C15: Exothermic and Endothermic Reactions
Energy changes and reaction profiles
📋 Key Definitions
Exothermic reaction: A reaction that releases energy to the surroundings, causing the temperature of the surroundings to increase. Energy is transferred from the reactants to the surroundings.
Endothermic reaction: A reaction that absorbs energy from the surroundings, causing the temperature of the surroundings to decrease. Energy is transferred from the surroundings to the reactants.
Activation energy: The minimum amount of energy that particles must have to react when they collide. Shown on reaction profiles as the energy "hill" from reactants to the peak.
Bond energy: The amount of energy needed to break a particular chemical bond, or the amount released when the bond forms. Breaking bonds is endothermic; making bonds is exothermic.
Bond Energy Calculation:
Energy change = energy absorbed (breaking bonds) − energy released (making bonds)
If the value is negative → exothermic
If the value is positive → endothermic
📝 Exothermic vs Endothermic
| Feature | Exothermic | Endothermic |
|---|---|---|
| Energy transfer | Released to surroundings | Absorbed from surroundings |
| Temperature change | Surroundings get warmer | Surroundings get cooler |
| Energy of products vs reactants | Products have less energy than reactants | Products have more energy than reactants |
| Energy change value | Negative (e.g. −571 kJ/mol) | Positive (e.g. +571 kJ/mol) |
| Everyday examples | Hand warmers, self-heating cans | Cold packs, sports injury packs |
| Chemical examples | Combustion, neutralisation, oxidation | Thermal decomposition, photosynthesis |
📈 Reaction Profiles
A reaction profile (energy level diagram) shows the energy of reactants and products, and the activation energy.
Reading reaction profiles:
- Exothermic: Products are at a lower energy level than reactants. The arrow showing energy change points downwards.
- Endothermic: Products are at a higher energy level than reactants. The arrow showing energy change points upwards.
- Activation energy: The arrow from reactants up to the peak of the curve. Both exothermic and endothermic reactions have activation energy.
Worked Example 1: Identifying reaction type from a profile
A reaction profile shows the products at a lower energy level than the reactants. Is this exothermic or endothermic?
Since the products have less energy than the reactants, energy has been released to the surroundings.
This is an exothermic reaction.
⚡ Bond Energy Calculations
Every chemical bond has a bond energy — the energy required to break it. When new bonds form, energy is released. The overall energy change depends on the balance between energy needed to break bonds and energy released when new bonds form.
Steps for bond energy calculations:
1. Calculate total energy required to break all bonds in the reactants
2. Calculate total energy released when all new bonds form in the products
3. Energy change = energy in (breaking) − energy out (making)
Worked Example 2: Bond energy calculation for H₂ + Cl₂ → 2HCl
Bond energies: H−H = 436 kJ/mol, Cl−Cl = 242 kJ/mol, H−Cl = 431 kJ/mol
Step 1: Energy to break bonds = H−H + Cl−Cl = 436 + 242 = 678 kJ/mol
Step 2: Energy released making bonds = 2 × H−Cl = 2 × 431 = 862 kJ/mol
Step 3: Energy change = 678 − 862 = −184 kJ/mol
Negative value → exothermic. More energy released than absorbed.
Worked Example 3: Bond energy calculation for CH₄ + 2O₂ → CO₂ + 2H₂O
Bond energies: C−H = 413, O=O = 498, C=O = 805, O−H = 464 (all kJ/mol)
Step 1: Energy to break bonds = (4 × C−H) + (2 × O=O) = (4 × 413) + (2 × 498) = 1652 + 996 = 2648 kJ/mol
Step 2: Energy released making bonds = (2 × C=O) + (4 × O−H) = (2 × 805) + (4 × 464) = 1610 + 1856 = 3466 kJ/mol
Step 3: Energy change = 2648 − 3466 = −818 kJ/mol
Negative → exothermic (combustion is always exothermic).
Worked Example 4: Endothermic bond energy calculation
For the reaction H₂ + I₂ → 2HI, calculate the energy change.
Bond energies: H−H = 436, I−I = 151, H−I = 299 (all kJ/mol)
Step 1: Energy to break bonds = 436 + 151 = 587 kJ/mol
Step 2: Energy released making bonds = 2 × 299 = 598 kJ/mol
Step 3: Energy change = 587 − 598 = −11 kJ/mol
This is slightly exothermic (very close to zero).
🔥 Common Exothermic Reactions
Worked Example 5: Exothermic reactions
Combustion: Fuel + oxygen → carbon dioxide + water. E.g. CH₄ + 2O₂ → CO₂ + 2H₂O. Releases lots of heat and light.
Neutralisation: Acid + alkali → salt + water. E.g. HCl + NaOH → NaCl + H₂O. Temperature of the solution rises.
Oxidation: E.g. adding water to calcium oxide: CaO + H₂O → Ca(OH)₂. Used in self-heating cans.
❄️ Common Endothermic Reactions
Worked Example 6: Endothermic reactions
Thermal decomposition: Heating calcium carbonate: CaCO₃ → CaO + CO₂. Heat must be continuously supplied.
Photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂. Energy from sunlight is absorbed.
Sports injury packs: Ammonium nitrate dissolving in water — absorbs heat, making the pack feel cold.
