Solutions, electrolyte vapour pressure

Experiments on sufficiently dilute solutions of non-electrolytes yield Henry s laM>, that the vapour pressure of a volatile solute, i.e. its partial pressure in a gas mixture in equilibrium with the solution, is directly proportional to its concentration, expressed in any units (molar concentrations, molality, mole fraction, weight fraction, etc.) because in sufficiently dilute solution these are all proportional to each other. 

These can be prepared by electrolytic oxidation of chlorates(V) or by neutralisation of the acid with metals. Many chlorates(VII) are very soluble in water and indeed barium and magnesium chlorates-(VII) form hydrates of such low vapour pressure that they can be used as desiccants. The chlorate(VII) ion shows the least tendency of any negative ion to behave as a ligand, i.e. to form complexes with cations, and hence solutions of chlorates (VII) are used when it is desired to avoid complex formation in solution. 

The electroneutrality condition decreases the number of independent variables in the system by one these variables correspond to components whose concentration can be varied independently. In general, however, a number of further conditions must be maintained (e.g. stoichiometry and the dissociation equilibrium condition). In addition, because of the electroneutrality condition, the contributions of the anion and cation to a number of solution properties of the electrolyte cannot be separated (e.g. electrical conductivity, diffusion coefficient and decrease in vapour pressure) without assumptions about individual particles. Consequently, mean values have been defined for a number of cases. 

When a 60 MW turbine at Hinkley A power station disintegrated in 1969 from stress corrosion cracking of a low pressure turbine disc (consequences shown in Plate 1) it was considered that Na H solutions were most probably involved (84) and it was soon found that if NaOH were the sole electrolyte present its maximum concentration (based on vapour pressure depression) was sufficient to have caused the cracking. However, it was also found that in mixtures it was only the free NaOH which led to rapid stress corrosion cracking. Considerations of acid gas solubility and solution thermodynamics showed that at the CO2 and acetate levels present it was most unlikely that free NaOH was present in sufficient quantity to be responsible for the Hinkley failure (85). 

Other properties of aqueous solutions have been studied, such as the density,13 vapour-pressure,14 boiling-point,15 molecular depression of the freezing-point,16 electric conductivity,17 electrolytic dissociation,18... 

The -maxima and minima on viscosity-composition curves are reminiscent of those on vapour pressure-composition curves of binary, mixtures. 5 The vapour pressures and viscosities are equal at some temperatures, say T and To, and T and To respectively. Then To/T—To7T =C(T —T), where C is a constant. A plot of TojT—To IT against T—T gives a straight line in many cases, both for vapour pressure and viscosity in other cases, the vapour pressure shows a minimum and the viscosity a maximum, and the vapour pressure a maximum and the viscosity a minimum. Prasad, 6 from the relation with vapour pressure deduced the equation rj =rjjrio= +ac, where c=conc. of non-electrolyte. The theoretical value of a is 0 00652 the observed values were glucose 0 44, fructose 0 44, sucrose 0 78, independent of temperature. According to Errera, the curves depend on the electric dipolarity of the liquids if both are nonpolar, the curve is concave to the composition axis whilst if both are polar, it is convex. Wolkowa found that the viscosity of a solution is approximately proportional to its heat of dilution. There seems to be no relation between the viscosity and surface tension of a mixture of acetic acid and water (cf. salt solutions, 13.VIII E). Mixtures of isomorphous substances obey an approximately linear relation. 

A quantitative test of these relationships is not yet possible, as the heats of dilution of solutions of non-electrolytes have not been determined, and the available data on the vapour pressures of moderately concentrated solutions are too inaccurate. 

It is thus possible in principle to determine the mean activity coefficient of the electrolyte by dividing the actual vapour pressure of the electrolyte over the solution by that over an ideal solution which is obtained by finding by extrapolation to very dilute solutions (c/. 21.27). In fact the methods outhned in the two following paragraphs are more precise and more easily applied. 

