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9.6: Hard and Soft Acids and Bases

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  • Origin of the Hard-Soft Acid-Base Principle

    One of the strengths of the Lewis acid-base concept is the readiness with which it illuminates the role that covalent and electrostatic interactions in acid base behavior, specifically through its ability to explain chemical interactions in terms of frontier orbitals and the interactions between charged groups as electrons are donated from a base to an acid. However, simply acknowledging the presence of such interactions does little to illuminate the degree to which each mode of explanation best explains the bonding in a given system? To what extent is a given adduct better described as held together by covalent bonds as opposed to ionic ones - e.g. better described as a molecule rather than an ion pair? Moreover, does it even matter, given that the orbitals of quantum mechanics result from the combination of electrons' wavelike behavior with their electrostatic attraction to nuclei in either case?1 These questions and more are addressed by one of the most important conceptual tools in contemporary inorganic chemistry, the Hard-Soft acid-base principle.

    The Hard-Soft acid-base principle (HSAB principle) stems from the recognition that some Lewis acids and bases seem to have a natural affinity for one another.* Consider the following:

    • Some metals are commonly found in nature as salts of chloride or as oxide ores while others are found in combination with sulfur. Geochemists even use the Goldschmidtt classification scheme to classify the halide and oxide formers as lithophiles and the sulfide formers as chalcophiles.
    • In living systems small highly charged metals ions like Fe3+ are usually found bound to N and O atoms while larger metals with lower charges such as Zn2+ are often found attached to at least one S atom. Similarly, metals prefer to bind to one coordination site over the other when forming complexes with ambidentate ligands. The most well-known instances involve complexes of cyanate and thiocyanate, which can coordinate metals through either the N or chalcogen atom. For instance, Cu2+ and Zn2+ form N-thiocyanato complexes in species like [Cu(NCS)2(py)2] and [Zn(NCS)4]2- while their larger cogeners Au3+ and Hg2+ preferentially forms S-thiocyanato complexes, giving speciles like [Hg(SCN)4]2-.

    Ambidentate ligands

    Ambidentate ligands possess multiple coordination sites through which a metal may bind. For instance, thiocyanate may coordinate metals (M) at either the S or N to give S-thiocyanato or N-thiocyanato complexes.


    • The solubility trends for the alklai metal halides and silver halides are opposite, even though both involve salts of formula M+X- (salts can be thought of as involving Lewis acid-base adduct formation between the anions and cations). Specifically, although the silver halides are all relatively insoluble in water, the very modest solubility they possess follows the order:

    X = F >> Cl > Br > I (for the solubility of AgX)

    In contrast, the much more ample solubility of the alkali metal halides** follows the opposite order. For example, the order for the lithium halides is

    X = F << Cl < Br << I (for the solubility of LiX)


    1. Albeit suitably fudged through such niceties as the Aufbau principle and other approximations.

    * Despite the fruitfulness of this observation, in general it is important to reduce the potential for observer bias by checking observations like these against compounds reported in the chemical literature and databases like the Inorganic Crystal Structure and Cambridge Crystallographic Databases.

    ** These are very soluble in water, to the point where some solutions are perhaps better described as solutions of water in the halide.

    † This can be predicted based on the relative hardness of BF3, BR3, and BH3 in the list of hard and soft acids. However, for those of you who may be confused as to why H is considered a better electron donor for the purposes of softening a Lewis acid center while alkyl groups are better electron donors for the purposes of stabilizing carbocations in organic chemistry, the dominant effect is the lower electronegativity of H relative to carbon (in CH3). The effect of electron donation due to hyperconjugation isn't as great for thermodynamically stable bases like BX3/BR3.

    Contributors and Attributions

    Stephen M. Contakes (Westmont College)