Equilibrium Chapter Notes | CBSE Class 11 Chemistry

Key Definitions in Equilibrium

Chemical Equilibrium

A state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant over time.

Dynamic Equilibrium

A type of chemical equilibrium where the forward and reverse reactions continue to occur at equal rates, meaning the system is active at a molecular level, but macroscopically no net change is observed.

Law of Mass Action

At a given temperature, the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants, each raised to the power equal to its stoichiometric coefficient in the balanced chemical equation.

Equilibrium Constant (K_c)

For a reversible reaction aA + bB ⇌ cC + dD, K_c = ([C]^c [D]^d) / ([A]^a [B]^b) at equilibrium, where [ ] denotes molar concentrations.

Le Chatelier's Principle

If a change of condition (temperature, pressure, concentration) is applied to a system in equilibrium, the system will shift in a direction that tends to counteract the change.

Conjugate Acid-Base Pair

An acid and a base that differ by a proton (H⁺). When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

pH

A measure of the acidity or alkalinity of a solution, defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH = -log[H⁺].

Buffer Solution

A solution that resists changes in pH upon the addition of small amounts of acid or base, typically composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.

Solubility Product (K_sp)

For a sparingly soluble ionic compound, it is the product of the molar concentrations of its constituent ions, each raised to the power of its stoichiometric coefficient in the balanced dissolution equation, at saturation.

Understanding the Dynamic Nature of Equilibrium

Chemical reactions can be broadly classified as reversible or irreversible. Irreversible reactions proceed in one direction until reactants are consumed. Reversible reactions, however, can proceed in both forward and reverse directions. When a reversible reaction starts, reactants combine to form products (forward reaction). As products accumulate, they begin to react to reform the reactants (reverse reaction). Initially, the rate of the forward reaction is high, and the rate of the reverse reaction is low. Over time, as reactants are consumed, the forward reaction rate decreases. Concurrently, as products accumulate, the reverse reaction rate increases.

Eventually, a state is reached where the rate of the forward reaction becomes exactly equal to the rate of the reverse reaction. This is the state of chemical equilibrium. At this point, the concentrations of reactants and products become constant and do not change further, provided external conditions remain unaltered. It's crucial to understand that equilibrium is not a static state where reactions stop. Instead, it's a dynamic state where both forward and reverse reactions continue to occur at equal and opposite rates. This means molecules are constantly being converted from reactants to products and vice-versa, but the net change in concentration is zero. This dynamic balance is fundamental to how many biological and industrial processes operate efficiently.

Must Remember: Key Concepts & Formulas

• Equilibrium is Dynamic: Both forward and reverse reactions continue at equal rates.
• Equilibrium Constant (K): Temperature-dependent. Does not depend on initial concentrations or catalyst.
• Relation between K_c and K_p: K_p = K_c (RT)^Δn_g, where Δn_g = (moles of gaseous products) - (moles of gaseous reactants). R is the gas constant, T is absolute temperature.
• Le Chatelier's Principle: Predicts shift: increase reactant/decrease product shifts right; increase product/decrease reactant shifts left. Increase pressure shifts to fewer moles of gas; decrease pressure shifts to more moles of gas. Increase temperature shifts to absorb heat (endothermic direction); decrease temperature shifts to release heat (exothermic direction).
• pH Scale: pH = -log[H⁺], pOH = -log[OH⁻], pH + pOH = 14 (at 298 K).
• Ionization Constants: K_a for weak acids, K_b for weak bases. K_a × K_b = K_w (for conjugate acid-base pairs).
• Hydrolysis: Salts of strong acid/weak base are acidic (e.g., NH₄Cl). Salts of weak acid/strong base are basic (e.g., CH₃COONa). Salts of strong acid/strong base are neutral (e.g., NaCl).
• Buffer Action: Weak acid/conjugate base or weak base/conjugate acid system resists pH changes. Henderson-Hasselbalch equation: pH = pK_a + log([Salt]/[Acid]) for acidic buffers.
• Solubility Product (K_sp): Higher K_sp means higher solubility for compounds with the same stoichiometry. Common ion effect decreases solubility of sparingly soluble salts.

Homogeneous vs. Heterogeneous Equilibrium

Worked Examples: Equilibrium Calculations

• {"title":"Example 1: Calculating K_c","bodyMarkdown":"Problem: For the reaction A(g) + B(g) ⇌ 2C(g), at equilibrium, [A] = 0.5 M, [B] = 0.2 M, and [C] = 0.8 M. Calculate K_c.\n\nSolution:\nThe equilibrium constant expression is K_c = [C]² / ([A][B])\nSubstitute the equilibrium concentrations:\nK_c = (0.8)² / (0.5 × 0.2)\nK_c = 0.64 / 0.10\nK_c = 6.4"}
• {"title":"Example 2: pH Calculation","bodyMarkdown":"Problem: Calculate the pH of a 0.01 M solution of HCl.\n\nSolution:\nHCl is a strong acid, so it completely dissociates:\nHCl(aq) → H⁺(aq) + Cl⁻(aq)\nTherefore, [H⁺] = 0.01 M = 1 × 10⁻² M\npH = -log[H⁺]\npH = -log(1 × 10⁻²)\npH = -(-2)\npH = 2"}

Exam Tip: Mastering Le Chatelier's Principle

When applying Le Chatelier's Principle, always analyze the stress applied and then determine the direction of shift that alleviates that stress. Pay close attention to:

1. Temperature changes: Identify if the reaction is endothermic (ΔH > 0) or exothermic (ΔH < 0). Heat can be treated as a reactant (endothermic) or product (exothermic).
2. Pressure changes: Only affect reactions involving gases. Count the total moles of gaseous reactants and products. The system shifts to reduce the number of gas moles if pressure is increased, and vice-versa.
3. Inert gas addition: If volume is kept constant, adding an inert gas does not change partial pressures of reactants/products, so equilibrium is unaffected. If pressure is kept constant (volume increases), partial pressures decrease, and the equilibrium shifts to the side with more moles of gas.

Misapplying these nuances is a common trap!

Practice Questions with Solutions

• Q1: What is the conjugate base of H₂S? A1: The conjugate base of H₂S is HS⁻ (formed by losing one proton).
• Q2: For an exothermic reaction at equilibrium, how will increasing the temperature affect the value of the equilibrium constant (K)? A2: For an exothermic reaction, increasing temperature shifts the equilibrium to the left (reactant side). This decreases the product concentration and increases reactant concentration, thus decreasing the value of K.
• Q3: Write the expression for K_p for the reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). A3: K_p = (P_NH₃)² / (P_N₂ × P_H₂)³.
• Q4: A catalyst is added to a reversible reaction at equilibrium. What effect does this have on the equilibrium position? A4: A catalyst increases the rate of both forward and reverse reactions equally. Therefore, it helps the system reach equilibrium faster but does not shift the equilibrium position or change the value of the equilibrium constant.

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