Ralph G. Pearson Northwestern University Evanston, Illinois 60201
Hard and Soft Acids and Bases, HSAB, Part I Fundamental principles
According to G. N. Lewis a base is an atom, molecule, or ion which has at least one pair of valence electrons which are not already being shared in a covalent bond. An acid is similarly a unit in which at least one atom has a vacant orbital in which a pair of electrons can be accommodated. The typical acidbase reaction is A + :B
—
A:B
(1)
The species A:B may be called a coordination compound, an adduct, or an acid-base complex. In fact, a wide variety of acid-base complexes exist under different names. The species A is usually called a Lewis acid, or a generalized acid, to avoid confusion with Brpnsted acids. A Lewis base, B, is identical with a Br0nsted base. Sidgwick suggested the terms electron acceptor, in place of Lewis acid, and electron donor, in place of base. These terms are widely used, particularly in Europe. Also certain types of weak generalized acidbase interactions are almost always discussed under the heading of donor-acceptor complexes. The disadvantage is that a different term, electron donor, is used to describe substances which are generally and conveniently called bases. It is also true that special The names are sometimes used for special categories. use of the term ligand in place of base when A in eqn. (1) is a metal ion is firmly established. Also in speaking of rates of reaction it is usual to call A an electrophile and to call B a nucleophile.* Probably the most important class of chemical reaction is the generalized acid-base reaction of eqn. (1). The easiest way to appreciate this is to consider the different kinds of acid-base complexes, A:B, that may be formed. For example, all metal atoms or ions are Lewis acids. They are usually found coordinated to several bases or ligands simultaneously since they are polyvalent. The combination may be electrically charged, in which case we have a complex ion formed. Also the combination may be electrically neutral in which case a normal inorganic molecule such as SnCL is formed. Most cations are Lewis acids and most anions are bases. Hence salts are automatically acid-base complexes. MgClj in the solid state consists of the acid Mg2+ coordinated to six neighboring Cl- ligands. In dilute solution the magnesium ion is coordinated to the 1
Editor's Note: This is a two part article. The second part, which discusses the theory? underlying the HSAB principle, will appear in the October issue. Hayek, E., Osiers. Chem.-Zent., 63, 109 (1962), has a general discussion of the problem of acid-base terminology. 1
base water forming the solvated ion. The chloride ion also forms an acid-base complex, via hydrogen bonding in which water molecules are Lewis acids.
Inorganic compounds, as solids, liquids, gases, or in solution, are examples of acid-base complexes. The same thing can be said for organic compounds. The method here is to mentally dissect, the organic molecule into two fragments, one of which is a Lewis acid and the other a base. For example, ethyl alcohol can be thought of as composed of the ethyl carbonium ion, All carbonium C2Hs+, and the hydroxide ion, OH-. ions are Lewis acids and the hydroxide ion is the base. Similarly, ethyl acetate, can be thought of as a complex of the acylium ion, CH3CO+, which is an acid, and the base ethoxide ion, OTLO-- Even a hydrocarbon can be broken down (conceptually) into an acid such as H+ and a carbanion, which is a base. Thus methane can be considered to be H+ and CH3- combined. It is equally true that CHi can be viewed as a combination of CH3 + and H-. Such ambiguity is, in fact, universal among both organic and inorganic molecules. While at first confusing, it turns out to be absolutely necessary to explain the variety of reactions undergone by these molecules. That is, methane sometimes reacts as if it were splitting into CH3+ and H-, and sometimes as if it were splitting into CH3- and H+. In addition, there are reactions in which it behaves as These arc redox, or free radical reacCH3' and H\ and do not concern us here. tions, however, In the case of ethyl acetate, the molecule can be viewed as an acid-base complex as explained above. It is also true that it is a Lewis acid and a Lewis base. Ethyl acetate acts as a Lewis base when it forms complexes through one of its oxygen atoms to the proton, It acts as a Lewis acid when it or other Lewis acids. adds bases, such as the hydroxide ion, to its carbonyl group. Such