NEET Chemistry - Chapter 5

Equilibrium

Original NEET equilibrium notes on dynamic equilibrium, equilibrium constants, Le Chatelier principle, acids and bases, pH, ionic product, buffers, hydrolysis, and solubility product.

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Concept Block

1. Chemical Equilibrium: Kc, Kp, and the Reaction Quotient Q

A reversible reaction reaches dynamic equilibrium when the forward and reverse rates become equal. Macroscopic properties (concentration, pressure, colour) remain constant but molecular exchange continues.

K_c=\frac{[C]^c[D]^d}{[A]^a[B]^b}\quad\text{for }aA+bB\rightleftharpoons cC+dD

K_p=K_c(RT)^{\Delta n_g}

Pure solids and pure liquids are excluded from $K$ expressions (their activity = 1).

The reaction quotient $Q$ uses the same expression as $K$ but with non-equilibrium concentrations:

$Q < K$ : reaction proceeds in the forward direction
$Q > K$ : reaction proceeds in the reverse direction
$Q = K$ : system is at equilibrium

NEET key facts:

K

depends only on temperature. A large

K

(>10

^3

) means the reaction goes nearly to completion. A small

K

(<10

^{-3}

) means very little product at equilibrium.

Concept Block

2. Le Chatelier's Principle and Equilibrium Shifts

When a system at equilibrium is disturbed, it responds to oppose the change and re-establish equilibrium at a new position.

Disturbance	Effect on Equilibrium Position	Effect on K
Add reactant	Shifts forward	No change
Add product	Shifts reverse	No change
Increase pressure ( $\Delta n_g < 0$ )	Shifts toward fewer gas moles (forward)	No change
Increase temperature (exothermic rxn)	Shifts reverse (absorbs heat)	$K$ decreases
Catalyst added	No shift	No change

Haber process: $N_2+3H_2\rightleftharpoons 2NH_3\;(\Delta H=-92\text{ kJ})$ . High pressure and low temperature favour NH $_3$ formation, but low temperature makes rate too slow — so a compromise of 400–500°C with a catalyst is used.

NEET trap: Adding an inert gas at constant volume does NOT shift equilibrium (concentrations unchanged). Adding inert gas at constant pressure decreases partial pressures of all species and shifts toward more gas moles.

Concept Block

3. Ionic Equilibrium: Acids, Bases, Ka, Kb, pH, and Ostwald's Law

Brønsted-Lowry definition: an acid donates H $^+$ , a base accepts H $^+$ . Every acid has a conjugate base and vice versa.

K_a=\frac{[H^+][A^-]}{[HA]},\quad K_b=\frac{[BH^+][OH^-]}{[B]}

pK_a=-\log K_a,\qquad K_a\times K_b=K_w=10^{-14}\text{ at 25°C}

pH=-\log[H^+],\qquad pH+pOH=14

Ostwald's Dilution Law (for weak acid at degree of dissociation $\alpha$ , concentration $C$ ):

K_a=\frac{C\alpha^2}{1-\alpha}\approx C\alpha^2\;(\alpha\ll1),\quad\alpha=\sqrt{K_a/C},\quad[H^+]=\sqrt{K_aC}

pH of solutions:

Strong acid $C$ M: $pH = -\log C$
Weak acid $C$ M: $pH = \frac{1}{2}(pK_a - \log C)$
Strong base $C$ M: $pOH = -\log C$ , then $pH = 14 - pOH$

NEET trap: Stronger acid → larger

K_a

→ smaller

pK_a

(more negative log). HCl (

K_a \to \infty

, completely ionised) vs CH

_3

COOH (

K_a=1.8\times10^{-5}

Concept Block

4. Buffer Solutions, Salt Hydrolysis, and Henderson-Hasselbalch Equation

A buffer resists large pH changes on adding small amounts of acid or base. An acidic buffer = weak acid + its conjugate base (as a salt). A basic buffer = weak base + its conjugate acid salt.

pH=pK_a+\log\frac{[\text{conjugate base}]}{[\text{weak acid}]}\qquad(\text{Henderson-Hasselbalch})

pOH=pK_b+\log\frac{[\text{conjugate acid}]}{[\text{weak base}]}

Salt hydrolysis: When salts dissolve, the ions may react with water, shifting the pH away from 7.

Salt type	Hydrolysis?	Solution pH
Strong acid + Strong base (NaCl)	No	7 (neutral)
Weak acid + Strong base (CH $_3$ COONa)	Anionic hydrolysis	> 7 (basic)
Strong acid + Weak base (NH $_4$ Cl)	Cationic hydrolysis	< 7 (acidic)
Weak acid + Weak base (CH $_3$ COONH $_4$ )	Both ions hydrolyse	Depends on $K_a$ vs $K_b$

Concept Block

5. Solubility Product (Ksp), Common Ion Effect, and Precipitation

For a sparingly soluble salt $A_xB_y \rightleftharpoons x A^{y+} + y B^{x-}$ :

K_{sp}=[A^{y+}]^x[B^{x-}]^y

For BaSO $_4$ : $K_{sp}=[Ba^{2+}][SO_4^{2-}]$ . If solubility is $s$ , then $K_{sp}=s^2$ . For $A_2B$ : $K_{sp}=4s^3$ .

Ionic product (IP): Same expression as $K_{sp}$ but with actual (non-equilibrium) concentrations.

$IP < K_{sp}$ : solution is unsaturated, more salt can dissolve
$IP = K_{sp}$ : solution is saturated (equilibrium)
$IP > K_{sp}$ : precipitation occurs until equilibrium is reached

Common ion effect: Adding a common ion decreases the solubility of a sparingly soluble salt. Example: adding NaCl to AgCl solution decreases Ag $^+$ ion concentration by driving equilibrium backward.

NEET trap: Only temperature changes the value of

K

(including

K_{sp}

). Adding a common ion shifts the equilibrium but doesn't change

K_{sp}

. Pressure changes don't affect ionic equilibria significantly (liquids are incompressible).

Practice Tests

5 Chapter Tests of 25 Questions Each

Each test is original, NEET-aligned, and answer-backed. Use them as sectional revision instead of a single long mock so your weak subtopics become easier to identify quickly.

Test 1: Chemical Equilibrium

Kc, Kp, Q, Le Chatelier principle, and equilibrium position.

Test 2: Acid-Base Basics

Ka, Kb, pH, pOH, conjugate pairs, and water ionisation.

Test 3: Buffers and Ksp

Buffer action, hydrolysis, precipitation, and common ion effect.

Test 4: Equilibrium Numericals

ICE tables, weak-acid approximations, pH numericals, and solubility calculations.

Test 5: Mixed NEET Drill

Shift logic, pH, K relations, and integrated chemical-plus-ionic equilibrium questions.

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