Kp, Partial Pressure & Changing Conditions (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching Chemistry, and this video is about Kp, partial pressure and how changing conditions affect them.

The Kp expression is built from partial pressures, and while changing the pressure shifts the position of equilibrium, it never changes the value of Kp — only temperature actually does that.

For gaseous equilibria, Kp is defined using the equilibrium partial pressures of the reactants and products, each raised to the power of its stoichiometric coefficient. Those partial pressures come from the mole fraction of each gas multiplied by the total pressure. Le Chatelier's Principle can then be used to predict how changes in temperature and pressure affect both the position of equilibrium and the value of Kp itself.

This video starts with how the Kp expression is written, then covers how partial pressure is calculated from mole fraction, and finishes with what actually changes Kp when conditions change — and what doesn't.

For a gaseous equilibrium, Kp is defined the same way as Kc but using partial pressures instead of concentrations, with each partial pressure raised to the power of its stoichiometric coefficient in the balanced equation. Solids and liquids are ignored in the Kp expression, since only gases have partial pressures. Square brackets are never used as these denote concentration. Kp is constant at a given temperature and only changes if the temperature itself changes.

The partial pressure of a gas in a mixture is the pressure it would have if it were alone in the container, and the total pressure is the sum of all the partial pressures. The mole fraction of a gas is the number of moles of that gas divided by the total number of moles of gas present, and multiplying the mole fraction by the total pressure gives the partial pressure of that gas. As a check, the mole fractions of all the gases present should add up to 1, and their partial pressures should add up to the total pressure.

Le Chatelier's Principle applies to gaseous equilibria just as it does to aqueous ones: a system at equilibrium shifts to counteract any change made to it. For a reaction that's exothermic in the forward direction, increasing the temperature pushes the equilibrium to the left, decreasing the ratio of products to reactants and so decreasing Kp. For a reaction that's endothermic in the forward direction, increasing the temperature increases Kp instead. Changing the pressure shifts the position of equilibrium, but it does not change the value of Kp — the shift itself is what restores it. A catalyst speeds up the forward and reverse reactions equally, so it affects neither the position of equilibrium nor the value of Kp.

Check your working: the mole fractions of all the gases should add up to 1, and their partial pressures should add up to the total pressure. Never use square brackets in a Kp expression — square brackets mean concentration, and imply a Kc expression instead.

Kp is built from the partial pressures of the gases in an equilibrium, each raised to the power of its stoichiometric coefficient. Partial pressure itself comes from the mole fraction of a gas multiplied by the total pressure. Changing the pressure shifts the position of equilibrium but leaves Kp unchanged, and a catalyst affects neither — only a change in temperature actually changes the value of Kp.