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INTRODUCTION

1.1 THE SIGNIFICANCE AND PHENOMENOLOGY OF IONS IN SOLUTION

Chemistry is for a large part conducted in solutions involving ions and such solutions are ubiquitous in nature. Oceans are vast aqueous solutions of salts, consisting mainly of sodium chloride, but other salts and minor components are also present in ocean water. Lakes, rivers, and brackish water are dilute solutions of ions and are essential to survival, since they provide drinking water and water for irrigation. Rain and other precipitates may remove ionic species from the atmosphere that arrived there as spray from oceans and seas or from human activities, for example, acid rain. Physiological fluids consist mostly of water in which colloidal substances, but also ions essential to their function, are dissolved.

It appears from the above that water is the only medium in which ions play a role, but this picture is too narrow because human endeavors utilize many other liquid media in which ions are present and have an active role. The manufacture of organic substances, as raw materials or intermediates in many industries, such as textiles, drugs, and food additives, generally involves reactions carried out in mixed aqueous-organic or completely nonaqueous liquid media in which ions participate. In chemical analysis, such media have long been of invaluable use, for instance in electroanalytical measurements or chromatographic separations. Industrial uses of nonaqueous media involving ions include solvent extraction in hydrometallurgy or in nuclear fuel reprocessing and nuclear waste disposal (Chapter 8).

The ions involved in these as well as other systems and applications interact with each other and with nonionized solutes that may be present. These interactions are of prime interest to the chemist, but the extent, intensity, and rate of proceeding of these interactions depend heavily on the solvent or solvent mixture present, a fact that is not always clearly recognized by the operator. The ion-solvent interactions should be understood in order to make the best use of the solutions of the ions, since it is the solvated ions that take part in the interactions of interest. If a free choice of the solvent or the solvent mixture to be used is possible, the most suitable one for the purpose should be selected on the basis of the knowledge available on the interactions that take place, bearing in mind also costs, ecology, and hazards. If the solvent is prescribed, this knowledge is still needed in order to select the proper reaction conditions or the additives that could be useful. So-called "bare" or nonsolvated ions occur in gas-phase reactions (Chapter 2) but not in condensed media, that is, in solutions. A seeming exception to this generalization is the use of room temperature ionic liquids (RTILs) as reaction media, where all the ions are surrounded by ions of the opposite charge sign rather than by a nonionic liquid medium. Whether the ions of RTILs are called "bare" or "solvated" is a semantic question. In common situations, which are the subject of this book, there is always an excess of a nonionic liquid medium in which the ions find themselves, the molecules of which surround the ions more or less completely, unless some other species, be it another ion (of opposite sign) or a solute molecule (a ligand) replaces some of the solvent molecules in the ionic solvation shell.

Ions cannot be added individually to any major extent to a solvent or a solution; it is always an electrolyte consisting of ions of both signs in a combination that makes the electrolyte electrically neutral, which is added to form a solution involving ions. Many commonly used and studied electrolytes are crystalline solids, such as NaCl or (C4H9)4NClO4. The electrostatic energy that holds the ions constituting such crystals together, the lattice energy that must be invested in order to separate the ions in the solution, is compensated by the solvation energy that is gained in the process of dissolution, with some effect also of the entropic changes encountered in the process. Some potential electrolytes are gaseous, for example, HCl, but they produce ions only on reaction with the solvent in which the covalent H-Cl bonds are broken and replaced with others to compensate for the energy involved. On the other hand, ions may leave the solution, if not individually then as a small combination of ions, in electrospray experiments, in which they are then monitored in the gas phase by mass spectrometry. The results of such experiments have some bearing on the state of the liquid ionic solutions, but this subject is outside the scope of this book.

It should be kept in mind that, connected with such ion solvation reactions with crystalline or gaseous electrolytes, a further reaction takes place, which is not always recognized, namely the breaking of some solvent-solvent molecular interactions, required to produce the space to accommodate the ions in the solution. It is the balance of all the (Gibbs) energies that have to be invested and those that are gained that determine the extent to which an electrolyte will dissolve in a given solvent (Chapter 4).

Some electrolytes are completely dissociated into "free," that is, solvated, positively charged cations, and negatively charged anions. Other electrolytes are only partly so dissociated, depending on the concentration and on the nature of the solvent. Some substances are ionogenic, in the sense that some dissociation into ions occurs only under specific conditions, and these include also so-called "weak electrolytes" such as many acids and essentially basic substances in aqueous solutions.

A solution of a single, individual ion in a very large amount of solvent could presumably be the basis of a study of ion solvation which is not encumbered by other interactions. This situation cannot be achieved in the laboratory but can be dealt with as a thought process and now for many years also in computer simulations. The results of the latter have by now consolidated into a large body of knowledge that is constantly not only extended but also improved by the level of sophistication that can nowadays be achieved in such simulations. Such results are incorporated into the discussions in the present book, where they are compared with laboratory experimental data obtained on electrolytes, extrapolated to infinite dilution (Chapters 4 and 5). At such high dilutions, each ion is surrounded by solvent molecules only and does not interact with the very remote ions of the opposite charge sign that must be present somewhere in the solution. Still, the allocation of the extrapolated values of the electrolyte properties to its constituent individual ions is a problem that must be solved.

For the purpose of only reducing the number of items in the properties list from the many electrolytes (combination of cations and anions), that have been derived at infinite dilution to the much smaller number of individual ions that constitute them, it is sufficient to employ the so-called "conventional" values. These are based on assigning to one ion, say the solvated hydrogen ion, an arbitrary value (generally zero) and rely on the additivity of individual infinite dilution ionic values to derive values of all other ions. The sum of the conventional ionic values, weighted according to the stoichiometric coefficients (the numbers of ions of each kind constituting the electrolyte), expresses correctly the infinite dilution property of the electrolyte. Within a given charge sign series of ions, say cations only, comparisons between conventional values of diverse ions can throw some light on the effects of the individual ionic properties, such as size and valency, but the cation and the anion series cannot be compared with each other.

The problem of assignment of the so-called "absolute" individual ionic values to these infinite dilution electrolyte data is solved mainly on the basis of chemical intuition (Chapter 4) that can be assisted by the results from computer simulations. Once individual ionic values of their properties in a given solvent (or solvent mixture) at a given thermodynamic state [temperature and pressure, usually specified as 298.15?K (25°C) and 0.1?MPa (less commonly now 1?atm?=?0.101325?MPa)] have been established, they may be compared with other properties of the ions (e.g., their sizes) or with theoretical expectations (models). The latter are the main incentives to obtaining the absolute values. Such comparisons and correlations provide insights into the ion-solvent interactions that take place and form the basis for understanding interactions of ions with other solutes, be they ionic themselves or nonionic.

There are some experimental measurements that can be made on solutions of ions that pertain directly to individual ions. These include transport properties, such as the ionic conductivities that are obtained from specific conductivities of electrolytes in conjunction with transport number measurements. Diffusivities of individual ionic species can also be measured by the use of isotopically labeled ions and should be compatible with the mobilities deduced from the ionic conductivities. Spectroscopic data can also in certain cases be due to individual ionic species, such as NMR chemical shifts and relaxation rates of the signals from appropriate nuclei (e.g., 7Li or 27Al). Such information may be used as a guide for the "chemical intuition" mentioned earlier needed for obtaining absolute...