The amount of chemical energy a substance can deliver is encoded in the bonds that hold its atoms together. During a chemical reaction, these bonds are broken and re‑formed, and the net energy change depends on the relative strengths of the bonds involved.
Atoms are linked by different kinds of bonds—covalent, ionic, metallic, and hydrogen—each carrying a characteristic amount of energy. Covalent bonds, formed by electron sharing, are typically the strongest and thus store the most energy (e.g., the O–H bonds in water). Ionic bonds, such as Na⁺–Cl⁻ in table salt, are weaker, while hydrogen bonds between water molecules are among the weakest.
In practice, a chemist records the quantities of reactants, the temperature, and the pressure before and after a reaction. Only the net change in bond energy matters: if the bonds in the products contain less energy than those in the reactants, heat is released (an exothermic process). Conversely, if the products possess more energy, the reaction absorbs heat from the surroundings (endothermic).
Exothermic reactions release heat, e.g., combustion of wood, where carbon and hydrogen react with oxygen to form CO₂ and H₂O. Endothermic reactions consume heat, such as dissolving NaCl in water, which slightly lowers the solution’s temperature.
Whether a reaction occurs on its own depends on the system’s free energy. Spontaneous reactions, like sodium metal reacting violently with water, proceed without external input. Nonspontaneous reactions, such as igniting gasoline, require an energy input (e.g., a spark) to cross an activation barrier.
Understanding these principles allows chemists to predict and control the energy flow in chemical processes.
