Every chemical reaction involves an energy change. Understanding whether energy is released or absorbed — and how much — is essential for everything from industrial chemistry to understanding how our bodies burn food.
9.1 Exothermic and Endothermic Reactions
Definition
Exothermic reaction
An exothermic reaction is one that releases energy in the form of heat or light. Another way of describing an exothermic reaction is that it is one in which the energy of the products is less than the energy of the reactants, because energy has been released during the reaction.
Definition
Endothermic reaction
An endothermic reaction is one that absorbs energy from the surroundings in the form of heat. The energy of the products is greater than the energy of the reactants (ΔH > 0). The temperature of the surroundings decreases during an endothermic reaction.
Definition
Enthalpy
Enthalpy is a measure of the total energy of a chemical system for a given pressure, symbol H.
Definition
Activation energy (Ea)
Activation energy (Ea) is the minimum energy that colliding reactant particles must possess for a reaction to occur. On an enthalpy diagram, Ea is the difference between the energy of the reactants and the peak of the energy barrier (the transition state). A catalyst lowers Ea without being consumed.
A potential energy diagram (reaction energy profile) shows how the potential energy of reactants and products changes during a reaction. The peak of the curve is the transition state — the activation energy (Ea) is the height of that peak above the reactants.
Exothermic vs Endothermic Reactions
| Property | Exothermic | Endothermic |
|---|---|---|
| Energy change | Releases energy (heat/light) | Absorbs energy (heat/light) |
| ΔH sign | ΔH < 0 (negative) | ΔH > 0 (positive) |
| Product energy | Lower than reactants | Higher than reactants |
| Temperature of surroundings | Increases | Decreases |
| Example | Combustion, neutralisation | Photosynthesis, dissolving NH₄NO₃ |
Exam Tip
Remember ΔH = H(products) − H(reactants). If products have less energy, ΔH is negative → exothermic. If products have more energy, ΔH is positive → endothermic.
Practice Question
A reaction has ΔH = −285 kJ·mol⁻¹. Is this exothermic or endothermic? What happens to the temperature of the surroundings?
(3 marks)
9.2 Bond Energy and ΔH Calculations
Definition
Bond energy
Bond energy is a measure of bond strength in a chemical bond. It is the amount of energy (in kJ·mol⁻¹) needed to break the chemical bond between two atoms.
Breaking bonds always REQUIRES energy (endothermic step). Forming bonds always RELEASES energy (exothermic step). The overall ΔH of a reaction is the difference: energy in (breaking) minus energy out (forming). If more energy is released than absorbed, the reaction is exothermic overall.
Formula
ΔH from bond energies
ΣE(bonds broken) = total bond energy absorbed to break all bonds in reactants (kJ·mol⁻¹); ΣE(bonds formed) = total bond energy released forming all bonds in products (kJ·mol⁻¹)
SI unit: kJ·mol⁻¹
Note
A catalyst lowers the activation energy (the peak on the PE diagram) by providing an alternative reaction pathway. It does NOT change the enthalpy of reactants or products, so ΔH is unaffected.
Worked Example
Calculate ΔH for H₂ + Cl₂ → 2HCl using bond energies: H–H = 436 kJ·mol⁻¹, Cl–Cl = 242 kJ·mol⁻¹, H–Cl = 431 kJ·mol⁻¹.
Given
- Bond energies: H–H = 436, Cl–Cl = 242, H–Cl = 431 kJ·mol⁻¹
- Equation: H₂ + Cl₂ → 2HCl
Find
ΔH
Solution
- 1Bonds broken: 1(H–H) + 1(Cl–Cl) = 436 + 242 = 678 kJ·mol⁻¹
- 2Bonds formed: 2(H–Cl) = 2 × 431 = 862 kJ·mol⁻¹
- 3ΔH = 678 − 862 = −184 kJ·mol⁻¹
Watch Out
A very common mistake: students subtract bonds formed MINUS bonds broken and get the wrong sign. Always use: ΔH = Σ(broken) − Σ(formed). A negative ΔH means exothermic.
Practice Question
Calculate ΔH for N₂ + 3H₂ → 2NH₃ using bond energies: N≡N = 945 kJ·mol⁻¹, H–H = 436 kJ·mol⁻¹, N–H = 391 kJ·mol⁻¹.
(6 marks)
Practice Question
On a potential energy diagram for an exothermic reaction, label: reactants, products, transition state, Ea (forward), Ea (reverse), and ΔH. Where would the products line move if a catalyst is added?
(5 marks)
9.3 Interpreting Energy Profiles
You need to be able to draw PE diagrams from scratch. For an exothermic reaction: draw the reactants higher on the y-axis, a peak in the middle, and products LOWER than reactants. For an endothermic reaction the products line is HIGHER than the reactants line.
What to label on a PE diagram
- y-axis: 'Potential energy / kJ·mol⁻¹'
- x-axis: 'Reaction coordinate' (or 'Progress of reaction')
- Reactants line (with label)
- Products line (with label)
- Transition state (peak)
- Ea (forward) — from reactants to peak
- Ea (reverse) — from products to peak
- ΔH — vertical distance between reactants and products
Exam Tip
In an exam: if the question says 'exothermic', draw products BELOW reactants and use a negative ΔH arrow going DOWN. If 'endothermic', products are ABOVE reactants and ΔH arrow goes UP.
Worked Example
A reaction has Ea(forward) = 150 kJ·mol⁻¹ and ΔH = −80 kJ·mol⁻¹. Calculate Ea(reverse).
Given
- Ea(forward) = 150 kJ·mol⁻¹
- ΔH = −80 kJ·mol⁻¹
Find
Ea(reverse)
Solution
- 1Ea(reverse) = Ea(forward) − ΔH
- 2Ea(reverse) = 150 − (−80)
- 3Ea(reverse) = 150 + 80 = 230 kJ·mol⁻¹
Practice Question
An endothermic reaction has Ea(forward) = 200 kJ·mol⁻¹ and ΔH = +50 kJ·mol⁻¹. A catalyst lowers Ea(forward) by 40 kJ·mol⁻¹. What is the new Ea(reverse) with the catalyst?
(5 marks)
Real World
Catalytic converters in cars use platinum and palladium catalysts to lower the activation energy of reactions that convert toxic CO and NOₓ exhaust gases into less harmful CO₂ and N₂.