Standard enthalpy of formation

In

For example, the standard enthalpy of formation of carbon dioxide is the enthalpy of the following reaction under the above conditions:

${\displaystyle {\ce {C(s, graphite) + O2(g) -> CO2(g)}}}$

All elements are written in their standard states, and one mole of product is formed. This is true for all enthalpies of formation.

The standard enthalpy of formation is measured in units of energy per amount of substance, usually stated in

All elements in their reference states (oxygen gas, solid carbon in the form of graphite, etc.) have a standard enthalpy of formation of zero, as there is no change involved in their formation.

The formation reaction is a constant pressure and constant temperature process. Since the pressure of the standard formation reaction is fixed at 1 bar, the standard formation enthalpy or reaction heat is a function of temperature. For tabulation purposes, standard formation enthalpies are all given at a single temperature: 298 K, represented by the symbol Δ_fH^⦵
_{298 K}.

For many substances, the formation reaction may be considered as the sum of a number of simpler reactions, either real or fictitious. The

For ionic compounds, the standard enthalpy of formation is equivalent to the sum of several terms included in the Born–Haber cycle. For example, the formation of lithium fluoride,

${\displaystyle {\ce {Li(s) + 1/2 F2(g) -> LiF(s)}}}$

may be considered as the sum of several steps, each with its own enthalpy (or energy, approximately):

1. H_sub, the
   standard enthalpy of atomization (or sublimation
   ) of solid lithium.
2. IE_Li, the
   first ionization energy
   of gaseous lithium.
3. B(F–F), the standard enthalpy of atomization (or bond energy) of fluorine gas.
4. EA_F, the electron affinity of a fluorine atom.
5. U_L, the lattice energy of lithium fluoride.

The sum of these enthalpies give the standard enthalpy of formation (Δ_fH) of lithium fluoride:

${\displaystyle \Delta H_{\text{f}}=\Delta H_{\text{sub}}+{\text{IE}}_{\text{Li}}+{\frac {1}{2}}{\text{B(F–F)}}-{\text{EA}}_{\text{F}}+{\text{U}}_{\text{L}}.}$

In practice, the enthalpy of formation of lithium fluoride can be determined experimentally, but the lattice energy cannot be measured directly. The equation is therefore rearranged to evaluate the lattice energy:^[3]

${\displaystyle -U_{\text{L}}=\Delta H_{\text{sub}}+{\text{IE}}_{\text{Li}}+{\frac {1}{2}}{\text{B(F–F)}}-{\text{EA}}_{\text{F}}-\Delta H_{\text{f}}.}$

The formation reactions for most organic compounds are hypothetical. For instance, carbon and hydrogen will not directly react to form methane (CH₄), so that the standard enthalpy of formation cannot be measured directly. However the standard enthalpy of combustion is readily measurable using bomb calorimetry. The standard enthalpy of formation is then determined using Hess's law. The combustion of methane:

${\displaystyle {\ce {CH4 + 2 O2 -> CO2 + 2 H2O}}}$

is equivalent to the sum of the hypothetical decomposition into elements followed by the combustion of the elements to form carbon dioxide (CO₂) and water (H₂O):

${\displaystyle {\ce {CH4 -> C + 2H2}}}$

${\displaystyle {\ce {C + O2 -> CO2}}}$

${\displaystyle {\ce {2H2 + O2 -> 2H2O}}}$

Applying Hess's law,

${\displaystyle \Delta _{\text{comb}}H^{\ominus }({\text{CH}}_{4})=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{\text{O}})]-\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4}).}$

Solving for the standard of enthalpy of formation,

${\displaystyle \Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4})=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{\text{O}})]-\Delta _{\text{comb}}H^{\ominus }({\text{CH}}_{4}).}$

The value of ⁠${\displaystyle \Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4})}$⁠ is determined to be −74.8 kJ/mol. The negative sign shows that the reaction, if it were to proceed, would be

It is possible to predict heats of formation for simple unstrained organic compounds with the heat of formation group additivity method.

The standard enthalpy change of any reaction can be calculated from the standard enthalpies of formation of reactants and products using Hess's law. A given reaction is considered as the decomposition of all reactants into elements in their standard states, followed by the formation of all products. The heat of reaction is then minus the sum of the standard enthalpies of formation of the reactants (each being multiplied by its respective stoichiometric coefficient, ν) plus the sum of the standard enthalpies of formation of the products (each also multiplied by its respective stoichiometric coefficient), as shown in the equation below:^[4]

${\displaystyle \Delta _{\text{r}}H^{\ominus }=\sum \nu \Delta _{\text{f}}H^{\ominus }({\text{products}})-\sum \nu \Delta _{\text{f}}H^{\ominus }({\text{reactants}}).}$

If the standard enthalpy of the products is less than the standard enthalpy of the reactants, the standard enthalpy of reaction is negative. This implies that the reaction is exothermic. The converse is also true; the standard enthalpy of reaction is positive for an endothermic reaction. This calculation has a tacit assumption of ideal solution between reactants and products where the enthalpy of mixing is zero.

For example, for the combustion of methane, ${\displaystyle {\ce {CH4 + 2O2 -> CO2 + 2H2O}}}$:

${\displaystyle \Delta _{\text{r}}H^{\ominus }=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{}{\text{O}})]-[\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4})+2\Delta _{\text{f}}H^{\ominus }({\text{O}}_{2})].}$

However ${\displaystyle {\ce {O2}}}$ is an element in its standard state, so that ${\displaystyle \Delta _{\text{f}}H^{\ominus }({\text{O}}_{2})=0}$, and the heat of reaction is simplified to

${\displaystyle \Delta _{\text{r}}H^{\ominus }=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{}{\text{O}})]-\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4}),}$

which is the equation in the previous section for the enthalpy of combustion ${\displaystyle \Delta _{\text{comb}}H^{\ominus }}$.

• When a reaction is reversed, the magnitude of ΔH stays the same, but the sign changes.
• When the balanced equation for a reaction is multiplied by an integer, the corresponding value of ΔH must be multiplied by that integer as well.
• The change in enthalpy for a reaction can be calculated from the enthalpies of formation of the reactants and the products
• Elements in their standard states make no contribution to the enthalpy calculations for the reaction, since the enthalpy of an element in its standard state is zero. Allotropes of an element other than the standard state generally have non-zero standard enthalpies of formation.

Thermochemical properties of selected substances at 298.15 K and 1 atm

• Calorimetry
• Thermochemistry