UNIT 2 CHEMISTRY · PEARSON AAQ · D1.3.1
Enthalpy Change
FREE CHEMISTRY REVISION · TOPIC 38
Enthalpy Change
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · D1.3.1
What you will learn
Define enthalpy change, identify the system and surroundings, and explain the signs of ΔH for exothermic reactions and endothermic processes. Apply the ideas to chemical reactions and physical state changes. Energy-level and reaction-profile diagrams are reserved for D1.3.2.
1. Enthalpy change: the system is your reference
The system is the part being studied: for example, the reacting chemicals or the water undergoing a state change. The surroundings are everything outside that chosen boundary, such as the solution used as a thermal reservoir, container and room. State the boundary before discussing where energy goes.
Enthalpy change, ΔH, is the enthalpy of the final system minus the enthalpy of the initial system. For a reaction, this is the enthalpy of the products minus that of the reactants, compared at the specified conditions. At constant pressure, when only pressure–volume work occurs, ΔH equals the heat transferred to the system. H is a state function: the change depends on the initial and final states, not the route between them.
Enthalpy has units of energy, such as J or kJ. Reaction enthalpy data are often expressed per mole of reaction as written, in kJ mol⁻¹. The amount and balanced equation therefore matter. If a question gives total heat for a particular sample, use the stated energy unit rather than automatically adding mol⁻¹.
Use ΔH = H(final) − H(initial). A positive value means an enthalpy increase; a negative value means an enthalpy decrease. Absolute enthalpy values are not normally needed. Physical states, amounts, temperature and pressure define the initial and final states, so they affect a quoted enthalpy change. Formal standard-enthalpy definitions are a later topic.
2. Exothermic reactions and processes
An exothermic change transfers energy as heat from the system to the surroundings at constant pressure. The system has lower final enthalpy, so ΔH is negative. Nearby surroundings often become warmer. “Energy is released” means transferred to the surroundings; energy has not been created or destroyed.
Chemical examples include combustion of fuels, many acid–alkali neutralisations and many oxidation reactions. A reaction can need ignition before releasing energy overall: starting the reaction and its net energy change are different ideas.
Physical exothermic processes include freezing a liquid and condensing a gas. A freezer removes heat from water while it freezes. During a pure substance’s state change at fixed pressure, heat transfer can occur while its temperature stays approximately constant.
3. Endothermic reactions and processes
An endothermic change transfers energy as heat into the system at constant pressure. The final enthalpy is higher, so ΔH is positive. If nearby surroundings supply the energy and it is not replaced, they can cool. With a heater maintaining temperature, an endothermic process can proceed without a falling measured temperature.
Chemical examples include thermal decomposition of calcium carbonate, CaCO₃(s) → CaO(s) + CO₂(g), and dissolving ammonium nitrate in water under typical cold-pack conditions. Dissolving is not always endothermic: its net enthalpy depends on the substance and conditions.
Physical endothermic processes include melting, boiling and evaporation. They require energy to overcome intermolecular attractions; they do not normally break covalent bonds within molecules. Evaporation can occur below the boiling point and can cool the liquid or surrounding surface.
Photosynthesis stores energy supplied by light. It is commonly described as an energy-absorbing overall chemical process, but light absorption is not heat transfer from a colder environment. Distinguish the mode of energy input when the context involves electrical work or radiation; the simple ΔH = heat relationship assumes constant pressure and only pressure–volume work.
4. Compare signs, transfers and observations
| Feature | Exothermic | Endothermic |
|---|---|---|
| Sign of ΔH | Negative | Positive |
| Direction of heat at constant pressure | System → surroundings | Surroundings → system |
| Final system enthalpy | Lower than initially | Higher than initially |
| Typical nearby temperature observation | Rises if heat is retained | Falls if heat is supplied without replacement |
| Chemical example | Fuel combustion | Thermal decomposition of CaCO₃ |
| Physical example | Condensation or freezing | Evaporation or melting |
Temperature and enthalpy are different quantities. Temperature reflects average particle kinetic energy; enthalpy change includes changes in the system’s energy associated with interactions as well as other contributions. Equal initial and final temperatures do not imply zero ΔH: a process can exchange heat with a reservoir or undergo a state change at constant temperature.