acid-base processes are important in the acid and base catalyzed hydrolyses of esters. It should be noted that a certain group of atoms is often designated as an acid or base even if it has no stable existence. A carbonium ion such as CH3+ is considered to be a Lewis acid because its structure shows that it can accept, a pair of electrons from a base. The breaking down of a molecule, such as CH3C1, into a methyl carbonium ion and a chloride ion is a purely conceptual process and lias nothing to do with the stability of CH3L The point is that most reactions of methyl chloride can be classified as being exchanges of the chloride ion by other bases, or of the methyl cation by other Lewis acids. Just as the proton does not exist free under ordinary circumstances, so it is likely that CILi+ does not ordinarily exist as a free species. Volume 45, Number 9, September J968
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581
This brings up the point that eqn. (1) is oversimplified. What actually occurs in most cases is the exchange reaction ArB' + A':B^A;B + A':B' (2)
In solution A' and B' are often solvent molecules. Thus, as already mentioned in the case of ions dissolved in water, most solute-solvent interactions can be classified as generalized acid-base reactions. A polar molecule will always have an electron rich, or basic site, and an electron poor, or acid site. Other kinds of acid-base complexes are the so-called charge transfer complexes which are responsible for the colors produced when many substances
are
mixed.
An example is iodine and the intense brown color it gives in solvents which are bases, such as water, alcohols, or ethers. Many charge transfer complexes are formed between Lewis acids which are unsaturated molecules with electron-withdrawing substituents, such as tetraeyanoethylene, and unsaturated molecules with electron donating substituents, such as toluene. These systems are called ir-acids and 7r-bases, respectively. Finally, free atoms and radicals containing electronegative elements act as Lewis acids and form complexes with a variety of bases. These complexes of free radicals cannot be isolated but they have a very great effect on the reactivity of the radicals. It is apparent that it is possible, in principle at least, to view the greater part of chemistry as examples of interaction of generalized acids and bases. This in turn means that any rules that can be developed concerning the stability of complexes A:B in reaction (1) will have very wide application and can be useful in many areas. Recently a rule has been suggested, “The Principle of Hard and Soft Acids and Bases,” or HSAB principle, which does seem to have value in understanding a wide variety of chemical phenomena (7).
Acid and Base Strength The concept of acid or base strength comes in at this point. In a qualitative way, what is meant by generalized acid or base strength is simply that a strong acid, A, and a strong base, B, will form a stable complex, A:B. A weaker acid and base will form a less stable complex. Operationally we may define generalized acid and base strength by competition experiments. Consider the acid-base substitution reactions A' -f- A:B
-*
A':B + A
(3)
B' + A:B
—
A:B' + B
(4)
If the reactions go as indicated, it means that A' is a stronger Lewis acid than A, and that B' is a stronger base than B. If it were possible to put all Lewis acids into an order of decreasing strength, and the same for all bases, then it would be possible to predict the stabilities of all possible acid-base complexes. That is, we could predict what chemical reactions would occur under various conditions, what compounds would be stable, etc. We might then expeet the equilibrium constant of eqn. (1) to be given by an equation such as log K
where 582
/
SA
is
a
=
SASB
strength factor for the acid and
Journal of Chemical Education
(5) SB is a
