Application of the [[First Law of Thermodynamics|first law]] upon a simple [[Combustion Process|combustion process]] will result in the following:
$Q+H_R=H_P$
Where $H_R$ denotes the [[Enthalpy|enthalpy]] of [[Chemical Reactant|reactants]] and $H_P$ denotes the enthalpy of [[Chemical Product|products]].
$H_R=\sum_Rn_i\bar{h}_i$
$H_P=\sum_Pn_e\bar{h}_e$
Since enthalpy is measured from an arbitrary reference point we can assume that HR=0. The following is measured when assuming zero reactant enthalpy for the formation of carbon dioxide:
$C+O_2=CO_2$
$\dfrac{Q}{n_{CO_2}}=\dfrac{H_{CO_2}}{n_{CO_2}}=−393522\quad\text{kJ/kmol}$
Since $Q$ is negative, the [[Reaction|reaction]] is releasing [[Heat|heat]] to the surrounding environment. The resulting measurement is termed the enthalpy of formation:
$\bar{h}_{f,CO_2}^0=−393522\quad\text{kJ/kmol}$
The enthalpy of formation is dependent on the product as well as the [[Pressure|pressure]] and [[Temperature|temperature]] at which the reaction takes place. Atmospheric conditions, superscript 0, are atmospheric pressure and 25 degrees Celsius.
$\bar{h}_{T,P}=\bar{h}_f^0+(\bar{h}_f−\bar{h}_f^0)_{T,P}$
$\bar{h}_{T,P}=\bar{h}_f^0+\Delta \bar{h}_{T,P}$
The enthalpy of formation for a [[Substance|substance]] is tabulated in any standard thermodynamics textbook.
A couple notes on enthalpy of formation:
- $\bar{h}_f^0=0$ for pure elements like $C$, $O_2$, $H_2$, $N_2$, $S$
- The enthalpy of formation for stable compounds is negative and for unstable compounds is positive