Electrolytes apparent dissociation

Fig. 1.2 Dependence of the dissociation degree a of a week electrolyte on molar concentration c for different values of the apparent dissociation constant K (indicated at each curve)... 

The excess electrolyte is often termed the indifferent electrolyte. From the practical point of view, solutions containing an indifferent electrolyte are very often used in miscellaneous investigations. For example, when determining equilibrium constants (e.g. apparent dissociation constants, Eq. 1.1.26) it is necessary only to indicate the indifferent electrolyte and its concentration, as they do not change when the concentrations of the reactants are changed. Moreover, the indifferent electrolyte is important in the study of diffusion transport (Section 2.5), for elimination of liquid... 

In this formula K m is the dissociation constant expressed solely by the equilibrium concentrations, according to the classical Guldberg-Waage interpretation of the law of mass action. This value is identical with the true thermodynamical dissociation constant Km in highly diluted solutions only, for which the mean activity coefficient y+w very nearly equals unity. In all other solutions K m is not a true constant, but it depends on the actual concentration and on the presence of additional electrolytes therefore, it is called the apparent dissociation constant, in contradistinction of the true dissociation constant. For concentration expressed in terms of molarity, a similar equation is valid-... 

In methanolic solution, [Rh2(DIPHOS)2] [BF4]2 apparently dissociates into mononuclear [Rh(DIPHOS)]" ions (presumably containing coordinated solvent), as demonstrated by (i) electrical conductance measurements which yielded a slope of — 270 ohm" corresponding to a 1 1 electrolyte (20) for a plot of equivalent conductance vs. Vconcn (ii) P NMR measurements which revealed only a single P signal (d, 2P, 8 80, /Rh-p = 203 Hz) and (iii) measurements on the equilibria for the formation of various 1 1 alkene and arene adducts of [Rh-(DIPHOS)] see below). When base (OMe" or a sterically hindered amine such as triethylamine) was added to a methanolic solution of [Rh(DIPHOS)] an irreversible (i.e., not reversed by addition of acid) yellow to red-brown color change was observed, to yield a new species, [Rh3(DIPHOS)3(OMe)2] (4, A.ax 445 nm, e .ax 3.3 X 10 cm" P... 

This last condition is fulfilled when the ionic concentrations are very low, as they are in fact in dilute solutions of weak electrolytes. The dissociation constants of substances such as weak organic acids can be determined by a combination of the formulae of Ostwald and Arrhenius, but the procedure is quite inadmissible for salts. Here the degree of dissociation is large. In fact the value of a is often indistinguishable from unity, and the mutual influences of the ions are considerable. They are calculable in principle by methods due to Debye and Hiickel, and operate differently on different properties. The procedure outlined on p. 276 allows the calculation of the activity coefiicients. In general the thermodynamic properties of the salt in solution correspond to those of a system with apparently incomplete dissociation, not because the concentrations of the ions are reduced by molecule formation but because the activity coefficients are lowered by mutual ionic influences. 

K is the true dissociation constant of the electrolyte and is independent of composition because of the characteristics of the It is only equal to the apparent dissociation constant... 

Since the apparent dissociation constants are usually reported as pK,(/ 0 7) and not as K (/ 0 7) values, the scattering of results and differences caused by different ionic media and temperatures are less evident in the hterature presentations than they actually are (Figs. 3.12, 3.13, 3.14 and 3.15). Significant differences in the apparent dissociation constants are observed with regard to apphed supporting electrolyte, the ionic strength and temperature. 

As the titration begins, mostly HAc is present, plus some H and Ac in amounts that can be calculated (see the Example on page 45). Addition of a solution of NaOH allows hydroxide ions to neutralize any H present. Note that reaction (2) as written is strongly favored its apparent equilibrium constant is greater than lO As H is neutralized, more HAc dissociates to H and Ac. As further NaOH is added, the pH gradually increases as Ac accumulates at the expense of diminishing HAc and the neutralization of H. At the point where half of the HAc has been neutralized, that is, where 0.5 equivalent of OH has been added, the concentrations of HAc and Ac are equal and pH = pV, for HAc. Thus, we have an experimental method for determining the pV, values of weak electrolytes. These p V, values lie at the midpoint of their respective titration curves. After all of the acid has been neutralized (that is, when one equivalent of base has been added), the pH rises exponentially. 

A study of the concentration dependence of the molar conductivity, carried out by a number of authors, primarily F. W. G. Kohlrausch and W. Ostwald, revealed that these dependences are of two types (see Fig. 2.5) and thus, apparently, there are two types of electrolytes. Those that are fully dissociated so that their molecules are not present in the solution are called strong electrolytes, while those that dissociate incompletely are weak electrolytes. Ions as well as molecules are present in solution of a weak electrolyte at finite dilution. However, this distinction is not very accurate as, at higher concentration, the strong electrolytes associate forming ion pairs (see Section 1.2.4). 

