Energy is the thread that runs through every reaction. Breaking the bonds in the reactants takes in energy, and making the bonds in the products gives out energy; whichever of the two is bigger decides whether the reaction warms its surroundings or cools them. That contrast is the difference between a hand warmer and an instant cold pack.
The first part of C5 is for everyone: exothermic and endothermic reactions, how to measure a temperature change, and how to read a reaction profile. The rest is Triple Chemistry only, looking at how chemical cells, batteries and fuel cells get a reaction’s energy out as electricity rather than heat.
Not all of C5 is for everyone — check what your tier needs:
- Everyone — exothermic and endothermic reactions and their everyday uses, the required practical on temperature changes, and reaction profiles and activation energy.
- Higher Tier H — calculating the overall energy change of a reaction from bond energies.
- Triple (Chemistry only) T — chemical cells and batteries, and hydrogen fuel cells. (The electrode half equations within these are Higher Tier.)
1Exothermic & 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 | Examples |
|---|---|---|---|
| Exothermic | Given out to surroundings | Increases (warms up) | Combustion (burning fuels), many oxidation reactions, neutralisation, displacement of a metal from its salt solution |
| Endothermic | Taken in from surroundings | Decreases (cools down) | Thermal decomposition, 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, the specification 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).
2Measuring 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.
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 the 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
3Reaction Profiles
For a reaction to happen, particles must collide — and collide with enough energy. The minimum energy that colliding particles must have for a reaction to occur is called the activation energy. A reaction profile (an energy level diagram) is a picture of the energy as the reaction proceeds: it shows the relative energies of the reactants and products, the activation energy, and the overall energy change.
The activation energy is the minimum amount of energy that reacting particles must have in order to react. On a reaction profile it is the “hill” the reactants must climb — the gap from the reactants up to the top of the curve.
Reading a profile
Two arrows do all the work, and they are always drawn the same way:
- The activation energy is measured from the reactants up to the top of the curve (the peak).
- The overall energy change — also called the enthalpy change (ΔH) — is measured from the reactants across to the products.
Whether the reaction is exothermic or endothermic — energy given out or taken in, first defined back in section 1 — is shown by where the products sit relative to the reactants:
Exothermic: the products are lower in energy than the reactants, so energy is given out to the surroundings — the enthalpy change ΔH is negative.
Endothermic: the products are higher in energy than the reactants, so energy is taken in from the surroundings — the enthalpy change ΔH is positive.
The overall energy change is the enthalpy change, written ΔH. The symbol Δ is the Greek capital letter delta, and in science it always means “change in” — so ΔH simply reads as the change in enthalpy (enthalpy being the energy stored in the chemicals). It is the difference between the energy of the products and the energy of the reactants. It is negative for an exothermic reaction (products lower than the reactants, so the energy fell) and positive for an endothermic one (products higher, so the energy rose). (The term enthalpy, the symbol ΔH and the +/− sign convention are A-level terminology, beyond the GCSE spec — a handy way to label the overall energy change, but at GCSE you only need to describe a reaction as exothermic or endothermic and give the direction of the energy change.)
Questions often mark three arrows on a profile and ask you to identify or calculate them:
- the activation energy — reactants up to the peak;
- the enthalpy change ΔH — reactants across to the products;
- the energy released as the products form — the peak down to the products.
They are linked by simple arithmetic: for an exothermic reaction, the drop from the peak to the products = the activation energy + the size of ΔH. So if you are given any two arrows you can find the third — for example, the overall energy change = (peak-to-products) − (activation energy).
- Label the axes: energy up the side, progress of reaction along the bottom.
- Draw a flat reactants level and a flat products level — products lower for exothermic, higher for endothermic.
- Join them with a curved line that goes up over a hump.
- Add a labelled arrow for the activation energy (reactants → peak) and one for the overall energy change (reactants → products).
🧪 Exam-style questions
The activation energy of a reaction is best described as…
On a reaction profile, the products are drawn at a higher energy level than the reactants. What does this tell you about the reaction?
