Equilibrium is among the highest-yield Physical Chemistry topics in NEET, consistently contributing 3–4 questions per year. It is divided into two interconnected domains: chemical equilibrium (behaviour of reversible reactions) and ionic equilibrium (acid-base chemistry, buffers, and solubility).

Chemical Equilibrium

A reversible reaction reaches dynamic equilibrium when the rate of the forward reaction equals the rate of the backward reaction. At this point, the concentrations of all species remain constant — but they are not necessarily equal, and molecules continue to interconvert at the molecular level.

The equilibrium constant Kc is defined by the law of mass action: for the reaction aA + bB ⇌ cC + dD,

$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

where brackets denote equilibrium molar concentrations. Pure solids and pure liquids are excluded (activity = 1). For gaseous reactions, Kp uses partial pressures and is related to Kc by:

$K_p = K_c(RT)^{\Delta n}$

where $\Delta n$ = (moles of gaseous products) − (moles of gaseous reactants) and R = 0.0821 L·atm/(mol·K).

The reaction quotient Q has the same mathematical form as K but uses instantaneous (non-equilibrium) concentrations. Comparing Q to K predicts the direction of reaction: Q < K → forward; Q > K → backward; Q = K → at equilibrium. Thermodynamically, K is linked to Gibbs energy: $\Delta G$° = −RT ln K. When K > 1, $\Delta G$° < 0 and products are favoured; when K < 1, $\Delta G$° > 0 and reactants are favoured.

Le Chatelier's principle states that a system at equilibrium subjected to a stress will shift to partially relieve that stress. The key stresses and their effects:

• Concentration: Adding reactant shifts equilibrium forward; adding product shifts it backward. K does not change.
• Pressure: Increasing pressure (decreasing volume) shifts equilibrium toward the side with fewer moles of gas ($\Delta n$). No effect when $\Delta n$ = 0. K does not change.
• Temperature: Increasing temperature favours the endothermic direction; K changes — increases for endothermic reactions, decreases for exothermic reactions. This is the ONLY factor that changes K.
• Catalyst: Provides an alternate pathway with lower activation energy for both forward and backward reactions equally. No shift in equilibrium position; no change in K. Only the rate of attaining equilibrium increases.
• Inert gas at constant volume: Partial pressures and concentrations of reacting species are unchanged → no effect. At constant pressure, adding inert gas expands volume, decreases concentrations, and shifts equilibrium toward the side with more moles of gas.

Ionic Equilibrium

Acid-Base Theories:

• Arrhenius: Acid produces $H^{+}$ in water; base produces $OH^{-}$. Limited to aqueous solutions.
• Bronsted-Lowry: Acid is a proton donor; base is a proton acceptor. Introduces conjugate acid-base pairs. Ka × Kb = Kw for any conjugate pair.
• Lewis: Acid accepts electron pair; base donates electron pair. Broadest definition; explains reactions like $BF_{3}$ + $NH_{3}$.

Water and pH: Kw = [$H^{+}$][$OH^{-}$] = 10^{-14} at 25°C; pH = −log[$H^{+}$]; pOH = −log[$OH^{-}$]; pH + pOH = 14 (at 25°C only). The neutral pH = 7 is valid only at 25°C; at higher temperatures Kw increases and neutral pH drops below 7.

Weak Acid Ionisation: Ka = Cα^{2} (where C = concentration, α = degree of dissociation). When α ≪ 1: [$H^{+}$] ≈ √(Ka × C).

Buffers and Henderson-Hasselbalch: A buffer is a solution of a weak acid and its conjugate salt (or weak base and conjugate salt) that resists pH change. The pH of an acidic buffer is given by:

$pH = pK_a + \log\frac{[\text{salt}]}{[\text{acid}]}$

When [salt] = [acid], pH = pKa — the buffer's most effective point. For basic buffers: pOH = pKb + log([salt]/[base]).

Salt Hydrolysis: The pH of a salt solution depends on the strengths of its parent acid and base. Strong acid + strong base → neutral (pH = 7). Strong acid + weak base → acidic (pH < 7). Weak acid + strong base → basic (pH > 7). Weak acid + weak base → pH = 7 + ½pKa − ½pKb.

Common Ion Effect: Adding an ion common to a weak electrolyte equilibrium suppresses its ionisation (Le Chatelier applied to ionic equilibrium). This dramatically reduces the solubility of sparingly soluble salts.

Solubility Product (Ksp): For a sparingly soluble salt AmBn ⇌ mA^n+ + nB^m−, Ksp = [A^n+]^m [B^m−]^n. Solubility s is calculated using the ICE method. Precipitation occurs when the ionic product Qsp exceeds Ksp. Selective precipitation (qualitative analysis) exploits differences in Ksp — the salt with the lower Ksp precipitates first.

Most Tested NEET Traps

• Catalyst never shifts equilibrium or changes K.
• Inert gas at constant volume has no effect.
• pH of 10^{-8} M HCl ≈ 6.98, not 8 (acid cannot be basic).
• pH = 7 is neutral only at 25°C.
• Ka × Kb = Kw (not Ka + Kb).
• Ksp of $Ag_{2}CrO_{4}$ = [$Ag^{+}$]^{2}[$CrO_{4}^{2-}$] — stoichiometric coefficients become exponents.