Skip to content
Dabour HubLEARN · PRACTICE · EXPLORESign inStudent

3.6Periodic Table and Energy

Chemical Equilibrium — OCR A Level Chemistry A H432 Revision Notes & Topical Questions

10 min read 18 parts Lesson 18 of 21

📚 OCR A Level Chemistry A H432 Hub / Topical Revision
🩺
100% Curriculum Specification Match Lesson Doctor Verified
The only platform with 100% curriculum specification match and verified interactive question banks.
0%read
Lesson summary

Complete PDF with figures and original page layoutPage 1 OCR A Level Chemistry A Your notes Chemical Equilibrium Contents Dynamic Equilibrium Le Chatelier’s Principle The Equilibrium Constant, Kc Page 2 Dynamic Equilibrium …

Part 1 of 10Reversible reactions and dynamic equilibrium

Reversible reaction
  • Reversible reaction the products can react to reform the reactants, shown by the ⇌ symbol
  • Some reactions instead go to completion and stop when the reactants are used up.
Dynamic equilibrium
Dynamic equilibrium the state in a closed system where the rate of the forward reaction equals the rate of the backward reaction and the concentrations are constant
  • Reactants and products are constantly reacting.
  • Equilibrium can be approached from the reactant side or the product side.
  • HI breaking down to H2 and I2 at the same rate as they recombine is an example.
Misconception
  • Misconception the concentrations at equilibrium are constant, not equal

Part 2 of 10Open and closed systems

Open and closed systems
Closed system
  • Closed system none of the reactants or products escape
  • CaCO3 ⇌ CaO + CO2 reaches equilibrium only if the CO2 cannot escape.
Open system
  • Open system matter and energy can be lost to the surroundings
  • If CO2 escapes, the CaCO3 keeps decomposing and the reaction goes to completion.
Solutions
  • Solutions a reaction wholly in solution can reach equilibrium in an open flask because only negligible material evaporates
Gases
  • Gases a reaction involving gases reaches equilibrium only in a closed system

Part 3 of 10Position of equilibrium and Le Chatelier principle

Position of equilibrium
Position of equilibrium the relative amounts of products and reactants in the equilibrium mixture
  • A shift to the left increases the concentration of reactants.
  • A shift to the right increases the concentration of products.
Le Chatelier principle
Le Chatelier principle when a change is made to a system in dynamic equilibrium, the position of equilibrium moves to counteract the change
  • It predicts the effect of changes in temperature, pressure or concentration.

Part 4 of 10Effect of concentration

Adding a substance
  • Adding a substance the equilibrium moves away from it to reduce its concentration
Removing a substance
  • Removing a substance the equilibrium moves towards it to replace it
Example
  • CH3COOH + C2H5OH ⇌ CH3COOC2H5 + H2O
  • Adding ethyl ethanoate or removing ethanol shifts the position left, forming more ethanoic acid and ethanol.
Dilution
  • Dilution adding water to Ce4+ + Fe2+ ⇌ Ce3+ + Fe3+ has no effect
  • Water dilutes all ions equally, so the ratio of reactants to products does not change.

Part 5 of 10Effect of pressure

Scope
  • Scope pressure only affects equilibria with gaseous reactants or products
Rule
  • Rule an increase in pressure shifts the equilibrium to the side with fewer gas molecules, and a decrease shifts it to the side with more gas molecules
  • N2O4(g) ⇌ 2NO2(g) shifts left on increasing pressure because there are fewer gas molecules on the left.
  • CaCO3(s) ⇌ CaO(s) + CO2(g) shifts left because only the right side contains gas.
  • 2NO2(g) ⇌ 2NO(g) + O2(g) shifts right on decreasing pressure because there are more gas molecules on the right.

Part 6 of 10Effect of temperature

Rule
  • Rule raising the temperature shifts the equilibrium in the endothermic direction, and lowering it shifts it in the exothermic direction
  • The system absorbs the added energy or releases energy to oppose the change.
Example
  • H2 + CO2 ⇌ H2O + CO with ΔH = +41.2 kJ mol−1 shifts right on heating because the forward reaction is endothermic
Deduction
  • Deduction if heating Ag2CO3(s) ⇌ Ag2O(s) + CO2(g) gives more CO2, the forward reaction is endothermic

Part 7 of 10Catalysts and equilibrium

  • Effect
    • a catalyst speeds up the forward and reverse reactions equally
    • The equilibrium is reached faster.
    • The position of equilibrium is not changed.

Part 8 of 10Investigating concentration changes: chromate and dichromate

Investigating concentration changes: chromate and dichromate
Equilibrium
  • Equilibrium chromate, CrO42− (yellow), and dichromate, Cr2O72− (orange), are interconverted and depend on H+ concentration
  • The different colours make any shift easy to see.
Adding dilute sulfuric acid
  • Adding dilute sulfuric acid increases [H+]
  • The equilibrium shifts right to reduce [H+].
  • The solution turns orange as more Cr2O72− forms.
Adding aqueous sodium hydroxide
  • OH− removes H+, H+(aq) + OH−(aq) → H2O(l)
  • The equilibrium shifts left to replace H+.
  • The solution turns yellow as more CrO42− forms.

Part 9 of 10Investigating temperature changes: cobalt chloride

Equilibrium
  • [Co(H2O)6]2+(aq) + 4Cl−(aq) ⇌ [CoCl4]2−(aq) + 6H2O(l)
  • The forward reaction is endothermic and the backward reaction exothermic.
  • Cobalt chloride dissolves to give a pink solution containing both complexes.
Heating
  • Heating the equilibrium shifts right, favouring the endothermic direction
  • The system takes in heat to oppose the temperature rise and the solution turns blue as more [CoCl4]2− forms.
Cooling
  • Cooling the equilibrium shifts left, favouring the exothermic direction
  • The system releases heat to oppose the fall and the solution turns pink as more [Co(H2O)6]2+ forms.

Part 10 of 10Haber process conditions

1
Reaction
  • N2(g) + 3H2(g) ⇌ 2NH3(g), ΔHr = −92 kJ mol−1
2
Pressure
  • Pressure high pressure shifts the equilibrium to the right, towards fewer gas molecules, and raises the rate
  • Closer particles give more successful collisions.
  • Very high pressures are expensive, so a compromise of 200 atm is used.
3
Temperature
  • Temperature the reaction is exothermic, so a low temperature shifts the equilibrium right and gives a higher yield
  • At a low temperature the gases lack kinetic energy to react, so equilibrium is reached too slowly.
  • A compromise of 400–450 °C is used.
  • A heat exchanger warms the incoming gases so they collide more frequently.

Key terms

TermMeaning
Dynamic equilibriumClosed system state where forward and backward rates are equal and concentrations are constant
Closed systemA system in which no reactants or products escape
Position of equilibriumThe relative amounts of products and reactants at equilibrium
Le Chatelier principleThe equilibrium moves to oppose any change made to it
Reversible reactionA reaction in which products can react to reform the reactants

Avoid these mistakes

1
Stating that concentrations are equal at equilibrium — they are constant, not equal
2
Saying a catalyst shifts the position of equilibrium — it only speeds up reaching it
3
Applying pressure effects to equilibria with no gas, or with equal moles of gas on both sides
4
Saying the equilibrium stops when it is reached — the forward and reverse reactions continue at equal rates

Last updated:

3 & IGCSE —

Kc

  • : Kc! (ΔH < 0) Kc. (ΔH > 0) Kc.
  • : ( ).
  • : Kc .