Understanding the Concept of Equilibrium

Equilibrium represents a state where the rate of the forward reaction equals the rate of the reverse reaction.
An equilibrium system can be identified by the use of a double arrow in the chemical equation.
The concentrations of the reactants and products remain constant at equilibrium (though they may not be equal).
Physical changes can also reach equilibrium- for instance, the rate of evaporation equals the rate of condensation in a closed system.
Equilibrium does not necessarily mean equal quantities of reactants and products; it represents equal rates of reaction.

Le Chatelier’s Principle

Le Chatelier’s principle is used to predict how a change in conditions affects the position of equilibrium.
An increase in temperature will shift the equilibrium to favour the endothermic reaction.
A decrease in temperature favours the exothermic reaction.
Increasing the pressure will shift the equilibrium to favour the side with fewer gas molecules.
Increasing the concentration of a reactant causes the equilibrium to shift to use up the added reactant.

The equilibrium constant Kc is calculated for a reaction at a particular temperature.
Kc does not change unless the temperature changes.
Kc is the ratio of the concentration of the products to the concentration of the reactants. Each concentration has an exponent that corresponds to the number of moles in the balanced chemical equation.
A larger Kc value (greater than 1) indicates that the products are favoured at equilibrium.
A smaller Kc value (less than 1) suggests that the reactants are favoured at equilibrium.

Catalysts speed up the rates of both the forward and reverse reactions.
They decrease the activation energy needed for the reactions.
Catalysts don’t change the position of the equilibrium, but they allow the system to reach equilibrium faster.