Understanding gaseous equilibria requires describing gas mixtures quantitatively, using mole fractions and partial pressures to relate composition to equilibrium expressions and chemical behaviour accurately.

Gaseous Mixtures in Chemical Equilibria

In many equilibria studied at A-Level, reactants and products are gases contained in the same system. Unlike solids and liquids, gases mix completely and occupy the entire volume of their container. As a result, the behaviour of each gas depends not only on its own amount, but also on the total amount of gas present.

To describe such systems precisely, chemists use mole fraction and partial pressure. These concepts allow the contribution of each gas to be quantified independently, which is essential for understanding equilibrium constants such as Kp.

Mole Fraction

The mole fraction expresses how much of a mixture is made up by a particular gas, relative to the total number of moles present.

Mole fraction has no units because it is a ratio. For a mixture of gases A, B, and C:

• The mole fraction of A depends only on the number of moles of A and the total moles present.

• The sum of the mole fractions of all gases in a mixture is always equal to 1.

This reflects the idea that the entire gas mixture is fully accounted for by its individual components.

Mole fraction is particularly useful because it is independent of temperature and pressure, provided the gases behave ideally. This makes it a convenient way to describe composition in changing conditions.

Partial Pressure

Each gas in a mixture behaves as if it alone occupies the container, exerting its own pressure. This individual contribution is known as the partial pressure.

Partial pressure is measured in pascals (Pa) or kilopascals (kPa), consistent with standard pressure units used in equilibrium calculations.

In a gas mixture:

• The total pressure is the sum of the partial pressures of all gases.

• Each gas contributes to the total pressure in proportion to its mole fraction.

This schematic shows gases A, B, and C in separate containers, each exerting its own partial pressure. When combined at the same conditions, the total pressure is the sum of the individual partial pressures, illustrating Dalton’s law. Source

Relationship Between Mole Fraction and Partial Pressure

Mole fraction and partial pressure are directly linked through the total pressure of the system.

This relationship shows that:

• A gas with a larger mole fraction exerts a larger partial pressure.

• If the total pressure changes, the partial pressures change proportionally, provided mole fractions remain constant.

This equation is fundamental when working with gaseous equilibria, as equilibrium constants depend on partial pressures rather than mole fractions alone.

Three equal-volume cylinders containing gases at different pressures are combined into one cylinder. The total pressure equals the sum of the individual pressures, demonstrating how partial pressures contribute to total pressure in a gas mixture. Source

Partial Pressure and Gaseous Equilibria

In equilibria involving gases, the position of equilibrium depends on the partial pressures of reactants and products, not their individual amounts or concentrations.

For example:

• Increasing the partial pressure of a gaseous reactant tends to favour the forward reaction.

• Decreasing the partial pressure of a product can shift equilibrium towards further product formation.

These effects are explained by Le Chatelier’s principle, but the quantitative description relies on partial pressures.

Importance for Kp Expressions

The equilibrium constant Kp is defined using partial pressures of gaseous species at equilibrium.

• Each gas appears in the Kp expression raised to the power of its stoichiometric coefficient.

• Partial pressures are used instead of concentrations because gases are compressible and their behaviour depends strongly on pressure.

Understanding how partial pressures are obtained from mole fractions and total pressure is therefore essential before attempting Kp calculations.

Ideal Gas Assumption

At A-Level, gases are assumed to behave as ideal gases unless stated otherwise. Under this assumption:

• Gas particles have negligible volume.

• There are no intermolecular forces between particles.

• The relationship between mole fraction and partial pressure holds exactly.

This allows mole fraction and partial pressure concepts to be applied consistently across different equilibrium systems without correction factors.

Practical Context and Terminology

When describing gaseous equilibria, correct terminology is essential:

• Always specify partial pressure, not simply “pressure”, when referring to an individual gas.

• Use mole fraction to describe composition rather than percentage, as it links directly to equilibrium expressions.

• Remember that solids and liquids do not have partial pressures and are excluded from gas-phase equilibrium descriptions.

Accurate use of mole fraction and partial pressure terminology underpins all further work on equilibrium constants and equilibrium position in gaseous systems.