OxfordAQA International A Level Chemistry
OxfordAQA International A Level Chemistry 3.1.12 Equilibrium Constant Kp for Homogeneous Systems Exam Questions
Free OxfordAQA International A Level Chemistry revision for 3.1.12 Equilibrium Constant Kp for Homogeneous Systems: what you need to know, common mistakes and exam-style questions with answers. Covers mole fractions, partial pressures, Kp expressions and units, and the effects of temperature and pressure on Kp.
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About OxfordAQA International A Level Chemistry 3.1.12 Equilibrium Constant Kp for Homogeneous Systems
This A2 physical chemistry topic, examined in Unit 4, extends the work on Kc from Unit 2 to gas-phase equilibria. You need to calculate mole fractions and partial pressures, write expressions for Kp for homogeneous gaseous systems, calculate Kp and its units, and explain how changes in temperature and pressure affect the position of equilibrium and the value of Kp. You also need to know that catalysts do not change Kp. Examiners often set multi-step questions where you start with initial moles, find equilibrium moles, convert to mole fractions and partial pressures, then calculate Kp.
What you need to know
- Mole fraction Mole fraction of a gas = moles of that gas / total moles of gas. The mole fractions in a mixture add up to 1.
- Partial pressure Partial pressure = mole fraction x total pressure. The partial pressures add up to the total pressure.
- Kp expression For aA(g) + bB(g) ⇌ cC(g) + dD(g), Kp = p(C)^c p(D)^d / p(A)^a p(B)^b. Only gases appear in the expression.
- Units Work out the units by cancelling the pressure units, for example kPa or Pa, in the expression.
- Changing conditions Only temperature changes Kp. Raising the temperature increases Kp if the forward reaction is endothermic. Changing pressure shifts the position of equilibrium but Kp is unchanged, and a catalyst has no effect on Kp.
Common mistakes in 3.1.12 Equilibrium Constant Kp for Homogeneous Systems
- Using moles rather than partial pressures in Kp. Convert each to mole fraction and then partial pressure first.
- Forgetting to recalculate the total moles at equilibrium before finding mole fractions.
- Saying Kp changes when total pressure changes. Kp depends only on temperature.
- Giving Kp without units, or with the wrong units, when units are required.
Try three 3.1.12 Equilibrium Constant Kp for Homogeneous Systems exam-style questions
1. At equilibrium in 2SO2 + O2 ⇌ 2SO3, there are 0.20 mol SO2, 0.10 mol O2 and 1.70 mol SO3 at a total pressure of 200 kPa. Calculate the partial pressure of SO3. (2 marks)
Show the mark scheme answer
Total moles = 2.00, so mole fraction of SO3 = 1.70 / 2.00 = 0.85; partial pressure = 0.85 x 200 = 170 kPa.
2. For N2O4(g) ⇌ 2NO2(g), the partial pressures are N2O4 = 60 kPa and NO2 = 40 kPa. Calculate Kp and give its units. (2 marks)
Show the mark scheme answer
Kp = 40^2 / 60 = 26.7; kPa.
3. The forward reaction for N2O4 ⇌ 2NO2 is endothermic. State and explain the effect of increasing temperature on Kp. (2 marks)
Show the mark scheme answer
Kp increases; equilibrium shifts in the endothermic (forward) direction, increasing partial pressure of NO2 relative to N2O4.
3.1.12 Equilibrium Constant Kp for Homogeneous Systems FAQs
What does OxfordAQA International A Level Chemistry 3.1.12 Equilibrium Constant Kp for Homogeneous Systems cover?
It covers mole fractions, partial pressures, Kp expressions and units, and the effects of temperature and pressure on Kp.
Which unit is 3.1.12 Equilibrium Constant Kp for Homogeneous Systems examined in?
It is examined in Unit 4, an International A2 paper covering organic and physical chemistry. The Unit 5 practical and synoptic paper can also assess it.
What is the method for a Kp calculation?
Find the equilibrium moles of each gas, add them to get total moles, calculate mole fractions, multiply by total pressure to get partial pressures, then substitute into the Kp expression and work out the units.