4Bond Energy Calculations Higher
For Higher Tier you have to explain where a reaction’s energy change comes from, and calculate it. It all comes down to bonds: during a reaction the bonds in the reactants must break, and new bonds form in the products. Whether the reaction turns out exothermic or endothermic — giving out or taking in energy overall, as we set out in section 1 — comes down to which of those two steps transfers more energy.
Breaking bonds takes energy in (endothermic). Making bonds gives energy out (exothermic). The overall energy change — the enthalpy change ΔH you met on the reaction profile in section 3 — is the difference between the two:
overall energy change = (energy in to break bonds) − (energy out making bonds)
If more energy is released making bonds than was taken in breaking them, the reaction is exothermic (a negative overall change). If more energy is needed to break bonds than is released making them, it is endothermic (a positive change).
The method
You will be given a balanced equation, the displayed (structural) formulae and a table of bond energies (sometimes called bond dissociation energies) in kJ/mol. Then:
- Count every bond broken in the reactants and add up their bond energies — this is the energy in.
- Count every bond made in the products and add up their bond energies — this is the energy out.
- Overall change = energy in − energy out. A negative answer means exothermic.
Worked example — the combustion of methane
Methane burns completely in oxygen. Rather than read the answer off, build it yourself: click every bond in the reactants and products below. Each click adds the energy taken in to break that bond, or given out when it forms — and the two totals decide whether the reaction is exo- or endothermic.
CH4 + 2O2 → CO2 + 2H2O
| Bond | Energy |
|---|---|
| C–H | 413 |
| O=O | 498 |
| C=O | 799 |
| O–H | 463 |
The overall change is negative, so the reaction is exothermic — as we’d expect for burning a fuel. About 802 kJ is given out for every mole of methane burned.
Click each bond in the molecules to count it. 0 of 12 bonds counted.
- Forgetting to multiply by the big numbers in the equation — 2O2 means two O=O bonds, and 2H2O means four O–H bonds.
- Counting double bonds as single — a C=O bond has its own bond energy; it is not two C–O bonds.
- Subtracting the wrong way round. It is always reactants (in) − products (out); a negative answer is exothermic.
🧪 Exam-style questions
Which statement about bonds is correct?
Hydrogen reacts with chlorine: H2 + Cl2 → 2HCl. Use the bond energies to calculate the overall energy change.Give your answer in kJ/mol, including the + or − sign.
Show answer
- Bonds broken (energy in): 1 × H–H + 1 × Cl–Cl = 436 + 242 = 678 kJ/mol. 1 mark
- Bonds made (energy out): 2 × H–Cl = 2 × 431 = 862 kJ/mol. 1 mark
- Overall change = 678 − 862 = −184 kJ/mol — negative, so exothermic. 1 mark
5Cells & Batteries Triple Only
An exothermic reaction normally releases its energy as heat. A cell is a clever way of getting that energy out as electricity instead: it contains chemicals that react to produce a voltage, which can push a current around a circuit. The simplest cell is just two different metals (the electrodes) dipped into an electrolyte — a liquid or solution containing ions that are free to move — and connected by a wire.
A simple cell: two different metals in an electrolyte. The bigger the difference in reactivity between the metals, the bigger the voltage.
The voltage of a simple cell depends on the type of electrodes (the two metals) and the electrolyte. The key rule for the metals:
- The bigger the difference in reactivity between the two metals, the bigger the voltage.
- Two of the same metal give a difference of zero, so the voltage is 0 V.
This means you can use cell voltages to put metals in order of reactivity — the pairs that are furthest apart in the reactivity series give the largest readings.
From cells to batteries
A single cell only gives a small voltage. A battery is simply two or more cells connected together in series, which adds their voltages up to give a greater total. Four 1.5 V cells in series, for example, make a 6.0 V battery.
Rechargeable or not?
- In a non-rechargeable cell or battery (such as an alkaline battery), the reaction stops when one of the reactants is used up. It cannot be recharged because the reaction is not reversible.
