reaction is valid only for equivalent amounts of acid and base and only for those cases where .water is not directly involved in the reaction. If water is involved, either as a reactant or as a product, the extent of reaction is dependent on the amount (initial concentration) of acid or base. For the reaction of a weak acid or weak base with water, the weU-known equation for the relationship between the fraction reacted, a , and the equilibrium constant may be derived,
where c is the amount (number of moles) of weak acid or weak base available to react per liter of solution. Similarly, for the reaction of a strone acid rH20', with u weak base or of a strong base (OH-) with a weak acid,
for equal initial "wnceutrations", c, of each of the acid and base. In either case the extent of reaction, a , is a function of the amount, c, for any particular value ofK. W. G. Baldwin C. E. Burchill University of Manitoba Winnipeg. MB, Canada R3T 2VI
idea that this constant is 55.5 from Starkey et al. (I), in spite of the fact that it was clearly refuted by Baldwin and Burchill (2).This value is obtained by expressing the wncentration of water for two sides of reaction 1in different ways. Thompson does it in a n attempt to apply, to an equimolar mixture of a strong acid and a weak base, the procedure that he previously used to calculate the extent of the reaction between a weak acid and a (not conjugated) weak base. Taking aniline as an example:
he obtains a concentration-independent value of 0.9994 for the extent of reaction 2. The correct approach to the solution of this problem is different. I t is very important for a student to understand the bidirectional (forward and backward) aspect of chemical reactions and realize that the reaction between a weak base and a strong acid is merely the reverse of the dissociation of the conjugated weak acid. In f a d reaction 2, if correctly interpreted, can he read: proton + weak base = conjugated weak acid (free and solvating water can he neglected, as usual). The equilibrium constant of this reaction is K = UKa=K& (where K, is the acidity constant of the weak acid and Kb the hydrolysis constant of the conjugated base), and not
To the Editor:
The interpretation of thermodynamic equilibrium constants by Baldwin and Burchill is quite proper and technically wrrect. The conventions employed by them would in fact make water and hydronium ion special cases. However. mv formulation of the acid-base reactions and the K, values Lsed are based upon the principle of the law of mass action, which almost invariably is the approach - used in lower undergraduate level instruction. The assertion by Baldwin and Burchill that K.(H20) = K, and Ka(H30+)= 1.00 can lead to erronious conclusious when comparing relative acidities of certain weak acids. Ethylene glycol (K. = 6 x 10-15) is a stronger acid than water, but if we choose K,as the K, value of water, then we reach t h e wrong conclusion about t h e relative acid strengths of these two compounds. This demonstrates the necessity to carefully state one's approach when doing calculations involving equilibrium constants, especially aqueous acid-base equilibria Ralph J. Thompson Eastern Kentucky University Richmond, KY 40475 The Extent of Acid-Base Reactions To the Editor:
Understanding the meaning of "completeness"of a reaction is very important for average first-year students (as well as for junior and senior ones !). To be effective, a discussion about the "extent of acid-base reactions" must be both simple and based on sound principles. This is not the case in Thompson's treatment of this topic [1990,67,2201, where severai misleading Rtatcments rnn he firund. It would be very difficult to explain to a student why the equilibrium constant of the rea&ion H30+(aq)+ HzO 2 HzO+H30t(aq)
(1)
should be different from unity. In fact, the equilihrium constant of reaction 1is equal to one, whether it be an activity or a concentration constant. Thompson has borrowed the
as obtained by Thompson. On the above grounds, the evaluation of the "completeness" of the reaction in a n equimolar solution of HC1 and aniline becomes trivial, because this solution is exactly equivalent (3) to a solution of the weak acid anilinium chloride at the same concentration, C. As a consequence, the "completeness" (a)of the reaction is related with the more widely used "dissociation degree" of the weak acid: "completeness" of forward reaction (2) formation degree of anilinium ion 1-(dissociation degree of anilinium ion) the dissociation degree is calculated from K, using a wellknown equation (4).Alternatively, the value of a is directly obtained from the equation
which is different from the one used by Thompson. The value of a depends on K. but also on the concentration of the solution (C) and is, for instance, 0.985 a t C = 0.1 M. Calculating with eq 3, the concentration for which a = 0.9994 (the "concentration independent" value obtained by Thompson) vields C = 64.6 M. As kegards the separation between carboxylic acids and phenols. also examined by Thompson, . . the mnditions of the problem are incorrectly stated by the author, who seems to consider separate equimolar solutions of HC03- and carboxilic acid and, respectively, H C 0 3 and phenol. In the real application, the ampholyte hydrogen carbonate ion is much more concentrated than the acids to be separated and acts as a buffer (although a poor one), fixing the solution pH at a value in the vicinity of 8.3 = 112 (pK1 + pKz) (the "pH of the ampholyte"). The ratio between the basic and acidic forms of either phenol or carboxylic acid, required to evaluate the "yield" of the separation, is correctly calculated by using the acidity constant of weak acids: Volume 69 Number 6 June 1992
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