Chemical reactions involve energy transfers to or from the surroundings. Exothermic reactions release heat energy and endothermic reactions absorb heat energy. Understanding bond energies allows you to calculate the overall energy change for a reaction.
Chemical reactions involve the breaking and making of chemical bonds:
The overall energy change of a reaction depends on the difference between the energy needed to break bonds and the energy released when new bonds form.
In an exothermic reaction, more energy is released when new bonds form in the products than is required to break bonds in the reactants. The excess energy is transferred to the surroundings as heat, causing the temperature to rise.
| Feature | Exothermic Reactions | Endothermic Reactions |
|---|---|---|
| Energy transfer | From reactants to surroundings | From surroundings to reactants |
| 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 |
| Overall energy change (ΔH) | Negative (e.g. -571 kJ/mol) | Positive (e.g. +572 kJ/mol) |
| Bond energies | Energy released (making bonds) > Energy needed (breaking bonds) | Energy needed (breaking bonds) > Energy released (making bonds) |
Common examples of exothermic reactions:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
This is an exothermic reaction. Energy is transferred to the surroundings as heat and light. The temperature of the surroundings increases. This is why methane (natural gas) is used as a fuel for heating and cooking.
In an endothermic reaction, more energy is required to break bonds in the reactants than is released when new bonds form in the products. The deficit in energy is taken from the surroundings as heat, causing the temperature to fall.
Common examples of endothermic reactions:
CaCO₃(s) → CaO(s) + CO₂(g)
This is an endothermic reaction. Energy must be continuously supplied (by heating) to keep the reaction going. If you remove the heat source, the reaction stops. The temperature of the surroundings would decrease if no external heat were supplied.
When citric acid is added to sodium hydrogencarbonate, the mixture fizzes (carbon dioxide is produced) and the temperature drops. This is an endothermic reaction — energy is taken in from the surroundings, making the container feel cold to the touch.
| Application | Type | How It Works |
|---|---|---|
| Hand warmers | Exothermic | Iron powder oxidises in air (rusting), releasing heat slowly over several hours |
| Self-heating cans | Exothermic | Calcium oxide reacts with water to form calcium hydroxide, releasing heat |
| Instant cold packs | Endothermic | Ammonium nitrate dissolves in water, absorbing heat from the surroundings |
| Sports injury packs | Endothermic | Endothermic dissolution process cools the injured area |
| Flameless ration heaters | Exothermic | Magnesium powder reacts with water (activated by adding water to the pouch) |
Reaction profiles show the energy changes that occur during a chemical reaction:
For an exothermic reaction profile:
For an endothermic reaction profile:
Even exothermic reactions need activation energy to get started. For example, methane does not spontaneously combust in air — a spark or flame is needed to provide the activation energy. Once the reaction starts, the energy released by the reaction keeps it going.
The combustion of methane is highly exothermic (releases a large amount of energy). However, a mixture of methane and oxygen at room temperature does not react because the molecules do not have enough energy to overcome the activation energy barrier. A spark or flame must be applied to provide the activation energy. Once the reaction begins, the heat released sustains the reaction.
If the result is negative, the reaction is exothermic. If the result is positive, the reaction is endothermic.
Common bond energies (in kJ/mol):
| Bond | Bond Energy (kJ/mol) | Bond | Bond Energy (kJ/mol) |
|---|---|---|---|
| C−H | 435 | O−H | 464 |
| C−C | 347 | O=O | 498 |
| C=C | 612 | H−H | 436 |
| C=O | 805 | N≡N | 945 |
| C−O | 358 | N−H | 391 |
| C−N | 286 | Cl−Cl | 243 |
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
Step 1: Identify bonds broken in the reactants
CH₄: 4 × C−H = 4 × 435 = 1740 kJ/mol
2O₂: 2 × O=O = 2 × 498 = 996 kJ/mol
Total energy to break bonds = 1740 + 996 = 2736 kJ/mol
Step 2: Identify bonds formed in the products
CO₂: 2 × C=O = 2 × 805 = 1610 kJ/mol
2H₂O: 4 × O−H = 4 × 464 = 1856 kJ/mol
Total energy released making bonds = 1610 + 1856 = 3466 kJ/mol
Step 3: Calculate overall energy change
ΔH = Energy broken − Energy made = 2736 − 3466 = -730 kJ/mol
The reaction is exothermic (ΔH is negative) because more energy is released making bonds than is needed to break bonds.
N₂(g) + 3H₂(g) → 2NH₃(g)
Step 1: Bonds broken
N₂: 1 × N≡N = 1 × 945 = 945 kJ/mol
3H₂: 3 × H−H = 3 × 436 = 1308 kJ/mol
Total energy to break bonds = 945 + 1308 = 2253 kJ/mol
Step 2: Bonds formed
2NH₃: 6 × N−H = 6 × 391 = 2346 kJ/mol
Total energy released making bonds = 2346 kJ/mol
Step 3: Calculate overall energy change
ΔH = 2253 − 2346 = -93 kJ/mol
The reaction is exothermic (ΔH is negative).
