Chemistry

Le Châtelier’s Principle

According to Le Châtelier’s Principle, when the temperature, pressure or concentration of a reaction in equilibrium is changed, the reaction shifts in the direction where the effect of these changes is reduced.

15.3.1. Effect of Temperature on equilibrium

Activity 15.3(b)

1. Explain the following terms

(a)Endothermic

(b)Exothermic

(c) Suggest how temperature affects the position of equilibrium.

When dealing with temperature, we distinguish exothermic and endothermic reactions. A change in the temperature of a system already in equilibrium could either shift the equilibrium to the right (favoring the forward reaction) or to the left(favoring the backward reaction). This depends on whether the forward reaction is exothermic or endothermic. Heat can be considered a reactant in an endothermic reaction and a product in an exothermic reaction. For a reversible reaction, when the forward reaction is exothermic, the enthalpy change is negative (ΔH < 0), then the backward reaction is endothermic and the enthalpy change is positive (ΔH > 0).

For exothermic forward reactions, an increase in temperature will cause the system to counter balance it by favouring the reaction that consumes heat, hence the backward reaction will be favoured or promoted. On the contray, if the temperature is decreased, the system reacts to produce more heat by favouring the forward reaction.

In this equilibrium the forward reaction is exothermic, and the backward reaction is endothermic. If the temperature is increased, the reverse reaction will be favoured and we say the equilibrium is displaced toward the left side. If the temperature is decreased, then the forward reaction is favoured and the equilibrium is displaced toward the right side.

For endothermic forward reaction, an increase in temperature favors forward reaction, while a decrease in temperature favors reverse reaction.

In summary:

Effect of temperature on Haber process

The industrial production of ammonia is based on the reaction, shown above, between nitrogen and hydrogen gases following the Haber process.

The Haber process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic.

The reaction of the production of ammonia is exothermic. According to the Le Châtelier’s Principle, since this reaction is exothermic and if only the factor temperature is considered ,lowering the temperature would thermodynamically favors the forward reaction, i.e. the formation of ammonia. Neverthless, working at low temperature will result in a low rate of the reaction (reaction kinetics), hence low yield or low production,

On the other hand, if you look at the state of reactants and products, you notice that at the reactants side, there are 4 gaseous molecules, whereas at the products side there are 2 gaseous molecules. If the pressure is increased, the system reacts by producing more products or less number of molecules; hence increasing pressure moves the equilibrium toward the formation of more products, i.e. less number of gaseous molecules.

Taking account of all those theoretical considerations, combined with financial and economical aspects of the operations, the Haber process adopts intermediate of compromise condtions: 400-450^oC, 150-300 atm and catalyst. The Haber process yields 15% of ammonia.

15.3.2. The effect of change in concentration on equilibrium

Activity 15.3(c):

Experiment to investigate the effect of changing concentration on equilibrium

Equipment/materials

If the concentration of either Fe³⁺ or SCN^- is increased, the equilibrium shifts to right to reduce the concentration of the added reactant. Similarly, if FeSCN²⁺(aq) is added to the equilibrium, the position of the equilibrium will shift to the left to reduce the concentration of FeSCN²⁺(aq) according to Le Châtelier’s principle.

Note that if silver nitrate solution is added to the equilibrium, it supplies silver ions (Ag⁺) which removes cyanide ions (SCN^-) and causes the equilibrium to shift from right to left to replace the cyanide ions.

According to Le Châtelier’s principle, if the concentration of reactants is increased, the reaction will shift in the forward direction. Adding a reactant or product, the equilibria shifts away from the increase in order to consume the part of the added substance. By removing a reactant or product, the equilibria shifts toward the decrease to replace part of the removed species.

Consider the following equilibrium reaction:

If we add either N₂ or H₂, the collisions between N₂ and H₂ are increased and more product NH₃ are formed. It means that the equilibrium shifts from left to right. However, if more NH₃ are added, the concentration of NH₃ is increased. As a result, some NH₃ decomposes and forms more reactants. The equilibrium shifts from right to left.

