Endothermic against exothermic reactions: the difference in one line, the comparison table, the energy level diagrams, 20 everyday examples of each, how to tell which is which from a temperature change or a bond-energy calculation, the mistakes that cost marks, and a printable sheet with practice

"Endothermic vs exothermic reactions" and its variants are searched in every English-speaking school system (the US, Canada, the UK and Australia between them about 3,000 times a month, with a traffic potential near 7,000 for the page that answers it well), because the pair is examined everywhere from Grade 7 to GCSE and the two words are so easy to swap. This page gives the difference in one line, the comparison table, the energy level diagrams with the three things to label, 20 examples of each kind, the two ways to decide (thermometer, bond energies), the curriculum placements and the mistakes, with a printable sheet and practice set.

Download the endothermic and exothermic sheet with practice (PDF) (A4, the table, the diagrams and the examples on page 1, the bond-energy method and the practice set with answers on page 2, free).

The difference in one line

An exothermic reaction transfers energy out to the surroundings, so the surroundings warm up. An endothermic reaction transfers energy in from the surroundings, so the surroundings cool down.

The word parts do the work: exo = out, endo = in, thermic = heat. The thing that trips children is which "side" the temperature belongs to: a thermometer in a beaker measures the surroundings (the solution), not the chemicals' stored energy. In an exothermic reaction the chemicals lose energy and the solution gains it, so the thermometer goes up.

The comparison table

The energy level diagrams

Both diagrams plot energy (vertical) against the progress of the reaction (horizontal), with three things to label:

1. Reactants on the left and products on the right, at their energy levels.
2. The activation energy (E a): the hump between them, the minimum energy the reactants need to start reacting. Every reaction has one, exothermic or not, which is why a match needs striking and petrol needs a spark.
3. ΔH, the overall energy change: the vertical gap between the reactants' level and the products' level.

Exothermic: the products sit below the reactants. The reaction climbs the hump, then drops to a lower level; the drop is the energy released. ΔH is negative (for example, burning methane: ΔH = −890 kJ/mol).

Endothermic: the products sit above the reactants. The reaction climbs the hump and settles higher; the rise is the energy absorbed. ΔH is positive (for example, thermal decomposition of calcium carbonate: ΔH = +178 kJ/mol).

A catalyst lowers the hump (a different path with a smaller activation energy) but does not move the reactants' or products' levels, so ΔH is unchanged. The two-diagram question with a catalyst drawn in is a GCSE staple.

20 exothermic examples

1. Burning fuels: wood, coal, natural gas, petrol, a candle (combustion).
2. Respiration in every living cell: glucose + oxygen → carbon dioxide + water + energy.
3. Neutralisation: an acid and an alkali, such as hydrochloric acid and sodium hydroxide (the beaker warms).
4. Rusting of iron (slow, but exothermic).
5. Hand warmers: iron powder oxidising in a sachet.
6. Self-heating cans: calcium oxide and water.
7. Setting of cement and concrete (hydration, warm for days in a large pour).
8. Thermite: aluminium and iron oxide, hot enough to weld rails.
9. Mixing water with concentrated sulfuric acid (always acid into water).
10. Dissolving sodium hydroxide pellets or calcium chloride in water.
11. The reaction of sodium or potassium with water.
12. Explosions: fireworks, dynamite, airbag inflation (sodium azide).
13. Glow sticks (chemiluminescence: energy out as light rather than heat).
14. Condensation of steam (a change of state; a steam burn is worse than a hot-water burn for this reason).
15. Freezing of water (energy released as ice forms).
16. Composting and the heat in a manure heap (microbial respiration).
17. Fermentation (slightly exothermic; a brewing vat warms).
18. Polymerisation, such as setting epoxy resin (the mix warms).
19. Catalytic converters in cars oxidising carbon monoxide.
20. Adding water to quicklime to make slaked lime (steams).

