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Hard and Soft Acids and Bases

The thermodynamic stability of a metal complex depends greatly on the properties of the ligand and the metal ion and on the type of bonding. Metal–ligand interaction is an example of a Lewis acid–base interaction. Lewis bases can be divided into two categories:

  • hard bases contain small, relatively nonpolarizable donor atoms (such as N, O, and F), and
  • soft bases contain larger, relatively polarizable donor atoms (such as P, S, and Cl).

Metal ions with the highest affinities for hard bases are hard acids, whereas metal ions with the highest affinity for soft bases are soft acids. Some examples of hard and soft acids and bases are given in Table \(\PageIndex{1}\). Notice that hard acids are usually cations of electropositive metals; consequently, they are relatively nonpolarizable and have higher charge-to-radius ratios. Conversely, soft acids tend to be cations of less electropositive metals; consequently, they have lower charge-to-radius ratios and are more polarizable. Chemists can predict the relative stabilities of complexes formed by the d-block metals with a remarkable degree of accuracy by using a simple rule: hard acids prefer to bind to hard bases, and soft acids prefer to bind to soft bases.

Table \(\PageIndex{1}\): Examples of Hard and Soft Acids and Bases
  Acids Bases
hard H+ NH3, RNH2, N2H4
Li+, Na+, K+ H2O, ROH, R2O
Be2+, Mg2+, Ca2+, VO2+ OH, F, Cl, CH3CO2
Al3+, Sc3+, Cr3+ CO32−
Ti4+ PO43−
soft BF3, Al2Cl6, CO2, SO3  
Cu+, Ag+, Au+, Tl+, Hg22+ H
Pd2+, Pt2+, Hg2+ CN, SCN, I, RS
GaCl3, GaBr3, GaI3 CO, R2S

Hard acids prefer to bind to hard bases, and soft acids prefer to bind to soft bases.

Because the interaction between hard acids and hard bases is primarily electrostatic in nature, the stability of complexes involving hard acids and hard bases increases as the positive charge on the metal ion increases and as its radius decreases. For example, the complex of Al3+ (r = 53.5 pm) with four fluoride ligands (AlF4) is about 108 times more stable than InF4, the corresponding fluoride complex of In3+ (r = 80 pm). In general, the stability of complexes of divalent first-row transition metals with a given ligand varies inversely with the radius of the metal ion, as shown in Table \(\PageIndex{2}\). The inversion in the order at copper is due to the anomalous structure of copper(II) complexes, which will be discussed shortly.

Table \(\PageIndex{2}\)
complex stability Mn2+ <   Fe2+ <   Co2+ <   Ni2+ <   Cu2+ >   Zn2+
ionic radius (pm) 83 78 74.5 69 73
74

Because a hard metal interacts with a base in much the same way as a proton, by binding to a lone pair of electrons on the base, the stability of complexes of hard acids with hard bases increases as the ligand becomes more basic. For example, because ammonia is a stronger base than water, metal ions bind preferentially to ammonia. Consequently, adding ammonia to aqueous solutions of many of the first-row transition-metal cations results in the formation of the corresponding ammonia complexes.

In contrast, the interaction between soft metals (such as the second- and third-row transition metals and Cu+) and soft bases is largely covalent in nature. Most soft-metal ions have a filled or nearly filled d subshell, which suggests that metal-to-ligand π bonding is important. Complexes of soft metals with soft bases are therefore much more stable than would be predicted based on electrostatic arguments.

The hard acid–hard base/soft acid–soft base concept also allows us to understand why metals are found in nature in different kinds of ores. Recall that most of the first-row transition metals are isolated from oxide ores but that copper and zinc tend to occur naturally in sulfide ores. This is consistent with the increase in the soft character of the metals across the first row of the transition metals from left to right. Recall also that most of the second- and third-row transition metals occur in nature as sulfide ores, consistent with their greater soft character.