Reaction Rates and Equilibrium (OCR A-Level Chemistry A): Revision Notes

Dynamic equilibrium and Le Chatelier's principle

Understanding reversible reactions

Most chemical reactions you have studied so far have been irreversible reactions - they proceed in one direction only until the reactants are used up. For example, when hydrogen gas is ignited in oxygen, water forms through a reaction that goes to completion:

$2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)}$

In this reaction, all the hydrogen reacts with oxygen, the reaction finishes, and only products remain.

However, many important chemical reactions are reversible reactions - they can proceed in both forward and reverse directions simultaneously. These reactions are particularly important in industrial processes.

A classic example is the Haber process used to manufacture ammonia:

$\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)}$

The $\rightleftharpoons$ symbol indicates that this reaction is reversible and can reach equilibrium. In this system, nitrogen and hydrogen combine to form ammonia (forward reaction), while simultaneously ammonia breaks down to reform nitrogen and hydrogen (reverse reaction).

What is dynamic equilibrium?

When a reversible reaction is carried out in a closed system (isolated from its surroundings), it can reach a state called dynamic equilibrium. Understanding this concept is essential for predicting and controlling chemical reactions.

Characteristics of equilibrium systems

The term "dynamic" is important here. Although concentrations stay constant, both reactions continue to occur. Reactant molecules are constantly converting to product molecules at the same rate that product molecules convert back to reactants. This creates a balance where the overall composition doesn't change, even though molecular conversions are still happening.

Closed systems

For equilibrium to be maintained, the system must be closed - it must be isolated from its surroundings. This means temperature, pressure, and concentrations cannot be affected by external factors. Any exchange of matter or uncontrolled energy with the surroundings would disrupt the equilibrium state.

Le Chatelier's principle

When an equilibrium system is disturbed by changing conditions, it responds in a predictable way. This behavior is described by Le Chatelier's principle.

In simpler terms, if you push an equilibrium system one way, it pushes back to oppose the change. The equilibrium position shifts to counteract the disturbance and establish a new equilibrium.

This diagram shows how an equilibrium system responds when disturbed. When extra reactant is added, the system shifts to the right (toward products) to minimize the change in reactant concentration. A new equilibrium is then established with different concentrations but with forward and reverse rates again equal.

Position of equilibrium

The position of equilibrium tells us the extent of a reaction - how much of the reactants have converted to products at equilibrium. We can describe the position as:

• Shifted to the right when products are favored (more products than reactants at equilibrium)
• Shifted to the left when reactants are favored (more reactants than products at equilibrium)

Changes in temperature, pressure (for gaseous reactions), or concentration can alter the position of equilibrium. Le Chatelier's principle helps us predict which direction the equilibrium will shift.

Effect of concentration changes on equilibrium

When you change the concentration of a reactant or product in an equilibrium system, the position of equilibrium shifts to counteract that change.

General principles

Investigating the chromate-dichromate equilibrium

An excellent demonstration of concentration effects uses the equilibrium between yellow chromate ions and orange dichromate ions:

$2\text{CrO}_4^{2-}\text{(aq)} + 2\text{H}^+\text{(aq)} \rightleftharpoons \text{Cr}_2\text{O}_7^{2-}\text{(aq)} + \text{H}_2\text{O(l)}$

The different colors make it easy to see when the equilibrium position changes.

Effect of temperature changes on equilibrium

Temperature changes affect equilibrium position differently than concentration changes because temperature affects the rates of both forward and reverse reactions. The direction of the equilibrium shift depends on whether the reaction is exothermic or endothermic.

Understanding enthalpy changes in reversible reactions

In any reversible reaction, the forward and reverse reactions have enthalpy changes ($\Delta H$) that are equal in magnitude but opposite in sign:

• If the forward reaction is exothermic ($\Delta H$ is negative), the reverse reaction is endothermic ($\Delta H$ is positive)
• If the forward reaction is endothermic ($\Delta H$ is positive), the reverse reaction is exothermic ($\Delta H$ is negative)

General principles for temperature effects

Investigating the cobalt chloride equilibrium

The equilibrium between two cobalt complexes provides a vivid demonstration of temperature effects. One complex is pink and the other is blue:

$[\text{Co(H}_2\text{O)}_6]^{2+}\text{(aq)} + 4\text{Cl}^-\text{(aq)} \rightleftharpoons \text{CoCl}_4^{2-}\text{(aq)} + 6\text{H}_2\text{O(l)}$

The forward reaction (pink to blue) is endothermic ($\Delta H$ is positive), meaning it absorbs heat energy. The reverse reaction (blue to pink) is exothermic ($\Delta H$ is negative), meaning it releases heat energy.

Summary of temperature effects

The table below summarizes how temperature changes affect equilibrium position for both exothermic and endothermic reactions:

Effect of pressure changes on equilibrium

Pressure changes only affect equilibrium systems that contain gases. Moreover, the equilibrium only shifts if there are different numbers of gaseous molecules on each side of the equation.

Understanding pressure and gas molecules

The pressure of a gas is directly proportional to its concentration. In the same container, having more moles of gas means higher concentration and higher pressure. Therefore, two moles of a gas would have twice the pressure of one mole of the same gas in the same container.

General principles for pressure effects

The nitrogen dioxide-dinitrogen tetroxide equilibrium

Brown nitrogen dioxide gas ($\text{NO}_2$) exists in equilibrium with colorless dinitrogen tetroxide ($\text{N}_2\text{O}_4$):

$2\text{NO}_2\text{(g)} \rightleftharpoons \text{N}_2\text{O}_4\text{(g)}$

Note that there are:

• 2 moles of gas on the left (reactant side)
• 1 mole of gas on the right (product side)

Effect of catalysts on equilibrium

A catalyst is a substance that speeds up the rate of a chemical reaction without being used up in the process. For equilibrium systems, catalysts have a special characteristic:

How catalysts work in equilibrium systems

A catalyst speeds up both the forward and reverse reactions equally. This means:

• The same equilibrium position is reached
• The same final concentrations are obtained
• The same ratio of products to reactants exists at equilibrium

The only difference is that equilibrium is reached more quickly.

Why catalysts are useful

Even though catalysts don't change the position of equilibrium, they are extremely valuable in industrial processes. Reaching equilibrium faster means:

• Products are formed more quickly
• Manufacturing time is reduced
• Production rates increase
• Economic efficiency improves

Industrial application: the Haber process

The Haber process demonstrates how Le Chatelier's principle is applied in real industrial chemistry to manufacture ammonia efficiently and economically.

The Haber process equation

$\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)} \qquad \Delta H = -92 \text{ kJ mol}^{-1}$

This reaction is exothermic in the forward direction - it releases energy when ammonia forms.

Applying Le Chatelier's principle

Using Le Chatelier's principle, we can predict the best conditions to maximize ammonia production:

Temperature considerations:

Since the forward reaction is exothermic, a low temperature would shift the equilibrium to the right (toward products), producing more ammonia. However, if the temperature is too low, the reaction rate becomes extremely slow, and equilibrium may never be reached.

Pressure considerations:

Looking at the equation, there are:

• 4 moles of gas on the left (1 N₂ + 3 H₂)
• 2 moles of gas on the right (2 NH₃)

A high pressure would shift the equilibrium to the right (toward fewer molecules), producing more ammonia. High pressure also increases concentration and reaction rate.

Catalyst use:

An iron catalyst is used to speed up the reaction so that lower temperatures can be used. This reduces operating costs while still achieving economically viable production rates.

Recycling unreacted materials

Key points about the Haber process

Remember!