Contents
1. Exothermic and Endothermic Reactions
A reaction that releases energy to the surroundings — the temperature of the surroundings increases. Energy is released when bonds form; more energy is released forming bonds in products than is absorbed breaking bonds in reactants. ΔH is negative.
A reaction that absorbs energy from the surroundings — the temperature of the surroundings decreases. More energy is absorbed breaking bonds in reactants than is released forming bonds in products. ΔH is positive.
| Exothermic examples | Endothermic examples |
|---|---|
| Combustion of fuels | Thermal decomposition of calcium carbonate |
| Neutralisation (acid + alkali) | Dissolving ammonium nitrate in water |
| Oxidation of metals (rusting) | Photosynthesis |
| Respiration | Cracking of alkanes |
2. Energy Profile Diagrams
An energy profile diagram shows how energy changes as reactants are converted to products.
- Activation energy: the energy barrier — the difference between the reactants energy level and the peak of the curve.
- Exothermic profile: products are at a lower energy level than reactants. The peak is above reactants; products are below reactants.
- Endothermic profile: products are at a higher energy level than reactants.
- With catalyst: the peak is lower (activation energy is reduced) — the reactants and products energy levels are unchanged.
On an energy profile, a catalyst lowers the peak of the curve only — the starting energy of reactants and the finishing energy of products stay exactly the same. Draw the catalysed curve as a lower, flatter hump.
3. Bond Energy Calculations
Breaking bonds requires energy (endothermic). Forming bonds releases energy (exothermic).
H₂ + Cl₂ → 2HCl. Bond energies: H−H = 436 kJ/mol; Cl−Cl = 242 kJ/mol; H−Cl = 431 kJ/mol.
Energy absorbed (bonds broken): 436 + 242 = 678 kJ/mol
Energy released (bonds formed): 2 × 431 = 862 kJ/mol
ΔH = 678 − 862 = −184 kJ/mol (exothermic)
4. Calorimetry
Calorimetry measures heat released or absorbed in a reaction using a known mass of water and a thermometer.
50 cm³ of acid reacts with alkali. Temperature rises from 21.0°C to 27.4°C. Find the heat released.
m = 50 g (assume density of solution ≈ 1 g/cm³)
ΔT = 27.4 − 21.0 = 6.4°C
Q = 50 × 4.2 × 6.4 = 1344 J = 1.344 kJ
Heat loss to the surroundings (use insulated cup); assuming density of solution = 1 g/cm³; thermometer precision. These are the standard improvements asked in Paper 5 questions.
- Exothermic: temperature rises, delta H negative, products lower energy than reactants. e.g. combustion, neutralisation, oxidation.
- Endothermic: temperature falls, delta H positive, products higher energy than reactants. e.g. thermal decomposition, photosynthesis.
- Activation energy = minimum KE needed for successful collision. Catalyst lowers activation energy.
- Bond breaking: ENDOTHERMIC (energy IN). Bond forming: EXOTHERMIC (energy OUT).
- Net delta H = energy to break bonds - energy released forming bonds. Negative = exothermic.
- Calorimetry: Q = mc(delta)T. Water c = 4.2 J/(g degC). Calculate moles then energy per mole.
5. Common Exam Traps
This is always true — never reversed. "Energy is released when bonds are broken" is always wrong. Energy is absorbed to break bonds and released when bonds form.
A negative ΔH means the products are at lower energy than the reactants — energy has been released to the surroundings. Students often say "ΔH is negative so it is endothermic" — this is the opposite of the truth.
A catalyst provides a different pathway but the overall energy change (ΔH) between reactants and products remains identical. Only the activation energy changes.
Key Terms — Flashcard Review
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Original study notes for Singapore students. Not affiliated with MOE, SEAB or Cambridge.