Equilibria

In many chemical reactions, reactants turn into products, and that is the end of the story. The reaction goes to completion. However, in many other chemical processes, the products can react together to reform the original reactants. This is where we enter the fascinating world of chemical equilibria.

Understanding how to control these reactions is vital for chemical industries, allowing us to maximise the yield of useful products like ammonia or sulfuric acid.

What you'll learn:

What a reversible reaction is and how to represent it.
The conditions required to establish a dynamic equilibrium.
Higher Tier: How to use Le Chatelier's Principle to predict how changing temperature, pressure, or concentration shifts the position of equilibrium.

Reversible Reactions

Most everyday reactions you have met so far, like burning magnesium or rusting iron, only go in one direction. These are irreversible. However, some reactions can go both forwards and backwards.

Definition

Reversible Reaction

A reversible reaction is a reaction in which the products can react together to reform the original reactants.

We represent a reversible reaction using a special double arrow: $⇌\rightleftharpoons$ .

For a general reversible reaction:

\text{Reactant A} + \text{Reactant B} \rightleftharpoons \text{Product C} + \text{Product D}

The reaction reading from left to right is the forward reaction.
The reaction reading from right to left is the reverse (or backward) reaction.

Reversing Reactions by Changing Conditions

You can often force a reversible reaction to go in a particular direction by altering the reaction conditions, such as temperature.

A classic example is the thermal decomposition of ammonium chloride:

\text{ammonium chloride (s)} \rightleftharpoons \text{ammonia (g)} + \text{hydrogen chloride (g)}

\text{NH}_4\text{Cl(s)} \rightleftharpoons \text{NH}_3\text{(g)} + \text{HCl(g)}

Heating solid ammonium chloride causes it to decompose into ammonia gas and hydrogen chloride gas (the forward reaction).
Cooling the gases causes them to react together again to reform the solid ammonium chloride (the reverse reaction).

Another standard example is hydrated copper(II) sulfate:

\text{hydrated copper(II) sulfate (s)} \rightleftharpoons \text{anhydrous copper(II) sulfate (s)} + \text{water (l)}

\text{CuSO}_4\cdot5\text{H}_2\text{O(s)} \rightleftharpoons \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(l)}

Heating blue hydrated copper(II) sulfate crystals drives off the water, leaving white anhydrous copper(II) sulfate.
Adding water to the white anhydrous copper(II) sulfate releases heat and turns the solid back to blue.

Dynamic Equilibrium

If we carry out a reversible reaction in a closed vessel, something special happens. At first, only the forward reaction takes place because there are no products. As the reactants are used up, the forward reaction slows down.

At the same time, as more products are made, they begin to react to form the reactants again. The reverse reaction speeds up. Eventually, the rate of the forward reaction becomes exactly equal to the rate of the reverse reaction. This state is called dynamic equilibrium.

Definition

Dynamic Equilibrium

Dynamic equilibrium is the state reached in a closed system when the rate of the forward reaction is exactly equal to the rate of the reverse reaction.

There are two vital requirements for dynamic equilibrium to exist:

It must occur in a closed system: A closed system is one where no reactants or products can escape, and no external substances can enter. If a gas escapes, the reverse reaction cannot happen.
It is "dynamic": The reaction has not stopped! Both the forward and backward reactions are happening at the exact same speed.

Rates of reaction over time establishing equilibrium

Common Mistake

Equal vs. Constant Concentrations

A very common exam error is stating that the concentrations of reactants and products are equal at equilibrium. This is rarely true! At equilibrium, the concentrations of reactants and products remain constant (unchanging), but they do not have to be equal to each other.

Analogy

The Two-Way Escalator

Imagine you are walking up a "down" escalator. If you walk up at the exact same speed that the escalator is moving down, you will stay in the exact same spot. To an observer, you look stationary (your position is constant), but you are still walking and the escalator is still moving (the process is dynamic).

Key Idea

The Balance of Rates

At dynamic equilibrium:

Rate of forward reaction = Rate of reverse reaction.
The concentrations of reactants and products stay constant.

Altering the Equilibrium (Higher Tier Only)

In a closed system, the relative amounts of reactants and products at equilibrium depend on the conditions of the reaction. We refer to this balance as the position of equilibrium.

If the position of equilibrium "lies to the right", it means there is a higher concentration of products than reactants.
If the position of equilibrium "lies to the left", it means there is a higher concentration of reactants than products.

To predict how the position of equilibrium will shift when we change the conditions, we use Le Chatelier’s Principle.

Definition

Le Chatelier's Principle

Le Chatelier's Principle states that if a system at dynamic equilibrium is subjected to a change in conditions, the position of equilibrium will shift to counteract and oppose that change.

1. Changing Temperature

To predict the effect of temperature, you must look at the enthalpy change ( $ΔH\Delta H$ ) of the reaction.

