Changing the Position of Equilibrium

A reversible reaction attains equilibrium when the opposing reactions occur at an equal rate. Therefore, if a reaction at equilibrium is subjected to some change it changes the rate of one reaction more than the other. By changing the concentration, temperature or pressure of any of the reactants or products, the position of equilibrium can be changed. The effect of these changes can be summarised by Le-Chatelier's Principle.

Le-Chatelier's Principle

If a system at equilibrium is subjected to a change in concentration, temperature or pressure, the equilibrium readjusts itself to counteract the effect of the applied change.

Equilibrium Constant and Extent of a Chemical Reaction

Consider a reversible reaction expressed by the following general equation:

where A and B are two substances which react to form new substances C and D and C and D react under the same conditions forming the original substances A and B.

Let us assume that the concentrations of these four substances respectively are [A], [B], [C] and [D] (square bracket denotes molar concentration) and the temperature remains constant throughout.

The equilibrium constant K for this equation is:

This expression is also known as Law of Chemical Equilibrium.

At a particular temperature, equilibrium constant has a definite value. When we express concentration in molL^-1, then the equilibrium constant is denoted by K_c.

Thus for a reversible reaction at equilibrium, the multiplication product of the concentrations of the products divided by the multiplication product of the concentrations of the reactants, each concentration term raised to the power which is the coefficient of the substance in the balanced chemical equation, is a constant for a given temperature.

Significance of Equilibrium Constant

The magnitude of equilibrium constant K is a measure of the extent to which the reversible reaction has proceeded at equilibrium. The larger the value of the equilibrium constant, the greater is the extent to which the forward reaction takes place and therefore the greater the concentration of the products at equilibrium. A small equilibrium constant would likewise indicate that the forward reaction takes place only a small extent and therefore only a small amount of the products is being formed at equilibrium.

Knowing the value of the equilibrium constant and the initial concentrations of the reacting substances, the concentrations of various substances at equilibrium can be easily calculated.

For the reaction at 298 K,

Very small value of K_c, implies that reactants N₂ and O₂ will be dominant species in the system at equilibrium.

For the reaction at 298 K,

Large value of K, indicates that at equilibrium, a major portion of nitrogen and hydrogen (mixed in the molar ratio 1:3) will be converted into ammonia.

Industrially ammonia is manufactured at a high pressure of 152000 to 760000 mm of Hg and at a higher temperature of 723 K as the reaction is very slow at room temperature.

The rate of a reaction increases by raising the temperature of the reactants. The value of equilibrium constant also depends on temperature.

Effect of Change of Concentration

Increasing the concentration of one of the reactants or decreasing the concentration of one of the products shifts the equilibrium in favor of the products. On the other hand, decreasing the concentration of one of the reactants or increasing the concentration of one of the products shifts the equilibrium in favor of reactants.

The effect of concentration on equilibrium state can be demonstrated by studying the following equilibrium, which exists in a solution of bromine in water.

Effect of Concentration on Equilibrium State

When sodium hydroxide solution is added to the above system at equilibrium, the solution becomes colorless indicating that the equilibrium has shifted to the right. Addition of OH^- ions removes H⁺ ions to form water.

In order to undo the effect of the change, equilibrium shifts in the forward direction producing colorless solution.

Addition of dilute sulphuric acid solution restores the reddish brown color indicating that the equilibrium has shifted to the left. The acid increases the concentration of H⁺ ions and shifts the equilibrium in backward direction. Due to the formation of bromine, reddish brown color is restored.

Effect of Temperature

A chemical equilibrium involves two opposing reactions. If one of them is exothermic, the other must be endothermic. If a system at equilibrium is subjected to change in temperature, the equilibrium adjusts itself to counteract the effect of the applied change. Thus, if the temperature is increased, the equilibrium shifts in favor of the process that absorbs heat i.e., the endothermic process. Similarly, a decrease in temperature favors exothermic process.

Effect of temperature on equilibrium position can be demonstrated by studying the following equilibrium:

Nitrogen dioxide (NO₂) gas is produced by heating conc. nitric acid (HNO₃) with copper chips and the gas evolved is collected in a test tube. The mouth of test tube is closed with a cork. The reddish gas in the test tube is actually an equilibrium mixture of nitrogen dioxide and dinitrogen tetroxide (N₂O₄). When the tube is placed in a beaker containing ice it is observed that the color of the gas mixture lightens indicating that the equilibrium has shifted towards left (exothermic direction). When the tube is placed in a beaker containing hot water, the gas mixture becomes darker in color, indicating that the equilibrium has shifted towards right (endothermic direction).

Effect of Temperature on Equilibrium

Effect of Pressure

The change of pressure has effect only on those equilibria which involve gaseous substances and proceed with a change in the number of moles of the gases. According to Le-Chatelier's principle, increase of external pressure should affect the equilibrium in such a way as to reduce the pressure. This implies that the equilibrium shifts in the direction which has smaller number of moles of the gaseous substances. This can be easily understood from the following equilibrium representing the dissociation of N₂O₄:

1 mole              2 mole

If the pressure is increased in this reaction at equilibrium, the equilibrium shifts towards left (because it has less number of moles of gases) resulting in the formation of more N₂O₄.