strength factor for the base. Equation (5) is not the only equation that might result, but it would be the simplest one that would correctly predict the direction of displacement reactions such as (3) and (4). Of course, SA and SB would be functions of the environment and the temperature. The “intrinsic” strengths would presumably refer to gas phase reactions. There would then be strong solvent corrections, but even so an equation such as eqn. (5) would be most useful since a series of SA and $B values would need to be determined once and for all in water at 25°C, and the stabilities of possible complexes would be known in that medium. SA and S* values for 100 acids and 100 bases would predict the stabilities of 10,000 complexes, for example.2 Unfortunately it is not possible to write down any universal order of acid or base strength (#). If a series of different bases, B', are used in testing against a fixed reference A:B, then the order of base strength that one gets is very much a function of the nature of the reference acid, A. When A is the chromic ion, one gets a different order from the case when A is the cuprous ion. In fact, the order is different for cupric ion and for cuprous ion so that even a change in oxidation state of the reference acid can have an effect on relative base strengths. We are forced to the conclusion that there is no straight-forward way of evaluating base strengths or acid strengths even if we were to agree that only gas phase data for reaction (1) be used. Of course, a useful scale of base strengths towards the special Lewis acid, the proton, in aqueous solution does exist. It must always be remembered that this scale is not valid for all Lewis acids. It can be useful in a general sense, however, in telling us that some bases such as CKR-, are very weak, and other bases, such as H are very strong. Such a rough ordering will usually be valid if two bases of widely different strengths are being compared. A comparable scale to order Lewis acids in aqueous solution has not been devised, partly because many Lewis acids are decomposed by water. For many others, a scale in which the hydroxide ion is a reference base could be used. In aqueous solution we could write both the reactions ,
H+{aq) + B(aq) ^ BH+(aq) OH-(aq) + A(aq) ^ AOH-(aq)
(6) (7)
Any arbitrary value of SA could be assigned to the proton. Let us pick 9.0, for example. This would lead to values of SB ranging from about 5 for a strong base 1 for a weak base such as I-. The such as CH3_ to value of SB for the hydroxide ion would then be (15.74/ —
1.75. The number 15.74 is the logio of the equilibrium constant for eqn. (6) when B is the hydroxide ion (55.5/1.0 X 10-u). This value for Sb of OH- then leads to values of SA of 5.9 and 8.6 for the aqueous Hg2+ and CH3+, respectively. These figures come simply from the acid ionization constants of Hg(H20)22+ and CH3OH2+ applied to reaction (7). Again, such numbers are useful
9)
=
2 One might ask at this point why standard free energies of formation are not used as the fundamental properties. The answer is that these involve prior knowledge of the stability of A:B. We are seeking a method that predicts the properties of A: B from a knowledge of A and B only, in a given medium.
in telling us that H+ and CHs+ are probably stronger acids than Hg2+. However, they cannot be combined with the values of *SB previously obtained to accurately predict the stabilities of CH3HgH20+, HgI(H20) + CH3I, or C2H3 in aqueous solution. They may however, give some rough idea of stabilities. Furthermore, there will be some Lewis acids sufficiently like the proton so that a Br0nsted relationship exists between the equilibrium constants, Ka, for the reaction A(aq) + B(aq) vi A: B(aq)
(8)
and the equilibrium constant for eqn. (6), Ka, Ka
GKaa
=
(9)
In such cases eqn. (5) will be valid since eqn, (9) is simply another form of the equation with a equal to Usually the simple eqn. (5) will not be adequate, and it would be logical to replace it with a more complex equation involving more parameters. That is, instead of having only one parameter, jS, characteristic of each acid and each base, it will be necessary to have at least two. Such an equation (not the only one possible) would be log K
=
SaSb
+
(10)
a\ As > Sb
Te Cl > Br > I S > Se >
P > As > Sb
Te Cl < Br < I S < Se
~
consider any Lewis acid, such as Cu+, NO+, or simply examine the literature to see if a complex such as Cul or CuF is more stable, or if Cu(PRB)2+ is more stable than Cu(NH3)2+. The term stable is ambiguous, as ordinarily used, and a strict definition would refer to the equilibrium constants for reactions in water such as wc
I2, wc
CuF(aq)
-+•
I (aq)
Cu(PRs)+(aq) + NH,(aq) ^
Cul(aq) + F
(17)
Cu(NH,)+(aq) + PRJflq)
(18)
Oftentimes the data is incomplete and a variety of interpolations must be made to draw a conclusion (1). Nevertheless, it is usually possible to conclude that a Lewis acid prefers either the hard bases of Table 3, or the soft bases. When this is done for a large number of Lewis acids, the results are as shown in Table 4. The entries in the
The Principle of Hard and Soft Acids and Bases or
» » > (N « < O « I.F
