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Exothermic reactions release heat (ΔH < 0); endothermic reactions absorb heat (ΔH > 0). The sign and magnitude of ΔH determine heat flow between system and surroundings.
From your study of thermochemistry and enthalpy, you know that enthalpy (H) is a state function measuring the heat content of a system at constant pressure, and from calorimetry you know how to measure the heat exchanged during a reaction. The classification of reactions as exothermic or endothermic applies these ideas to a simple but powerful question: does energy flow out of the reacting system into the surroundings, or into the system from the surroundings?
In an exothermic reaction, the products have lower enthalpy than the reactants — the system has released energy, typically as heat. The enthalpy change ΔH is negative because the final state is lower in energy than the initial state. You can feel this directly: the reaction vessel gets warm. Combustion is the classic example — burning methane (CH₄ + 2O₂ → CO₂ + 2H₂O) releases 890 kJ per mole because the bonds in CO₂ and H₂O are collectively stronger (lower energy) than the bonds in CH₄ and O₂. The excess energy escapes as heat. Neutralization reactions (mixing a strong acid with a strong base) are also exothermic: the formation of water from H⁺ and OH⁻ releases about 56 kJ/mol.
In an endothermic reaction, the products have higher enthalpy than the reactants — the system has absorbed energy from its surroundings. ΔH is positive, and the reaction vessel feels cold. Dissolving ammonium nitrate (NH₄NO₃) in water is a familiar example — this is the chemistry behind instant cold packs. The energy required to break apart the crystal lattice exceeds the energy released when ions are hydrated, so the net process absorbs heat. Photosynthesis is another endothermic process: plants use sunlight to drive the conversion of CO₂ and H₂O into glucose and O₂, storing the sun's energy in chemical bonds.
The sign convention is critical and often trips students up: ΔH is defined from the system's perspective. When the system *loses* heat (exothermic), ΔH is negative — even though the surroundings are gaining that heat. Think of it as the system's energy account balance: a withdrawal (energy leaving) is negative, a deposit (energy entering) is positive. An energy diagram makes this visual — exothermic reactions show products sitting below reactants on the enthalpy axis, with the vertical gap equal to |ΔH|. Endothermic reactions show products above reactants. This framework connects directly to your upcoming work with bond energies and Hess's Law, where you will calculate ΔH by tracking the energy cost of breaking bonds versus the energy released in forming new ones.