3.14CH1 · Section 3: Physical chemistry

Energetics

Key learnings

Exothermic vs endothermic and the sign of ΔH
An exothermic reaction transfers energy to the surroundings, so their temperature rises — combustion, neutralisation and most displacement reactions. An endothermic reaction takes energy in from the surroundings, so their temperature falls — thermal decomposition and many dissolving reactions. The enthalpy change carries the sign: exothermic $\Delta H$ is negative, endothermic $\Delta H$ is positive (measured per mole, units $\text{kJ/mol}$ ).
Q = mcΔT for heat, then ΔH = −Q ÷ n per mole
Heat energy change: $Q = m\,c\,\Delta T$ , where $m$ = mass of water/solution (g), $c = 4.2\ \text{J/g}\,^{\circ}\text{C}$ , $\Delta T$ = temperature change. $Q$ is in joules — ÷1000 for kJ. The molar enthalpy change is $\Delta H = -\dfrac{Q}{n}$ per mole of reactant. E.g. 0.050 mol of fuel warms 50 g water by 20 °C: $Q = 50 \times 4.2 \times 20 = 4200\ \text{J} = 4.2\ \text{kJ}$ , so $\Delta H = -4.2 \div 0.050 = -84\ \text{kJ/mol}$ (exothermic).
Bond energies: ΔH = broken − formed
Breaking bonds needs energy (endothermic); making bonds releases energy (exothermic). Overall $\Delta H = \sum(\text{bonds broken}) - \sum(\text{bonds formed})$ . For $\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$ (H–H 436, Cl–Cl 242, H–Cl 431): broken $= 436 + 242 = 678$ ; formed $= 2 \times 431 = 862$ ; $\Delta H = 678 - 862 = -184\ \text{kJ/mol}$ (exothermic). If breaking costs more than forming releases, $\Delta H$ is positive (endothermic).

Common mistakes

Drawn from real examiner reports.

Doing broken − formed the wrong way round

The formula is broken minus formed, never formed minus broken. Reversing it flips the sign: for $\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$ the answer is $-184\ \text{kJ/mol}$ , but reversing gives $+184$ and wrongly makes the reaction look endothermic. A wrong sign then forces a wrong energy level diagram, losing more marks by error carried forward.

June 2024 Paper 2C Q7(c) — a common error was giving +184 rather than -184; with a positive value the energy diagram should then have been drawn as endothermic.

Forgetting a bond energy's coefficient

Count each bond with the correct multiplier from the balanced equation. In $\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$ there are two H–Cl bonds formed, so the H–Cl bond energy is multiplied by 2. Candidates forget the $\times 2$ on H–Cl, or wrongly double the H–H and Cl–Cl bonds that occur only once. List the bonds with their counts before adding.

June 2024 Paper 2C Q7(c) — candidates either multiplied the first two bond energies twice, or forgot to multiply the H-Cl bond energy by two.

Both bond breaking and making 'need energy'

Bond breaking is endothermic (needs/absorbs energy); bond making is exothermic (releases energy). Calling both processes "energy needing" scores zero. To explain why a reaction is exothermic, compare the two: more energy is released making the product bonds than is taken in breaking the reactant bonds. Learn the pairing: breaking in, making out.

June 2023 Paper 2CR Q7(b)(ii) — most candidates did not distinguish bond breaking (endothermic) from bond making (exothermic).

Restating the sign is not an explanation

Writing " $\Delta H$ is negative, so it is exothermic" or "it releases energy, so it is exothermic" simply restates the result — examiners give it no explanation marks. A proper explanation must compare the bond energies: more energy is released forming the product bonds than is absorbed breaking the reactant bonds. Always turn the conclusion into a bonds comparison.

June 2023 Paper 2CR Q7(b)(ii) — stating only that energy was released and so the reaction was exothermic did not score, as it was not an explanation.

