Acids and Bases and their Ionisations

Strong acid : 100% dissociation
HClO₄, H₂SO₄, HNO₃, HI, HBr
HCl + H_{2}O \rightarrow H^{+} + Cl^{-}

Strong base :
NaOH, KOH, RbOH, CsOH, Ba(OH)₂

Weak acid : Which undergoes partial dissociation

Hydrogen ion concentration :

K_{a} = \frac{C{\alpha} \ C{\alpha}}{C - C{\alpha}} = \frac{C\alpha^{2}}{1 - \alpha}
K_a = Cα₂ (α < < < 1, 1 − α = 1)
\alpha = \sqrt{\frac{K_{a}}{C}}
[H^{+}] = C\alpha = C. \sqrt{\frac{K_{a}}{C}} = \sqrt{K_{a}.C}
[H^{+}] = C\alpha = \sqrt{K_{a}.C}

Arrhenius acid - base theory :
Strong acid : Which produce more no. of H⁺ (HClO₄, H₂SO₄, HCl)
Weak acid : Which produce less no. of H⁺(CH₃COOH, HCN, H₂S)
Strong base : Which produce more no. of OH⁻ (NaOH, KOH)

Bronsted - Lowry acid - base theory :
Acid : Proton donor [HCl , H₂SO₄, CH₃COOH ......]
Base : Proton acceptor[NaOH, KOH, NH₃ .......]
Salt : Neither proton acceptor nor proton donor. [C₆H₆, CCl₄ ..... aprotic]

Conjugate acid - base pair :
A Bronsted - Lowry acid - base pair which differ by only one proton is called conjugate acid - base pair.
Acid − H⁺ → conjugate base
Base + H⁺ → conjugate acid.

Ionic product of water : (K_W)
H_{2}O + H_{2}O \rightleftharpoons H_{3}O^{+}+ OH^{-}
K = \frac{[H_{3}O^{+}][OH^{-}]}{[H_{2}O]^{2}}
K{[H_{2}O]^{2}} = [H_{3}O^{+}][OH^{-}]
K_{W} = [H_{3}O^{+}][OH^{-}]

Kw → acid, ΔT = 25°C
K_w = 1.0 × 10^-14 mol² lit^-2
[H⁺][OH^-] = 10^-14, [H⁺] = 1.0 × 10^-7 mol/lit

Degree of dissociation :
\alpha = \frac{1}{55.5 \times 10^{7}} = \frac{10^{-7}}{55.5} = \frac{10^{-7}}{\left(\frac{1000}{18}\right)}
α = 1.8 × 10^-9
% dissociation = 1.8 × 10^-7 , \alpha = \frac{10^{-7}}{55.5}