H₂(g) + Cl₂(g) → 2HCl(g)
Step 1: Bonds broken
H₂: 1 × H−H = 1 × 436 = 436 kJ/mol
Cl₂: 1 × Cl−Cl = 1 × 243 = 243 kJ/mol
Total energy to break bonds = 436 + 243 = 679 kJ/mol
Step 2: Bonds formed
2HCl: 2 × H−Cl = 2 × 432 = 864 kJ/mol
Total energy released making bonds = 864 kJ/mol
Step 3: Calculate overall energy change
ΔH = 679 − 864 = -185 kJ/mol
The reaction is exothermic (ΔH is negative).
The thermal decomposition of hydrogen bromide: 2HBr(g) → H₂(g) + Br₂(g)
Step 1: Bonds broken
2HBr: 2 × H−Br = 2 × 366 = 732 kJ/mol
Total energy to break bonds = 732 kJ/mol
Step 2: Bonds formed
H₂: 1 × H−H = 1 × 436 = 436 kJ/mol
Br₂: 1 × Br−Br = 1 × 193 = 193 kJ/mol
Total energy released making bonds = 436 + 193 = 629 kJ/mol
Step 3: Calculate overall energy change
ΔH = 732 − 629 = +103 kJ/mol
The reaction is endothermic (ΔH is positive) because more energy is needed to break bonds than is released when new bonds form.
For example, the C−H bond energy is given as 435 kJ/mol, but the actual energy needed to break each C−H bond in methane is slightly different. The first C−H bond breaks more easily than the last one. The value 435 kJ/mol is an average across many different molecules containing C−H bonds.
When you calculate the energy change for a reaction using mean bond energies, you may get a slightly different result from the experimentally measured value. This is because:
Where:
25 cm³ of hydrochloric acid is added to 25 cm³ of sodium hydroxide in a polystyrene cup. The temperature rises from 20°C to 27°C.
m = 25 + 25 = 50 g (total volume of solution)
c = 4.2 J/g°C
ΔT = 27 − 20 = 7°C
Q = 50 × 4.2 × 7 = 1470 J = 1.47 kJ
This is the energy released by the reaction and transferred to the solution.
For example, the thermal decomposition of hydrated copper(II) sulfate:
CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -93 kJ/mol
The forward reaction (making ammonia) is exothermic: ΔH = -93 kJ/mol.
The reverse reaction (decomposing ammonia) is endothermic: ΔH = +93 kJ/mol.
The energy absorbed in the reverse reaction exactly equals the energy released in the forward reaction. This is a consequence of the conservation of energy.
1. Explain why breaking bonds is always endothermic and making bonds is always exothermic.
Breaking bonds requires energy input because you need to overcome the attractive forces between the bonded atoms. This means energy is taken in from the surroundings, so it is endothermic. Making bonds releases energy because when atoms come together and form bonds, the system moves to a more stable (lower energy) state, and the excess energy is released to the surroundings, so it is exothermic.
2. Calculate the overall energy change for the combustion of hydrogen: 2H₂(g) + O₂(g) → 2H₂O(g). Bond energies: H−H = 436 kJ/mol, O=O = 498 kJ/mol, O−H = 464 kJ/mol.
Bonds broken: 2 × H−H = 2 × 436 = 872 kJ/mol; 1 × O=O = 498 kJ/mol. Total = 872 + 498 = 1370 kJ/mol.
Bonds made: 4 × O−H = 4 × 464 = 1856 kJ/mol. (2H₂O has 4 O−H bonds)
ΔH = 1370 − 1856 = -486 kJ/mol. The reaction is exothermic.
3. In a calorimetry experiment, 50 cm³ of HCl at 21.0°C was added to 50 cm³ of NaOH at 21.0°C. The final temperature was 34.5°C. Calculate the energy released. (c = 4.2 J/g°C)
m = 50 + 50 = 100 g. ΔT = 34.5 − 21.0 = 13.5°C. Q = 100 × 4.2 × 13.5 = 5670 J = 5.67 kJ.
4. Draw a reaction profile for an endothermic reaction. Label the reactants, products, activation energy, and overall energy change.
The reaction profile should show: the reactants at a lower energy level than the products. A curve rising from the reactants to a peak (transition state) then falling to the products. The activation energy arrow points upward from the reactants to the peak. The overall energy change (ΔH) arrow points upward from the reactants to the products. ΔH is positive.
5. Explain why a catalyst increases the rate of a reaction by referring to activation energy and reaction profiles.
A catalyst provides an alternative reaction pathway with a lower activation energy. On a reaction profile, the catalysed pathway has a lower peak than the uncatalysed pathway. Because the activation energy is lower, more particles have sufficient energy to react when they collide, so a greater proportion of collisions are successful. This increases the rate of reaction. The catalyst does not change the overall energy change (ΔH) of the reaction.
