Last updated : September 3, 2024

Ionic Equilibrium : pH of Acid & Base, Dissociation Constant, Degree of Ionisation, Salt Hydrolysis, Buffer Solution, Solubility Product

Acid Base Theory

Arrhenius Concept

Acid : H⁺ donor in aqueous solution.

Base : OH^- donor in aqueous solution.

Bronsted-Lowery Concept

Acid : H+ donor.

Base : H+ acceptor.

Acid Conjugate-base pair and
Base Conjugate-acid pair

Lewis Concept

Acid - Lone pair acceptor.
(i) Having vacant orbital (ii) Electron deficiency

Example :

Base - Lone pair donor.
(i) min 1 lone pair present (ii) Octate complete

Example :

pH of a Solution

Ionic Product of Water (K_w)

pH of Strong Acid and Strong Base

if [H⁺]/[OH^-] concentration from Acid/Base is more than 10^-6, [H⁺]/[OH^-] concentration from water (10^-7) can be ignored!

Q. pH of 10^-3M H₂SO₄ | and pH of 10^-4M NaOH solutions?

Q. pH of 10^-12M HCl | and 10^-8M NaOH?

pH of Mixture Solution

pH of Weak Acid and Weak Base

Suppose - HA is a weak monoprotic acid

Common Ion Effect

Degree of Dissociation (α) of weak acid or weak base decreases when it is mixed with strong acid or strong base respectively.

Salt Hydrolysis

Acid + Base = Salt

On the basis of nature of Acid and Base, Salts are of 4 types :

Strong acid + Strong base : pH = 7 (No hydrolysis)
HCl + NaOH = NaCl + H₂O
Weak acid + Strong base : pH > 7 (Anionic hydrolysis)
CH₃COOH + NaOH = CH₃COONa + H₂O
Strong acid + Weak base : pH < 7 (Cationic hydrolysis)
HCl + NH₄OH = NH₄Cl + H₂O
Weak acid + Weak base : pH depends on pK_a & pK_b (both Anionic & Cationic hydrolysis)
CH₃COOH + NH₄OH = CH₃COONH₄ + H₂O

Only Conjugate ion of Weak acid or Weak base participate in hydrolysis!

Hydrolysis of CH₃COONa, is a salt of weak acid (CH₃COOH) and strong (NaOH) base :

Buffer Solution

Acidic buffer - Weak acid + Salt of its Conjugate ion
Basic buffer - Weak acid + Salt of its Conjugate ion

Composition of Buffer Solution

Solution = Weak acid + Salt of its Conjugate ion

The resulting buffer solution contains

Solution = Weak acid + Strong base

C₁ > C₂ (If equal Salt Hydrolysis take place)
The resulting buffer solution contains

pH of Acidic Buffer

Using concentration of H⁺ ion from buffer solution

pH of Basic Buffer

Buffer Range

\begin{array}{rcl} pH & = & {pK}_{a} \pm 1 \\ log \frac{[Salt]}{[Acid]} & = & - 1 to + 1 \\ \frac{[Salt]}{[Acid]} & = & \frac{1}{10} to 10 \end{array}

Buffer Capacity

Buffer Capacity = \frac{Conc. of S.A, S.B Added}{Change in pH of Buffer solution}

Solubility

Total amount of springly soluble (S <<< 1) solute in 1 litre of solution at saturated (equilibrium) condition is known as Solubility.

Unit of Solubility (S) = $\frac{gram}{Litre}$

Solubility Product (K_sp)

Solubility of salt AB = S

Common ion Effect in Solubility

In presence of common ion, solubility of salt decreases due to Le-chatelier’s principle.

Solubility (S) = \frac{K_{sp}}{Conc. of common ion of strong electrolyte}

Ionic Product (K_ip)

K_ip > K_sp : Backward reaction - form ppt
K_ip < K_sp : Forward reaction - soluble
K_ip = K_sp : Backward reaction - at equilibrium

Selective Precipitation

Case - 1
Number of Salts = n
Number of ions in each salt = same
K_sp = different
Case - 2
Number of Salts = n
Number of ions in each salt = different
K_sp = same
Case - 3
Number of Salts = n
Number of ions in each salt = different
K_sp = different
Use Solubility formula to find order of solubility and ppt formation
Solubility min = first ppt
Solubility max = last ppt

# pH of Acid Base # Dissociation Constant # Degree of Ionisation # Salt Hydrolysis # Buffer Solution # Solubility Product # Physical Chemistry