Ionic Equilibrium

Concepts of Acids and Bases

Arrhenius Concept:
According to Arrhenius, an acid is a substance that contains hydrogen and releases hydrogen ions (H+) in aqueous solution. A base is a substance that produces OH(aq).

$\large HCl_{(aq)} \rightarrow H^+_{(aq)} + Cl^-$

$\large NaOH_{(aq)} \rightarrow H^+_{(aq)} + OH^-_{(aq)}$

In aqueous solution NH3 forms NH4OH which furnishes OH(aq),

The strength of an acid is defined in terms of concentration of H+(aq) that is present in the aqueous solution of a given concentration of the acid.

Similarly the strength of a base depends upon the relative concentration of OH−(aq) in aqueous solution of the base.  In neutralisation reaction, H + and OH ions combine to form water.

$\large H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2 O$

Bronsted Lowry’s Protonic Concept :
Bronsted acid is a proton donor and Bronsted base is proton acceptor. The reaction of an acid with a base involves the transfer of proton from the acid to the base. Acids and bases exist in a solution in a state of dynamic equilibrium. The conjugate base of a Bronsted acid (HA) is the base that is formed when the acid has donoted a proton. The conjugate acid of a Bronsted base is the acid that is formed when the base has accepted a proton.

For a conjugate acid base pair HA – A in aqueous solution.

HA ↔ H+ + A;                        Ka

A + H2O ↔ HA + OH         Kb

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on adding ,       H2O ↔H+ + OH

$\large K_a \times K_b = [H]^+[OH]^- = K_w = 10^{-14} (at \; 25^o C)$

That is strong acid has weak conjugate base and weak acid has strong conjugate base and  also lower the pKa value, higher the acid strength and lower the base strength of conjugate base.

Lewis Acids and Bases:
Lewis acid is a species that is capable of accepting a pair of electrons to form covalent bond and Lewis base is a substance that is capable of donating an electron pair to form a covalent bond.

Some examples of Lewis acid: BF3, Cu2+, Fe3+, SO2, SnCl2

Some examples of Lewis bases: F , NH3, CN , OH etc.