Proton-Donors in the Electronic Theory of Acids and Bases
T
INTRODUCTION
:
HE exact of the Bronsted and Lewis . relatlonshxp ' theories of ac~dsand bases has not yet been fully clarsed. In view of the fact that there appears to be a tendency to imply a nonexistent antagonism between the two theories, i t seems advisable a t this time to give a full explanation of their relation to each other. To followers of Lewis, it is obvious that there is no conflict between the two theories. Lewis himself in the first statement of his theory (1) pointed out that the electronic theory of acids and bases includes as special cases both the proton-donor and the solventsystem theories. This has been restated several times since then (2, 3, 4). Yet it must be admitted that there is some ground for confusion. I t is not sufficient to say that since all acids are electron-pair acceptors, the proton as a powerful electron-pair acceptor is a special case among all the substances which show similar activity toward bases. The adherents of the Lewis theory should be able to demonstrate exactly how the type equation of the Bronsted theory fits into the broader picture. This paper represents an attempt to do just that.
W. F. LUDER Northeastern University, Boston, Massachusetts
electronic theory of acids and bases is founded are (5):
I I1 I11 IV
Neutralization Titration with Indicators Displacement Catalysis
The theoretical explanation (1,2, 3 , P ) of this experimental behavior is that an acid is capable of accepting a share in a lone electron-pair from a base to form a coordinate bond. A base donates a share in a lone electron-pair to the acid. The formation of the coordinate bond is the important step in all neutralization reactions:
HMROGEN ACIDS
The four experimental criteria upon which the
acid
base
In some cases formation of the coordinate bond is followed by ionization so that the product is a salt. For example, the compound CaHsN:AlBr3 seems to be a typical salt. Both pyridine and aluminum bromide are covalent compounds. When they are mixed, the white precipitate appears to have all the characteristics of a. salt. The electrical stress produced by the formation of the coordinate bond must result in the ionization of at least one of the bromine atoms:
acid
'
hse!
In other cases, the formation of the coordinate bond and the resulting ionization must be regarded as simultaneous. A few years ago, when the "hydrogen bridge" was regarded as involving 2-covalent hydrogen, the reaction between hydrogen chloride and water might have been written to correspond with equation
(4):
base
'acid (5)
The existence of the hypothetical intermediate addition compound in which the hydrogen bridge between the hydrogen chloride and water molecules involves 2covalent hydrogen is now regarded as unlikely. Prohably it would be better to represent the formation of the coordinate bond as taking place simultaneously with ionization as follows:
base
acid
This is exactly the same as the Bronsted theory pictures snch a reaction. To designate acids which owe their acidity to the proton, ~ e w i ssuggested the term hydrogen acid. This is onlv one examale of the manv which could be given to show that the electronic theory of acids and bases includes the proton theory as a s$ecial case. The simultaneous coordination and ionization pictured by the Lewis theory is equivalent to the proton-transfer mechanism of the Bronsted theory. How, then, is it possible to maintain that hydrogen acids need special consideration in the Lewis terminology? Apparently the explanation of the confusion lies in the fact that the Bronsted theory as ordinarily presented is confined to only one of the four experimental ~
~
criteria of acid-base behavior: namely, displacement. DISPLACEMENT
The type equation of the Bronsted theory actually represents the displacement of a weaker hase by a stronger one: HCI acid,
+ HOH base
-
HsOfl acid2
+ base, CI-I
(7)
The chloride ion is a very weak base combined with a strong acid (the proton) to form hydrogen chloride. In this reaction, a stronger base, water, has displaced the weaker base, chloride ion, from its combination with the proton. For every snch displacement reaction (or competition of bases for the proton) there exists an equilibrium constant. Therefore, the extent of the reaction depends upon concentration as well as the relative "strength" of the two bases. For example, even weak bases such as ammonia or acetate ion will displace a strong hase (the hydroxyl ion) from water to some extent:
+ HOH P NH4+' + 011-' acidz acid, base* OAc-I + HOH F-' HOAc + OH-'
NH3 base,
(8) (9)
For both displacement reactions, the equilihrium point ordinarily lies far to the left. It is this undue emphasis of the Bronsted theory upon displacement which seems to he responsible for the idea that hydrogen acids need special consideration. Displacement is only one of the four "phenomenological criteria" of acids and bases. Yet some writers on the Bronsted theory have emphasized it so strongly as to give the impression that the other three--even neutralization-are inconsequential. In order to show that special treatment is unnecessary for hydrogen acids, we need only demonstrate that the displacement reaction represented hy the typical Bronsted equation is not unique. The same kind of reaction may take place between many other acids and bases. A detailed consideration of one example should be sufficient illustration. When the addition compound formed by the neutralization of boron trichloride and acetone is added to pyridine, displacement of the weaker base (acetone) by the stronger base (pyridine) takes place (5). This reaction is exactly analogous to the displacement of the weakly basic chloride ion from combination with the proton when hydrogen chloride is added to water:
acid,
bases
acidz /CHr
:6::c 'CH, base,
(10)
acid,
bash
acids
base,
Both boron trichloride and the proton are primary acids (5) which, to form the acid, compounds shown above, have been neutralized by the weak bases acetone and chloride ion, respectively. However, in both cases stronger bases are able to displace the weaker ones from combination with the two acids. All four acids (in both equations) are secondary acids (5). Many more examples of displacement reactions of exactly the same type could be given. Some of them are responsible for basic catalysis and have been dealt with previously (4, 6, 7). One familiar example is interesting since it has been the subject of some misunderstanding: Ag(NH&+' acrd,
+ OH-'-
AgOH acid2
basen
+ 2NHs base,
(12)
The silver ion is a fairly strong acid which will combine with many bases including ammonia and hydroxyl ion. The reaction represented by the above equation is analogous to any Briinsted type equation, for example: NH4+' acid,
+ OH-'
bash
-
HOH acid?