As with solubility, Kow is a function of the presence of electrolytes and for dissociating chemicals it is a function of pH. Accurate values can generally be measured up to about 107, but accurate measurement beyond this requires meticulous technique. A common problem is the presence of small quantities of emulsified octanol in the water phase. The high concentration of chemical in that emulsion causes an erroneously high apparent water phase concentration. 

In the approximate treatment of the conductance of weak electrolytes, the decrease in A is treated as resulting only from changes in the degree of dissociation, a. On this basis, it can be shown that an apparent degree of dissociation a can be obtained from... 

When the apparently penta-coordinated diarsine complexes just described are dissolved in solvents more polar than nitrobenzene, they tend to dissociate into halide ions and bivalent cations, thus becoming 2 1 electrolytes (119). The effect is most marked with the platinum compounds. It has been shown that solvation effects might be less with platinum than with palladium, and so, other things in the equilibria being equal, it can also be concluded that the bonding of further ligands by a square-planar complex is much weaker with platinum than with palladium. Square-planar nickel complexes are of course the most ready to take up further ligands. 

Electrical Conductance of Aqueous Solutions of Ammonia and Metal Hydroxides. Check the electrical conductance of 1 W solutions of sodium hydroxide, potassium hydroxide, and ammonia. Record the ammeter readings. Arrange the studied alkalies in a series according to their activity. Acquaint yourself with the degree of dissociation and the dissociation constants of acids and bases (see Appendix 1, Tables 9 and 10). Why is the term apparent degree of dissociation used to characterize the dissociation of strong electrolytes ... 

The great increase in complexity in solution thermodynamics which occurs when a salt is dissolved to substantial concentration in a mixture of two liquid components becomes fully apparent in the realization that the liquid phase so created is a concentrated solution of an electrolyte whose degree of dissociation is a function of the relative proportions of the other two components present, and... 

According to modem theory, many strong electrolytes are completely dissociated in dilute solutions. The freezing-point lowering, however, does not indicate complete dissociation. For NaCl, the depression is not quite twice the amount calculated on the basis of the number of moles of NaCl added. In the solution, the ions attract one another to some extent therefore they do not behave as completely independent particles, as they would if they were nonelectrolytes. From the colligative properties, therefore, we can compute only the "apparent degree of dissociation" of a strong electrolyte in solution. 

There are some cases where a reaction, that is, the formation or dissolution of a chemical bond, is involved along with ion exchange phenomena (Helfferich, 1983). Examples of this are acid-base neutralization, dissociation of weak electrolytes in solution or weak ionogenic groups in ion exchangers, complex formation, or combinations of these (Table 5.2). With some of these, very low apparent D in ion exchangers have been noted. 

The determination of the molecular weight by depression of the freezing-point indicates some electrolytic dissociation.8 A value of 93 has been found for the molecular weight.9 The acid apparently... 

Recent surface force measurements revealed a similar trend (20). Comparing steam-treated to flame-treated silica sheets using site-dissociation/site-binding model, a decrease in silanol surface sites and apparent decrease in average pKa was observed upon heat treatment. Furthermore, a repulsive force other than double-layer and van der Waals forces was observed 15 A from the surface. This repulsion was attributed to hydration of the surface and was found to be independent of surface treatment and electrolyte concentration. In Bums treatment, an arbitrary plane of shear was introduced to provide a best model fit (l 3). A value of 9 A from the surface for the plane of shear was determined from electro-osmosis measurements. 

When this is the case, the heat of reaction must be quite independent of the nature of the anion and of the cation, aa these are not affected by the reaction. This is clearly true for nitric and hydrochloric acids with all the bases given in the table. For sulphuric and carbonic acids, however, the conditions for the validity of the theory are apparently not fulfilled. In the first case, the heat of dilution of sulphuric acid amounts to 2000 cal., and this amount must be subtracted from the figure given in the table, as it is evolved when the alkali and acid are mixed. In the second case, carbonic acid is so weak an acid that it is practically undissociated. The heat necessary for the dissociation into ions therefore uses up part of the heat of neutralisation. From the table it follows that the electrolytic dissociation of J mol. HgCOg requires 13700 — 10200 = 3500 calories. The constant heat of neutrahsation 13700 cal. is the heat of ionisation of water, i.e, the quantity of heat required for the dissociation of water, and liberated on the combination of its ions. 

The factor i only occurs in solutions which are good conductors of electricity, and in 1887 Arrhenius succeeded in explaining these apparent deviations from the simple laws by his electrolytic dissociation theory. The molecules of an electrolyte are broken up to a greater or less extent into their free ions, even when the solution is not conducting a current of electricity. Thus we have the equation HCl H - - CL... 

Firstly, ion exchange resins when hydrated generally dissociate to yield equivalent amounts of oppositely charged ions. Secondly, as with conventional aqueous acid or alkali solutions, resins in their acid or base forms may be neutralized to give the appropriate salt form. Finally, the degree of dissociation can be expressed in the form of an apparent equilibrium constant (or pK value) which defines the electrolyte strength of the exchanger and is usually derived from a theoretical treatment of pH titration curves. ... 

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