# Calculating an Equilibrium Constant, Kp, with Partial Pressures

$$K_p$$ is the equilibrium constant calculated from the partial pressures of a reaction equation. It is used to express the relationship between product pressures and reactant pressures. It is a unitless number, although it relates the pressures.

## Introduction

In calculating Kp, the partial pressures of gases are used. The partial pressures of pure solids and liquids are not included. To use this equation, it is beneficial to have an understanding of partial pressures and mole fractions.

Partial Pressures: All of the partial pressures add up to the total pressure, as shown in the equation (Dalton's law)

$P_{total} = P_A + P_B + ... \tag{1}$

Individual gases maintain their respective pressures when combined. However, their individual pressures add up to the total pressure in the system

Mole Fractions: Although mole fractions can be used for non-gases, this description is referring specifically to those involving gases. This fraction expresses the moles of an individual gas compared to the total moles. It is used for reactions in which more than one gas is involved. In general the equation for finding the mole fraction of a gas (represented by A) is

$x_A = \dfrac{mol_A}{mol_{total}}\tag{2}$

Note

If the mole fraction is multiplied by 100, it becomes a percent expressing the mole amount for the individual gas.

When combined with the total pressure, this will equal the partial pressure (for gas A). The equation is thus

$P_A = X_A \times P_{total}\tag{3}$

## Calculations

Partial pressures are used for calculating the pressure constant. Consider this general equation for a gas-phase reaction:

$aA + bB \rightleftharpoons cC + dD \tag{4}$

The formula uses the partial pressures, PA, PB, PC, and PD, raised to exponents equal to their respective coefficients in the chemical equation. Kp is calculated as follows:

$K_p = \dfrac{P_C^cP_D^d}{P_A^aP_B^b} \tag{5}$

Note

This expression is similar to $$K_c$$; however, $$K_c$$ is calculated from molar concentrations. To prevent confusion, do not use brackets around partial pressures.

$$K_p$$ can also be obtained from $$K_c$$ (the concentration equilibrium constant), temperature, change in moles, and the gas constant, as shown:

$K_p = K_c(RT)^{\Delta{n}} \tag{6}$