- A rechargeable cell or battery (such as a lithium-ion battery) can be recharged because passing an external current through it reverses the reactions, turning the products back into reactants.
This part of the course is deliberately light-touch. You need that a cell is two different metals (electrodes) in an electrolyte whose chemicals react to produce electricity, that a bigger difference in reactivity gives a bigger voltage, and that a cell goes flat once a reactant is used up — unless an external current can reverse the reaction (rechargeable). You don’t need the detail of what happens at each electrode: in AQA’s words, students “do not need to know details of cells and batteries other than those specified.” The Higher-Tier note below gives an optional first look, but the full picture — oxidation and reduction at the electrodes, electrode potentials and so on — is electrochemistry you’ll meet at A‑level.
In a simple cell the reaction is redox. The more reactive metal loses electrons (it is oxidised) and is the negative electrode; the less reactive metal gains them (it is reduced). For a zinc–copper cell:
Zn → Zn2+ + 2e− (oxidation — the negative electrode)
Cu2+ + 2e− → Cu (reduction — the positive electrode)
🧪 Exam-style questions
Which combination would produce a voltage greater than zero?
A single cell produces 1.5 V. How many of these cells must be connected in series to make a 12 V battery?
Show answer
- Cells in series add their voltages, so divide the target by one cell’s voltage: 12 ÷ 1.5 = 8 cells. 2 marks
An alkaline battery is non-rechargeable. Why does it eventually stop working?
6Hydrogen Fuel Cells Triple Only
A fuel cell is different from an ordinary cell: instead of running down when its chemicals are used up, it is continuously supplied with fuel from outside. In a hydrogen fuel cell, hydrogen and oxygen (from the air) are fed in, the hydrogen is oxidised electrochemically, and this produces a voltage (a potential difference). The only product is water.
hydrogen + oxygen → water 2H2 + O2 → 2H2O
The diagram below runs as an interactive. Choose an acidic or an alkaline electrolyte and watch the cell work: each side has an entry tube at the top and an exit tube below — hydrogen and oxygen are fed in, the gas that doesn’t react leaves again, and the product water leaves with it.
Choose an electrolyte to start the fuel cell — then watch the ions cross the electrolyte, electrons drive the lamp, and water leave at the bottom.
In the animation the lamp is only there to show that the circuit is complete and that a current is flowing. The fuel cell itself is just a power supply — like a battery — so it could drive any electrical component instead: a motor to turn a car’s wheels, a phone to charge, a heater, and so on. The lamp simply makes the current easy to see.
Back in electrolysis we said the anode was positive — so why is it the negative terminal here? The trap is assuming that “anode” means “the positive electrode”. It doesn’t. Anode just means the electrode where oxidation happens (and cathode the one where reduction happens). Whether that electrode turns out positive or negative depends on whether electricity is being used or generated.
In electrolysis, a power supply forces electrons onto the cathode. So:
- the cathode is negative — it receives electrons from the power supply;
- the anode is positive — electrons are pulled away from it by the power supply.
At the anode, negative ions lose those electrons — for example, in a chloride:
2Cl− → Cl2 + 2e−
Losing electrons is oxidation, so oxidation happens at the anode.
In a cell, nothing forces the electrons: the cell generates electricity, because the chemical reaction itself pushes electrons around the circuit. Hydrogen is oxidised at the anode:
H2 → 2H+ + 2e−
Those electrons are produced at the anode, so the anode becomes the negative terminal. They then flow through the wire to the positive cathode — the very current that lights the lamp above.
You will not be expected to use the terms “anode” and “cathode” for electrochemical cells at GCSE — but if the labels ever seem to flip, this is why.
The reaction is redox: hydrogen is oxidised and oxygen is reduced. With an acidic electrolyte:
Anode (oxidation): 2H2 → 4H+ + 4e−
Cathode (reduction): O2 + 4H+ + 4e− → 2H2O
Add them together and the H+ and e− cancel, leaving the overall 2H2 + O2 → 2H2O.
