Board exam focus — Equilibrium (CBSE & HBSE)

Equilibrium is the state where the rates of forward and backward processes become equal and macroscopic properties stay constant. This chapter covers equilibrium in physical and chemical processes, the law of mass action, equilibrium constants Kc and Kp (and their relation), Le Chatelier's principle, and a deep dive into ionic equilibrium — acid-base theories (Arrhenius, Bronsted-Lowry, Lewis), ionization, Ka/Kb, pH, buffers, the common ion effect and solubility product Ksp. It is both conceptual and numerically demanding (especially pH problems).

Physical & Chemical Equilibrium, Law of Mass Action, Kc, Kp & Le Chatelier

Nature of Equilibrium

At equilibrium, the rate of the forward reaction equals the rate of the backward reaction, so concentrations of reactants and products remain constant. It is dynamic (both reactions continue) and is attained only in a closed system.

Physical Equilibria

Law of Mass Action & Equilibrium Constant

For a general reaction aA + bB ⇌ cC + dD, the equilibrium constant in terms of concentration is:

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

In terms of partial pressures (for gases):

$$K_p = \frac{(p_C)^c (p_D)^d}{(p_A)^a (p_B)^b}$$

Relation Between Kp and Kc

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

where Δn = (moles of gaseous products) − (moles of gaseous reactants).

• If Δn = 0 ⇒ Kp = Kc.

Characteristics of K

• K is constant at a given temperature; changes only with temperature.
• Pure solids and pure liquids are omitted (their concentration is constant).
• For the reverse reaction, K' = 1/K. If coefficients are multiplied by n, K becomes Kⁿ.

Reaction Quotient (Q) and Predicting Direction

Q has the same form as K but uses non-equilibrium concentrations:

• Q < K ⇒ forward reaction proceeds (more products form).
• Q = K ⇒ at equilibrium.
• Q > K ⇒ reverse reaction proceeds.

Le Chatelier's Principle

If a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the equilibrium shifts to counteract (partly undo) the change.

Trap: A catalyst speeds up both forward and reverse reactions equally — it does NOT change K or the equilibrium position, only the time to reach it.

Ionic Equilibrium: Acids, Bases, Ka/Kb and pH

Acid-Base Theories

Conjugate acid-base pairs (Bronsted): differ by one proton. e.g., HCl/Cl⁻ and H₂O/H₃O⁺. A strong acid has a weak conjugate base.

Lewis examples: BF₃, AlCl₃, H⁺ are Lewis acids; NH₃, H₂O, OH⁻ are Lewis bases.

Ionization of Water & Kw

Water self-ionizes: 2H₂O ⇌ H₃O⁺ + OH⁻.

$$K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \text{ at } 298\,K$$

• Neutral: [H⁺] = [OH⁻] = 10⁻⁷ M.
• Acidic: [H⁺] > 10⁻⁷; Basic: [H⁺] < 10⁻⁷.

pH Scale

$$pH = -\log[H^+], \quad pOH = -\log[OH^-]$$ $$pH + pOH = 14 \text{ (at 298 K)}$$

• pH < 7 acidic, = 7 neutral, > 7 basic.

Strength of Weak Acids/Bases

For a weak acid HA ⇌ H⁺ + A⁻ with concentration c and degree of ionization α:

$$K_a = \frac{c\alpha^2}{1-\alpha} \approx c\alpha^2 \text{ (for small } \alpha)$$

Ostwald's dilution law: α = √(Kₐ/c) — degree of ionization increases on dilution.

Also, [H⁺] = √(Kₐ·c) for a weak monoprotic acid, and pKₐ = −log Kₐ. Lower pKₐ ⇒ stronger acid.

For a conjugate pair: Kₐ × K_b = Kw, so pKₐ + pK_b = 14.

Trap: Strong acids (HCl, HNO₃, H₂SO₄) ionize completely (α ≈ 1); pH is found directly from concentration. Weak acids need Ka.

Buffers, Common Ion Effect & Solubility Product (Ksp)

Buffer Solutions

A buffer resists change in pH on adding small amounts of acid or base.

• Acidic buffer: weak acid + its salt (e.g., CH₃COOH + CH₃COONa). pH < 7.
• Basic buffer: weak base + its salt (e.g., NH₄OH + NH₄Cl). pH > 7.

Henderson-Hasselbalch equations: $$pH = pK_a + \log\frac{[\text{salt}]}{[\text{acid}]}$$ $$pOH = pK_b + \log\frac{[\text{salt}]}{[\text{base}]}$$

Maximum buffer capacity occurs when [salt] = [acid], i.e., pH = pKₐ.

Common Ion Effect

The suppression of ionization of a weak electrolyte by adding a strong electrolyte that has a common ion.

• Example: adding CH₃COONa to CH₃COOH provides extra CH₃COO⁻, shifting CH₃COOH ⇌ H⁺ + CH₃COO⁻ to the LEFT ⇒ less ionization, higher pH.
• This principle underlies buffer action and salt analysis (e.g., adding NH₄Cl before NH₄OH in qualitative analysis).

Salt Hydrolysis (brief)

• Salt of strong acid + strong base (NaCl): neutral, pH ≈ 7.
• Salt of weak acid + strong base (CH₃COONa): basic, pH > 7.
• Salt of strong acid + weak base (NH₄Cl): acidic, pH < 7.

Solubility Product (Ksp)

For a sparingly soluble salt AₓBᵧ ⇌ xAⁿ⁺ + yBᵐ⁻:

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

For a salt of solubility S (mol/L):

• AB type (e.g., AgCl): Ksp = S²
• AB₂ or A₂B type (e.g., CaF₂, Ag₂CrO₄): Ksp = 4S³
• AB₃ type: Ksp = 27S⁴

Predicting Precipitation (Ionic Product, Q)

Compare the ionic product (Qsp) with Ksp:

• Qsp < Ksp ⇒ unsaturated, no precipitate (more solute dissolves).
• Qsp = Ksp ⇒ saturated solution (equilibrium).
• Qsp > Ksp ⇒ supersaturated ⇒ precipitation occurs.

Common ion effect on solubility: adding a common ion (e.g., Cl⁻ to an AgCl solution) decreases the salt's solubility — exploited in gravimetric analysis.

Frequently asked questions

Are these Equilibrium notes free?

Yes — the Equilibrium notes for Chemistry (Class 11) on Siksha Sarovar are completely free to read, with no account required.

Do these notes follow CBSE and HBSE?

Yes. The Equilibrium notes are NCERT-aligned and include guidance for both CBSE and Haryana Board (HBSE), with important questions and MCQs for revision.

What does the Equilibrium chapter cover?

Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Equilibrium.