C6.3 - Equilibria

What is a reversible reaction?

A reaction that can go forwards (reactants → products) and backwards (products → reactants) $

Symbol for reversible reaction

⇌ $

Example of reversible reaction

NH₄Cl(s) ⇌ NH₃(g) + HCl(g) $

What happens when ammonium chloride is heated?

Decomposes to ammonia + hydrogen chloride gases $

What happens when ammonia + hydrogen chloride cool?

Reform solid ammonium chloride $

What is dynamic equilibrium?

When forward and backward reaction rates are equal in a closed system $

Conditions for equilibrium

Closed system; forward and backward reactions occur simultaneously $

What happens to concentrations at equilibrium?

They remain constant (but not necessarily equal) $

Dynamic equilibrium diagram

[DRAW: forward rate decreases, backward rate increases, lines meet at equilibrium] $

Open vs closed system equilibrium

Open system: no equilibrium; Closed system: equilibrium can form $

Iodine equilibrium example

I₂(s) ⇌ I₂(g) $

Open vs closed iodine diagram

[DRAW: open tube vapour escapes; closed tube vapour + solid in equilibrium] $

Le Chatelier’s Principle

If a change is made to a system at equilibrium, the equilibrium shifts to oppose the change $

Effect of increasing concentration of reactants

Shifts right (more products formed) $

Effect of decreasing concentration of reactants

Shifts left (more reactants formed) $

Effect of increasing concentration of products

Shifts left $

Effect of decreasing concentration of products

Shifts right $

Concentration table

[DRAW: increase reactants → right; decrease reactants → left] $

Effect of temperature increase

Shifts to endothermic direction $

Effect of temperature decrease

Shifts to exothermic direction $

Exothermic forward reaction: increase temp

Shifts left $

Exothermic forward reaction: decrease temp

Shifts right $

Endothermic forward reaction: increase temp

Shifts right $

Endothermic forward reaction: decrease temp

Shifts left $

Temperature table

[DRAW: ↑T → endothermic; ↓T → exothermic] $

Effect of pressure (gases only)

Equilibrium shifts to side with fewer moles of gas when pressure increases $

Effect of decreasing pressure

Shifts to side with more moles of gas $

Pressure only affects

Reactions where gas moles differ on each side $

Pressure table

[DRAW: ↑P → fewer moles; ↓P → more moles] $

Haber process equilibrium

N₂ + 3H₂ ⇌ 2NH₃ (ΔH = –92 kJ/mol) $

Compromise between rate and yield

Effect of increasing temperature on Haber

Shifts left (endothermic), decreases ammonia yield $

Effect of decreasing temperature on Haber

Shifts right (exothermic), increases ammonia yield $

Effect of increasing pressure on Haber

Shifts right (fewer moles), increases ammonia yield $

Effect of decreasing pressure on Haber

Shifts left (more moles), decreases ammonia yield $

What does a catalyst do to equilibrium?

Speeds up forward + backward reactions equally; no change in equilibrium position $

Why catalysts don’t shift equilibrium

They lower activation energy for both directions equally $