Commonly, it is accepted that a compound is an Arrhenius acid if it liberates hydrogen ions as H+ in water. In the next step, these H+ ions combine with water molecules to form hydronium ions (H3O+). The two steps can be summarised to say that Arrhenius acids are compounds that form hydronium ions, provided water is the solvent.
A Bronsted acid, HA, is a proton donor (with ‘proton’, in this context, a hydrogen ion, H+) in any aqueous solvent. In the example below, HNO3 splits in aqueous ammonia to liberate H+ and NO3- ions.
The loss of the proton results in the formation of the conjugate base of the acid, A- (NO3-). Conversely, the acquisition of a proton by a base, B (here NH3), results in the formation of its conjugate acid, HB+ (NH4+).
So, Bronsted and Lowry expanded on the Arrhenius definitions, defining acids as proton donors and bases as proton acceptors. They also introduced the concept of conjugate acid-base pairs.
An acid is a strong acid if it effectively and fully deprotonates in a solution, whereas a weak acid is only partially deprotonated at normal concentrations.
This acid strength is measured by the dissociation constant Ka; strong acids have higher Ka values, reflecting higher H+ concentrations. For a weak acid, the acidity constant, Ka, has a value of Ka<<1.
G.N. Lewis further expanded the scope and defined acid from the perspective of the electrons. A Lewis acid is an electron-pair acceptor (therefore, a base is an electron-pair donor; when they both combine, an adduct, or an acid-base complex, is formed). This definition also brought within its scope the “standard” Bronsted-Lowry acid-base reactions. A Bronsted-Lowry acid must donate a proton, which means it has an electron-deficient site (the proton) capable of accepting an electron pair, making it a type of Lewis acid. The formation of conjugate acid (HB+) proves just that!
In addition to Lowry-Bronsted acids, Lewis acids include transition-metal ions, central atoms with incomplete octets, and central atoms with available extra d-orbitals to accommodate an incoming electron pair.
Hard acids and soft acids are some of the terms associated with Lewis acids. It is important to realize that hard/soft considerations have nothing to do with acid or base strength. An acid or a base may be hard or soft and also be either weak or strong.
A hard and soft acid (or base) is indicative of its polarizability, that is, its innate ability to undergo electron cloud distortion under a set of applied conditions and therefore form temporary dipoles. Hard acids (and bases) are small, compact, and nonpolarizable. Soft acids (and bases) are larger and have a more diffuse distribution of electrons.
This polarizability of a Lewis acid (or base) plays a role in its reactivity and in which species it associates in reactions. A hard Lewis acid tends to bond strongly to a hard base; a soft acid is a Lewis acid that tends to bond strongly to a soft base. Hard acid – hard base combinations are predominantly ionic, whereas soft acid – soft base combinations are largely covalent.
Learning about Acids and Bases is crucial to understanding reactivity in Organic Chemistry. Beyond organic chemistry, they play a vital role in biological functions such as digestion and blood pH balance, as well as in numerous everyday products, from cleaning supplies and antacids to foods and fertilizers.
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Quiz on Acids in Organic Chemistry - Arrhenius, Bronsted-Lowry, Lewis (set 1)
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