+ NHs
base,
which may be evaluated by combining:
(13)
A number of statements appear in the literature to the effect that "in dilute aqueous solution silver ions do not combine with hydroxyl ions . . whereas they do combine with ammonia." In view of the fact that the solubility product for silver hydroxide and the instability constant for the ammonia complex are, respectively, 2 X 10-5 and 6 X it is difficult to understand how such statements ever came to be made. Silver ions and ammonia molecules combine until equilibrium is reached when their concentrations are low. Silwr ions and hydroxyl ions also combine until equilibrium i s reached when the concentrations of both ions are low. If the two ions have been added in equal amounts, equilibrium is attained when each has a concentration equal to 1.4 X lo-& M. NOWif in the other solution, the concentrations of ammonia and silver ion in equilibrium with the complex ion are adjusted to have comparable molarities, i , e., 1.4 X 10-4 M, the concentration of the complex is 0.47 X lo-&M. In such a solution, only about one-fourth of the silver ions have combined to form the complex! Thus under comparable conditions, a large proportion of the silver ions do not combine with ammonia molecules. Perhaps the cause of the misunderstanding lies in the method by which we are accustomed to observe the preparation of the silver-ammonia complex. When we add excess ammonia solution to silver nitrate solution, the initial precipitate dissolves, forming the complex ion. This apparently means that ammonia is a much stronger base toward silver ion than is hydroxyl ion. But we must remember that the silver
.
hydroxide precipitate is formed first. When the first few drops of ammonia solution are added, the ammonia is so dilute that i t is largely ionized t o form hydroxyl ions which coordinate* with the silver ions to form the hydroxide. But by the time the precipitate is completely dissolved, the amount of ammonia added is enough so that if the silver ions were not present, the concentration of NH3 would be roughly 100 times the concentration of OH-' ion. Actually if concentrated sodium hydroxide solution is now added to the solntion, the ammonia is displaced and the precipitate returns! This is what one would expect when mass action effects are taken into account. The effects observed are similar whether the acid is the proton or the silver ion. The displacement of one base by another depends not only upon their relative strength, but upon concentration' factors also. An idea of the relative basic strengths of ammonia and hydroxyl ions toward silver ion, can be obtained from equation (12). The equilibrium constant for equation (12) is:
with [Ag+l] X [OH-'] = 2 X
lo-$
We see that the constant has a value of 3, indicating that ammonia is about as strong a base as hydroxyl ion toward silver ion. I t is worth while noting a t this point that in equations (lo), (12). and (13) the acid1 and acida compounds are amphoteric in exactly the same sense. The addition compounds of boron trichloride with acetone and with pyridine are amphoteric because they contain both acidic and basic constituents. They can react both ways because stronger bases will displace the weaker bases and stronger acids will displace the boron trichloride. In the same way, silver hydroxide in equation (12) and water in equation (13) are also amphoteric. They appear as acids in these equations, but both may act as bases because of the presence of the hydroxyl group. It. i s obvious that the terms "acid" and "base" are not absolute but relatiwe terms to be used with reference to a particular reaction. One more point about the role of displacement in the two theories should be discussed. Not only does the Bronsted theory overemphasize displacement a t the
* Since silver hydroxide actually appears as silver oxide when attempts are made to isolate it, we can not be positive that a coordinate bond is formed between silver and hydroxyl ions. However, there are two reasons for believing that one is formed: (1) Other hydroxides of similar nature (small positive ion) are often not ionic even in the solid state. (2) Silver iodide does not crystallize into an ionic lattice, although the chloride and bromide do.
