context is that it should not undergo redox reactions with the solvated...