Section 1: The Equilibrium Concept

Equilibrium in chemistry describes a dynamic steady state in reversible reactions. Unlike a static balance, both forward and backward reactions continue at equal rates. Key consequences: concentrations are constant but not zero; equilibrium can be reached from either direction; the position of equilibrium (described by K) depends only on temperature.

The equilibrium constant Kc encodes the ratio of product to reactant concentrations (each raised to stoichiometric coefficients) at equilibrium. Its value is fixed at a given temperature. Increasing the value of K means products are more favoured; decreasing K means reactants are favoured. The reaction quotient Q uses the same expression with instantaneous concentrations to predict which direction a system will shift.

Section 2: Kp, Kc, and the $\Delta n$ Formula

For reactions involving gases, Kp (in terms of partial pressures) is more convenient than Kc. Since partial pressure P = CRT (from PV = nRT), substituting into the equilibrium expression yields Kp = Kc(RT)^$\Delta n$. Calculating $\Delta n$ correctly is critical: count only gaseous moles, not solids or liquids.

Example: $N_{2}$ + $3H_{2}$ ⇌ $2NH_{3}$ → $\Delta n$ = 2 − 4 = −2; Kp < Kc (since RT > 1 at normal temperatures, raising to −2 makes Kp smaller than Kc).

Section 3: Le Chatelier's Principle

Le Chatelier's principle is both conceptually important and heavily NEET-tested. The principle applies to all equilibria — chemical and ionic. Key points:

• Only temperature changes K.
• Concentration changes and pressure changes shift position but leave K unchanged.
• Catalyst: no shift, no K change — only rate of attainment increases.
• Inert gas differentiation: constant V (no effect) vs constant P (shift toward more gas moles).

Section 4: Acid-Base Theories and pH

Three theories govern acid-base behaviour with increasing generality: Arrhenius (aqueous only) → Bronsted-Lowry (any proton transfer) → Lewis (any electron-pair transfer). NEET predominantly uses Bronsted-Lowry for pH problems and conjugate pair relationships, and Lewis for identifying acids in inorganic/organic contexts.

pH calculations follow from the relevant equilibrium expression: strong acids give [$H^{+}$] directly; weak acids require [$H^{+}$] = √(Ka·C); buffers use Henderson-Hasselbalch. Kw = 10^{-14} at 25°C links pH and pOH. Salt hydrolysis produces acidic or basic solutions depending on the relative strengths of parent acid and base.

Section 5: Solubility Product and Common Ion Effect

Ksp is an equilibrium constant for the dissolution of sparingly soluble salts. Its magnitude determines solubility in pure water (using ICE tables). The common ion effect — an application of Le Chatelier — reduces solubility when a common ion is present in solution. This principle is the basis of selective precipitation in qualitative analysis. The critical test: comparing ionic product Qsp to Ksp to decide whether precipitation will occur.