Checking up 15.3.2

1.Explain Le Châtelier’s principle

2. What factors did Fritz Haber employ to determine the changing positions of equilibrium reactions?

3. Determine what would happen to the following reactions if the concentration of any one of the reactants is increased or decreased?

15.3.3. Effect of Pressure

In which direction—toward reactants or toward products—does the reaction shift if the equilibrium is stressed by each change?

(a)H₂ is added.

(b)NH₃ is added.

(c )NH₃ is removed.

According to Le Châtelier’s principle, an increase in pressure favours the reaction in the direction where the volume of reactants is reduced, or less molecules of gas are formed, and a decrease in pressure favours the reaction in the direction where the volume of reactants is increased, or more molecules of gas are formed.

Example

Consider the reaction:

Because the pressure of gases is related directly to the concentration by P = n/V, changing the pressure by increasing/decreasing the volume of a container will disturb an equilibrium system. In the above example, the volume of reactants is 4 units while the volume of products is 2 units. Therefore, according to Le Châtelier’s principle an increase in pressure of this reaction will favor the forward reaction to form fewer ammonia while a decrease in pressure of the reaction will favor the backward reaction to form more nitrogen and hydrogen.

To obtain much ammonia in the equilibrium mixture, a high pressure of 200 atmospheres is needed.

The effect of pressure can be summarized by the graph indicated below

High Pressure gives a good yield of ammonia as indicated from the graph above, at 400 atmosphers the yield of ammonia is 70%

Higher pressure increases the rate of reaction

However, the higher the pressure used, the higher the cost of the equipment needed to withstand the pressure.

The higher the pressure the higher the electrical energy costs for pumps to produce the pressure.

A moderately high pressure of between 150 – 300 atmospheres is used.

Checking up 15.3.3

Predict whether each of the following reactions are favored by high or low pressures?

15.3.4. The effect of a catalyst on equilibrium

Activity 15.3.4

1. What is an enzyme?

2. What is a catalyst? Name the catalyst used in the Haber process and contact process

3. What would happen if the enzymes involved in the digestion of food were not present?

4. Most of the metabolic processes in the body are controlled by enzymes. What would happen to these metabolic processes if the enzymes were missing?

The function of a catalyst is to speed up the reaction by lowering the activation energy. The catalyst lowers the activation energy of the forward reaction and reverse reaction to the same extent. Adding a catalyst doesn’t affect the relative rates of the two reactions and therefore the catalyst has no effect on the equilibrium system. But the catalyst helps the system to reach the equilibrium more quickly. The catalyst does not appear in the overall equation of the reaction.

Practical and financial aspects: In industry, all the above factors must be considered, taking in account not only the theoritical advantages but also their costs and risks. That is why for example the manufacture of ammonia is based on a compromise of conditions that guaranty profitability: not too high pressure between 150-300 atm. Not too low temperature between 400-450^oC, plus a catalyst.

15.4. End unit assessment

1. Consider the following equilibrium reaction:

a. What happens if we increase the concentration of N₂ by adding more N₂?

b. What happens if we increase the concentration of H₂ by adding more H₂?

c. What happens if we increase the concentration of NH₃ by adding more NH₃?

d. What happens if we decrease the concentration of N₂ by removing some N₂?

e. What happens if we decrease the concentration of H₂ by removing some H₂?

f. What happens if we decrease the concentration of NH₃ by removing some NH₃?

2. Read the passage and answer the questions that follow:

carbon dioxide and formation of Stalactites and Stalagmates in limestone caves

In addition to being a component of the atmosphere, carbon dioxide also dissolves in the water of the oceans. The dissolving process can be described by the following equations

In nature, surface water often becomes acidic because atmospheric carbon dioxide dissolves in it; this acidic solution can dissolve limestone.