6. Calculate the energy change for the reaction: H₂(g) + F₂(g) → 2HF(g). Bond energies: H−H = 436, F−F = 158, H−F = 562 kJ/mol.
Bonds broken: 1 × H−H = 436 kJ/mol; 1 × F−F = 158 kJ/mol. Total = 436 + 158 = 594 kJ/mol.
Bonds made: 2 × H−F = 2 × 562 = 1124 kJ/mol.
ΔH = 594 − 1124 = -530 kJ/mol. The reaction is exothermic.
Use a polystyrene cup as a calorimeter to reduce heat loss. Add a measured volume of one solution to the cup and record its initial temperature. Add a measured volume of the second solution, stir, and record the highest (or lowest) temperature reached. Calculate the temperature change.
For exothermic reactions (e.g. neutralisation of HCl with NaOH), the temperature increases. For endothermic reactions (e.g. citric acid with sodium hydrogencarbonate), the temperature decreases. Record results in a table showing initial temperature, final temperature, and temperature change.
Control variables: the volumes and concentrations of the solutions, the type of cup, whether the solution is stirred, and the starting temperature. The independent variable is the type of reaction. The dependent variable is the temperature change.
To improve accuracy: use a lid on the cup to reduce heat loss, insulate the cup further, stir thoroughly to distribute heat evenly, and take repeated readings to calculate a mean.
Energy change = sum of bond energies broken − sum of bond energies made. If the result is negative, the reaction is exothermic (more energy released making bonds than needed to break them). If positive, it is endothermic (more energy needed to break bonds than released making them).
Always draw out the displayed formula of each substance to count every bond correctly. For example, CH4 has 4 C−H bonds, and 2H2O has 4 O−H bonds. Do not count bonds that appear in both reactants and products — they cancel out. A common mistake is forgetting that bond energies are per mole of bonds, not per mole of substance.
The vertical axis shows energy, the horizontal axis shows reaction progress. The activation energy is the difference between the reactants and the peak. The overall energy change is the difference between reactants and products. For exothermic reactions, products are lower than reactants; for endothermic, products are higher.
Wrong: Exothermic reactions are always fast Correct: The speed of a reaction and its energy change are different things. Exothermic means energy is released to the surroundings, but it can happen slowly. Rusting is exothermic but very slow. Conversely, some endothermic reactions happen quickly, like dissolving ammonium nitrate in water. Rate depends on activation energy; energy change depends on bond strengths
Wrong: Breaking bonds releases energy Correct: Breaking bonds ALWAYS requires energy input (it is endothermic). Energy is needed to overcome the attractive forces between bonded atoms. Making bonds ALWAYS releases energy (it is exothermic). A reaction is exothermic overall only because MORE energy is released making new bonds than is needed to break the original bonds
Wrong: Exothermic reactions always feel hot Correct: Exothermic reactions transfer energy to the surroundings, causing a temperature rise. However, very slow exothermic reactions (like rusting) release heat so gradually you cannot feel any temperature change
Explain, using bond energy calculations, why the combustion of methane is exothermic.
During the combustion of methane, bonds in the reactants (CH4 and O2) must first be broken, which requires energy input. The bonds broken are 4 C−H bonds (4 × 435 = 1740 kJ/mol) and 2 O=O bonds (2 × 498 = 996 kJ/mol). Total energy required to break bonds = 2736 kJ/mol. [2 marks]
When new bonds form in the products (CO2 and H2O), energy is released. The bonds made are 2 C=O bonds (2 × 805 = 1610 kJ/mol) and 4 O−H bonds (4 × 464 = 1856 kJ/mol). Total energy released making bonds = 3466 kJ/mol. [2 marks]
The overall energy change = energy to break bonds − energy released making bonds = 2736 − 3466 = −730 kJ/mol. The negative value shows the reaction is exothermic because more energy is released when new bonds form (3466 kJ) than is needed to break the original bonds (2736 kJ). The excess 730 kJ/mol is transferred to the surroundings as heat, which is why combustion of methane can be used as a fuel for heating. [2 marks]
Given the following bond energies in kJ/mol: N≡N = 945, H−H = 436, N−H = 391. Calculate the enthalpy change for the reaction N2(g) + 3H2(g) → 2NH3(g). Classify the reaction as exothermic or endothermic and explain your reasoning.
Bonds broken: N≡N = 1 × 945 = 945 kJ/mol; H−H = 3 × 436 = 1308 kJ/mol. Total broken = 945 + 1308 = 2253 kJ/mol. Bonds made: N−H = 6 × 391 = 2346 kJ/mol (2NH3 has 6 N−H bonds). Energy change = 2253 − 2346 = −93 kJ/mol. The reaction is exothermic because the energy change is negative. More energy is released making the 6 N−H bonds (2346 kJ) than is needed to break 1 N≡N and 3 H−H bonds (2253 kJ). The 93 kJ/mol of excess energy is transferred to the surroundings as heat. This is why the Haber process releases heat and must be cooled to maintain optimal conditions.
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