Class 11 Chemistry Chapter 7 Equilibrium Revision Notes

Understanding Equilibrium: The Dynamic Balance

Equilibrium represents a state in a reversible process where the rate of the forward reaction equals the rate of the reverse reaction. At equilibrium, the concentrations of reactants and products remain constant over time, but the reactions do not stop; rather, they continue simultaneously at equal rates. This is known as dynamic equilibrium.

Key characteristics of equilibrium include:

• Dynamic Nature: Reactions continue in both directions. It's not a static state where everything stops.
• Reversibility: Equilibrium can only be established in reversible reactions, denoted by a double arrow (⇌).
• Attainable from Either Side: Equilibrium can be reached whether you start with reactants or products.
• Constant Macroscopic Properties: Measurable properties like concentration, pressure, temperature, and color remain constant at equilibrium.
• Catalyst Effect: A catalyst helps achieve equilibrium faster but does not alter the equilibrium state or the equilibrium constant. It speeds up both forward and reverse reactions equally.
• Closed System: Chemical equilibrium is generally established in a closed system to prevent loss or gain of substances.

Equilibrium can be broadly classified into:

1. Physical Equilibrium: Involves a change of state (e.g., solid ⇌ liquid, liquid ⇌ gas, dissolution of solids or gases in liquids) where physical properties are in balance.
2. Chemical Equilibrium: Involves chemical reactions where reactants and products are in dynamic balance. This is the primary focus of the chapter.

Key Equilibrium Terms

Equilibrium Constant (Kc/Kp)

A value that expresses the ratio of product concentrations/partial pressures to reactant concentrations/partial pressures at equilibrium, each raised to the power of their stoichiometric coefficients. It indicates the extent to which a reaction proceeds.

Le Chatelier's Principle

States that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that counteracts the change to re-establish equilibrium.

Conjugate Acid-Base Pair

Two species that differ from each other by a single proton (H+). An acid, after donating a proton, forms its conjugate base; a base, after accepting a proton, forms its conjugate acid.

Buffer Solution

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

Solubility Product (Ksp)

The equilibrium constant for the dissolution of a sparingly soluble salt. It represents the product of the concentrations of its ions in a saturated solution, each raised to the power of their stoichiometric coefficients.

pH

A measure of the acidity or basicity of an aqueous solution, defined as the negative logarithm (base 10) of the hydrogen ion (H+ or H3O+) concentration: pH = -log[H+].

Ionization Constant (Ka/Kb)

The equilibrium constant for the ionization of a weak acid (Ka) or a weak base (Kb) in water, indicating its strength. Higher values mean stronger acids/bases.

Chemical Equilibrium: Law of Mass Action & Equilibrium Constant

The Law of Mass Action, proposed by Guldberg and Waage, states that at a constant temperature, the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants, with each concentration term raised to the power of its stoichiometric coefficient in the balanced chemical equation.

For a general reversible reaction: aA + bB ⇌ cC + dD

The equilibrium constant (Kc) is expressed as:

Kc = ([C]^c [D]^d) / ([A]^a [B]^b)

where [A], [B], [C], [D] are molar concentrations at equilibrium.

When reactants and products are gases, the equilibrium constant can also be expressed in terms of partial pressures (Kp):

Kp = (P_C^c P_D^d) / (P_A^a P_B^b)

The relationship between Kp and Kc is given by:

Kp = Kc(RT)^Δn

where:

• R is the gas constant (0.0831 L bar mol⁻¹ K⁻¹ or 0.0821 L atm mol⁻¹ K⁻¹).
• T is the absolute temperature in Kelvin.
• Δn is the change in the number of moles of gaseous products minus the number of moles of gaseous reactants (Δn = (c+d) - (a+b)).

Characteristics of Equilibrium Constant (K):

• Temperature Dependent: K changes only with temperature.
• Independent of Initial Concentrations: K is constant for a given reaction at a given temperature, regardless of initial concentrations.
• Predicts Extent of Reaction: A large K (>10³) indicates products are favored; a small K (<10⁻³) indicates reactants are favored; K ≈ 1 means comparable concentrations of reactants and products.
• Reaction Reversal: If a reaction is reversed, the new K is the reciprocal of the original K.
• Multiplying by a Factor: If a reaction is multiplied by a factor 'n', the new K is (original K)ⁿ.
• Adding Reactions: If two reactions are added, the K for the overall reaction is the product of their individual K values.

Factors Affecting Equilibrium: Le Chatelier's Principle

1. Effect of Change in Concentration — If the concentration of a reactant is increased, the equilibrium shifts in the forward direction (to consume the added reactant). If the concentration of a product is increased, the equilibrium shifts in the reverse direction (to consume the added product).
2. Effect of Change in Pressure (for gaseous reactions) — If pressure is increased (by decreasing volume), the equilibrium shifts towards the side with fewer moles of gas. If pressure is decreased (by increasing volume), the equilibrium shifts towards the side with more moles of gas. If Δn = 0, pressure has no effect.
3. Effect of Change in Temperature — For an endothermic reaction (ΔH > 0), increasing temperature shifts equilibrium in the forward direction (absorbs heat). For an exothermic reaction (ΔH < 0), increasing temperature shifts equilibrium in the reverse direction (consumes heat). Decreasing temperature has the opposite effect.
4. Effect of Adding Inert Gas — Adding an inert gas at constant volume has no effect on equilibrium (as partial pressures and concentrations of reacting gases remain unchanged). Adding an inert gas at constant pressure increases the total volume, effectively decreasing partial pressures of reacting gases, thus shifting equilibrium towards the side with more moles of gas (similar to decreasing pressure).
5. Effect of Catalyst — A catalyst does not affect the position of equilibrium or the value of the equilibrium constant. It only speeds up the attainment of equilibrium by lowering the activation energy for both forward and reverse reactions equally.

