Chemical energetics deals with the energy change between reacting chemicals and their surroundings. An exothermic reaction transfers energy to the surroundings, so the temperature of the mixture rises and ΔH is negative. An endothermic reaction takes in energy from the surroundings, so the temperature falls and ΔH is positive.
It is the shortest topic in the Sec 3 Pure Chemistry syllabus, but its language reappears in questions on combustion of fuels, neutralisation, respiration and rusting. Energy profile diagrams link it to speed of reaction. Bond energy calculations are a guaranteed calculation question in O-Level Paper 2, worth 3–4 marks.
What these notes cover
- •Exothermic and endothermic reactions
- •Sign convention for ΔH
- •Named syllabus examples and their equations
- •Energy profile diagrams
- •Activation energy
- •Effect of a catalyst
- •Bond energy calculations
- •Hydrogen as a fuel
- •The hydrogen–oxygen fuel cell
- •Common mistakes and exam technique
- •Summary
Key definitions
- Exothermic reaction
- A reaction that transfers energy, often as heat, to the surroundings. The surroundings gain energy, so the temperature rises and ΔH is negative.
- Endothermic reaction
- A reaction that takes in energy, often as heat, from the surroundings. The surroundings lose energy, so the temperature falls and ΔH is positive.
- Enthalpy change (ΔH)
- The energy change of a reaction, measured in kJ/mol. It is taken from the reaction's point of view, so energy leaving the chemicals is counted as negative.
- Activation energy (Eₐ)
- The minimum energy that colliding particles must possess for a reaction to occur.
- Bond energy
- The energy needed to break one mole of a particular covalent bond in the gaseous state, measured in kJ/mol. The same amount of energy is released when one mole of that bond forms.
Key formulas
Sign convention
Exothermic: ΔH < 0; Endothermic: ΔH > 0
A temperature rise in the mixture means ΔH is negative; a temperature fall means ΔH is positive.
Enthalpy change from an energy profile
ΔH = energy of products − energy of reactants
Drawn as a vertical arrow from the reactants' level to the products' level: pointing down for exothermic, up for endothermic.
Activation energy from an energy profile
Eₐ = energy at peak − energy of reactants
Always measured from the reactants' level to the top of the peak, never from the horizontal axis or the products.
Enthalpy change from bond energies
ΔH = total energy absorbed breaking bonds − total energy released forming bonds
Multiply each bond energy by the bonds per molecule and the coefficient in the equation. A negative answer means exothermic; a positive answer means endothermic.
Ionic equation for neutralisation
H⁺(aq) + OH⁻(aq) → H₂O(l)
Applies to any strong acid–strong alkali neutralisation once the spectator ions are crossed out. Neutralisation is exothermic.
Combustion of hydrogen
2H₂(g) + O₂(g) → 2H₂O(l), ΔH = −572 kJ
Highly exothermic, and the only product is water. A hydrogen–oxygen fuel cell has the same overall reaction.
Worked example
Nitrogen reacts with hydrogen to form ammonia: N₂(g) + 3H₂(g) → 2NH₃(g). Using the bond energies N≡N = 944 kJ/mol, H–H = 436 kJ/mol and N–H = 391 kJ/mol, calculate ΔH and state whether the reaction is exothermic or endothermic.
- 1List the bonds: the reactants contain one N≡N bond and 3 × 1 = 3 H–H bonds; each NH₃ molecule has three N–H bonds and there are 2 molecules, so the products contain 2 × 3 = 6 N–H bonds.
- 2Energy absorbed breaking bonds = 944 + (3 × 436) = 944 + 1308 = 2252 kJ.
- 3Energy released forming bonds = 6 × 391 = 2346 kJ.
- 4ΔH = energy absorbed − energy released = 2252 − 2346 = −94 kJ/mol.
- 5ΔH is negative, so the forward reaction of the Haber process is exothermic.
Watch out: A coefficient multiplies every bond in that molecule: 2NH₃ contains 2 × 3 = 6 N–H bonds, not 3. Forgetting this is the most common way to lose all the marks on this question type.
Common mistakes
Getting the sign of ΔH backwards
Wrong: The reaction gives out heat, so ΔH is positive because energy is produced.
Right: ΔH is measured from the system's point of view. An exothermic reaction loses energy to the surroundings, so ΔH is negative; an endothermic reaction gains energy from the surroundings, so ΔH is positive.
Saying that breaking bonds releases energy
Wrong: Energy is released when the bonds in the fuel break.
Right: Bond breaking always absorbs energy and bond forming always releases energy. A fuel releases energy overall because the energy released forming bonds in the products is greater than the energy absorbed breaking bonds in the reactants.
Saying a catalyst lowers ΔH
Wrong: The catalyst lowers the energy change of the reaction, or makes the reaction more exothermic.
Right: A catalyst provides an alternative reaction pathway with a lower activation energy. The energy levels of the reactants and products do not change, so ΔH is unchanged.
Questions students ask
How do you tell if a reaction is exothermic or endothermic from the temperature change?
If the temperature of the reaction mixture rises, energy is being released to the surroundings, so the reaction is exothermic and ΔH is negative. If the temperature falls, energy is being absorbed from the surroundings, so the reaction is endothermic and ΔH is positive. Quote the temperature change first, then link it to the direction of energy transfer.
How do you calculate enthalpy change using bond energies?
Use ΔH = total energy absorbed breaking bonds − total energy released forming bonds. List every bond in the reactants and products, multiply each bond energy by the number of bonds including the coefficients, total each side and subtract. Then state the classification from the sign: negative means exothermic, positive means endothermic.
Where is activation energy measured on an energy profile diagram?
Activation energy is measured from the energy level of the reactants up to the top of the peak, never from the horizontal axis or from the products. Measuring from the products to the peak gives the activation energy of the reverse reaction instead.
What are the advantages and disadvantages of hydrogen as a fuel?
Burning hydrogen produces only water, so no CO₂ is released at the point of use, and it releases about 143 kJ per gram, far more than methane or petrol. However, hydrogen is a flammable, very low-density gas that is difficult and hazardous to store. Producing it by electrolysis is also expensive and energy-intensive, so it is only as clean as the electricity used.
In the full PDF
- •Labelled energy profile diagrams for exothermic, endothermic and catalysed reactions, plus a checklist for a full-mark sketch
- •Further worked examples: thermometer readings, the cold pack, and bond energy calculations for H₂ + Cl₂, the combustion of methane and the formation of NO
- •An energy-per-gram comparison of hydrogen and methane, and a model 4-mark 'evaluate' answer on hydrogen as a fuel
- •The remaining misconceptions, including writing heat into equations and misreading a hot container
- •Exam-technique sentence frames for bond-energy explanations, catalysts and the definition of activation energy
- •Section summaries and a full-topic summary