CBSE Class 10 Chemistry Chemical Equilibrium Notes Set 12

Reversible Reactions and Equilibrium

In a reversible reaction the products can react to form the reactants again. When the rates of the forward and backward reactions become equal in a closed system, a state of chemical equilibrium is reached: the concentrations stay constant although both reactions continue (dynamic equilibrium). Example: N₂ + 3H₂ ⇌ 2NH₃.

Law of Mass Action and Equilibrium Constant

The rate of a chemical reaction is proportional to the product of the active masses (molar concentrations) of the reactants, each raised to the power of its coefficient. For aA + bB ⇌ cC + dD the equilibrium constant is K = [C]ᶜ[D]ᵈ / ([A]ᵃ[B]ᵇ). K depends only on temperature; a catalyst does not change it. A large K means that the products are favoured.

Le Chatelier's Principle

If a change of concentration, pressure or temperature is applied to a system at equilibrium, the system shifts in the direction that opposes the change.
1. Adding a reactant (or removing a product) shifts the equilibrium to the right.
2. Raising the pressure favours the side with fewer gas moles; if the moles are equal, pressure has no effect.
3. Raising the temperature favours the endothermic direction; lowering it favours the exothermic direction.
4. A catalyst only helps to reach equilibrium sooner.

Industrial Applications

Haber process: N₂ + 3H₂ ⇌ 2NH₃ + heat. High pressure (about 200 atm), moderate temperature (about 450°C) and an iron catalyst with molybdenum as promoter.
Contact process: 2SO₂ + O₂ ⇌ 2SO₃ + heat. High pressure, about 450°C and vanadium pentoxide catalyst.

Ionic Product of Water and pH

Water ionises very slightly: H₂O ⇌ H⁺ + OH⁻. The ionic product K_w = [H⁺][OH⁻] = 10⁻¹⁴ mol²/L² at 25°C. In pure water [H⁺] = [OH⁻] = 10⁻⁷ M.
pH = −log [H⁺]. Acids have pH below 7, neutral solutions 7 and bases above 7. pH + pOH = 14. Example: 0.0001 M HCl has pH 4; 0.0001 M NaOH has pOH 4 and pH 10.

Chemical Kinetics

Chemical kinetics deals with the rate of a reaction, the change in concentration of a reactant or product per unit time. The rate decreases with time as reactants are used up. Rate increases with concentration, with temperature (roughly doubling for each 10°C rise), with surface area of solids and with a catalyst. The velocity constant k is the rate when each reactant has unit concentration, and K = k(forward)/k(reverse).