Exothermic and Endothermic ReactionsCambridge IGCSE Chemistry: Revision notes
Section 1
What is the difference between exothermic and endothermic reactions?
Chemical reactions transfer thermal energy between the reaction and its surroundings. This transfer determines whether a reaction is exothermic or endothermic.
| Type | Energy Transfer | Temperature Change | Definition |
|---|---|---|---|
| Exothermic | Transfers energy TO surroundings | Surroundings warm up (ΔT increases) | Energy is released; the system loses energy |
| Endothermic | Takes energy FROM surroundings | Surroundings cool down (ΔT decreases) | Energy is absorbed; the system gains energy |
Common examples:
- Exothermic: combustion, neutralisation reactions, respiration, condensation
- Endothermic: melting ice, evaporation, dissolution of some salts, photosynthesis
The key distinction is the direction of energy flow: exothermic reactions release energy (felt as heat), whilst endothermic reactions absorb energy (causing cooling).
Think of exothermic reactions as a person giving away money (energy leaves the system), making their bank balance (temperature) go down, whilst endothermic reactions are like withdrawing money from a shared bank (energy leaves surroundings), reducing the shared balance (cooling the surroundings).
Students often confuse 'the system gets hot' with 'exothermic'. In an exothermic reaction, the surroundings warm up. If a reaction vessel becomes hot to touch, it is exothermic.
Section 2
What is enthalpy change and how is it represented?
Enthalpy change (ΔH) is the term used to describe the transfer of thermal energy during a chemical reaction. It is a measure of the overall energy change of the reaction.
Sign convention:
- Negative ΔH (ΔH < 0) = exothermic reaction (energy released)
- Positive ΔH (ΔH > 0) = endothermic reaction (energy absorbed)
Units: Enthalpy change is measured in kJ/mol (kilojoules per mole of reaction).
Examples:
- Combustion of methane: ΔH = −890 kJ/mol (exothermic; negative sign)
- Thermal decomposition of calcium carbonate: ΔH = +178 kJ/mol (endothermic; positive sign)
The magnitude (size) of ΔH tells you how much energy is released or absorbed. A larger negative value means more energy is released; a larger positive value means more energy is absorbed.
Examiners expect you to write ΔH with the correct sign and units. Always write 'ΔH = −890 kJ/mol', not just '890 kJ/mol'. The sign is part of the answer.
Section 3
What is activation energy and how is it shown on reaction pathway diagrams?
Activation energy (Ea) is the minimum energy that colliding particles must possess in order for a reaction to occur. Even if a reaction is exothermic overall, particles must have enough kinetic energy to break existing bonds and start the reaction.
Key points about activation energy:
- Every reaction has an activation energy barrier that must be overcome
- Particles with less energy than Ea will not react, no matter how many times they collide
- Catalysts lower the activation energy (making reactions faster) but do not change ΔH
- Activation energy is always positive and is measured from the energy level of the reactants
How to identify Ea on a reaction pathway diagram:
- Identify the energy level of the reactants (starting point)
- Identify the highest point on the curve (the activation energy peak)
- Activation energy is the vertical distance from reactants to the peak
How to identify ΔH on a reaction pathway diagram:
- Find the energy level of the reactants (starting point)
- Find the energy level of the products (ending point)
- ΔH is the vertical distance from reactants to products
- If products are lower: ΔH is negative (exothermic)
- If products are higher: ΔH is positive (endothermic)
On a reaction pathway diagram, Ea is always measured from the reactants to the peak (transition state), not from the products. If you measure from the wrong starting point, your answer will be incorrect.
For an exothermic reaction shown on a pathway diagram: reactants at 400 kJ/mol, peak at 550 kJ/mol, products at 350 kJ/mol. Ea = 550 − 400 = 150 kJ/mol. ΔH = 350 − 400 = −50 kJ/mol (negative = exothermic).
Section 4
How do bond breaking and bond making explain enthalpy changes?
