2.2:
Enthalpy and Heat of Reaction
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.
2.2:
Enthalpy and Heat of Reaction
Combustion, commonly known as burning, is a reaction in which a substance reacts with an oxidizing agent, which in most cases is molecular oxygen, to liberate energy in the form of heat, light, or sound. The heat of combustion is also known as the enthalpy of combustion. The energy released when one mole of a substance undergoes complete combustion at constant pressure is called molar heat of combustion. Combustion reactions are exothermic; that is, they release energy, and their ΔH sign convention is negative.
In 1772, French chemist Antoine Lavoisier, discovered that the products of burnt sulfur weighed more than the initial mass of the reactant. He postulated that sulfur combined with air, which resulted in the increased weight. Later, Joseph Priestley's discovery of "oxygen" in 1774, as a component of air, led Lavoisier to believe that sulfur combined with oxygen in the air, leading to an increase in its mass. He concluded that combustion means combining with oxygen. In other words, sulfur underwent combustion.
Examples of combustion reactions include the burning of hydrocarbon fuels like natural gas and coal. In the case of combustion reactions involving hydrocarbons, the amount of energy released varies depending on the type of fuel undergoing combustion.
For example, the combustion of natural gas, methane (CH4), given by the reaction:
generates less heat energy than that of butane (C4H10), given by the reaction:
Thus the number of oxygen molecules required to combust the hydrocarbon and the number of molecules of each product formed depend on the hydrocarbon composition.
The heat of combustion governs the relative stability of branched hydrocarbons with the same molecular formula. The difference in structure arises due to methyl groups attached at different positions along the hydrocarbon chain. The amount of heat energy released decreases with increasing branching, where the highly branched 2,2-dimethylhexane generates low energy compared to octane. Hence, unbranched octane is less stable than its branched counterpart.