Equilibrium I
Quick Notes
- Reversible reactions can proceed in both forward and reverse directions.
- At dynamic equilibrium:
- Forward rate = backward rate
- Concentrations of reactants and products remain constant
- If conditions change, the system shifts to oppose the change:
- Increased temperature: favours endothermic direction
- Decreased temperature: favours exothermic direction
- Increasing concentration of something shifts equilibrium to favour direction that uses it as a reactant
- Increased pressure (gases only): favours direction that produces fewest moles of gas
- Decreased pressure: favours direction that produces most moles of gas
- Industrial conditions often use compromises between high yield and fast rate.
- Kc (equilibrium constant) expresses ratio of product to reactant concentrations at equilibrium:
- For a reaction: aA + bB ⇌ cC + dD
Full Notes
Reversible reactions can go forward (reactants → products) and backward (products → reactants).
Example The Haber Process (Ammonia Production)
- Forward reaction: N2 + 3H2 → 2NH3
- Reverse reaction: NH3 → N2 + H2
Dynamic equilibrium is reached in a closed system when rate of the forward reaction equals the rate of the backward reaction.
- The concentrations of all species remain constant (though not necessarily equal).
It is dynamic because both reactions continue, but there is no overall change in concentrations.
The ‘position’ of equilibrium refers to the relative amounts of reactants and products. More products mean equilibrium lies to the right.
Effect of a change in temperature, concentration or pressure
When changes are made to a system at equilibrium, the rates of forward and reverse reactions change unequally.
This means one direction occurs at a higher rate than the other (the 'favoured' direction), until a new equilibrium is reached.
We can predict the effect of changing conditions on the position of equilibrium as the system responds to oppose the change using Le Chatelier’s Principle:
Changing Temperature
Increasing temperature favours the endothermic direction (+ΔH).
Decreasing temperature favours the exothermic direction (−ΔH).
Example In the Haber Process
- Forward reaction is exothermic (−ΔH).
- Increasing temperature shifts equilibrium left, reducing NH3 yield.
- Decreasing temperature shifts equilibrium right, increasing NH3 yield.
Changing Concentration
Increasing reactant concentration shifts equilibrium right (more products).
Increasing product concentration shifts equilibrium left (more reactants).
Example
- Adding more N2 in the Haber Process shifts equilibrium right, producing more NH3.
Changing Pressure (for Gaseous Equilibria)
Increasing pressure shifts equilibrium in the direction that produces the fewest moles of gas.
Decreasing pressure shifts equilibrium in the direction that produces the most moles of gas.
If gas moles are equal on both sides, pressure has no effect.
Example Haber Process
- 4 moles (N2 + 3H2) ⇌ 2 moles (NH3)
- The forward direction produces the fewest moles of gas (4 to 2), meaning higher pressure shifts equilibrium right, increasing NH3 yield.
Industrial Equilibrium: Yield vs Rate
Industrial conditions are chosen as compromises between high yield and fast rate.
Example Haber Process (N2 + 3H2 ⇌ 2NH3)
- Low temperature favours forward (exothermic) direction → high yield (but too slow).
- High pressure favours forward direction → high yield (but expensive/safety issues).
- Compromise: 450 °C and 200 atm
Equilibrium Constant, Kc
Kc shows the position of equilibrium for reactions in homogeneous equilibrium.
General formula for Kc:
For a reaction:
Kc =
where:
- [A], [B], [C], [D] = equilibrium concentrations (mol dm⁻³)
- a, b, c, d = balancing numbers
If...
- Kc > 1
This means equilibrium favours products - there is a greater proportion of products in the mixture at equilibrium compared to reactants (forward direction favoured) - Kc < 1
This means equilibrium favours reactants - there is a greater proportion of reactants in the mixture at equilibrium compered to products (reverse direction favoured)
Solids aren’t ever included in Kc expressions and if water is a solvent, it also isn’t included (even if it is also a reactant or product).
How to Calculate Kc
CH3COOH + C2H5OH ⇌ CH3COOC2H5 + H2O
Equilibrium concentrations:
[CH3COOH] = 0.20 mol dm⁻³
[C2H5OH] = 0.20 mol dm⁻³
[CH3COOC2H5] = 0.40 mol dm⁻³
[H2O] = 0.40 mol dm⁻³
- Kc = [CH3COOC2H5][H2O] ÷ [CH3COOH][C2H5OH]
- = (0.40 × 0.40) ÷ (0.20 × 0.20)
- = 4.0
Since Kc > 1, equilibrium favours the products. This means in the equilibrium mixture there is a higher concentration of products (ester and water) compared to reactants.
Remember concentrations in Kc are equilibrium values. Always check if you need to calculate equilibrium concentrations from given data first.
Effect of Changing Conditions on Kc
Changing conditions can cause a change to the value of Kc for a given reaction.
Temperature:
- Kc changes with temperature because temperature affects equilibrium position.
- If the forward reaction is endothermic (+ΔH), increasing temperature increases Kc.
- If the forward reaction is exothermic (-ΔH), increasing temperature decreases Kc.
Concentration:
- Kc does not change with concentration or pressure
- The system shifts equilibrium to restore the original Kc
Catalysts:
- Catalysts do not change Kc.
- They speed up how fast equilibrium is reached but do not affect equilibrium position.
Summary
- Dynamic equilibrium is when forward and reverse reaction rates are equal in a closed system.
- Changing temperature, pressure, or concentration shifts equilibrium position according to Le Chatelier’s principle.
- Kc indicates equilibrium position. Kc > 1 favours products, Kc < 1 favours reactants.
- Kc only changes with temperature, not with catalysts, concentration, or pressure.