Chemical Equilibrium

Some reactions can go forwards and backwards. At equilibrium, the forward and reverse changes happen at the same rate, so the amounts stay constant even though particles keep reacting.

Key ideas

• Reversible reaction: goes both ways, shown by $\rightleftharpoons$.
• Closed system: no substances enter or leave.
• Dynamic equilibrium: in a closed system, the rate forward = rate backward; concentrations are constant but reactions continue.
• Equilibrium does not mean equal amounts; it means equal rates.

Changing the position of equilibrium

When conditions change, the system shifts to oppose the change (Le Chatelier’s idea).

• Concentration: adding more of a reactant makes equilibrium shift to use it up; removing a product pulls the reaction to make more of it.
• Temperature: treat heat like a substance. If the forward reaction is exothermic (releases heat), increasing temperature favors the reverse (endothermic) direction; decreasing temperature favors the exothermic direction.
• Pressure (gases only): increasing pressure shifts to the side with fewer gas molecules; decreasing pressure shifts to the side with more gas molecules.
• Catalyst: speeds up reaching equilibrium by making both directions faster equally; it does not change the equilibrium position.

Everyday and industrial examples

• Cobalt(II) chloride paper: pink when hydrated, blue when dried; adding water or heat shifts the equilibrium and changes color.
• Hydrated copper(II) sulfate: blue crystals lose water on heating to form white anhydrous powder; adding water reverses it.
• Haber process (ammonia):

[ \n\text{N}_2(g) + 3\,\text{H}_2(g) \rightleftharpoons 2\,\text{NH}_3(g) \n]

• More pressure favors ammonia (fewer gas molecules on right). Lower temperature favors ammonia if formation is exothermic, but too low makes it slow, so a moderate high temperature and iron catalyst are used.
• Contact process (making SO₃):

[ \n2\,\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\,\text{SO}_3(g) \n]

Common misconceptions

• “At equilibrium, reactions stop.” False: they continue at equal rates.
• “Equal amounts on both sides.” Not required; only rates are equal.
• “Catalysts change equilibrium position.” They do not; they only help reach equilibrium faster.
• Open systems do not reach dynamic equilibrium because substances can escape.