GCSE Combined Science Rates & equilibrium revision

The mean rate of a reaction is how much stuff changes divided by how long it takes: rate = quantity changed ÷ time. A reaction filling a syringe with 72 cm³ of gas in 45 s runs at 72 ÷ 45 = 1.6 cm³/s. Rates carry units — cm³/s for gas volume, g/s for mass lost, mol/s for moles — and forgetting the unit drops the mark.

Collision theory explains every factor in one breath: particles react only when they collide with at least the activation energy. Temperature up → particles move faster → collisions more frequent *and* more energetic, so many more succeed. Concentration or pressure up → more particles in the same volume → more collisions per second. Surface area up → a powder exposes far more particles than one lump → more collisions at the surface. A catalyst plays a different game: it offers a route with a lower activation energy, so the same collisions succeed more often — and it emerges unchanged, with exactly the mass it went in with.

Pick the method from the reaction. Gas made? Collect it in a gas syringe (or over water). Gas escaping from an open flask? Follow the mass loss on a balance. A precipitate clouding a solution? Time a disappearing cross under the beaker. All three plot quantity against time — and all three curves tell the same story: steepest at the start (reactants plentiful), flattening as a reactant runs out, perfectly flat when the reaction is over.

The mean rate covers a whole interval; the rate at one moment needs a tangent. Draw the line that just kisses the curve at your chosen time, form a large triangle, and divide rise by fall. A tangent at 15 s that climbs 33 cm³ over the next 15 s gives a rate of 33 ÷ 15 = 2.2 cm³/s at that instant. Once the curve is flat, the gradient — and the rate — is zero.

Trilogy’s rates practical compares how fast a reaction runs at different concentrations, by either the disappearing-cross or the gas-syringe method. Disappearing cross: draw a cross on paper, stand a beaker of sodium thiosulfate solution over it, add hydrochloric acid, start the stopwatch, and stop when the cloudy precipitate hides the cross. Repeat with the same volumes at different concentrations.

The independent variable is the concentration; the dependent variable is the time for the cross to vanish (or the gas volume after a fixed time); control variables include the total volumes used, the mass and surface area of any solid, and the temperature. A shorter time means a faster rate, so results are compared as 1/time: if doubling the concentration roughly doubles 1/time, rate is proportional to concentration.

Some reactions run in both directions. Heat blue hydrated copper(II) sulfate and it releases water vapour, leaving white anhydrous powder; add water and the blue returns, with the tube warming. The symbol ⇌ marks a reversible reaction. Forward and reverse reactions can run at the same time, and in a closed system (nothing escapes) they eventually run at the same rate — dynamic equilibrium. “Dynamic” means both reactions keep going; “equilibrium” means the amounts of reactants and products no longer change.

(Higher) Change the conditions and the balance tips. Le Chatelier-style rule: the equilibrium shifts to oppose the change. Add more of a reactant → the system consumes it, shifting to the products. Increase the pressure on a gas equilibrium → the system shifts to the side with fewer gas molecules, because fewer molecules lowers the pressure. For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), four gas molecules sit on the left and two on the right, so extra pressure pushes the equilibrium right, towards more ammonia. A catalyst, note, speeds both directions equally — it reaches equilibrium faster without moving its position.