Enthalpy, abbreviated H, equals the sum of internal energy, abbreviated E or U, and the product of pressure and volume.
The change in enthalpy, ΔH, is expressed as the difference between the enthalpies of the products and the reactants. At constant temperature and pressure, ΔH is equal to the amount of heat energy exchanged between the system and the surroundings, or the heat of the reaction.
When ΔH is positive, reactions absorb heat and are endothermic. If ΔH is negative, reactions release heat and are exothermic.
Combustion is an example of an exothermic process, where a substance burns in the presence of an oxidant like atmospheric oxygen to release energy in the form of heat.
The heat released is quantified as the molar heat of combustion, which is the amount of heat energy released on burning one mole of a substance.
When a hydrocarbon burns, the carbon and hydrogen from the fuel combine with molecular oxygen to produce water and carbon dioxide, along with the release of energy.
The value of the heat of combustion for a hydrocarbon increases with the number of carbon atoms in the chain since more carbon is available for burning and more bonds undergo changes.
For example, the combustion of methane, a single-carbon compound, generates less heat energy than that of butane, which has four carbon atoms.
The heat of combustion is a critical way to determine the relative stability of hydrocarbons with the same molecular formula but different structures.
Consider the heats of combustion of octane, 2-methylheptane, and 2,2-dimethylhexane.
These compounds have the same number of carbon atoms, but the methyl groups are attached at different positions in each molecule.
Octane has the largest heat of combustion. As the branching increases the ΔH decreases, suggesting that branching increases the stability of a hydrocarbon.