Every reversible reaction is exothermic (gives out heat) in one direction, and endothermic (takes in heat) in the opposite direction.
The energy values are equal and opposite. For example, if the forward reaction releases $kJ/mol92\text{ kJ/mol}$ of energy ( $kJ/mol\Delta H = -92\text{ kJ/mol}$ ), the reverse reaction must absorb $kJ/mol92\text{ kJ/mol}$ ( $kJ/mol\Delta H = +92\text{ kJ/mol}$ ).

According to Le Chatelier's Principle:

If you increase the temperature, the system tries to cool things down. It does this by shifting the equilibrium in the endothermic direction to absorb the extra heat.
If you decrease the temperature, the system tries to heat things up. It does this by shifting the equilibrium in the exothermic direction to release more heat.

2. Changing Pressure (Gases Only)

This rule only applies to reactions involving gases. Pressure in a gas is caused by gas molecules colliding with the walls of the container. The more gas molecules there are in a given space, the higher the pressure.

According to Le Chatelier's Principle:

If you increase the pressure, the system tries to decrease the pressure. It does this by shifting the equilibrium to the side with fewer gas molecules (fewer moles of gas).
If you decrease the pressure, the system tries to increase the pressure. It does this by shifting the equilibrium to the side with more gas molecules (more moles of gas).

If both sides of the balanced equation have the same number of moles of gas, changing the pressure has no effect on the position of equilibrium.

3. Changing Concentration

According to Le Chatelier's Principle:

If you increase the concentration of a reactant, the system tries to use it up. The equilibrium shifts to the right (making more products).
If you decrease the concentration of a product (for example, by continuously removing it as it forms), the system tries to make more of it. The equilibrium shifts to the right.

Tip

Catalysts and Equilibrium

Adding a catalyst speeds up the rate of both the forward and reverse reactions by the exact same amount. Therefore, a catalyst has no effect on the position of equilibrium. It simply helps the reaction reach dynamic equilibrium faster.

Worked Example: Predicting Equilibrium Shifts

Let's apply these rules to a real industrial reaction: the Haber process, which is used to manufacture ammonia.

Example

Predicting the effect of temperature and pressure changes

The chemical equation for the synthesis of ammonia is:

\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)} \quad (\Delta H = -92\text{ kJ/mol})

Predict and explain the effect on the yield of ammonia ( $NH3\text{NH}_3$ ) when you:

Increase the temperature.
Increase the pressure.

Step-by-step solution:

Analyze the thermal properties of the reaction: Look at the sign of $ΔH\Delta H$ . Because $kJ/mol\Delta H = -92\text{ kJ/mol}$ (negative), the forward reaction is exothermic. This means the reverse reaction must be endothermic.
Apply the temperature change rule: An increase in temperature causes the system to try to lower the temperature. It will shift the equilibrium position in the endothermic direction (to the left). Therefore, the yield of ammonia ( $NH3\text{NH}_3$ ) will decrease.
Count the moles of gas on each side of the equation: On the reactant (left) side, there is $N21\text{ mole of }\text{N}_2$ and $H23\text{ moles of }\text{H}_2$ , making a total of $moles4\text{ moles}$ of gas. On the product (right) side, there are $moles2\text{ moles}$ of gas ( $NH3\text{NH}_3$ ).
Apply the pressure change rule: An increase in pressure causes the system to try to lower the pressure. It does this by shifting to the side with fewer moles of gas. Since there are fewer moles on the right ( $moles2\text{ moles}$ compared to $moles4\text{ moles}$ ), the equilibrium shifts to the right. Therefore, the yield of ammonia ( $NH3\text{NH}_3$ ) will increase.

Exam technique

In the exam

Always check the state symbols: Pressure changes only affect reactants and products that are in the gaseous state $(g)\text{(g)}$ . If a substance is solid $(s)\text{(s)}$ or liquid $(l)\text{(l)}$ , ignore it when counting moles of gas for pressure shifts!
Identify the exothermic direction first: Before answering any question about temperature, write down which direction is exothermic and which is endothermic. It prevents silly mix-ups under pressure.
Distinguish between rate and yield: A high temperature speeds up a reaction (increases rate) but can sometimes decrease the amount of product made at equilibrium (decreases yield). Make sure you read the question carefully to see which one is being asked.

Self review

Check yourself

State the two conditions required for a chemical system to reach dynamic equilibrium.
Explain why a catalyst does not change the position of dynamic equilibrium.
Higher Tier: For the reversible reaction $H2(g)+I2(g)⇌2HI(g)\text{H}_2\text{(g)} + \text{I}_2\text{(g)} \rightleftharpoons 2\text{HI(g)}$ , explain why changing the overall pressure has no effect on the position of equilibrium.

