Chemical Equilibruim – SS2 Chemistry Lesson Note (Second Term)

Equilibrium can be defined as the state of a system where there is no observable or detectable change in the properties of the system with time. Examples of systems in equilibrium are:

1. a balanced see-saw
2. a saturated solution of NaCl

Behavioural objectives

By the end of the lesson, students should be able to:

• distinguish between static and dynamic equilibrium;
• state the properties of a system in equilibrium;
• state Le Chatelier’s principle and use it to predict the effect of temperature, concentration, pressure and a catalyst on the position of equilibrium;
• explain the conditions used in the Haber process;
• state the law of mass action and write expressions for Kc and Kp.

What are static and dynamic equilibrium?

Static equilibrium is attained when a system in equilibrium is in a state of rest or stationary state. An example of this is a balanced see-saw.

Dynamic equilibrium can occur during a physical change or a reversible chemical change. Dynamic equilibrium involving a physical change is called physical equilibrium, while that involving a chemical change is called chemical equilibrium.

A system is in a state of dynamic equilibrium when the forward and backward reactions take place at the same rate. Example:

N₂O₄(g) ⇌ 2NO₂(g)

Equilibrium in a reversible reaction

A reversible reaction proceeds in both forward and backward directions under suitable conditions. For example:

NH₄Cl(s) ⇌ NH₃(g) + HCl(g)

A reversible reaction is in dynamic equilibrium when the forward and backward reactions occur at the same rate, so that there is no net change in the concentrations of the reactants or products.

What are the properties of a system in equilibrium?

1. A chemical reaction at equilibrium is in a dynamic state, with the rates of the forward and backward reactions equal.
2. The equilibrium position can be reached from the reactants alone or from the products alone.
3. For equilibrium to be established, the system must be closed.

What factors affect the position of equilibrium?

The position of equilibrium depends on the following conditions:

1. Temperature
2. Concentration
3. Pressure (for gases)

A change in any of these factors upsets the balance of the system and shifts the position of the equilibrium. Le Chatelier studied these factors and their effect on chemical systems in equilibrium, and stated Le Chatelier’s principle.

Le Chatelier’s principle states that if a chemical system is in equilibrium and one of the factors involved in the equilibrium is altered, the equilibrium position will shift to neutralise the effect of the change.

Importance of Le Chatelier’s principle in the chemical industry

It helps to:

• define optimum conditions for a chemical process
• reduce undesirable reversibility
• predict the effect of altered factors on the equilibrium

Effect of temperature

For an endothermic reaction, an increase in temperature shifts the position of equilibrium to the right. That is, the forward reaction is favoured, while a decrease in temperature favours the backward reaction. The reverse is the case for exothermic reactions.

Effect of concentration

In a system at equilibrium, if the concentration of the reactants is increased or that of the products is decreased, the forward reaction is favoured. That is, the equilibrium position shifts to the right. If the concentration of the products is increased or that of the reactants is decreased, the backward reaction is favoured and the equilibrium position shifts to the left.

Effect of pressure

The pressure of a gas is proportional to the number of moles. An increase in the number of moles implies an increase in pressure and vice versa. If the forward reaction proceeds with a decrease in the number of moles (pressure), an increase in pressure favours the forward reaction. The reverse is the case if the forward reaction proceeds with an increase in the number of moles (pressure).

In other words, when high pressure is applied, the equilibrium shifts to the side with fewer moles. When low pressure is applied, the equilibrium shifts to the side with more moles.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

1 mole + 3 moles ⇌ 2 moles

An increase in pressure favours the forward reaction, while a decrease in pressure favours the backward reaction.

Effect of a catalyst

A catalyst does not change the position of equilibrium. It increases the rate of both the forward and backward reactions by lowering the activation energy of the reaction. The equilibrium position is therefore attained faster when a catalyst is present.

What is the Haber process?

The Haber process is an industrial process for the manufacture of ammonia. Ammonia, NH₃, is an important chemical for the manufacture of fertilisers. During the process, the following equilibrium is established:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -96 kJ mol^-1

1 mole + 3 moles ⇌ 2 moles

Choice of temperature

The forward reaction is exothermic, so Le Chatelier’s principle shows that the greatest yield of ammonia would be obtained at low temperatures. However, if the temperature is too low, the time taken to reach equilibrium is very long. A temperature of 500 °C gives a reasonable yield of ammonia.

Choice of pressure

The forward reaction leads to a decrease in the number of moles (pressure), so Le Chatelier’s principle shows that a high yield of ammonia would be obtained at high pressure. A pressure of 25 atm gives a satisfactory yield.

Choice of catalyst

Iron is used as a catalyst in the Haber process. It is effective in speeding up the rate at which equilibrium is reached. It is used in the form of small pellets, which increase the surface area available for contact.

What are the equilibrium constant (K) and the law of mass action?

The law of mass action states that, at constant temperature, the rate of a reaction is proportional to the active masses of each of the reactants.

The active mass of a reactant is the concentration of the substance raised to the power of its coefficient (number of moles) in the equation of the reaction under study.

Note: Active mass = molar concentration × activity coefficient. In most reactions, the active masses of the reactants may be taken as their molar concentrations.

Example: aA + bB → products

Thus, r ∝ [A]^a and r ∝ [B]^b, where [ ] represents the concentration in mol dm^-3.

r ∝ [A]^a × [B]^b

r = k[A]^a[B]^b, where k is called the rate constant.

Equilibrium constant, Kc

Applying the law to the reversible reaction represented by the equation below:

aA(aq) + bB(aq) ⇌ cC(aq) + dD(aq)

Kc = ([C]^c [D]^d) / ([A]^a [B]^b)

where Kc is called the equilibrium constant.

Equilibrium constant, Kp

For reactions involving gases, the equilibrium constant is expressed in terms of the partial pressures of the gases involved in the reaction. For the general gaseous reaction:

aA(g) + bB(g) ⇌ cC(g) + dD(g)

Kp = (P_C^c × P_D^d) / (P_A^a × P_B^b)

where Kp is the equilibrium constant for the gaseous reaction, and P_A, P_B, P_C and P_D are the partial pressures of gases A, B, C and D respectively.

What does the value of K tell us?

• For a given reaction, the concentrations of solids are ignored.
• A Kc or Kp value greater than one shows that product formation is favoured at equilibrium.
• A Kc or Kp value less than one shows that reactant formation is favoured at equilibrium.
• A Kc or Kp value equal to one shows that neither reactants nor products are favoured.

Evaluation

1. Define equilibrium and state the difference between static and dynamic equilibrium.
2. State Le Chatelier’s principle.
3. For the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), state the effect of increasing the pressure on the position of equilibrium.
4. What is the effect of a catalyst on the position of equilibrium?
5. Write the expression for Kc for the reaction aA(aq) + bB(aq) ⇌ cC(aq) + dD(aq).

Assignment

1. State three properties of a system in equilibrium.
2. Explain why a temperature of 500 °C is used in the Haber process instead of a very low temperature.
3. State the law of mass action and write the expression for Kp for the reaction aA(g) + bB(g) ⇌ cC(g) + dD(g).