Openings formed in the limestone as the calcium carbonate dissolves

Slight cooling of the water saturated with carbon dioxide can reduce the solubility of the carbon dioxide. The position of the equilibrium shifts, resulting in the precipitation of calcium carbonate. This precipitate , the statalctite is formed immediately when the seeping water comes in contact with air currents in a cave stalagmites form on the floors of the caves in the same way.

a) Suggest why the balance between CO₂(g) in the atmosphere and CO₂(aq) in the oceans cannot be regarded as a true dynamic equilibrium.

The system is not closed.

b) Based on equations (1) and (2), explain the likely effect of the increasing con-centration of atmospheric carbon dioxide on the pH of water at the ocean sur-face.

c) Use Le Châtelier’s principle to explain why slight cooling of the water saturated with carbon dioxide will result in the precipitation of calcium carbonate.

3. The hydrogen used in the Haber process is made by the following reaction:

Discuss how the yield of hydrogen in the process is affected by changing the pressure, changing the temperature and using a catalyst.

(For this question, you are required to give answers in paragraph form.)

4. Read the following passage and answer the questions that follow.

Chlorine for disinfection Chlorine is used in water treatment for disinfection.

When chlorine is added to pure water, hypochlorous acid (HOCl) and hydrochloric acid (HCl) are formed.

The principal disinfecting action of aqueous chlorine is due to hypochlorous acid formed.

Hypochlorous acid is dissociated into hydrogen ions and hypochlorite ions

The concentration of hypochlorous acid and hypochlorite ions in chlorine water depend on the pH of water. Instead of using chlorine gas, some plants apply sodium hypochlorite or calcium hypochlorite in water. Sodium hypochlorite completely dissociate in water to form sodium ions and hypochlorite ions. In solution, the hypochlorite ions hydrolyze to form the disinfectant hypochlorous acid according to the following equation.

a) State Le Châtelier’s principle.

b) Use Le Châtelier’s principle to explain how the pH of the chlorinated water will affect the concentrations of hypochlorous acid and hypochlorite ions in the water.

c) The hypochlorous acid produced in a solution of sodium hypochlorite can react further to produce small amount of chlorine according to the following equation:

What will happen to the concentration of chlorine if a little sodium hydroxide solution is added to asodium hypochlorite solution? Explain your answer.

5. In the Contact process, sulphur dioxide is catalytically oxidized to sulphur trioxide according to the equation:

The following table summarizes several possible conditions for the preparation of sulphur trioxide:

Which of the cases would represent?

(a) The theoretical conditions for obtaining the maximum yield of sulphur trioxide?Explain your answer.

b) the most economical conditions for the industrial preparation of sulphur trioxide? Explain your answer.

6. Ethanol is manufactured by catalytic hydration of ethene:

a) The reaction represented by the above equation can reach a position of dynamic equilibrium..State two features of a system that is in dynamic equilibrium.

(b)The following table lists the percentage conversion of ethene using excess steam at various reaction conditions used in industry.

i) State and explain the effect of increasing the pressure on the percentage conversion.

ii) Deduce the sign of the enthalpy change for the forward reaction. Explain your answer.

iii) The equation for the formation of ethanol shows that equal numbers of moles of ethene and steam are required. In industry however excess steam is used.

Suggest why excess steam is used.

7. The following reaction is exothermic.

In which direction does the equilibrium shift as a result of each change?

a. Removing the hydrogen gas

b. Increasing the pressure of gases in the reaction vessel by decreasing the volume

c. Increasing the pressure of gases in the reaction vessel by pumping in argon gas while keeping the volume of the vessel constant

d. Increasing the temperature

e. Using a catalyst

8. The reaction reaches dynamic equilibrium in a closed vessel. The forward reaction is exothermic. The reaction is catalyzed by V₂O₅.

a. Explain the term dynamic equilibrium.

b. What will happen to the position of equilibrium when:

i. Some sulfur trioxide, SO₃, is removed from the vessel?

ii. The pressure in the vessel is lowered?

iii. More V₂O₅ is added?

iv. The temperature of the vessel is increased?

c. State Le Châtelier’s principle.

(d) Use Le Châtelier’s principle to explain what will happen to the position of equilibrium in the reaction when the concentration of hydrogen is increased.