The apparent strength of a protonic acid is dependent on the solvent in which the acid is dissolved. When all the acids in the acid chart which are stronger than H3O+ ion
(i.e., the acids above H3O+ acids) are added to H2O, they donate as proton to H2O to H3O+ ion and appear to have equal strength, since all these acids are levelled to the strength of H3O+ ion which is left in solution and is common to all such solutions.
This phenomenon viz. The strength of all the acids becomes equal to that of H3O+ ion is called leveling effect of the solvent, water, and here water is called a leveling solvent for all these acids.
In aqueous solution all very strong bases like Na+H–, Na+NH2–, Na+OC2H5– are levelled to the strength of OH– ion, for they react completely with H2O to produce OH– ions.
The solvent in which complete proton-transfer occurs are called levelling solvents.
In other words, the solvent in which the solute is ~100% ionised, are called levelling solvents.
Since HF and HCl both are ~ 100% ionised in liquid NH3 to give ~100% NH4+ ions, these appear to be of equal strength and liq. NH3 acts as a levelling solvent for HF and HCl. In H2O, HF is only partially ionised, whereas HCl and HBr are ~ 100% ionised.
Thus H2O is a differentiating solvent for HF, but for HCl and HBr it is a leveling solvent. Several mineral acids are partially ionised in glacial CH3COOH medium because CH3COOH is a poor proton-acceptor but rather a better proton donor.
CH3COOH, therefore, acts as a differentiating solvent towards the mineral acids. But, for bases, CH3COOH act as a leveling solvent.
Utility of Bronsted Concept
(i) It defines acids and bases in terms of the substances themselves and not in terms of their ions in aqueous solution. Thus HCl is an acid because of the fact that it can give a H+ ion.
HCl <—> H+ + Cl–
(ii) The Bronsted concept recognises that acid-base behaviour is neither restricted to, nor depends on, any particular, solvent
iii) This concept is useful in accounting for the hydrolysis of salt solution. When a salt is dissolved in water, an unbalance in the concentration of the solvent cation (H3O+) and anion (OH–) will result, if the salt cation and anion differ in their proton-donor and proton-acceptor properties towards H2O.
This point can be illustrated by considering the aqueous solution of FeCl3. Aqueous hydrated ferric ion, [Fe(H2O)x]3+, exceeds the proton-acceptor ability of Cl– ion and a considerable excess of H3O+ ion is produced in the solution, making FeCl3 acidic.
FeCl3 + xH2O <—> [Fe(H2O)x]3+ + 3Cl–
Aqueous solution of a Na2CO3 is alkaline in character, because the proton acceptor ability of ion exceeds the proton-donor ability of hydrated sodium ion, [Na(H2O)x]+.
2xH2O + Na2CO3 <—> 2[Na(H2O)x]+ + CO32-
i) This concept lays excessive emphasis on the proton – transfer. Although it is true that most common acids are protonic in nature, yet there are many which are not. It fails to explain the acidic character of many substances which do not contain any proton (e.g. SO3, AlCl3).
ii) There are a number of acid-base reactions in which no proton transfer takes place, e.g.
SO2 + SO2 <—> SO2+ + SO32-
Acid1 + Base2 <—-> Acid2 + Base1
Thus the protonic definition cannot be used to explain the reactions occuring in non-protonic solvents such as COCl2,SO2, N2O4 and BrF3.