Pressure effects on reversible gas-phase equilibria

The rate and extent of chemical change • Reversible reactions and equilibrium

Flashcards

Test your knowledge with interactive flashcards

Key distinction between total pressure and partial pressure effects

Click to reveal answer

Only changes in partial pressures of reactive gases affect equilibrium position; total pressure changes matter only insofar as they change those partial pressures.

Key concepts

What you'll likely be quizzed about

Fundamental definition and cause→effect rule

Equilibrium occurs when the forward and reverse reaction rates are equal and concentrations (or partial pressures for gases) remain constant. Le Chatelier’s principle predicts the direction of the shift when external conditions change. Cause: pressure increase (typically by reducing volume). Effect: equilibrium shifts toward the side with fewer gas molecules to reduce pressure. Cause: pressure decrease (typically by increasing volume). Effect: equilibrium shifts toward the side with more gas molecules to increase pressure.

Counting gaseous moles and deciding shift direction

Only gaseous species count when comparing moles on each side of the equation. Cause: unequal total moles of gas between reactants and products. Effect: the side with fewer gaseous moles becomes favored when pressure increases, while the side with more gaseous moles becomes favored when pressure decreases. Limiting factor: identical numbers of gaseous moles on both sides produce no shift when pressure changes.

Role of Kp and Qp in predictions (higher-tier)

Kp remains constant at a fixed temperature. Interpretation of data uses the reaction quotient Qp, calculated from current partial pressures, and comparison with Kp. Cause: Qp > Kp. Effect: equilibrium shifts in the direction that reduces product partial pressures (toward reactants). Cause: Qp < Kp. Effect: equilibrium shifts to increase product partial pressures (toward products). Pressure changes alter partial pressures and therefore change Qp; comparing Qp and Kp predicts the direction of the shift quantitatively.

Effects of adding an inert gas and practical limitations

Adding an inert gas at constant volume increases total pressure but leaves partial pressures of reacting gases unchanged, so the equilibrium position remains unchanged. Cause: addition of inert gas at constant pressure (system allowed to expand). Effect: partial pressures of reactive gases change, possibly causing a shift. Limiting factors include whether the pressure change occurs by volume change or by adding inert gas, and whether non-gaseous species are present. Solids and pure liquids do not change the equilibrium position when pressure changes because their activities remain effectively constant.

HT data-interpretation steps for pressure changes

Step 1: Identify which species are gaseous and count gas moles on each side of the equation. Step 2: Determine how pressure changes (increase by decreasing volume, decrease by increasing volume, or addition of inert gas). Step 3: Predict qualitative shift using the mole-count rule and Le Chatelier’s principle. Step 4: If numerical data are provided, calculate Qp from partial pressures before and after the change and compare with Kp to make a quantitative prediction. Step 5: State the expected change in concentrations/pressures and whether the system moves toward products or reactants.

Key notes

Important points to keep in mind

Count only gaseous species when deciding direction of shift.

Increased pressure favours the side with fewer gas moles; decreased pressure favours the side with more gas moles.

Solids and liquids do not change equilibrium position in response to pressure changes.

Addition of inert gas at constant volume does not alter equilibrium position.

Kp remains constant at fixed temperature; use Qp to assess direction after a pressure change.

Pressure changes alter collision frequency and partial pressures, producing a shift until Qp equals Kp again.

Catalysts change the rate of reaching equilibrium but not the equilibrium position.

Small pressure changes produce small shifts; large shifts require significant pressure differences or large differences in gas mole numbers.

Always state assumptions about whether temperature remains constant when predicting equilibrium changes.

When given numerical data, calculate Qp from partial pressures for a quantitative prediction.