Energy transfer: exothermic and endothermic reactions
Energy is never made or destroyed in a reaction. It moves between the chemicals and the surroundings (the water, the beaker, the air).
- Exothermic: energy goes out to the surroundings, so the temperature goes up. Examples: burning fuels (combustion), most oxidation reactions, neutralisation of an acid by an alkali, respiration.
- Endothermic: energy comes in from the surroundings, so the temperature goes down. Examples: thermal decomposition (like heating calcium carbonate), citric acid with sodium hydrogencarbonate, photosynthesis, dissolving ammonium nitrate.
Uses: hand warmers and self-heating cans (exothermic); sports cold packs (endothermic).
How to test: mix the chemicals in an insulated cup (a polystyrene cup with a lid), measure the start and highest (or lowest) temperature, and compare. Keep the amounts the same to make it fair.
Energy in changes of state
Changes of state are physical changes, but they also take in or give out energy:
- Take energy in (endothermic): melting, boiling/evaporating, subliming. Sweat evaporating cools your skin.
- Give energy out (exothermic): freezing, condensing. Steam burns badly because it gives out extra energy when it condenses on skin.
While a substance is changing state its temperature stays the same; the energy is used to separate (or is released by bringing together) the particles.
Reaction profiles and activation energy
A reaction profile is a graph of energy (up) against the progress of the reaction (across).
- Reactants are drawn on the left, products on the right.
- Exothermic: products are lower than reactants. The difference is the energy given out.
- Endothermic: products are higher than reactants. The difference is the energy taken in.
- The hump is the activation energy: the smallest energy particles need when they collide for the reaction to start. A match needs a spark (activation energy) even though burning is exothermic.
A catalyst gives a lower hump but does not change the start and end levels.
Bond energies: calculating the energy change
In a reaction, old bonds break and new bonds form.
- Breaking a bond takes energy in (endothermic step).
- Making a bond gives energy out (exothermic step).
ΔH = (total energy to break bonds in reactants) − (total energy released making bonds in products)
If more energy is released making bonds than is needed to break them, ΔH is negative: exothermic. If it is the other way round, ΔH is positive: endothermic.
Method: (1) draw every bond in each molecule, (2) count each type, (3) multiply by its bond energy, (4) add up the breaking side and the making side, (5) subtract. Units: kJ/mol.
Try it: a kitchen test
Put a spoon of baking soda in a glass, add a spoon of lemon juice or vinegar and feel the glass: it gets slightly cooler (endothermic). Then wet the back of your hand and blow on it: it feels cold because evaporation takes energy from your skin (an endothermic change of state). Predict first, then feel, and note which got colder.
Key formulas and definitions
- Exothermic: energy out, temperature of surroundings rises, ΔH negative
- Endothermic: energy in, temperature of surroundings falls, ΔH positive
- ΔH = Σ(bonds broken) − Σ(bonds made)
- Activation energy = energy from reactants to the top of the hump
- Units of bond energy and ΔH: kJ/mol
Worked examples
1. The temperature of an acid–alkali mixture goes from 21 °C to 29 °C. Is the reaction exothermic or endothermic?
The temperature rose by 8 °C, so energy was given out to the surroundings: exothermic.
2. Citric acid and sodium hydrogencarbonate solution drop from 20 °C to 14 °C. What type of reaction is this, and where did the energy go?
Temperature fell by 6 °C, so it is endothermic. Energy moved from the water (surroundings) into the chemicals to break bonds.
3. H₂ + Cl₂ → 2HCl. Bond energies: H–H 436, Cl–Cl 243, H–Cl 432 kJ/mol. Find ΔH.
Broken: 436 + 243 = 679 kJ. Made: 2 × 432 = 864 kJ. ΔH = 679 − 864 = −185 kJ/mol. Negative, so exothermic.
4. 2H₂ + O₂ → 2H₂O. H–H 436, O=O 498, O–H 464 kJ/mol. Find ΔH.
Broken: 2 × 436 + 498 = 1370 kJ. Made: each H₂O has 2 O–H bonds, so 4 × 464 = 1856 kJ. ΔH = 1370 − 1856 = −486 kJ/mol (exothermic).
5. CH₄ + 2O₂ → CO₂ + 2H₂O. C–H 413, O=O 498, C=O 805, O–H 464 kJ/mol. Find ΔH.
Broken: 4 × 413 + 2 × 498 = 1652 + 996 = 2648 kJ. Made: 2 × 805 + 4 × 464 = 1610 + 1856 = 3466 kJ. ΔH = 2648 − 3466 = −818 kJ/mol. Methane burning is strongly exothermic.
6. 2HBr → H₂ + Br₂. H–Br 366, H–H 436, Br–Br 193 kJ/mol. Find ΔH and say what type of reaction it is.
Broken: 2 × 366 = 732 kJ. Made: 436 + 193 = 629 kJ. ΔH = 732 − 629 = +103 kJ/mol. Positive, so endothermic.
Common mistakes
- Thinking endothermic means "cold chemicals give out cold". Cold is not given out; energy is taken in from the surroundings, so they cool.
- Saying breaking bonds releases energy. Breaking bonds always takes energy in.
- Subtracting the wrong way: ΔH = bonds broken − bonds made, not the other way round.
- Forgetting to multiply by the number of molecules and the number of bonds in each molecule (e.g. 2H₂O has 4 O–H bonds).