Sec 4 Combined Science (Chemistry) Chemical Energetics notes

Chemical reactions always involve an energy change alongside the substance change. The direction of flow decides the name: an exothermic reaction transfers energy (usually as heat) to the surroundings; an endothermic reaction takes energy in from the surroundings. The thermometer in the reaction mixture is the detector: rising for exothermic, falling for endothermic.

Watch the direction of the words: the system is the reacting mixture; the surroundings are everything else — the solution, the cup, the air, your hand. Heat always flows from the hotter to the cooler, so an exothermic reaction leaves the surroundings warmer, and an endothermic one leaves them cooler.

Combustion is the classic: burning methane (CH₄ + 2O₂ → CO₂ + 2H₂O) pours heat into the room — that is the point of a Bunsen burner or a gas stove. Neutralisation is quietly exothermic too: mix equal volumes of sodium hydroxide solution and hydrochloric acid, both starting at 25.0 °C, and the mixture warms — 25.0 → 31.5 °C is a typical classroom result (a 6.5 °C rise).

Metals with dilute acids release heat as well as hydrogen: magnesium ribbon in hydrochloric acid makes the tube noticeably warm (Mg + 2HCl → MgCl₂ + H₂). And quicklime slaking — adding water to calcium oxide — turns water visibly hot as CaO + H₂O → Ca(OH)₂: so much energy is released that self-heating cans are built on exactly this reaction.

• All combustions (natural gas, wax, wood) — exothermic.
• All neutralisations (acid + alkali) — exothermic.
• Reactive metal + acid — exothermic (and fizzing).
• Slaking quicklime (CaO + water) — exothermic.

Dissolve ammonium nitrate crystals in water and the solution turns noticeably colder — the dissolving takes in heat from the water. No new substance forms (dissolving is a physical change), but the energy change is real: the syllabus names this as the standard endothermic example, and it is what cold packs use.

Two chemical examples: mixing citric acid solution with sodium hydrogencarbonate fizzes (carbon dioxide) and cools — an endothermic reaction you can watch thermometre. And thermal decomposition: heating calcium carbonate (CaCO₃ → CaO + CO₂) needs a continuous, strong flame. Stop heating and the decomposition stops — the reaction keeps taking heat in, which is the definition of endothermic.

The standard method: (1) measure a known volume of the first solution in an insulated (polystyrene) cup and record its temperature; (2) add the second reactant, stir, and record the temperature every 30 seconds; (3) identify the highest temperature for an exothermic change or the lowest for an endothermic one. The temperature change is the difference from the starting value.

• Insulation matters: a polystyrene cup with a lid keeps heat from escaping (or entering) so the reading is honest.
• Stir so the temperature is even throughout before reading.
• Read the thermometer at eye level, to 0.5 °C (or better with a digital probe).
• The bigger the temperature change, the more energy was transferred — same volumes assumed.

Hand warmers and self-heating cans are exothermic chemistry packaged for use: the can warms its contents by mixing water with quicklime (calcium oxide), whose slaking releases a large amount of heat; some hand warmers work by crystallising a supersaturated solution or by the slow rusting of iron — both release heat.

Sports cold packs are the endothermic twin: a inner pouch of ammonium nitrate crystals (or a similar salt) bursts into water, and the dissolving pulls heat from the injured ankle, numbing it. One pack, one direction of energy flow — the same chemistry as the thermometer experiment, sized for first aid.