Thermochemistry

Before we can talk about energy changing, we have to draw a line around the bit we care about. The system is the part of the universe under study — usually the reaction mixture in a beaker. Everything else that can exchange energy or matter with it is the surroundings. The imaginary line between them is the boundary, and thermochemistry is really the bookkeeping of what crosses it.

• System — the reaction (or sample) we are studying.
• Surroundings — the water, air, beaker and bench around it.
• Boundary — the real or imaginary surface energy and matter pass through.

A thermos keeping your tea hot for hours is the closest everyday thing to an isolated system — it tries to let neither heat nor matter escape.

Every substance stores chemical energy. The enthalpy (H) is that heat content at constant pressure (H = U + PV). We can never read H directly, so we always track the change: ΔH = H(products) − H(reactants).

• Internal energy (U) — the total kinetic + potential energy stored in a system.
• Enthalpy (H) — the heat content at constant pressure.
• Exothermic reaction — gives heat OUT; the surroundings warm up; ΔH is negative.

When you light a campfire or shake a hand-warmer, the fuel turns into products that hold less energy than the starting materials. The leftover energy escapes as heat, so the thermometer climbs and ΔH comes out negative. Combustion and acid–base neutralisation are classic exothermic reactions.

The opposite happens in an endothermic reaction: the products store more energy than the reactants, so the reaction has to absorb heat from its surroundings to proceed. The surroundings lose that heat and cool down, which is why ΔH is positive.

• Endothermic reaction — takes heat IN; the surroundings cool down; ΔH is positive.

Snap an instant cold-pack and it turns icy in your hand — the salt dissolving inside soaks up heat from your skin, which is exactly how it soothes a sprain. Photosynthesis and thermal decomposition are endothermic too.

An enthalpy diagram plots energy on the vertical axis against reaction progress. Think of the reaction as a ball rolling downhill: it starts high (reactants) and settles into a lower, more stable valley (products). The height it drops is the energy given out — the size of ΔH.

• Exothermic diagram — products lie below reactants; the ball rolls down; ΔH negative.
• Endothermic diagram — products lie above reactants; you'd have to push the ball uphill; ΔH positive.

To compare reactions fairly we quote standard enthalpy changes — measured at 298 K, 1 atm, per mole. Common ones are formation (ΔH°_f), combustion (ΔH°_c, how we rate fuels), neutralisation (≈ −57 kJ) and solution.

Because H is a state function, Hess's law says the total ΔH depends only on the start and end — not the route. Picture a mountain hike: the total height you gain to the summit is the same whether you climb straight up or wander there by a detour. That lets chemists find a dangerous or slow ΔH by adding up easier steps.

How do we actually measure the heat? With a calorimeter: burn a known sample of food (or fuel) and let it warm a known mass of water in an insulated vessel. The water's temperature climbs, and the heat it absorbed is:

This is literally how a food scientist finds the Calories on a label. Drag the sliders in the live panel → to set the mass of water and its temperature change, and watch the heat work itself out.

Every reaction is really old bonds breaking and new bonds forming. Think of bonds as LEGO bricks: snapping two apart costs energy (you have to pull), while clicking two together releases energy (they snap home). So:

If the new bonds release more than the old ones cost, ΔH is negative — exothermic. That is why burning a fuel gives out energy overall: the strong new bonds in CO₂ and H₂O hand back more than was spent breaking the fuel apart.

Underneath it all sits the first law of thermodynamics — energy is conserved, never created nor destroyed:

1. System / surroundings / boundary; open, closed, isolated.
2. Internal energy, enthalpy, ΔH = H(products) − H(reactants).
3. Exothermic (−ΔH, warms) vs endothermic (+ΔH, cools).
4. Enthalpy diagrams; standard ΔH changes; Hess's law (path-independent).
5. Calorimetry q = mcΔT; ΔH from bond energies; first law ΔU = q + w.