Dropping the minus sign on an exothermic ΔH

An exothermic reaction has a negative $\Delta H$ — the minus sign is part of the answer and is dropped repeatedly. If the temperature rose, or the bonds-formed total exceeds the bonds-broken total, $\Delta H$ must carry a minus sign, not just its magnitude. Marks are also lost for not giving the value to the required significant figures.

June 2024 Paper 1CR Q6(b)(iii) — the negative sign for an exothermic reaction is dropped repeatedly, and answers are not always given to the required significant figures.

Using the fuel or solid mass in Q = mcΔT

In $Q = m\,c\,\Delta T$ , $m$ is the mass of the water or solution that is heated — never the mass of the fuel burned or the solid reactant added. The energy released is transferred into the water, so its mass and $c = 4.2\ \text{J/g}\,^{\circ}\text{C}$ are used. Using the fuel mass, or forgetting that $Q$ comes out in joules (divide by 1000 for kJ), are common slips.

June 2024 Paper 1CR Q6(b)(i) — in Q = mcΔT candidates must use the mass of water, not the fuel.

Exam tips

Bond-energy calc: list broken, list formed
(1) Balance the equation. (2) List each bond broken with its count and total them. (3) List each bond formed with its count and total them. (4) $\Delta H = \text{broken} - \text{formed}$ . (5) State the sign and $\text{kJ/mol}$ : negative exothermic, positive endothermic.
Calorimetry: Q = mcΔT, then ΔH = −Q/n
Find $Q = m\,c\,\Delta T$ using the mass of the solution and $c = 4.2$ , then convert J to kJ (÷1000). Find the moles $n$ of the limiting reactant. Then $\Delta H = -\dfrac{Q}{n}$ , keeping the minus sign when the temperature rose (exothermic).
Energy diagram: draw Ea from the reactants
On a reaction profile, draw the activation energy $E_a$ from the reactant level up to the top of the hump — not from the products. Draw the $\Delta H$ arrow between reactant and product levels: down for exothermic, up for endothermic. Mislabelled arrows lose marks.
Show working; keep the sign and units
Method marks reward the route, so show every line of working even if the arithmetic slips. Keep full calculator figures through the intermediate steps and round only at the very end. Finish with the sign and units — an exothermic $\Delta H$ is negative and in $\text{kJ/mol}$ .

Revision notes

Energetics

Exothermic reaction — a reaction that transfers energy to the surroundings, so the temperature of the surroundings rises (e.g. combustion, neutralisation). $\Delta H$ is negative.
Endothermic reaction — a reaction that takes in energy from the surroundings, so the temperature of the surroundings falls (e.g. thermal decomposition, many dissolving reactions). $\Delta H$ is positive.
Enthalpy change ( $\Delta H$ ) — the heat energy change of a reaction, measured per mole, at constant pressure (units $\text{kJ/mol}$ ).
Bond breaking is endothermic (energy must be put in); bond making is exothermic (energy is released).

Core relationships:

$Q = m\,c\,\Delta T \qquad \Delta H = -\frac{Q}{n} \qquad \Delta H = \sum(\text{bonds broken}) - \sum(\text{bonds formed})$

In $Q = mc\Delta T$ : $m$ = mass of the water/solution (g), $c = 4.2\ \text{J/g}\,^{\circ}\text{C}$ for water, $\Delta T$ = temperature change ( $^{\circ}\text{C}$ ). $Q$ is in joules — divide by 1000 for kJ.

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Energetics

Key learnings

Common mistakes

Exam tips

Revision notes

Energetics

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Energetics

Key learnings

Common mistakes

Exam tips

Revision notes

Energetics

Key facts and definitions

Worked example — calorimetry (statements 3.3 and 3.4)

Worked example — bond energies (statement 3.7C)

A chain of reasoning you can reuse

Diagram labelling checklist — energy level (reaction profile) diagrams (statement 3.5C)

Don't mix these up

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