In a solution of an electrolyte, it is often necessary to have a detailed knowledge of the species present. New ions or uncharged molecules resulting from interactions in the solution may behave quite differently from the constituent ions of the electrolyte. Some properties of the solutions will be profoundly affected, and the chemist, in order to understand these phenomena, will require to know the nature of the species present. There are a number of formidable difficulties in the analysis of such systems and, during the past forty years or so, a great deal of work has been done on the problem. The equilibrium properties of electrolyte solutions and the way in which ion-pair and complex formation can be detected and quantitatively studied are of primary importance. Although the application of new physical and chemical methods has produced significant contributions in this field, the information obtained from measurements of a system at equilibrium is to some extent limited, and in studying the phenomenon it is desirable to know the relevant kinetic parameters. Without this understanding, it is sometimes impossible to sketch the actual reaction mechanism by which the system approaches equilibrium. In general, we may regard the elucidation of the structure of an electrolyte solution as a difficult problem which requires as many independent lines of attack as possible.
A good deal of success in the study of molecules in the gas phase prompted chemists to attempt to build up theories of solutions in an analogous way. The classical theories regarded the solvent as merely providing space in which the solute particles moved and interactions between the ions and the solvent molecules were neglected. This assumption can be questioned on the basis of even the most elementary electrostatic considerations.
As long ago as 1833, both Faraday and Daniell concluded that electrolysis took place through the transport of electricity by mobile charged particles or ions which were discharged at the electrodes. These ions were produced simply by dissolving the electrolyte in the solvent and so the concept of bond-breaking in the molecules of electrolyte was first established. It is now realized that the energy required for such a process comes from the solvation of the ions. When the ions are introduced into solution, they interact with solvent molecules and a considerable heat of solvation may be involved. In order to understand such concepts, it is necessary to have a more detailed picture of the structure of the solvent molecules. Although non-aqueous solutions are of considerable interest, much of the work has been done in aqueous system and these continue to be of paramount importance.
Translate the text in written form. Mind the Infinitive Constructions.
Ideal Solutions The concept of an ideal or perfect gas was found to be very useful, so it is advantageous to introduce the idea of a perfect liquid solution. Among the criteria set up for an ideal solution is its exact obedience of Raoultís law. This law relates the magnitude of certain properties common to all solutions to the effect produced on the properties by a change in the composition of the solutions. The vapour pressure, boiling points, freezing points and osmotic pressure of solutions have been seen to be properties that are dependent upon the number and not the kind of the molecules concerned. Such properties are said to be colligative and a solution is considered to be ideal if its colligative properties are capable of exact mathematical treatment in terms of the laws describing these properties.
There proved to be deviations of solutions from Raoult`s law. For example, the molecular weights obtained by the freezing-point method are only approximate and not exact. There appeared to be difference between observed and calculated values for osmotic pressure. The extent to which a solution deviates from the ideal state is believed to depend upon the nature of the solution concerned. In general, an increase in dilution causes all solutions to approach ideality. In many cases the upper limit of the concentration has been estimated to be about 0.01 molar, while for others the limiting concentration may be either much higher or considerably lower.