The overall enthalpy change of a reaction can be explained by considering the energy required to break bonds and the energy released when new bonds form.
Bond breaking:
- Breaking bonds requires energy (endothermic process)
- This energy input is always positive in energy terms
- The energy required depends on the type and number of bonds broken
Bond making:
- Making bonds releases energy (exothermic process)
- This energy release is always negative in energy terms
- The energy released depends on the type and number of bonds formed
Overall enthalpy change:
- ΔH = Energy to break bonds + Energy released from making bonds
- ΔH = (Energy in) + (Energy out)
- If more energy is released making bonds than needed to break bonds: ΔH is negative (exothermic)
- If more energy is needed to break bonds than released making bonds: ΔH is positive (endothermic)
In summary:
- Exothermic reactions: stronger/more bonds are made than broken
- Endothermic reactions: stronger/more bonds must be broken than are made
Think of reactions like a house renovation: breaking down walls costs money (energy in = endothermic), whilst building new walls and installing fixtures earns you rental income (energy out = exothermic). If the rental income exceeds the demolition cost, you make a profit (exothermic; ΔH negative).
Students sometimes think bond breaking releases energy. It does not—breaking bonds always requires energy input. This is why the equation uses 'energy of bonds broken' as a positive (energy in).
Section 5
How do you calculate enthalpy change using bond energies?
Formula for calculating ΔH using bond energies:
ΔH = Total energy of bonds broken − Total energy of bonds made
Or equivalently:
ΔH = Energy required to break reactant bonds − Energy released when forming product bonds
Step-by-step method:
- Identify all bonds in the reactants and products from the chemical equation
- Count the number of each type of bond (e.g. C–C, C–H, O=O)
- Use a bond energy table to find the energy per mole for each bond type
- Calculate total energy of bonds broken: number of each bond × bond energy (sum all bonds in reactants)
- Calculate total energy of bonds made: number of each bond × bond energy (sum all bonds in products)
- Apply the formula: ΔH = (bonds broken) − (bonds made)
- Interpret the sign: negative = exothermic; positive = endothermic
Worked example:
For the reaction: H₂ + Cl₂ → 2HCl
- Bonds broken: 1 × H–H (432 kJ/mol) + 1 × Cl–Cl (244 kJ/mol) = 676 kJ/mol
- Bonds made: 2 × H–Cl (432 kJ/mol) = 864 kJ/mol
- ΔH = 676 − 864 = −188 kJ/mol (exothermic)
This means the reaction releases 188 kJ/mol of energy.
Always show your working clearly: list bonds broken, list bonds made, multiply by coefficients from the balanced equation, then subtract. Examiners award marks for method, not just the final answer.
For CH₄ + 2O₂ → CO₂ + 2H₂O: Bonds broken in reactants = 4(C–H) + 2(O=O); Bonds made in products = 2(C=O) + 4(O–H). Use bond energy values to calculate total energy for each, then ΔH = (broken) − (made).
Must Know
- Exothermic reactions release thermal energy to surroundings (ΔH negative; temperature rises); endothermic reactions absorb thermal energy from surroundings (ΔH positive; temperature falls)
- Enthalpy change (ΔH) is measured in kJ/mol with a sign that indicates reaction type: negative for exothermic, positive for endothermic
- Activation energy (Ea) is the minimum energy colliding particles must have to react; it is shown on reaction pathway diagrams as the height from reactants to the peak of the curve
- On a reaction pathway diagram, ΔH is the vertical distance from reactants to products (negative if products are lower; positive if products are higher)
- Bond breaking requires energy (endothermic); bond making releases energy (exothermic); ΔH depends on whether more energy is needed to break bonds or more is released making bonds
- To calculate ΔH using bond energies: ΔH = (sum of bond energies broken) − (sum of bond energies made); negative result = exothermic, positive = endothermic
That's the notes covered.
Carry on to the next subtopic.