Reactions that give heat and reactions that drink it
You have met this in daily life without naming it. A hand-warmer pouch gets hot; an instant cold-pack for a sprain turns icy. Both are chemistry releasing or absorbing energy as heat, and the study of that heat is thermochemistry. The core split is simple. An exothermic reaction *releases* heat to its surroundings — the flask and your hand warm up. An *endothermic* reaction *absorbs* heat from its surroundings — the flask and your hand go cold as the reaction pulls energy in.
Now connect this to the enthalpy we built last time. At constant pressure the heat of a reaction *is* its enthalpy change, ΔH, and the sign tells the whole story. An exothermic reaction has a *negative* ΔH: the products hold less enthalpy than the reactants, and the leftover spills out as heat. An endothermic reaction has a *positive* ΔH: the products hold more, so the reaction must drink in energy from the surroundings to climb to that higher level. Negative means heat out and warmer surroundings; positive means heat in and colder surroundings.
Putting a number on it: the enthalpy of reaction
We want more than "warm" or "cold" — we want a number. The enthalpy of reaction, written ΔH_rxn, is the heat released or absorbed when the reaction proceeds by the amounts written in its balanced equation, at constant pressure. It carries units of energy per mole of reaction, typically kilojoules per mole (kJ/mol). Burning one mole of methane, for instance, releases about 890 kJ, so its ΔH_rxn is roughly −890 kJ/mol — negative, telling you instantly that combustion is strongly exothermic.
Two honest cautions keep this number trustworthy. First, ΔH_rxn is tied to the *exact* balanced equation: double all the coefficients and you double ΔH, because you are now reacting twice as much. Second, it depends on the physical *states* of everything — making liquid water versus steam gives genuinely different values, because turning that water to vapour itself costs energy. This is why careful tables always pin down conditions and states, and why we will soon meet standard conditions to keep everyone comparing the same thing.
How much heat does it take to warm something?
Before we can measure a reaction's heat, we need one more idea: how much heat warming a substance actually takes. The heat capacity of an object is the heat needed to raise its temperature by one degree — a big swimming pool has an enormous heat capacity, a teaspoon of water a tiny one. To compare materials fairly, we use specific heat: the heat needed to raise *one gram* of a substance by one degree. Water's specific heat is famously high (about 4.18 joules per gram per degree), which is why oceans moderate climate and why water is the perfect liquid to trap a reaction's heat.
These two ideas are really the same one at different scales. Heat capacity belongs to a *particular object* (this exact pool); specific heat belongs to a *material* (water in general), and multiplying specific heat by mass gives back the object's heat capacity. The working formula that ties it all together is q = m · c · ΔT — heat equals mass times specific heat times the temperature change. That short equation is the engine of the experiment we are finally ready to meet.
Calorimetry: a thermometer becomes an energy meter
Here is the beautiful trick. We cannot measure heat directly, but we *can* read a temperature change off a thermometer. Calorimetry turns that temperature change into a heat measurement. The idea is pure First Law: run the reaction inside a well-insulated container of water, and any heat the reaction releases has nowhere to go but into that water. Measure how much the water warms, and q = m · c · ΔT converts the temperature rise straight into joules. The reaction's heat is captured, weighed, and read out — by little more than a foam cup, some water, and a thermometer.
- Run the reaction in insulated water and record the water's temperature before and after to get ΔT.
- Compute the heat absorbed by the water with q = m · c · ΔT, using the water's mass and its specific heat.
- Flip the sign: heat gained by the water equals heat lost by the reaction, so q_reaction = −q_water. At constant pressure this q_reaction is ΔH.
- Divide by the moles of reactant consumed to report the answer per mole — that is your measured enthalpy of reaction.