Some hydrogen fuel cells use an alkaline electrolyte (containing OH− ions) instead of an acidic one. The overall reaction is the same, but the half equations are written using OH−:
Anode (oxidation): H2 + 2OH− → 2H2O + 2e−
Cathode (reduction): O2 + 2H2O + 4e− → 4OH−
It is called an alkaline fuel cell simply because its electrolyte contains OH− ions.
Oxygen is the thing that does the oxidising (the oxidising agent), so in a fuel cell oxygen is itself reduced — never say oxygen is oxidised. And when comparing hydrogen with a fossil fuel, the marks come from the application point: using hydrogen instead of petrol or methane would reduce the rate of climate change, because the only product is water (no carbon dioxide).
Fuel cells vs rechargeable batteries
A favourite exam question asks you to evaluate hydrogen fuel cells against rechargeable batteries — for example to power a car. There are good points on both sides:
| Hydrogen fuel cells — for | Hydrogen fuel cells — against |
|---|---|
| Refuel in minutes (no long recharge); can give a greater range; lighter than a large battery for long journeys; only product is water, so no pollutants at the point of use; no toxic battery chemicals to dispose of | Hydrogen is hard to store and is highly flammable/explosive; few hydrogen filling stations; making hydrogen needs a lot of energy and often uses fossil fuels; expensive to produce |
🧪 Exam-style questions
Balance the overall equation for the reaction in a hydrogen fuel cell.Type a balancing number in each box (leave it as 1 if none is needed), then press Check. Any correct set of numbers is accepted.
Show answer
2H₂ + O₂ → 2H₂O
- Two water molecules need 4 H and 2 O, so you need 2H2 and 1O2. The numbers are 2 : 1 : 2. 1 mark
Which is an advantage of a hydrogen fuel cell over a rechargeable battery for powering a car?
A car manufacturer is choosing between a hydrogen fuel cell and a rechargeable lithium-ion battery to power a new car. Evaluate the use of a hydrogen fuel cell compared with a rechargeable battery.This is an extended-response question. Plan points on both sides and finish with a justified conclusion, then reveal the mark scheme to compare.
Show answer
How it is marked (levels):
- Level 3 (5–6): a clear judgement, supported by a good range of correct reasons on both sides, logically linked.
- Level 2 (3–4): some logically linked reasons; perhaps a simple judgement.
- Level 1 (1–2): relevant points, but not linked together.
Points in favour of the fuel cell: fast to refuel (no long recharge); greater range; lighter than a large battery; only product is water, so no pollutants at the point of use; no toxic chemicals to dispose of at end of life; gives a steady voltage.
Points in favour of the battery: uses energy more efficiently; cheaper to buy and to recharge; charging points are far more common than hydrogen stations; hydrogen is hard to store and is flammable/explosive; hydrogen is often made from fossil fuels using lots of energy.
To reach Level 3 you must end with a justified conclusion — for example, “the fuel cell is better for long-distance driving because of its range and fast refuelling, provided the hydrogen is made using renewable energy.”
- Exothermic: energy out to the surroundings, temperature up, products lower in energy (combustion, oxidation, neutralisation). Endothermic: energy in, temperature down, products higher in energy (thermal decomposition).
- Required practical: measure temperature change in a polystyrene cup with a lid; the cup insulates to reduce energy transfer.
- Displacement & data skills: displacement reactions are exothermic and a more reactive metal gives a bigger temperature rise; plot type of metal as a bar chart; uncertainty = ½ × range; a glass beaker instead of a polystyrene cup is a systematic error.
- Activation energy: the minimum energy needed to react. On a reaction profile it is reactants → peak; the overall energy change — the enthalpy change ΔH — is reactants → products (down/negative for exo, up/positive for endo).
- Bond energies H: breaking bonds is endothermic, making bonds is exothermic. Overall change = energy in (broken) − energy out (made); negative = exothermic.
- Cells T: two different metals in an electrolyte; bigger reactivity difference = bigger voltage. A battery is cells in series. Non-rechargeable reactions are not reversible.
- Fuel cells T: hydrogen + oxygen → water, producing electricity; evaluate against rechargeable batteries.