Strong Electrolytes Which Complex

The above treatment applies directly to strong electrolytes which completely dissociate to their constituent ions. The factor W in equation (8.17) accounts for the type of salt Involved. In Chapter VI, we saw that certain salts complex. For example, the strong electrolyte reaction, (Case 1) 

a different Ionic strength results in each case and the W factor would have to be adjusted accordingly. 

The problem highlighted in Case 3 above becomes even more significant in the case of weak electrolytes. This can be seen, in the case of COz. where  

The distribution among C0 2(aq), HCOf and CO is strictly determined by the simultaneous solution of the constituent equilibrium electroneutrality and material balance equations for the system. The total ionic strength is thus obtained only by solving this set. In order to handle weak electrolytes and stUl utilize the above framework, it would be necessary to  

2) Define the apparent molal volumes for all species (e.g. C02(aq), CO, HCO3) 

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The apparent acid strength of boric acid is increased both by strong electrolytes that modify the stmcture and activity of the solvent water and by reagents that form complexes with B(OH) 4 and/or polyborate anions. More than one mechanism may be operative when salts of metal ions are involved. In the presence of excess calcium chloride the strength of boric acid becomes comparable to that of carboxyUc acids, and such solutions maybe titrated using strong base to a sharp phenolphthalein end point. Normally titrations of boric acid are carried out following addition of mannitol or sorbitol, which form stable chelate complexes with B(OH) 4 in a manner typical of polyhydroxy compounds. EquiUbria of the type ... 

In aqueous electrolyte solutions the molar conductivities of the electrolyte. A, and of individual ions, Xj, always increase with decreasing solute concentration [cf. Eq. (7.11) for solutions of weak electrolytes, and Eq. (7.14) for solutions of strong electrolytes]. In nonaqueous solutions even this rule fails, and in some cases maxima and minima appear in the plots of A vs. c (Eig. 8.1). This tendency becomes stronger in solvents with low permittivity. This anomalons behavior of the nonaqueous solutions can be explained in terms of the various equilibria for ionic association (ion pairs or triplets) and complex formation. It is for the same reason that concentration changes often cause a drastic change in transport numbers of individual ions, which in some cases even assume values less than zero or more than unity. 

When an acid in solution is exactly neutralized with a base the resulting solution corresponds to a solution of the salt of the acid-base pair. This is a situation which frequently arises in analytical procedures and the calculation of the exact pH of such a solution may be of considerable importance. The neutralization point or end point in an acid-base titration is a particular example (Chapter 5). Salts may in all cases be regarded as strong electrolytes so that a salt AB derived from acid AH and base B will dissociate completely in solution. If the acid and base are strong, no further reaction is likely and the solution pH remains unaffected by the salt. However if either or both acid and base are weak a more complex situation will develop. It is convenient to consider three separate cases, (a) weak acid-strong base, (b) strong acid-weak base and (c) weak acid-weak base. 

The zinc halides behave in dilute solutions entirely as strong electrolytes, but in more concentrated solutions this is otherwise probably it is here a question of the direct formation of higher complexes, for example, [ZnX4]2-, the dissociation of which depends very much more on the concentration than for [ZnX]+. 

It can be seen that the added complexity of ion association is likely to make any simple model of ion-ion interactions very difficult to apply without a number of ad hoc assumptions concerning ionic radii. This is particularly true for ionic strengths in excess of 0.01 M or for low-dielectric-constant media. However, a further difficulty is raised by the problem of the nature of an ion pair. If we consider the simple case of univalent ions A+ B forming an ion pair, it is possible to picture the pair as varying in character from one in which the charges remain separated by the sum of the ionic radii of A+ + B to a molecule in which A and B form a covalent bond, not necessarily even polar in character. Nor is it necessarily true that a given species will behave the same in different solvents. If there is a tendency to covalent bond formation, then it is quite possible that the polarity of the A—B bond will depend on the dielectric constant of the solvent. Covalently bound molecules which ionize are considered as weak electrolytes, and they are not treated by the methods of Bjerrum, which are meant for strong electrolytes. The differences may not always be clear, but the important interactions for the weak electrolyte are with the solvent, and these we shall consider next. 

