Reaction 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?
Bond 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