It has already been pointed out that we may take the vapour pressure of a substance as a measure of its active mass, and therefore of its activity In the case of homogeneous reactions in gases it may be recalled that the law of mass action has been fully verified, the terms employed in the expression for K being simply the concentration terms of the various constituents In the case of gases, and therefore in the case of vapours at low concentrations, the concept of activity becomes identical with concentration On this basis the relative activities of the undissociated molecules of an electrolyte will be correctly measured at various concentrations of the solution by determining the values of the partial pressure of the molecules in the vapour in equilibrium with the solutions In the majority of cases the un-dissociated molecules of strong electrolytes are not sufficiently volatile to become measureable in this way There are, however, a few electrolytes in which this is possible, namely, aqueous solutions of HC1, HBr, and HI... 

It is known from physical chemistry that the equilibrium vapour pressure is smaller over solutions than over pure water. In the case of ideal solutions this vapour pressure decrease is proportional to x0, the mole fraction of the solvent (Raoult s law). If the solution is real, the interaction of solvent and solute molecules cannot be neglected. For this reason a correction factor has to be applied to calculate the vapour pressure lowering. This correction factor is the so-called osmotic coefficient of water (g ). We also have to take into account that the soluble substance dissociates into ions, forming an electrolyte. 

Electrolyte solutions are solutions which can conduct electricity. Colligative properties such as the lowering of the vapour pressure, depression of the freezing point, elevation of the boiling point and osmotic pressure all depend on the number of individual particles present in solution. They thus give information about the number of particles actually present in solution. For some solutes it is found that the number of particles actually present in solution is greater than would be expected from the formula of the compound. 

Isopiestic Method.— The isopiestic method can be used to determine provided only one of the components is volatile. The unknown solution and reference solution containing the same volatile component but different involatile components are placed in dishes on a metal block in an airtight enclosure. The volatile component distils between the two solutions until their compositions are such that they have the same chemical potential of volatile component. After equilibration the two solutions are analysed. As the isopiestic method is a comparative method it is necessary to know the variation of vapour pressure with composition for the reference solution. The technique has the advantage that no pressure measurements are required. The method has been frequently used in the study of electrolytes and involatile non-electrolytes in aqueous solutions. Corneliussen et a/. have used the method to study polymer solutions. An apparatus suitable for measurements with organic mixtures has been described by Harris and Dunlop. The method is suitable only when there is no possibility of distillation between the involatile component in the reference solution and the involatile component in the unknown solution. 

The colligative properties of solutions, e.g. osmotic pressure, boiling point elevations, freezing point depression and vapour pressure reduction, depend on the effect of solute concentration on the solvent aetivity. The ehemieal potential, yt, of a non-electrolyte in dilute solution may be expressed by... 

Calculated from the equation ln(H) = 11.04 —5196/T + 0.03998 1, where I is the ionic strength in molarity and T is the temperature in degree kelvin. This equation was obtained by linear regression of the data from Imakawa, H., Chemical reactions in the chlorate manufacturing electrolytic cell (part 1) the vapour pressure of hypochlorous add on its aqueous solution, J. Electrochem. Soc. Jpn., 18, 382, 1950 and Imagawa, H., Studies on chemical reactions of the chlorate cell (part 2) the vapour pressure of hypochlorous acid on its mixed aqueous solution with sodium chlorate, J. Electrochem. Soc. Jpn., 19, 271,1951. 

Imakawa, H., Chemical reactions in the chlorate manufacturing electrolytic cell (part 1) the vapour pressure of hypochlorous acid on its aqueous solution, J. Electrochem. Soc. Jpn., 18, 382, 1950. 

The impact of these liquid phase reactions on the phase equilibrium properties is thus an increased solubility of NH3, CO2, H2S and HCN compared with the one calculated using the ideal Henry s constants. The reason for the change in solubility is that only the compounds present as molecules have a vapour pressure, whereas the ionic species have not. The change thus depends on the pH of the mixture. The mathematical solution of the physical model is conveniently formulated as an equilibrium problem using coupled chemical reactions. For all practical applications the system is diluted and the liquid electrolyte solution is weak, so activity coefficients can be neglected. 

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