Exothermic & Endothermic Reactions
Energy is never created or destroyed — it is only transferred. The total amount of energy in the universe is exactly the same after a reaction as it was before; this is the law of conservation of energy. What changes is where the energy is. We sort reactions into two types depending on which way the energy flows between the chemicals (the system) and everything around them (the surroundings).
An exothermic reaction transfers energy to the surroundings, so the temperature of the surroundings increases. The products store less energy than the reactants.
An endothermic reaction takes in energy from the surroundings, so the temperature of the surroundings decreases. The products store more energy than the reactants.
The quickest way to tell them apart in the lab is with a thermometer: if the mixture (and its surroundings) warms up, the reaction is exothermic; if it cools down, it is endothermic.
| Type | Energy flow | Temperature of surroundings |
|---|---|---|
| Exothermic | Given out to surroundings | Increases (warms up) |
| Endothermic | Taken in from surroundings | Decreases (cools down) |
Examples of each type:
- Exothermic — combustion (burning fuels), many oxidation reactions, neutralisation, and displacement of a metal from its salt solution.
- Endothermic — thermal decomposition, and the reaction of citric acid with sodium hydrogencarbonate.
Everyday uses
Examiners often ask you to evaluate a use of an exothermic or endothermic reaction — so it helps to have real examples ready:
- Hand warmers and self-heating cans use exothermic reactions to release heat on demand.
- Sports injury cold packs use an endothermic change to draw heat away from a sprain or bruise.
Exothermic and endothermic are defined by the temperature change of the surroundings, not by whether the reaction “feels hot”. An exothermic reaction releases energy, so the surroundings get hotter. Don’t say “the energy is lost” — energy is conserved; it is transferred to the surroundings.
Choose a reaction type to see where the energy goes — and why the solution, not the chemicals, is what heats up or cools down.
For everyone, AQA limits this to measuring a temperature change. You do not need to calculate energy changes here — that only comes in for Higher Tier, using bond energies (section 4).
🧪 Exam-style questions
A reaction is endothermic. What happens to the temperature of the surroundings? Tick (✓) one box.
A hand warmer uses a chemical reaction to keep your hands warm. Explain, using the word surroundings, why the hand warmer feels hot.
Show answer
- The reaction is exothermic, so it transfers energy to the surroundings. 1 mark
- The temperature of the surroundings (your hand) increases, so it feels hot. 1 mark
Do not accept “the energy is lost” — energy is conserved; it is transferred to the surroundings.
Which of these reactions is endothermic? Tick (✓) one box.
Measuring Temperature Changes
This introduces Required Practical 4: investigating the variables that affect the temperature change in reacting solutions. The same method works for many reactions named in the course — an acid with a metal, an acid with a carbonate, a neutralisation, or a displacement reaction.
A reliable basic method:
- Stand a polystyrene cup inside a beaker (for support) and add a measured volume of the first solution.
- Measure and record the starting temperature.
- Add the second reactant, put a lid on, stir, and record the highest (or, for endothermic, the lowest) temperature reached.
- The temperature change = final temperature − starting temperature.
Variables: the independent variable is what you investigate (e.g. the type or concentration of a reactant); the dependent variable is the maximum temperature change; the control variables are the volume of solution, the insulation (same polystyrene cup and lid) and the same thermometer. The headline result is that a bigger temperature change means more energy is transferred per unit of reactant.
The polystyrene cup and lid are the key pieces of kit: polystyrene is a good insulator, so it reduces the energy transferred between the mixture and the surroundings. That makes your temperature change more accurate. For an exothermic reaction the insulation slows energy being lost; for an endothermic one it slows energy being gained — so for an endothermic reaction, write that it “prevents energy being gained”, not that it “prevents heat loss”.
Variables
As with any investigation, be clear about which variable does which job:
- Independent (the one you change) — e.g. the mass of metal added, the concentration of acid, or the type of metal.
- Dependent (the one you measure) — the temperature change.
- Control (kept the same) — e.g. the volume and concentration of the solution, and the starting temperature.
Always state the actual quantity you measure out: the mass (in g) of a solid, or the volume (in cm³) of a solution. The vague word “amount” is not accepted in the exam — it could mean mass, volume, concentration or number of moles, so it earns no mark.
A classic graph plots the temperature change against the mass or volume of one reactant you add to a fixed quantity of the other. Up to the point where the reactants are in their reacting proportions, the temperature change increases as you add more. What happens after that depends on what you added — and explaining it is where the marks are:
- add a solid and the line levels off (a plateau): the other reactant is now used up, so the extra solid just sits there unreacted and makes no further difference;
- add a solution and the line peaks then dips: the excess cold solution cools and dilutes the mixture, so the reading falls back.