Key Concepts in Ionic Equilibrium

• Acids and Bases: Defined by Arrhenius (H+/OH- in water), Brønsted-Lowry (proton donors/acceptors), and Lewis (electron pair acceptors/donors). Brønsted-Lowry is most commonly used for equilibrium.
• Strong vs. Weak Electrolytes: Strong electrolytes ionize completely in solution (e.g., HCl, NaOH); weak electrolytes ionize partially (e.g., CH3COOH, NH4OH).
• Ionization of Weak Acids/Bases: Quantified by Ka (acid dissociation constant) and Kb (base dissociation constant). Higher Ka/Kb means stronger acid/base. For a conjugate pair, Ka * Kb = Kw (ion product of water = 1.0 x 10⁻¹⁴ at 298 K).
• Common Ion Effect: The suppression of the dissociation of a weak electrolyte by the addition of a strong electrolyte containing a common ion. This principle is crucial for buffer action and precipitation.
• Hydrolysis of Salts: Salts formed from strong acid/weak base, weak acid/strong base, or weak acid/weak base hydrolyze in water to produce acidic, basic, or neutral solutions, respectively. The extent of hydrolysis is given by the hydrolysis constant (Kh).
• Buffer Solutions: Resist pH change. Acidic buffers (weak acid + its salt with a strong base) work around pH < 7. Basic buffers (weak base + its salt with a strong acid) work around pH > 7. pH of a buffer can be calculated using the Henderson-Hasselbalch Equation: pH = pKa + log([Salt]/[Acid]) or pOH = pKb + log([Salt]/[Base]).
• Solubility Product (Ksp): For a sparingly soluble salt AxBy ⇌ xA^y+ + yB^x-, Ksp = [A^y+]^x * [B^x-]^y. If Qsp > Ksp, precipitation occurs. If Qsp < Ksp, no precipitation. If Qsp = Ksp, the solution is saturated.
• Factors Affecting Solubility: Common ion effect decreases solubility. pH can affect solubility of salts of weak acids/bases.

Solved Mini-Examples

• {"title":"Example 1: Calculating Kc","description":"For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g), at equilibrium, concentrations are [N2] = 0.5 M, [H2] = 1.5 M, and [NH3] = 0.2 M. Calculate Kc.","solution":"Kc = [NH3]^2 / ([N2] [H2]^3)\nKc = (0.2)^2 / (0.5 (1.5)^3)\nKc = 0.04 / (0.5 * 3.375)\nKc = 0.04 / 1.6875\nKc ≈ 0.0237 M⁻²"}
• {"title":"Example 2: pH Calculation of a Weak Acid","description":"Calculate the pH of a 0.1 M CH3COOH solution if Ka = 1.8 x 10⁻⁵.","solution":"Let x be the concentration of H+ ions.\nCH3COOH ⇌ CH3COO⁻ + H⁺\nInitial: 0.1 M 0 0\nEquilibrium: 0.1-x M x x\nAssume x << 0.1 (weak acid approximation).\nKa = [CH3COO⁻][H⁺] / [CH3COOH] = x*x / (0.1-x) ≈ x^2 / 0.1\n1.8 x 10⁻⁵ = x^2 / 0.1\nx^2 = 1.8 x 10⁻⁶\nx = sqrt(1.8 x 10⁻⁶) = 1.34 x 10⁻³ M\npH = -log[H⁺] = -log(1.34 x 10⁻³) ≈ 2.87"}

Exam Strategy: Mastering Equilibrium

Equilibrium is a scoring chapter but often involves tricky conceptual questions and calculations. For numerical problems, always write down the balanced chemical equation and clearly define initial, change, and equilibrium concentrations/pressures. Pay close attention to units of Kc and Kp. When applying Le Chatelier's Principle, visualize the 'stress' and then the 'relief' the system provides. Remember that catalysts only affect the rate, not the equilibrium position. For ionic equilibrium, mastering the Henderson-Hasselbalch equation and conditions for buffer solutions is key. Practice problems involving the common ion effect and solubility product (Ksp) thoroughly, as these are frequent traps. Don't forget the sign convention for Δn in Kp = Kc(RT)^Δn and ensure temperature is always in Kelvin.

Practice Questions with Solutions

• Q: What are the characteristics of a system at dynamic equilibrium? A: Constant macroscopic properties (concentration, pressure, temp), forward and reverse reaction rates are equal, and it's attainable from either direction.
• Q: How does increasing temperature affect the equilibrium of an exothermic reaction? A: Increasing temperature for an exothermic reaction will shift the equilibrium in the reverse direction, favoring reactants, to consume the added heat.
• Q: Define a buffer solution and give an example of an acidic buffer. A: A buffer solution resists significant changes in pH upon addition of small amounts of acid or base. An acidic buffer consists of a weak acid and its salt with a strong base, e.g., CH3COOH and CH3COONa.
• Q: For the reaction 2SO2(g) + O2(g) ⇌ 2SO3(g), what is the relationship between Kp and Kc? A: Kp = Kc(RT)^Δn. Here, Δn = (2) - (2+1) = -1. So, Kp = Kc(RT)⁻¹ or Kp = Kc/RT.

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