20 endothermic examples

1. Photosynthesis: carbon dioxide + water + light energy → glucose + oxygen.
2. Thermal decomposition of calcium carbonate (limestone) to lime and carbon dioxide in a kiln.
3. Thermal decomposition of copper carbonate (green to black, a school experiment).
4. Instant cold packs: ammonium nitrate dissolving in water.
5. Dissolving ammonium chloride or potassium chloride in water (the beaker cools).
6. Citric acid and sodium hydrogencarbonate (the sherbet reaction, a classic cooling experiment).
7. Baking: the reactions in bread and cake batter take in heat from the oven.
8. Cooking an egg (the proteins absorb energy to denature).
9. Melting ice; evaporating sweat (changes of state that cool the surroundings).
10. Electrolysis of water into hydrogen and oxygen (electrical energy in).
11. Charging a rechargeable battery.
12. Cracking long-chain hydrocarbons in an oil refinery.
13. The Haber process's reverse: decomposing ammonia.
14. Nitrogen and oxygen forming nitrogen monoxide in a lightning strike or an engine.
15. Producing hydrogen by steam reforming of methane.
16. Converting ozone from oxygen in the upper atmosphere (ultraviolet energy in).
17. Dissolving sugar in water (slightly endothermic).
18. Barium hydroxide with ammonium chloride (cold enough to freeze a wet board to the flask: the demonstration).
19. Making quicklime in a lime kiln (industrial version of number 2).
20. Nuclear-style phrasing aside, every "we must keep heating it" reaction in the lab.

How to decide which is which

With a thermometer. Record the temperature before and after. Up: exothermic. Down: endothermic. In the GCSE required practical the reaction runs in a polystyrene cup with a lid, to keep the energy in the solution long enough to measure.

With bond energies. Breaking bonds costs energy; making bonds releases it.

ΔH = (energy to break all the bonds in the reactants) − (energy released making all the bonds in the products)

Worked: hydrogen burning, 2H₂ + O₂ → 2H₂O. Break 2 H–H (2 × 436 = 872 kJ) and 1 O=O (498 kJ): 1,370 kJ in. Make 4 O–H (4 × 464 = 1,856 kJ) out. ΔH = 1,370 − 1,856 = −486 kJ. Negative, so exothermic: more energy came out making bonds than went in breaking them.

Worked: nitrogen and oxygen, N₂ + O₂ → 2NO. Break N≡N (945) and O=O (498): 1,443 in. Make 2 N=O (2 × 607 = 1,214) out. ΔH = 1,443 − 1,214 = +229 kJ. Positive, endothermic, which is why it only happens in lightning and hot engines.

(Bond energies are averages; the values above are the ones GCSE data sheets give.)

The mistakes that cost marks

• Saying an exothermic reaction "has more energy". The surroundings gain energy; the chemicals lose it.
• Reading the thermometer as the chemicals' energy rather than the surroundings'.
• Drawing the exothermic diagram with products above reactants (the arrow for ΔH must point down).
• Forgetting the activation energy hump, or drawing it only on the endothermic diagram. Both have one.
• Letting a catalyst change ΔH on the diagram. It lowers the hump only.
• Getting the bond-energy subtraction backwards (it is broken minus made; a negative answer is exothermic).
• Calling respiration endothermic because it happens inside the body. It releases energy: exothermic.
• Treating "needs heating to start" as endothermic. A match needs striking but burning is exothermic; the test is the overall energy change, not the start.

Where it sits in each curriculum

Practice (answers in the PDF)

1. A student mixes two solutions and the temperature drops from 21 °C to 14 °C. Endothermic or exothermic? Where did the energy go?
2. Sketch the energy level diagram for the reaction in question 1 and label the reactants, products, activation energy and ΔH.
3. Classify each: rusting, photosynthesis, a cold pack, a glow stick, melting butter, neutralisation.
4. Methane burns: CH₄ + 2O₂ → CO₂ + 2H₂O. Bonds: C–H 413, O=O 498, C=O 805, O–H 464 kJ/mol. Calculate ΔH and classify the reaction.
5. A catalyst is added to an exothermic reaction. What changes on the energy level diagram, and what does not?
6. Why does a match need to be struck if burning releases energy?

Download the sheet and the practice set (PDF)

Where to get help

A 1-on-1 online chemistry class with a tutor who teaches NGSS middle and high-school science, GCSE and A-level chemistry, CBSE Class 9 to 12 and the Australian Years 9 to 12 every week: energy changes taught with a thermometer and a diagram before a definition, and the bond-energy method practised until the subtraction goes the right way every time. The chemistry page lists the courses; the free 30-minute first class is a real lesson on this week's topic. The partner pages are specific heat and Q = mcΔT and acid-base catalysis.

Questions parents ask

See how we teach this, 1-on-1 online →