Chemical Equilibrium

• Irreversible reaction — runs one way only and goes to completion (burning, precipitation). Written with a single arrow →.
• Reversible reaction — runs both ways; products can re-make reactants. Written with the double arrow ⇌, e.g. N₂ + 3H₂ ⇌ 2NH₃.

Picture a busy bridge between two towns. Cars stream across in both directions at once. A reversible reaction is exactly that traffic: the forward trip builds products while the return trip rebuilds reactants, and neither side ever empties.

• Chemical equilibrium — in a closed system, the state where the forward rate equals the reverse rate, so all concentrations stay constant.

Equilibrium is dynamic, not frozen. Think of a crowded dance floor: couples form and break up constantly, partners are swapped every minute — yet the number of people dancing stays the same. Underneath, everything is moving; from outside, nothing seems to change. A reaction at equilibrium is the same: forward and reverse never stop, they just cancel.

• Reached only in a closed system; it is dynamic.
• Concentrations stay constant (but not necessarily equal).
• It can be reached from either direction.
• A catalyst only speeds the approach — it never changes the final position.

Why does the balance hold? Imagine a bank account where deposits and withdrawals happen at equal rates. Money flows in and out all day, yet the displayed balance never moves. Equilibrium is reached the instant the forward "deposit" rate matches the reverse "withdrawal" rate — and it can settle there whether you started rich or broke.

The law of mass action says the rate of a reaction is proportional to the product of the molar concentrations of the reactants. Applied to both directions, it gives the equilibrium constant, the single number that fixes the balance.

Products go on top, reactants below, each raised to its coefficient. Kc is fixed by temperature alone; pure solids and liquids are left out (their concentration cannot change).

For gas reactions we can measure partial pressures instead of concentrations, giving Kp:

The magnitude of K is a shop's stock report. Walk into the "products shop": if K ≫ 1 the shelves are packed with products; if K ≪ 1 they are nearly empty and the reactants' warehouse is full; if K ≈ 1 both are about half-stocked.

• Le Chatelier's principle — if a stress (change of concentration, pressure or temperature) is applied to a system at equilibrium, it shifts in the direction that opposes (relieves) the stress.

Picture a see-saw resting level. Drop a weight on the left and it tips — but the system answers by moving mass to the right until a new balance is found. Add more reactant and the reaction shifts forward to "use it up"; add product and it shifts back. The system always pushes back against whatever you did.

Imagine a crowded room you can shrink. Squeeze it (raise the pressure) and the crowd relieves the crush by moving to the side that takes up less space — the side with fewer moles of gas. For N₂ + 3H₂ ⇌ 2NH₃, four moles become two, so squeezing favours ammonia. Move the slider to feel it.

Temperature works on the heat itself. Heat is a "reactant" for the endothermic side: add heat and the equilibrium shifts the endothermic way; cool it and it shifts exothermic.

Le Chatelier says: use high pressure (4 moles → 2) and low temperature (exothermic) to maximise NH₃. But low temperature is far too slow, so industry uses a compromise: ~200 atm, ~450 °C, an iron catalyst. The Contact process (2SO₂ + O₂ ⇌ 2SO₃, V₂O₅ catalyst) makes sulphuric acid by the very same reasoning.

1. Reversible vs irreversible; the ⇌ sign (the two-way bridge).
2. Dynamic equilibrium: equal forward & reverse rates (dance floor, bank balance).
3. Kc expression; Kp = Kc(RT)^Δn; the size of K.
4. Le Chatelier — concentration, pressure, temperature (the see-saw, the squeeze).
5. Haber & Contact processes (conditions & reasons).