Heat and enthalpy

Heat is energy moving because of a temperature difference. How much heat it takes to warm something depends on how much there is and what it's made of: q = m × c × ΔT, where c is the specific heat capacity. Water's is unusually high, 4.18 J per gram per kelvin. Warming 500 g of water by 20 K takes 500 × 4.18 × 20 = 41,800 J, or 41.8 kJ.

Reactions take in or give out heat because bonds hold energy. Breaking a bond always takes energy; making one always releases it. If the bonds made are stronger than the bonds broken, the reaction releases heat to its surroundings: it's exothermic, like burning fuel. If not, it absorbs heat and is endothermic, like the cold pack that gets cold when you squeeze it.

The heat a reaction gives or takes at constant pressure is its enthalpy change, ΔH, usually per mole of reaction as written. Exothermic reactions have a negative ΔH (the chemicals lose energy); endothermic ones positive. Burning one mole of methane has ΔH = −890 kJ/mol when the water ends up as liquid.

Heats of reaction are measured in a calorimeter: run the reaction in a known amount of water and watch its temperature. The water gains what the reaction gives out. Mix acid and base so the solution's 100 g warms by 6.0 K, and the reaction released 100 × 4.18 × 6.0 = 2508 J.

Changes of state have enthalpies too. Melting ice takes 6.01 kJ/mol, and boiling water takes 40.7 kJ/mol at 100 °C (about 2257 kJ per kg), which is why steam burns are so bad: condensing steam dumps all that energy into your skin.