Sec 4 Chemistry Chemical Energetics notes

Every reaction shuffles energy. In an exothermic reaction the mixture releases heat to the surroundings: the thermometer reading rises and ΔH, the enthalpy change per mole, carries a negative sign. Combustion, neutralisation, respiration and the setting of concrete are all exothermic — your body is a warm chemical reactor.

An endothermic reaction does the opposite: it absorbs heat, the thermometer falls, and ΔH is positive. Photosynthesis runs on sunlight; thermal decompositions need continuous heating — stop the Bunsen flame and calcium carbonate stops decomposing: CaCO₃ → CaO + CO₂.

Drawn in words: an energy level diagram is a vertical energy axis with two horizontal shelves — reactants on one, products on the other — joined by a curve that humps upward before descending (or climbing) to the product shelf. The vertical drop or climb between shelves is ΔH; the height of the hump above the reactants is the activation energy, Eₐ.

Exothermic: product shelf below reactant shelf, ΔH arrow pointing down, energy released to the surroundings. Endothermic: product shelf above, ΔH arrow pointing up, energy taken in. Label both shelves with formulae and the ΔH value with units (kJ/mol) — examiners award the marks piece by piece.

Activation energy is the minimum energy a collision must carry before bonds can rearrange — the entry fee for reaction. Collisions carrying less than Eₐ simply bounce. A catalyst offers an alternative route with a smaller hump: more collisions clear the lower barrier, so the reaction speeds up — but the two shelves never move, so ΔH is unchanged and the catalyst emerges unaltered.

That last point is the classic trap: a catalyst changes rate, never yield or ΔH. Combustion needs a spark for the same reason — methane and oxygen coexist peacefully until the initial Eₐ is paid, then the exothermic reaction funds itself.

Bonds are chemical duct tape: pulling them apart absorbs energy (endothermic); snapping them together releases energy (exothermic). So ΔH = total energy to break all bonds − total energy released forming all bonds. A negative answer means the bonds formed were stronger (more, or more energetic) than the bonds broken.

For methane combustion, CH₄ + 2O₂ → CO₂ + 2H₂O, using C–H = 413, O=O = 498, C=O = 803 and O–H = 464 kJ/mol: break 4(413) + 2(498) = 2648; form 2(803) + 4(464) = 3462; ΔH = −814 kJ/mol. Values from average bond energies are approximate — the measured value is about −890 kJ/mol — because average energies ignore each molecule’s exact neighbourhood. State this limitation for the evaluation mark.

One subtlety worth a mark: bond-energy arithmetic only works directly when all species are gaseous. For solutions and solids you would also need energy changes of dissolving and lattice breaking — beyond the O-Level calculation, but worth knowing why the numbers drift.

A fuel releases heat when burned, and a good fuel is cheap, plentiful, safe to store, easy to ignite and burns without toxic products. Complete combustion of a hydrocarbon needs plenty of oxygen: CH₄ + 2O₂ → CO₂ + 2H₂O. Limited air supply switches the chemistry to incomplete combustion: 2CH₄ + 3O₂ → 2CO + 4H₂O (and/or soot, C). Carbon monoxide binds haemoglobin 200 times more strongly than oxygen — the silent killer of badly ventilated heaters.

Hydrogen is the dream fuel from the products side: 2H₂ + O₂ → 2H₂O gives only water — no CO₂ to warm the planet, no SO₂ to acidify rain — and it packs about 3× the energy per kilogram of petrol. The catch is supply: most industrial hydrogen still comes from reforming natural gas, which emits CO₂; only electrolysis of water powered by renewable electricity makes it genuinely green. It is also very flammable and hard to store (lightest gas, leaks, needs high pressure or cooling).