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Science Class 11 - Equilibrium Notes

Comprehensive study notes for Class 11 - Equilibrium olympiad preparation

Equilibrium

Welcome to the chapter on Equilibrium for Class 11. In this chapter, you will learn about chemical equilibrium, the dynamic nature of equilibrium, the law of mass action, equilibrium constants, and how equilibrium is affected by changes in conditions. By the end of this chapter, you will be able to solve problems related to equilibrium and understand its importance in chemistry and real life.

Key Concepts

  • Chemical Equilibrium: A state in a reversible reaction where the rates of the forward and backward reactions are equal and the concentrations of reactants and products remain constant.
  • Dynamic Equilibrium: Both forward and backward reactions continue to occur, but there is no net change in the concentrations of reactants and products.
  • Law of Mass Action: The rate of a chemical reaction is proportional to the product of the concentrations of the reactants, each raised to the power of their stoichiometric coefficients.
  • Equilibrium Constant (Kc, Kp): A number that expresses the ratio of the concentrations (or partial pressures) of products to reactants at equilibrium.

Types of Equilibrium

  • Physical Equilibrium: Involves physical processes, e.g., evaporation of water in a closed vessel.
  • Chemical Equilibrium: Involves chemical reactions, e.g., the reaction between nitrogen and hydrogen to form ammonia.

Law of Mass Action and Equilibrium Constant

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

The equilibrium constant (Kc) is given by:
Kc = [C]c[D]d / [A]a[B]b

For reactions involving gases, the equilibrium constant can also be expressed in terms of partial pressures (Kp).

Characteristics of Equilibrium

  • It is dynamic in nature.
  • It can be attained only in a closed system.
  • The properties of the system remain constant at equilibrium.
  • The equilibrium constant is constant at a given temperature.

Le Chatelier’s Principle

If a system at equilibrium is disturbed by a change in concentration, temperature, or pressure, the system shifts in a direction that tends to reduce the disturbance and restore equilibrium.

  • Change in Concentration: Adding more reactant shifts equilibrium towards products.
  • Change in Temperature: Increasing temperature favors the endothermic direction.
  • Change in Pressure: Increasing pressure favors the side with fewer moles of gas.

Applications of Equilibrium

  • Industrial processes like the Haber process for ammonia and the Contact process for sulfuric acid.
  • Understanding biological systems, such as oxygen transport in blood.

Practice Questions

  1. Define chemical equilibrium and give one example.
  2. Write the expression for Kc for the reaction: 2SO2(g) + O2(g) ⇌ 2SO3(g).
  3. What will happen to the equilibrium if more reactant is added to a system at equilibrium?
  4. Explain Le Chatelier’s Principle with an example.
  5. How does temperature affect the value of the equilibrium constant?

Challenge Yourself

  • Calculate the equilibrium concentrations for the reaction: H2(g) + I2(g) ⇌ 2HI(g), given initial concentrations and Kc.
  • Describe how equilibrium is established in a closed vessel containing water at 100°C.

Did You Know?

  • The concept of equilibrium is not only used in chemistry but also in economics and biology!
  • At equilibrium, the rate of the forward reaction equals the rate of the backward reaction, but the reactions do not stop.

Glossary

  • Equilibrium: A state where the concentrations of reactants and products remain unchanged with time.
  • Dynamic: Always changing, but with no net change overall.
  • Le Chatelier’s Principle: A rule to predict the effect of a change in conditions on equilibrium.
  • Equilibrium Constant: A value that expresses the ratio of products to reactants at equilibrium.

Answers to Practice Questions

  1. Chemical equilibrium is the state in a reversible reaction where the rates of the forward and backward reactions are equal. Example: N2(g) + 3H2(g) ⇌ 2NH3(g).
  2. Kc = [SO3]2 / ([SO2]2 [O2])
  3. The equilibrium will shift towards the products to use up the added reactant.
  4. If the concentration of a reactant is increased, the system shifts to produce more products (e.g., in the Haber process, adding more N2 or H2 increases NH3 production).
  5. Temperature changes can increase or decrease the value of the equilibrium constant depending on whether the reaction is exothermic or endothermic.

Understanding equilibrium helps you master chemistry and solve real-world problems!

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