Don’t be thrown if a question draws the plateau sloping gently downward rather than dead flat — once the reaction is over, the warm solution slowly cools toward room temperature, so the reading drifts down. The mark is still for spotting that the reaction stops adding heat once the limiting reactant is used up.
Add more of one reactant to a fixed quantity of the other and the temperature change rises at first, up to the reacting-proportions point. After that the two cases split: a solid is simply left over, so the line levels off; excess solution cools and dilutes the mix, so it dips.
Try the four reactions for yourself. The simulator below runs the required practical: choose a reaction, pick your reactants, add them to the insulated cup and read the temperature change off the thermometer — then decide whether it is exothermic or endothermic. Run several and the results build into a comparison table.
🧪 Run the required practical — temperature change
Pick one of the four reactions named in the specification, choose your reactants, then add them to the polystyrene cup and watch the thermometer. Record the temperature change and decide: exothermic or endothermic?
Choose a reaction and its reactants above, then press Run experiment.
| Reaction | Final / °C | ΔT / °C | Result |
|---|
Every run starts from 21.0 °C (a control variable). Change one reactant at a time — the independent variable — and watch how ΔT responds.
Displacement reactions and reactivity
A more reactive metal will displace a less reactive metal from a solution of its salt. Adding zinc to copper sulfate solution is a classic example:
Zn + CuSO4 → ZnSO4 + Cu
Displacement reactions are exothermic, so the temperature of the solution rises. On a reaction profile they look like the exothermic diagram in section 3 — products lower than reactants. A common exam slip is to call them endothermic: if the temperature went up, the reaction released energy, so it must be exothermic.
- The greater the difference in reactivity between the added metal and the metal in the salt, the more energy is released — so the larger the temperature rise.
- A metal less reactive than the one already in solution gives no reaction and no temperature change (for example copper added to zinc sulfate, or copper added to copper sulfate).
- This lets you put metals in order of reactivity from their temperature rises — the bigger the rise, the more reactive the metal.
To place an unknown metal in the series: add the same mass (and surface area) of it to the same volume and concentration of copper sulfate solution at the same starting temperature — the control variables — record the temperature rise, and compare it with the rises for the known metals.
When the variable you change is the type of metal, that is a categoric variable, so the results are plotted as a bar chart, not a line graph. Label both axes with the quantity and its unit, choose a scale that fills at least half the grid, and draw bars of equal width with gaps between them.
Reading the chart: the more reactive the metal, the bigger the temperature rise — and copper, which cannot displace itself from the solution, gives none.
Errors and uncertainty
- Random errors scatter readings either side of the true value — for example slightly misreading the thermometer or misjudging the highest temperature. Repeating the experiment and taking a mean reduces their effect.
- A systematic error shifts every reading the same way. Using a glass beaker instead of a polystyrene cup is systematic: glass is a poorer insulator, so energy is lost to the surroundings every time and the temperature change always comes out too small.
- The uncertainty in a mean is often estimated as half the range: uncertainty = (highest valid value − lowest valid value) ÷ 2, quoted as mean ± uncertainty.
Don’t mix that up with a temperature–time graph of a single run, where the across-the-bottom axis is time rather than the mass or volume you added. Here the temperature shoots up to a maximum as the reaction happens, then slowly falls back: once the reaction has finished no more energy is released, so the warm mixture loses heat to the cooler surroundings.
Temperature against time within a single experiment: a sharp rise to a maximum as the reaction gives out energy, then a slow fall as the warm mixture loses heat to the surroundings.
🧪 Exam-style questions
A student repeated a temperature-change experiment four times. Ignoring the anomalous result, the four valid temperature rises were 5.6 °C, 5.8 °C, 5.9 °C and 5.7 °C. Calculate the mean temperature rise.
Show answer
- Add the four valid readings: 5.6 + 5.8 + 5.9 + 5.7 = 23.0. 1 mark
- Divide by 4: 23.0 ÷ 4 = 5.75 °C (5.8 °C to 2 s.f.). 1 mark
- Note: the anomaly is left out before averaging — dividing all five readings by five would be wrong.
A student dissolves sodium nitrate in water (an endothermic change) in a glass beaker. Which improvement would make the temperature reading most accurate?
Equal masses of four metals are each added, in turn, to separate samples of copper sulfate solution. Which metal would give the largest temperature rise?
Four valid temperature rises were recorded: 5.4 °C, 5.7 °C, 5.5 °C and 5.6 °C. Estimate the uncertainty in the mean as half the range.
Show answer
- Range = highest − lowest = 5.7 − 5.4 = 0.3 °C.
- Uncertainty = range ÷ 2 = 0.3 ÷ 2 = ±0.15 °C. 1 mark