5. Worked explanations and sign conventions
Example 1 — a neutralisation warms its surroundings
Acid and alkali react in a cup and the nearby solution temperature rises. The reaction transfers heat to its surroundings, so it is exothermic and its reaction ΔH is negative. The temperature rise is evidence of heat release under these conditions. State which material acts as the thermal surroundings when interpreting the observation.
Example 2 — a cold pack
When ammonium nitrate dissolves, the pack cools as heat is taken from nearby material into the dissolving system. The dissolution is endothermic, with positive ΔH. It does not “release cold”: energy transfers into the system from its surroundings.
Example 3 — a total enthalpy change
At constant pressure with only pressure–volume work, a specified sample transfers 12.0 kJ as heat to its surroundings. The system’s enthalpy change is −12.0 kJ. If that same sample instead absorbs 12.0 kJ under those conditions, ΔH is +12.0 kJ. These are sample totals, so no mol⁻¹ is added.
Example 4 — calculate a molar value from supplied heat
A reaction releases 15.0 kJ when 0.250 mol of reaction occurs as written, under the same constant-pressure assumptions. Molar enthalpy change = −15.0/0.250 = −60.0 kJ mol⁻¹. For 0.500 mol at the same conditions, the total change is (−60.0) × 0.500 = −30.0 kJ. Scaling amount changes the total, not the molar value.
Example 5 — reverse a state change
Suppose supplied data give H₂O(l) → H₂O(g), ΔH = +44.0 kJ mol⁻¹ at specified conditions. Reversing exactly those states and conditions gives H₂O(g) → H₂O(l), ΔH = −44.0 kJ mol⁻¹. Vaporisation absorbs energy; condensation releases the same magnitude for the reverse change. The supplied value is illustrative data for the stated comparison, not a universal boiling-point value.
Example 6 — boiling at constant temperature
Water at its boiling point continues to absorb energy while its temperature stays nearly constant at fixed pressure. Vaporisation is endothermic, so ΔH is positive. The energy changes the physical state and intermolecular separation rather than simply increasing average kinetic energy. A flat temperature reading does not mean no energy transfer.
Example 7 — combustion needs a spark
A fuel–oxygen mixture may need a spark to start. Once reacting, it releases heat overall, so it is exothermic with negative ΔH. An initial activation barrier does not determine the sign of the overall enthalpy change. Detailed profiles will be treated separately.
6. Bonds, attractions and the overall balance
Breaking a chemical bond requires energy; forming a chemical bond releases energy. A chemical reaction involves both. It is exothermic if the overall energy released by forming new bonds and interactions exceeds the energy required to break the original ones. It is endothermic when the required energy is greater. Do not say that breaking fuel bonds releases the combustion energy.
For melting and boiling a simple molecular substance, the molecules generally remain intact while intermolecular attractions are overcome. For freezing and condensation, stronger intermolecular association releases energy. Distinguish chemical bonds within molecules from attractions between them.
The sign of ΔH does not by itself tell you the rate, whether a reaction needs a catalyst, or whether it is thermodynamically feasible under all conditions. Nor does adding a catalyst change ΔH for the same initial and final states. Those are distinct questions.
Common misconceptions
- Exothermic means positive because temperature rises: ΔH refers to the system, which loses energy, so its sign is negative.
- Endothermic means the system releases cold: energy transfers into the system.
- Any heated process is exothermic: the heater can supply energy to an endothermic process.
- All dissolving is endothermic: dissolution can be either endothermic or exothermic.
- Bond breaking releases energy: bond breaking requires energy; bond formation releases it.
- Boiling breaks water molecules apart: normal boiling overcomes intermolecular attractions, not O–H bonds.
- No temperature change means ΔH = 0: heat can be exchanged at constant temperature, including during state changes.
- A negative ΔH guarantees a fast reaction: net energy change and kinetic barrier are distinct.
Exam tips
- Define the system before describing heat flow.
- State the transfer direction, the sign of ΔH and a suitable observation.
- Use “system loses/gains energy” rather than vague “it gets energy”.
- Distinguish total energy units from energy per mole.
- For reverse processes at identical conditions, reverse the sign.
- Specify physical states and conditions when interpreting enthalpy data.
- Explain a temperature observation with energy transfer; do not use temperature and enthalpy interchangeably.
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