GCSE Chemistry: Energy & Rates — Quick Burst Practice

Inside GCSE Chemistry: Energy & Rates — Quick Burst Practice

decide whether reactions are exothermic or endothermic, read a reaction profile for activation energy, work out an overall energy change from bond energies (energy in − energy out), find a mean rate as quantity ÷ time, and use collision theory to explain what speeds a reaction up. Five quick rounds, all checked on screen. About 10 minutes.

Learning objectives

• energy changes — exothermic reactions (energy out, surroundings warm) versus endothermic reactions (energy in, surroundings cool), with hand warmers and cold packs as examples;
• reaction profiles and activation energy — the activation-energy hump and the position of products for exothermic and endothermic changes;
• bond-energy calculations — energy in to break bonds − energy out to make bonds, worked for H₂ + Br₂ → 2HBr [higher];
• rate of reaction — the mean-rate formula (quantity ÷ time) and using a tangent to read the rate at an instant;
• collision theory — how enough energy and frequent collisions drive rate, catalysts, and reversible reactions reaching equilibrium.

Activities in this book

1. Round 1 — Exothermic or Endothermic?

   an exothermic reaction transfers energy out to the surroundings, so they warm up — combustion, neutralisation and oxidation are exothermic, and a hand warmer is a handy example. An endothermic reaction takes energy in, so the surroundings cool down — thermal decomposition and a sports cold pack are endothermic.

2. Round 2 — Profiles & Activation Energy

   a reaction profile plots energy against the progress of the reaction. The activation energy is the minimum energy colliding particles need to react — the height of the hump. For an exothermic reaction the products sit lower than the reactants; for an endothermic reaction they sit higher.

3. Round 3 — Bond-Energy Sum

   overall energy change = energy in to break bonds − energy out to make bonds. Add up the bonds broken in the reactants, add up the bonds made in the products, then subtract. A negative answer means energy was released overall, so the reaction is exothermic.

4. Round 4 — Mean Rate of Reaction

   the mean rate of a reaction = quantity of reactant used or product formed ÷ time taken. You can measure gas volume (cm³/s) or mass lost (g/s). On a curved rate graph, draw a tangent to find the rate at one instant — its gradient is the rate.

5. Round 5 — Collision Theory & Equilibrium

   collision theory says particles react only when they collide with enough energy (the activation energy). Higher temperature, higher concentration or pressure, larger surface area and a catalyst all speed a reaction up. A catalyst is not used up. A reversible reaction (⇌) in a closed system reaches dynamic equilibrium.

Suggestions for teaching

• Fix the direction of energy first. Drill exothermic (out, warmer) against endothermic (in, cooler) with concrete examples before profiles or sums.
• Sketch the profile. Label reactants, products and the activation-energy hump so the exothermic/endothermic shape is automatic.
• Lay out the bond-energy method the same way every time. Energy in (bonds broken) minus energy out (bonds made); a negative answer means exothermic.
• Rehearse mean rate as a division with units (cm³/s or g/s), then link the graph gradient and tangent to rate at an instant.
• Tie every rate idea to collisions — enough energy and frequent, successful collisions — and finish on reversible reactions and equilibrium.