expense of the other experimental criteria, but it considers only one kind of displacement: namely, the displacement of one base by a stronger one. This is true because it admits of only one primary acid, the proton. The other kind of displacement seems equally important in chemistry, i . e., the displacement of one acid by a stronger one. One example should he sufficient illustration : Si01 acid,
+ NaCO. base2
-
NaBi08 base,
+ acid. C02
The silicon dioxide displaces the weaker carbon dioxide from combination with the base sodium oxide. THE RELATIONSHIP BETWEEN A/B AND O/R REACTIONS In one other respect at least, the limited nature of the Bronsted theory has led to confusion. There is a formal analogy between hydrogen acids as protondonors and reducing agents as electron-donors. This implies an actual relationship between acids and reducing agents which does not exist. When their experimental properties are considered in the light of the Lewis theory (8) it is evident that both acids and oxidi~ingagents are electrophilic: i . e., acids accept a share in a lone electron-pair while oxidizing agents gain electrons outright. Conversely bases and reducing agents are electrodotic (8). Since this relationship has been presented in THIS JOURNAL previously (8) only one point need be made in connection with the subject of this paper. As mentioned in the preceding section, amphoteric behavior is implicit in both the Bronsted and Lewis theories of acids and bases. For example, water behaves as a base toward acetic acid, but as an acid toward ammonia. Water is amphoteric in the usual sense, but it may also act either as an oxidizing agent or a reducing agent. It oxidizes metallic sodium, but reduces fluorine. Evidently, an extension of the meaning of the word "amphoteric" is necessary to include the idea of the relativity bf oxidizing and reducing agents as well as acids and bases (Table 1). TABLE 1 AMPHOTERICBEHAVIOR oa WATER ElcctroPhilic HOH NHZ OH-' NH4+l acid 2HOH 2Na 20H-I f 2Nat' HZ oxidant . Electrodotic HOH SO2 H+' 4- HS08C1. base 2HOH 2F2 4H+I 4F-' 0 2 reductant ~
+ + + -
~
+
+
+
+
+
Obviously the terms electro~hilic and are no more absolute than the terms acid and base. ~h~~ apply to given substances only as they are reacting in a particular chemical reaction. Many other substances besides water can behave in all four ways (acid, base, oxidant, or reductant) dependmg upon the other substances involved. A few examples not pre-
viously mentioned will be discussed here. Hydrogen chloride is ordinarily thought of as a strong acid, hut it behaves as a base toward sulfur trioxide or stannic chloride (2). On the other hand, hydrogen chloride is an oxidizing agent toward active metals, but a reducing agent toward strong oxidizing agents such as permanganate. Hydroquinone is similar to water and hvdrocen chloride in its am~hotericbehavior. It can &have both as an acid and a n oxidizing agent tow?nd hases and reducing agents because of the presence of the labile protons. Or it may act both as a base and a reducing agent (particularly in water) because of the hydroquinone ion. Hydroquinone ion is readily oxidized to quinone. The familiar reactions of the closely related quinone again illustrate the need for an extension of our ideas of amphoteric behavior. Quinone as an oxidizing agent is electrophilic, but as a base, i t is electrodotic. Stannous ion, as far as its more familiar reactions are concerned, is amphoteric in oxidationreduction reactions, but not in general to be thought of as amphoteric in acid-base reactions (one would expect its basic properties to be extremely weak). These examples should be sufficient to illustrate the necessity of eradicating from our minds the idea that very many substances can be classified arbitrarily as electrophiiic or electrodotic. The terms are very useful, but we must not forget that they are relative terms applying to a substance only as it behaves in the particular reaction under consideration. This relativity in terminology is unavoidable. Whether we like i t or not, we must accept a certain amount of relativity in chemistry as well as in physics! CONCLUSION
1. No conflict exists between the Bronsted and Lewis theories of acids and bases, because the electronic theory includes hydrogen acids in such a way that no special treatment of them is required. 2. The typical equation of the Bronsted theory represents only one of the four experimental criteria upon which the Lewis theory is based. 3. Amphoteric behavior is implicit in both theories, hut is more widespread than indicated even by the Lewis theory. Many substances may be electrophilic (acids or oxidants) in some reactions, but electrodotic (bases or reductants) in other reactions. The terms electrophiiic and electrodotic are relative and should he used only with a particular reaction in mind. LITERATURE CITED
(1) LEWIS,G . N., "Valence and the Structure of Atoms and Molecules," The Chemical Catalog Company, h e . , New York, 1923. W. F., Chem. Revs., 27, 517 (1940). (2) LUDEK, (3) LWDER, W. F.,W. S. MCGUIRE, AND S. ZUPPANTI, J. CHEY. Eouc., 20, 34.1 (1943). (4) LUDER. W. F.,AND S. ZWTANTI, Chem. Reus.. 34,345 (1944). (5) LEWIS.G. N., J. Franklin Inst., 226, 293 (1944). (6) LUDER, W. F., AND S. ZUPPANTI, J. Am. Chem. SOL, 66,524 (1944). (7) ZUFFANTI, S., AND W. F.LUDER, J. CHEM.EDUC., 21, 485 (1944). (8) LUDER, w. F., ibid., 19, 24 (1942).
