The THEORY of ACIDS and BASES
in ANALYTICAL CHEMISTRY1 LOUIS P. HAMMETT Columbia University, New York City
T
HE subject of inorganic analytical chemistry can be presented as a series of unconnected empirical procedures; as a practical art instead of a branch of science. The position that it should be taught in this way has, however, been difficult to defend ever since Ostwald showed how brilliantly the Arrhenius theory of ionization illuminated the meaning of the practices which empirical investigation had developed. It has now few supporters, and the dual principle that the theory of ionization is the best way of understanding analytical chemistry, and that analytical chemistry offers the best introduction to the theory of ionization is generally accepted. Granted that analytical chemistry and the theory of ionization should be taught together, the next question is, what theory of ionization? The original t h e o j of Arrhenius, modified according to Bjerrum and Debye for the case of strong electrolytes, is competent to account for all of the phenomena of the classical inorganic analytical chemistry of aqueous solutions; most of us who are teaching analytical chemistry now were brought up in it; it is comfortable and familiar and therefore seems easier than the proton transfer theory of acids and bases, the modification which we are now urged to adopt; and the need for the new theory bas arisen from the study of the kinetics of acid and base catalyzed reactions and the properties of nonCaqueoussolutions, phenomena which seem to have little t o d o with analytical chemistry. Why then, runs a common argument, should we alter our instruction to include the new theory? This would he a good argument if analytical chemistry stood alone and independent of the rest of chemistry, a means and an end in itself. But this is not so; analytical chemistry is a part of chemistry, and instruction in analytical chemistry is part of a curriculum whose aim is to impart the maximum understanding of chemical phenomena. Given two theories, both satisfactory for analytical chemistry, there can be no question that the one which is most generally useful in chemistry as a whole is the one we should teach. Any other choice points toward a sterile classicism which is dangerous both for the future of analytical chemistry and for the future of the science as a whole. There can be little doubt which is the most useful
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Contribution to thr Symposium on Thcoricr and Teaching of Acids and R a w conducted by the Divisions of Physical Chem. irlrv and Chemical l i d t ~ a t i o nat thr ninetv-sevmrb mcctina - of the-A. C. S.. Baltimore. Md., April 4. 1939.-
theory, for the Arrhenius theory can take account only of reactions in a single solvent, which is, of course, always water, and only of equilibrium systems in that solvent. The proton transfer theory, on the other hand, can tell the beginning student why hydrogen chloride and sodium chloride are so much alike when they are dissolved in water and so different under any other conditions. It can offer the organic chemist the opportunity of explaining why he uses sulfuric acid or zinc chloride or sodium amide as condensing agents in one or another reaction. I hope i t points toward the day when our physical chemical tbeory is so good that an organic chemist cannot afford, as he now can, to be ignorant of it
Indeed one can hope for advances in analytical chemistry itself if we can escape from the vicious circle which trains analytical chemists only in the theory of dilute aqueous solutions because this is the field of analytical chemistry and consequently sees to i t that this remains the only field of analytical chemistry. We have already seen2 a new and valuable analytiPa1 method for the determination of %minoacids grow out of Conant's and Hall's theoretical investigations of acid-base systems in the solvent acetic acid, and there is every reason to expect that analytical chemists properly trained in theory will develop many more metbods utilizing nonaqueous solvents. The possibilities -of sulfuric acid, sulfur dioxide, ammonia, or hydrogen fluoride are so far absolutely untouched. The teaching of an adequate but limited theory in an elementary course might be justified if it were really simpler and more readily assimilable than the more general theory. There is, however, no evidence in my experience that this is the case, and I have taught both theories. Of course, the new theory is difficult to those of us trained in the older theory, but the difficulties we meet do not exist for students who have never met the older theory. Even now my own tongue stumbles over the new terminology, while it slips glibly off the lips of my students, both undergraduate and graduate; but I lived fifteen years with the old theory and the old names before I met the new ones, and the habits of fifteen years are hard to break. There are even some ways in which phenomena important in analytical chemistry can be more simply presented in terms of the new terminology than was ever possible with the old. If, for instance, one establishes NADEAU AND BRANCEEN. 3. Am. Chem. Soc., 57.1363 (1935).
the idea that the tenn acid includes ions like ammonium ion as well as neutral molecules, that the term base also includes ions as well as neutral molecules, and that any conjugate acid-base system comes to an equilibrium with the solvent, thus Acid
+ HnO F Base + OHs+
only a few words are required to account for all the properties of indicators and of buffer solutions, and to lay the foundation for the discussion of acid-base titrations. Thus an indicator is a conjugate acid-base system in which the color of the acid and base are different. The acid may be neutral and the base a negative ion as with 9-nitrophenol, the acid may be positively charged and the base neutral as with aminoazobenzene, or the acid may be a singly charged negative ion and the base a doubly charged one as with phenolphthalein; in any case
of acetic acid plus acetate ion, of H2P04- plus HPO&or any other. In any case [OH8+]= K . [ A c i d ] / [ B a s e ]
10H-
1 =! &X [Base] K. IAridl
The buffer property depends upon the fact that a small amount of added or adventitious acid or base can only convert a small fraction of the butrer acid to base or vice versa. Consequently i t can change the buffer ratio and the concentrations of oxonium or hydroxyl ions only slightly. If the acid is OHs+ and the base HaO, as in a solution of a strong acid in water, the buffer property of resisting change in acidity is present in fact and appears from the equation, although the solution is not usually called a buffer. It will be noted that these explanations do not involve any mention of hydrolysis, as the older ones did. As a matter of fact hydrolysis as a distinct phenomena vanishes in the new treatment. The hydrolysis reaction NH4+
+ H1O = N H I + OHs+
is no diierent in principle from the ionization Consequently a colorimetric determination of the ratio H G H a O z f HIO G H a O 1 - OH*[Acid]/ [Base] in a solution of known pH measures the and both are described by the same equation value of K., and, with this known, a colorimetric measurement in any solution determines the pH of -the solution. A number of obvious corollaries follow. At the half-way point in the color change, i. e., when rAcid] = [Base] then [OHa+] = K. and pH = pK.. The range with a the fraction hydrolyzed or ionized and K. the of usefulness of an indicator is the region of pH on acidity constant of the acid-base system. either side of the value of pK. for which neither [Acid] In the same way the course.of the titration of an acid nor [Base] is too small for measurement. This is about with sodium hydroxide, with ammonia, or with sodium all one needs to know about indicators in order to use carbonate may be described by identical equations, the them; the next refinement would be to consider the bases being hydroxyl ion, ammonia, and carbonate ion, effect of the ionic strength of the medium on the value respectively, in the three cases. This is very satisfactory of K,, an effect which differs according to the charge because it has always been obvious that the three titratype. The question why the acid and base have differ- tions are similar in fact, and we used to have to apoloent colors is an interesting problem in organic chemical gize and say that the similarity'was only apparent. theory but has nothing to do with t t e use of the indicaThese are only samples drawn almost a t random, but tor. they do show that the new theories and terminologies Similarly, a buffer solution is a solution containing an in acid-base theory may simplify rather than complicate acid and its conjugate base both a t relatively high cou- the treatment of the most familiar phenomena of anacentrations. I t may be of the type of NH4+ plus NH,, lytical chemistry.
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