CHEM-GUIDE: Bronsted lowrry concept of acids and bases

In 1923, J.N.Bronsted and T.M.Lowry proposed a more general theory known as the Bronsted-Lowry proton transfer theory. According to this concept, any hydrogen containing species (a molecule, a cation or an anion), which is capable of donating one or more protons to any other substance, is called an acid.

Any species (molecule, cation or anion), which is capable of accepting one or more protons from an acid, is called a base.

Thus, according to the Bronsted-Lowry concept, an acid is a proton-donor, and a base is a proton-acceptor.

The reaction of an acid with a base involves transfer of a proton from the acid to the base. So, an acid and a base should be present simultaneously in any system. The extent of an acid-base reaction is governed not only by the proton-donating ability of the acid, but also by the proton-accepting tendency of the base. Acids and bases classified on the basis of this concept are termed as Bronsted acids and bases.

In this reaction, HCl donates its one proton to become Cl^-, and H₂O accepts one proton to become H₃O⁺. Thus, HCl is Bronsted acid and H₂O is a Bronsted base. For the reverse reaction, H₃O⁺is able to transfer its proton to Cl^-. So, H₃O⁺is a Bronsted acid and Cl^- is a Bronsted base.

Every acid must form a base on donating its proton, and every base must form an acid on accepting a proton. The base that is produced when an acid donates its proton is called the conjugate base of the acid. The acid that is produced when a base accepts a proton is called the conjugate acid of the base. The above reaction can be written as

In this Cl^- is the conjugate base of the acid HCl and H₂O is the conjugate base of the acid H₃O⁺. The conjugate acid differs from conjugate base by one proton. A pair of an acid and a base which differ from one another by a proton constitute a conjugate acid base pair. Thus,

Although the Bronsted-Lowry concept of acids and bases is better than the Arrhenius concept, it cannot account for the acidic and basic character of compound not containing hydrogen. For example, acidic nature of oxides such as CO₂, SO₂ etc., and the basic nature of the compounds of the type CaO, Na₂O etc.

Relative Strengths of Conjugate Acid-Base Pairs

A stronger Bronsted acid will have a higher tendency for donating a proton to the base, hence would tend to exist as its conjugate base. The conjugate base so formed will have very little tendency to pick up a proton, hence would act as a weak base. Thus, there exists an interdependent relationship between the strengths of an acid and its conjugate base.

'The stronger an acid, weaker is its conjugate base, and stronger a base, weaker is its conjugate acid.'

In water, HCl acts as a strong acid. As the reverse reaction occurs to a very small extent Cl^- ion acts as a weak base. So, the anion of a strong acid is a weak base. Similarly, the cation of a weak base acts as a strong acid, e.g., NH₄⁺ ion in water is a strong acid.

Relative strengths of common conjugate acid-base pairs

Amphoteric substances

Substances which can act as an acid as well as a base are called amphoteric substances. For example, Al(OH)₃, Zn(OH)₂ behave both as acids and as bases in their reactions. The substance, which can donate or accept a proton to act as an acid or a base is termed amphiprotic. For example, the species such as HSO₄ and HCO₃ are amphiprotic, because of the following reactions.

Water is the most common solvent showing a unique behaviour. It can act as an acid as well as a base. In reactions like,

it acts as an acid, while it behaves like a base in a reaction of the type,

The dual role of water molecule may thus be represented by the equation,

Water is thus known as an amphiprotic solvent. Solvents which accept protons are called protophillic, while those which give up protons are called protogenic. Solvents which neither donate nor accept protons are called aprotic solvents. Thus, liquid ammonia is a protophillic solvent, while acetic acid is protogenic.

Relative Strength of Bronsted Acids and Bases

According to Bronsted-Lowry theory, an acid is a proton donor and base is a proton acceptor. Therefore the strength of an acid or a base is determined by its tendency to lose or gain protons. A strong acid is a substance which loses a proton easily to a base.

Relative strength of acids

The relative tendency of acids to transfer a proton to a common base, generally water is expressed as the relative strength of bronsted acids. For example, HCl has a higher tendency to transfer a proton to H₂O than CH₃COOH. This means HCl is a stronger acid than acetic acid. Thus,

Similarly, a base having a higher tendency to accept a proton is stronger. For example, ammonia accepts a proton more readily from water molecule than a water molecule from another water molecule, and hence ammonia is more basic than water.

The ability of an acid to lose a proton (acid strength) is described by its acid ionization constant. The larger the value of the acid ionization constant Ka, higher is the concentration of H₃O⁺ in the solution, stronger is the acid. Knowing the ionization constants of acids one can get the relative strengths of different acids at a particular temperature.

Since the ionization constant of HF is larger than the ionization constant of CH₃COOH, hydrofluoric acid (HF) is a stronger acid than acetic acid, (CH₃COOH).

Relative strength of bases

The ability of a base to accept a proton (basic strength) is described by its base ionization constant (Kb). The larger the value of Kb, higher is the concentration of OH^-, stronger is the base.

For example,