Standard enthalpy change of combustion

The standard enthalpy of combustion is the enthalpy change when one mole of a reactant completely reacts with oxygen under standard thermodynamic conditions (although experimental values are usually obtained under different conditions and subsequently adjusted). By definition, combustion reactions are always exothermic and so enthalpies of combustion are always negative, although the values for individual combustions may vary.

The most common way of calculating the enthalpy change of combustion (or formation) is by using a Hess cycle or by using numerical based bond enthalpies. It is commonly denoted as $\Delta H ^{\circ} _{\mathrm{comb}}$ or $\Delta H ^{\circ}_{\mathrm{c}}$. When the enthalpy required is not a combustion, it can be denoted as $\Delta H ^{\circ} _{\mathrm{total}}$. Enthalpies of combustion are typically measured using bomb calorimetry, and have units of energy (typically kJ); strictly speaking, the enthalpy change per mole of substance combusted is the standard molar enthalpy of combustion (which typically would have units of kJ mol⁻¹).

Comparisons

Hydrocarbons

The standard enthalpy change of hydrocarbons vary depending on their molecular size. For instance, the standard enthalpy change of Methane (CH₄) is -890.3 kJ mol⁻¹, whereas the standard enthalpy change Hexane (C₆H₁₄) is -4163.0 kJ mol⁻¹. This is due to the presence of more bonds in Hexane compared to Methane. Methane has only 4 bonds between its single carbon and four Hydrogen atoms. Hexane on the other hand, has 5 carbon-carbon bonds and 14 carbon-hydrogen bonds. Therefore, there is more bond making and bond breaking during the combustion of a larger hydrocarbon compared to a smaller hydrocarbon, thus the enthalpy change is larger for each mole of the substance.

Alcohols and alkanes

For alcohols and alkanes containing the same number of carbon atoms, e.g. Methane (CH₄) and Methanol (CH₃OH), the standard enthalpy change of the alkane would be greater than the alcohol ($\Delta H ^{\circ}_{\mathrm{c}}$ Methane = -890.3 kJ mol⁻¹, $\Delta H ^{\circ}_{\mathrm{c}}$ Methanol = -726.0 kJ mol⁻¹).

Since the complete combustion of both of these carbon compounds produce carbon dioxide (CO₂) and water (H_2O), there is more bond-breaking and bond-making when Methane is burnt. The presence of an -OH bond on Methanol means that there is less bond-breaking and bond-making to produce water compared to Methane.

External links

• NIST Chemistry WebBook