The arguments in favor of such a scheme are (juite strong but relatively complex and depend on a number of auxiliary data, kinetic, thermodynamic, and structural-chemical. The large activity effects on ions and Kion of weak electrolytes in the low-dielectric media have been quantitatively considered, and it appears that the active species in the system are not ions but rather ion pairs. This makes it likely that the same is true of some of the other work done in glacial acetic acid which has been probably incorrectly interpreted as ionic. 

The standard potentials of other metals, such as alkali metals, alkaline earth metals, aluminum, or titanium are so negative that these metals caimot be deposited from aqueous solutions. By choosing electrolytes which do not contain protons and strong complexing groups, it is possible to electrolytically deposit such metals which are not obtainable from aqueous electrolytes. Nonaqueous electrolytes may be either fused salts or solutions of metal compounds in organic or inorganic solvents. 

The Macinnes convention leads to = Tci = 7 kci, We can now compute individual ion activity coefficients from their mean values measured in solutions of strong electrolytes using y Kci values as our starting point. (In the ideal strong electrolyte, cations and anions are unassociated with each other and thus do not form complexes [see Chap. 3].) It is important to remember that all such calculations must be done with y values for KCl and other salts measured at the same ionic strength, which is not the same molality except for monovalent-monovalent salts. 

Pitzer s formulation offers a satisfactory and desirable way to model strong electrolyte activity coefficients in concentrated and complex mixtures. When sufficient experimental data are available, one can make calculations which are considerably more accurate than those presented in this paper. Attaining high accuracy requires not only experimentally-based parameters but also that one employ third virial coefficients and additional mixing terms and include explicit temperature dependencies for the various parameters. 

A student has prepared a cobalt complex that has one of the following three structures [Co(NH3)5]Cl3, [Co(NH3)5Cl]Cl2, or [Co(NH3)4Cl2]Cl. Explain how the student would distinguish between these possibilities by an electrical conductance experiment. At the student s disposal are three strong electrolytes— NaCl, MgCl2, and FeCl3—which may be used for comparison purposes. 

The subject of this chapter is ions which form complexes. As contrasted with strong electrolytes, discussed in Chapters IV and V, complexing species form ionic and/or molecular intermediates. A classic example is orthophosphoric acid which dissociates in water to form, among others ... 

Use of the Meissner family of curves as presented in Chapter IV, Figure (4.6), in order to correlate the reduced activity coefficients of strong electrolytes. Meissner found that the reduced activity coefficients of strong electrolytes fell into a pattern which he formalized. If it can be assumed that the reduced activity coefficients of strong electrolytes follow this pattern, then deviation from these curves may be said to indicate that the species in question does not completely dissociate or may be forming complexes. If activity coefficients from experimental data are available they may be plotted on the Meissner chart. Severe deviation from the curves, such as the crossing of lines, may be taken to indicate complex formation. However, there would be no indication of what the complexes might be. 

The most important message of this chapter is that nearly all electrolytes form one or more complexes over some range of their phase space with water and/or other compounds in water. For many species the conditions over which this occurs and the attendant quantities make it unreasonable to ignore the phenomena. Trying to bend strong electrolyte models to handle species which form complexes is often as invalid as using conventional VLB formulations to handle electrolyte solutions, ie. ignoring ion formation. 

The metabolism of certain aerobic bacteria produces strong acids, which can accelerate corrosion. The best known example are bacteria of the Thiobacillus family that are able to oxidize sulfides or sulfur compounds into sulfuric acid. Certain organic acids produced by bacteria have the ability to form chelate complexes with dissolved metal ions, changing the thermodynamic conditions for corrosion (Chapter 2). In some cases the chelates may precipitate with electrolyte cations and form a film. 

The parameter reflects interactions of the cation-constituent with the solvent and with itself. 1 likewise for the anion-constituent, while 1 is a measure of the degree of coupling between the motions of cation- and anion-constituent. The 1 approach has been most actively fostered by Don Miller (TtO 111. 115). An alternative method of representation, introduced by Gerhard Hertz (116.1171, is based on velocity correlation coefficients. There are six of these for a binary electrolyte with self-diffusion coefficients being required to evaluate them. Sets of vcc values have recently been published for concentrated solutions of electrolytes which exhibit strong complexation like Cdl (118). The main problem in calculating these turned out to be a deficiency in suitable transference number information ... 

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