Equilibria

Understanding how to control these reactions is vital for chemical industries, allowing us to maximise the yield of useful products like ammonia or sulfuric acid.

What you'll learn:

What a reversible reaction is and how to represent it.
The conditions required to establish a dynamic equilibrium.
Higher Tier: How to use Le Chatelier's Principle to predict how changing temperature, pressure, or concentration shifts the position of equilibrium.

Reversible Reactions

Most everyday reactions you have met so far, like burning magnesium or rusting iron, only go in one direction. These are irreversible. However, some reactions can go both forwards and backwards.

Definition

Reversible Reaction

A reversible reaction is a reaction in which the products can react together to reform the original reactants.

We represent a reversible reaction using a special double arrow: $⇌\rightleftharpoons$ .

For a general reversible reaction:

\text{Reactant A} + \text{Reactant B} \rightleftharpoons \text{Product C} + \text{Product D}

The reaction reading from left to right is the forward reaction.
The reaction reading from right to left is the reverse (or backward) reaction.

Reversing Reactions by Changing Conditions

You can often force a reversible reaction to go in a particular direction by altering the reaction conditions, such as temperature.

A classic example is the thermal decomposition of ammonium chloride:

\text{ammonium chloride (s)} \rightleftharpoons \text{ammonia (g)} + \text{hydrogen chloride (g)}

\text{NH}_4\text{Cl(s)} \rightleftharpoons \text{NH}_3\text{(g)} + \text{HCl(g)}

Heating solid ammonium chloride causes it to decompose into ammonia gas and hydrogen chloride gas (the forward reaction).
Cooling the gases causes them to react together again to reform the solid ammonium chloride (the reverse reaction).

Another standard example is hydrated copper(II) sulfate:

\text{hydrated copper(II) sulfate (s)} \rightleftharpoons \text{anhydrous copper(II) sulfate (s)} + \text{water (l)}

\text{CuSO}_4\cdot5\text{H}_2\text{O(s)} \rightleftharpoons \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(l)}

Heating blue hydrated copper(II) sulfate crystals drives off the water, leaving white anhydrous copper(II) sulfate.
Adding water to the white anhydrous copper(II) sulfate releases heat and turns the solid back to blue.

Dynamic Equilibrium

Definition

Dynamic Equilibrium

Dynamic equilibrium is the state reached in a closed system when the rate of the forward reaction is exactly equal to the rate of the reverse reaction.

There are two vital requirements for dynamic equilibrium to exist:

It must occur in a closed system: A closed system is one where no reactants or products can escape, and no external substances can enter. If a gas escapes, the reverse reaction cannot happen.
It is "dynamic": The reaction has not stopped! Both the forward and backward reactions are happening at the exact same speed.

Rates of reaction over time establishing equilibrium

Common Mistake

Equal vs. Constant Concentrations

Analogy

The Two-Way Escalator

Key Idea

The Balance of Rates

At dynamic equilibrium:

Rate of forward reaction = Rate of reverse reaction.
The concentrations of reactants and products stay constant.

Altering the Equilibrium (Higher Tier Only)

In a closed system, the relative amounts of reactants and products at equilibrium depend on the conditions of the reaction. We refer to this balance as the position of equilibrium.

If the position of equilibrium "lies to the right", it means there is a higher concentration of products than reactants.
If the position of equilibrium "lies to the left", it means there is a higher concentration of reactants than products.

To predict how the position of equilibrium will shift when we change the conditions, we use Le Chatelier’s Principle.

Definition

Le Chatelier's Principle

Le Chatelier's Principle states that if a system at dynamic equilibrium is subjected to a change in conditions, the position of equilibrium will shift to counteract and oppose that change.

1. Changing Temperature

To predict the effect of temperature, you must look at the enthalpy change ( $ΔH\Delta H$ ) of the reaction.

Every reversible reaction is exothermic (gives out heat) in one direction, and endothermic (takes in heat) in the opposite direction.
The energy values are equal and opposite. For example, if the forward reaction releases $kJ/mol92\text{ kJ/mol}$ of energy ( $kJ/mol\Delta H = -92\text{ kJ/mol}$ ), the reverse reaction must absorb $kJ/mol92\text{ kJ/mol}$ ( $kJ/mol\Delta H = +92\text{ kJ/mol}$ ).

According to Le Chatelier's Principle:

If you increase the temperature, the system tries to cool things down. It does this by shifting the equilibrium in the endothermic direction to absorb the extra heat.
If you decrease the temperature, the system tries to heat things up. It does this by shifting the equilibrium in the exothermic direction to release more heat.

2. Changing Pressure (Gases Only)

According to Le Chatelier's Principle:

If you increase the pressure, the system tries to decrease the pressure. It does this by shifting the equilibrium to the side with fewer gas molecules (fewer moles of gas).
If you decrease the pressure, the system tries to increase the pressure. It does this by shifting the equilibrium to the side with more gas molecules (more moles of gas).

