Reaction Enthalpy

Reaction enthalpy: ΔH = ΣΔHf°(products) − ΣΔHf°(reactants)

The enthalpy of reaction ΔH is the heat a reaction releases or absorbs when pressure stays constant: ΔH = H_products − H_reactants. A negative ΔH means the reaction is exothermic and gives off heat, as combustion and neutralization do. A positive ΔH means it is endothermic and pulls heat in, like photosynthesis or ice melting. Hess's law tells us ΔH behaves as a state function, so the total ΔH along a multi-step path adds up to the same value no matter which route you take. That is why tables of standard enthalpies of formation ΔHf° (measured at 25°C, 1 atm) are so handy: subtract the reactants from the products and you have the ΔH for any reaction. Take the combustion of methane, CH₄ + 2 O₂ → CO₂ + 2 H₂O, which releases ΔH = −890 kJ/mol. That much heat per mole is exactly why natural gas earns its keep as a fuel.

Applications

Engineers lean on ΔH when designing chemical processes and keeping them safe, estimating the heat load a reactor will throw off and sizing the cooling jackets to handle it. It shows up when fuels get ranked by heat of combustion (gasoline around 46 MJ/kg, hydrogen 142 MJ/kg, coal about 24 MJ/kg), inside self-heating military MREs that run on the CaO + H₂O reaction, in bomb calorimetry that measures the caloric content of food, in explosives work, and in the thermal management of batteries.

FAQ

Why use standard enthalpy of formation? By convention, a pure element in its reference state is assigned ΔHf° = 0. With everything anchored to that baseline, the tabulated ΔHf° values give you a reaction's ΔH on paper, no calorimeter required.

Does ΔH depend on temperature? A little, yes. Kirchhoff's law captures it: ΔH(T) = ΔH(T_ref) + ∫ΔCp dT. For reactions running near room temperature, the shift is small enough to ignore most of the time.

What's the difference between ΔH and ΔU? They are linked by ΔH = ΔU + Δ(PV). ΔU is the change in internal energy, while at constant pressure ΔH is what you actually measure as heat exchanged. The gap between them stays small until a reaction starts producing or consuming gases, and then it can get substantial.