Chemical equilibrium is a state when the quantities of reactants and products in a reaction remain the same. In other words, the forward and reverse reaction are occurring simultaneously. This form of equilibrium can also be regarded as ‘dynamic equilibrium’ as forward and backward reactions are constantly occurring; the reaction does not stop. Equilibrium reactions are always shown by double headed arrows.

Equilibrium Law (Equilibrium Constant)

The equilibrium law states that a system in equilibria follows a formula involving the concentration of both reactants and products. This formula equates to the equilibrium constant which is denoted as K.

(n is the number of mole of products and m is the number of mole of reactants.)

Units for K is determined by the number of particles in the numerator and denominator.

Eg) 

The K value tells us the ‘extent’ of a reaction

• K < 10^-4; small K value which means that mixture consists of mainly reactants. Reactants in reactions with small K values are usually either a weak acid/base.
• 10^-4 < K < 10⁴; significant amounts of products and reactants
• K > 10⁴; large K value which indicates that the mixture consists of mainly products. Reactants in reactions with large K values are usually either a strong acid/base.

Non-reversible reactions

Certain reactions are not able to reach equilibrium, instead they continue their forward reaction until it’s complete. Reasons for this include:

• Products such as gases escape from the system and thus no backward reaction can be observed
• There are only a minute amount of reactant in the system.

Le Chatelier’s Principle

Le Chatelier’s principle states that if an equilibrium system is subjected to a change, the system will adjust itself to partially oppose the effect of the change.

Factors that affect the position of the equilibrium include:

Adding/Removing reactants or products

Addition of reactants leads to the formation of more products (net forward reaction).
Addition of products leads to the formation of more reactants (net backwards reaction).

By adding or removing either products or reactants, the concentration of the reactant/product will either increase or decrease vertically on a concentration vs time graph at the point of the addition/reduction. After that point, the system will try to adjust for the change so that the altered product/reactant’s concentration will approach its original concentration.

Changing pressure

Change in equilibrium by changing the pressure of a system is dependent on the mole ratio of reactants/product in a reaction’s overall balanced equation. The effects of changing pressure can be illustrated through an example of:

If pressure were to increase, the system will aim to decrease the number of molecules/moles as stated by the stoichiometric equation. The mole ratio of reactants to products in the example above is 4:2. Hence the system will favour a net forward reaction to try to reduce the number of moles.

If the mole ratio is the same for both reactants and products then a change in pressure will not affect the equilibrium.

Increasing pressure will increase the concentration of both reactants and products initially. Similarly, decreasing pressure will decrease the concentration of both.

Adding an inert gas increases pressure within a system but doesn’t shift the equilibrium because the concentration of each individual species remains unchanged.

Dilution

Dilution, better known as adding water to a system and thus decreases the initial concentration of all species. It is similar to decreasing pressure as the system will favour the reaction that will produce more particles.

Changing temperature

The effect that changing the temperature has on a system depends on the whether the reaction is exothermic or endothermic and hence the value of ΔH.

In an exothermic reaction, energy is loss and can be thought of as a product of the forward reaction. By increasing the temperature of this system, we are adding more energy and hence; more product to the system. The system will try to correct this change by favouring a net backward reaction. The same concept can be applied to an endothermic reaction.

See Also