Combustion Reaction Calculator

Complete combustion of an organic fuel CxHyOz sends every carbon atom to CO2 and every hydrogen atom to H2O, so the product coefficients follow immediately: x moles of carbon dioxide and y/2 moles of water per mole of fuel. Counting oxygen atoms on the product side gives 2x + y/2, and subtracting the z atoms the fuel already carries leaves an oxygen requirement of x + y/4 - z/2 moles of O2. Methane, with x = 1 and y = 4, needs 2; octane, with x = 8 and y = 18, needs 12.5, which is why the textbook version is written as 2C8H18 + 25O2 to keep every coefficient whole. Oxygenated fuels demand less oxygen for the same reason ethanol, C2H6O, needs only 3 moles where the comparable hydrocarbon C2H6 needs 3.5. The heat released can be estimated from standard enthalpies of formation, since the products dominate the balance: about 393.5 kJ per mole of CO2 and 285.8 kJ per mole of liquid H2O, less the enthalpy of formation of the fuel itself. Using the liquid-water value gives the higher heating value; if the water leaves as vapour, roughly 44 kJ per mole is never recovered and the practical figure is the lower heating value, which is what engine and boiler calculations normally use. Air rather than pure oxygen is the usual oxidant, and at 23.2 percent oxygen by mass about 4.3 kilograms of air accompany each kilogram of O2 demanded. Real burners run with excess air, typically 5 to 10 percent for natural gas and 20 percent or more for solid fuels. The stoichiometric equation also ignores nitrogen oxides, which form from atmospheric N2 above roughly 1300 degrees C.

Combustion was explained backwards for the better part of a century. The phlogiston theory of Johann Joachim Becher and Georg Ernst Stahl, developed between roughly 1667 and 1703, held that burning substances released a fire principle, a picture that never reconciled with the observation that metals gain weight as they burn. Carl Wilhelm Scheele prepared oxygen around 1771 and Joseph Priestley independently in 1774, but it was Antoine Lavoisier who, through careful weighings during the 1770s, identified combustion as chemical combination with a component of the air, which he named oxygen in 1777. Once compositions and atomic weights were settled, writing a balanced combustion equation and reading the oxygen requirement off it became straightforward arithmetic. Quantifying the heat came next. Lavoisier and Pierre-Simon Laplace measured heats of combustion with an ice calorimeter in the early 1780s. Germain Hess established in 1840 that the total heat of a reaction is independent of the path taken, which allowed enthalpies of combustion to be assembled from tabulated enthalpies of formation instead of being measured fuel by fuel. Marcellin Berthelot introduced the sealed bomb calorimeter in the 1880s, providing the constant-volume measurements from which modern heating values descend, and the distinction between higher and lower heating value, turning on whether the product water condenses, belongs to the engineering practice of the same period. Ernest Mallard and Henry Le Chatelier's flame propagation studies of the 1880s, prompted by French coal-mine explosions, linked stoichiometry to flammability limits and ignition, a framework still used in furnace and engine design.