Class 10 Chemistry — Chapter 1: Chemical Equilibrium
Sindh Board — Complete solved notes and answers for Chapter 1 (definitions, short & long question solutions).
Important Definitions
- Chemical Equilibrium: The state in a reversible reaction when the rate of forward reaction equals the rate of backward reaction and concentrations of reactants and products remain constant.
- Reversible Reaction: A reaction that proceeds in both directions. Example: N2(g) + 3H2(g) ⇌ 2NH3(g).
- Irreversible Reaction: A reaction that goes only in one direction (e.g., combustion of fuel).
- Dynamic Equilibrium: At equilibrium both forward and backward reactions continue but at equal rates, so no net change is observed.
- Le-Chatelier’s Principle: If a system at equilibrium is disturbed (by change in concentration, pressure, or temperature), the system shifts to oppose the disturbance.
- Catalyst: A substance that speeds up the rate of a reaction without being consumed; it decreases the time to reach equilibrium but does not change equilibrium position or constant.
Solved Short Questions
Q1: Define chemical equilibrium.
It is the state when the rate of forward and backward reactions become equal and the concentrations of reactants and products remain constant in a closed system.
Q2: Difference between reversible and irreversible reactions?
Reversible reactions can proceed in both directions (e.g., Haber process). Irreversible reactions proceed only in one direction (e.g., burning of coal).
Q3: Write the Haber process equation.
N2(g) + 3H2(g) ⇌ 2NH3(g). Industrial conditions: iron catalyst, high pressure (∼150–300 atm), moderate temperature (400–500 °C).
Q4: Why is equilibrium called a dynamic state?
Because both forward and backward reactions continue to occur at equal rates; molecules are still reacting though the overall concentrations remain constant.
Q5: Effect of temperature on an exothermic reaction at equilibrium.
Increasing temperature shifts the equilibrium toward the reactants (left) for an exothermic reaction, reducing product yield. Lowering temperature favors product formation (right).
Q6: Examples of homogeneous and heterogeneous equilibrium.
Homogeneous: N2(g) + 3H2(g) ⇌ 2NH3(g). Heterogeneous: CaCO3(s) ⇌ CaO(s) + CO2(g).
Solved Long Questions
Q1: Characteristics of chemical equilibrium.
- Equilibrium is reached only in a closed system.
- It is a dynamic state where forward and backward rates are equal.
- Concentrations of reactants and products remain constant (but not necessarily equal).
- Equilibrium position can be shifted by changing concentration, pressure (for gases), or temperature.
- Catalysts speed up equilibrium attainment but do not change equilibrium position.
Q2: State Le-Chatelier’s principle and explain with examples.
Principle: If a system at equilibrium is disturbed by change in concentration, pressure or temperature, it shifts to minimize the disturbance.
Examples:
- Concentration: In N2 + 3H2 ⇌ 2NH3, adding more H2 shifts equilibrium right to produce more NH3.
- Pressure: For gases, increasing pressure favors the side with fewer moles. In Haber process (4 mol gas → 2 mol gas), higher pressure favors ammonia (products).
- Temperature: For exothermic reactions (heat as product), raising temperature shifts equilibrium left (toward reactants).
Q3: Factors affecting equilibrium.
- Concentration: Changing concentration shifts the position to oppose the change.
- Pressure: Affects equilibria involving gases; increase in pressure favors fewer gas moles.
- Temperature: Changes equilibrium constant; increasing temperature favors endothermic direction.
- Catalyst: Increases rate of both forward and backward reactions equally; no change in equilibrium composition.
Q4: Industrial importance — Haber process & Contact process (brief).
Haber process: N2 + 3H2 ⇌ 2NH3. Produces ammonia for fertilizers. Optimized by moderate temperature, high pressure and iron catalyst to balance rate and yield.
Contact process (summary): For H2SO4 manufacture — SO2 → SO3 (catalysed by V2O5) then converted to H2SO4. Equilibrium and conditions are chosen for best yield and rate.
Q5: Homogeneous vs Heterogeneous equilibrium (difference).
Homogeneous: All reactants and products are in the same phase (e.g., all gases).
Heterogeneous: Reactants/products are in different phases (e.g., solid + gas). Solids and pure liquids do not appear in the expression for Kc.
Quick Revision Points
- Equilibrium occurs in closed systems only.
- At equilibrium: rateforward = ratebackward.
- Catalyst affects rates, not equilibrium constant.
- For exothermic reactions, lower temperature favours product formation.
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