Why Combustions Are Always Exothermic, Yielding about 418 kJ per Mole of O2

Abstract

The strongly exothermic nature of reactions between molecular oxygen and all organic molecules as well as many other substances is explained in simple, general terms. The double bond in O2 is much weaker than other double bonds or pairs of single bonds, and therefore the formation of the stronger bonds in CO2 and H2O results in the release of energy, which is given off as heat or increases thermal motion. This explains why fire is hot regardless of fuel composition. The bond energies in the fuel play only a minor role; for example, the total bond energy of CH4 is nearly the same as that of CO2. A careful analysis in terms of bond enthalpies, counting double bonds as two bonds to keep the total number of bonds unchanged, gives the heat of combustion close to -418 kJ/mol (i.e., -100 kcal/mol) for each mole of O2, in good agreement (±3.1%) with data for >500 organic compounds; the heat of condensation of H2O, -44 kJ/mol, is also included in the analysis. For 268 molecules with ≥ 8 carbon atoms, the standard deviation from the predicted value is even smaller, 2.1%. This enables an instant estimate of the heat of combustion simply from the elemental composition of the fuel, even for a complex mixture or unknown molecular structure, and explains principles of biofuels production. The analysis indicates that O2, rather than fuels like octane, H2, ethanol, or glucose, is the crucial "energy-rich" molecule; we briefly explain why O2 is abundant in air despite its high enthalpy.