❓ Practice Questions
Q1: Foundation Define exothermic and endothermic reactions.
Q2: Foundation Classify each of the following as exothermic or endothermic: combustion, thermal decomposition, neutralisation, photosynthesis.
Q3: Higher For the reaction N₂ + 3H₂ → 2NH₃, calculate the energy change. Bond energies: N≡N = 945, H−H = 436, N−H = 391 (all kJ/mol).
Q4: Foundation Describe how the reaction profile of an exothermic reaction differs from that of an endothermic reaction.
Q5: Higher Explain, in terms of bond breaking and bond making, why combustion reactions are exothermic.
✅ Answers
- Exothermic: releases energy to surroundings, temperature increases. Endothermic: absorbs energy from surroundings, temperature decreases.
- Combustion = exothermic. Thermal decomposition = endothermic. Neutralisation = exothermic. Photosynthesis = endothermic.
- Energy in (breaking): N≡N + 3(H−H) = 945 + (3 × 436) = 945 + 1308 = 2253 kJ/mol. Energy out (making): 6 × N−H (because 2NH₃ has 6 N−H bonds) = 6 × 391 = 2346 kJ/mol. Energy change = 2253 − 2346 = −93 kJ/mol (exothermic).
- Exothermic: products at lower energy level than reactants, energy change arrow points down. Endothermic: products at higher energy level, energy change arrow points up. Both have activation energy (the hill from reactants to the peak).
- In combustion, the bonds formed in the products (C=O and O−H) release more energy than was needed to break the bonds in the reactants (C−H and O=O). Since more energy is released making bonds than absorbed breaking bonds, the reaction is exothermic.
🎯 Exam Tips
- In bond energy calculations, always count the number of each type of bond carefully — check the formulae
- A negative energy change means exothermic; a positive value means endothermic
- Breaking bonds is always endothermic (energy in); making bonds is always exothermic (energy out)
- On reaction profiles, the activation energy arrow always goes upwards from the reactants line
- Remember that both exothermic AND endothermic reactions need activation energy to get started
- When asked to sketch a reaction profile, label the axes (energy on y-axis, reaction progress on x-axis), the activation energy, and the overall energy change
🔢 Maths Skills
Mathematical Skills
Bond energy calculations: 1) Calculate total energy to break all bonds in reactants (energy in). 2) Calculate total energy released making all bonds in products (energy out). 3) Energy change = energy in − energy out. Negative = exothermic; positive = endothermic.
Example: H₂ + Cl₂ → 2HCl. Energy in = 436 + 242 = 678 kJ/mol. Energy out = 2 × 431 = 862 kJ/mol. ΔH = 678 − 862 = −184 kJ/mol (exothermic).
⚠️ Common Misconceptions
Watch Out!
Breaking bonds releases energy. Wrong: breaking bonds releases energy Correct: breaking bonds requires energy (endothermic); making bonds releases energy (exothermic)
Exothermic reactions are always fast. Wrong: exothermic reactions are always fast Correct: the speed of a reaction (rate) is independent of the energy change — an exothermic reaction can be slow (e.g. rusting) and an endothermic reaction can be fast
✍️ 6-Mark Question
Extended Answer
6 marks: Explain exothermic and endothermic reactions using bond energies.
In a chemical reaction, bonds in the reactants must first be broken, which requires energy (endothermic). New bonds then form in the products, which releases energy (exothermic). If the energy released when new bonds form is greater than the energy needed to break the original bonds, the reaction is exothermic — more energy is given out than taken in, so the surroundings get warmer and ΔH is negative. If the energy needed to break bonds is greater than the energy released making new bonds, the reaction is endothermic — more energy is taken in than given out, the surroundings get cooler and ΔH is positive. For example, combustion is exothermic because the C=O and O−H bonds formed in CO₂ and H₂O release more energy than was needed to break the C−H and O=O bonds in the reactants.
Mark scheme: 1 mark for bond breaking requires energy; 1 mark for bond making releases energy; 1 mark for exothermic = more energy released than absorbed; 1 mark for endothermic = more energy absorbed than released; 1 mark for correct sign of ΔH; 1 mark for combustion example.
📊 AO3: Analyse & Evaluate
Analysis and Evaluation
A student investigates the reaction between hydrogen and chlorine: H₂ + Cl₂ → 2HCl. Bond energies: H−H = 436, Cl−Cl = 242, H−Cl = 431 kJ/mol.
Question: Calculate the enthalpy change. Is the reaction exothermic or endothermic? A student says "the H−Cl bond is weaker than the H−H bond so the reaction must be endothermic." Evaluate this statement.
Answer: Energy in = 436 + 242 = 678 kJ/mol. Energy out = 2 × 431 = 862 kJ/mol. ΔH = 678 − 862 = −184 kJ/mol (exothermic). The student's statement is incorrect because the overall energy change depends on ALL bonds broken and formed, not just one comparison. Although H−Cl is weaker than H−H, two H−Cl bonds are formed per molecule, and the Cl−Cl bond is much weaker than H−H, so the total energy released making bonds exceeds the total energy needed to break bonds.
📝 Exam Questions by Topic
🎬 Video Resources
Ready to ace your GCSE Combined Science exams?
Get the best revision books and guides to boost your grades.