If both sides of the balanced equation have the same number of moles of gas, changing the pressure has no effect on the position of equilibrium.

3. Changing Concentration

According to Le Chatelier's Principle:

If you increase the concentration of a reactant, the system tries to use it up. The equilibrium shifts to the right (making more products).
If you decrease the concentration of a product (for example, by continuously removing it as it forms), the system tries to make more of it. The equilibrium shifts to the right.

Tip

Catalysts and Equilibrium

Worked Example: Predicting Equilibrium Shifts

Let's apply these rules to a real industrial reaction: the Haber process, which is used to manufacture ammonia.

Example

Predicting the effect of temperature and pressure changes

The chemical equation for the synthesis of ammonia is:

\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)} \quad (\Delta H = -92\text{ kJ/mol})

Predict and explain the effect on the yield of ammonia ( $NH3\text{NH}_3$ ) when you:

Increase the temperature.
Increase the pressure.

Step-by-step solution:

Analyze the thermal properties of the reaction: Look at the sign of $ΔH\Delta H$ . Because $kJ/mol\Delta H = -92\text{ kJ/mol}$ (negative), the forward reaction is exothermic. This means the reverse reaction must be endothermic.
Apply the temperature change rule: An increase in temperature causes the system to try to lower the temperature. It will shift the equilibrium position in the endothermic direction (to the left). Therefore, the yield of ammonia ( $NH3\text{NH}_3$ ) will decrease.
Count the moles of gas on each side of the equation: On the reactant (left) side, there is $N21\text{ mole of }\text{N}_2$ and $H23\text{ moles of }\text{H}_2$ , making a total of $moles4\text{ moles}$ of gas. On the product (right) side, there are $moles2\text{ moles}$ of gas ( $NH3\text{NH}_3$ ).
Apply the pressure change rule: An increase in pressure causes the system to try to lower the pressure. It does this by shifting to the side with fewer moles of gas. Since there are fewer moles on the right ( $moles2\text{ moles}$ compared to $moles4\text{ moles}$ ), the equilibrium shifts to the right. Therefore, the yield of ammonia ( $NH3\text{NH}_3$ ) will increase.

Exam technique

In the exam

Always check the state symbols: Pressure changes only affect reactants and products that are in the gaseous state $(g)\text{(g)}$ . If a substance is solid $(s)\text{(s)}$ or liquid $(l)\text{(l)}$ , ignore it when counting moles of gas for pressure shifts!
Identify the exothermic direction first: Before answering any question about temperature, write down which direction is exothermic and which is endothermic. It prevents silly mix-ups under pressure.
Distinguish between rate and yield: A high temperature speeds up a reaction (increases rate) but can sometimes decrease the amount of product made at equilibrium (decreases yield). Make sure you read the question carefully to see which one is being asked.

Self review

Check yourself

State the two conditions required for a chemical system to reach dynamic equilibrium.
Explain why a catalyst does not change the position of dynamic equilibrium.
Higher Tier: For the reversible reaction $H2(g)+I2(g)⇌2HI(g)\text{H}_2\text{(g)} + \text{I}_2\text{(g)} \rightleftharpoons 2\text{HI(g)}$ , explain why changing the overall pressure has no effect on the position of equilibrium.

Equilibria

What you'll learn:

Reversible Reactions

Reversing Reactions by Changing Conditions

Dynamic Equilibrium

Altering the Equilibrium (Higher Tier Only)

1. Changing Temperature

2. Changing Pressure (Gases Only)

3. Changing Concentration

Worked Example: Predicting Equilibrium Shifts

Recap questions

You've reached the end

Equilibria

What you'll learn:

Reversible Reactions

Reversing Reactions by Changing Conditions

Dynamic Equilibrium

Altering the Equilibrium (Higher Tier Only)

1. Changing Temperature

2. Changing Pressure (Gases Only)

3. Changing Concentration

Worked Example: Predicting Equilibrium Shifts

Recap questions

You've reached the end

Equilibria

What you'll learn:

Reversible Reactions

Reversing Reactions by Changing Conditions

Dynamic Equilibrium

Altering the Equilibrium (Higher Tier Only)

1. Changing Temperature

2. Changing Pressure (Gases Only)

3. Changing Concentration

Worked Example: Predicting Equilibrium Shifts

Recap questions

You've reached the end

Equilibria Revision Guide

Equilibria

What you'll learn:

Reversible Reactions

Reversing Reactions by Changing Conditions

Dynamic Equilibrium

Altering the Equilibrium (Higher Tier Only)

1. Changing Temperature

2. Changing Pressure (Gases Only)

3. Changing Concentration

Worked Example: Predicting Equilibrium Shifts

Recap questions

You've reached the end

Equilibria Revision Guide