The standard enthalpy of formation ($$\Delta_fH^\circ_{298}$$) for methane, CH$$_4$$ is $$-74.9$$ kJ mol$$^{-1}$$. In order to calculate the average energy given out in the formation of a C-H bond from this it is necessary to know which one of the following?

To calculate the average $$\text{C}-\text{H}$$ bond energy from the standard enthalpy of formation of methane ($$\Delta_f H^\circ$$), one must convert the elements from their standard states into isolated gaseous atoms using Hess's Law. This requires knowing both the enthalpy of sublimation of carbon, $$\text{C(graphite)} \rightarrow \text{C}(g)$$, and the bond dissociation energy of hydrogen, $$\text{H}_2(g) \rightarrow 2\text{H}(g)$$.

Combining these values allows to determine the total atomization energy of methane ($$\Delta H_{\text{atomization}}$$). Dividing this total energy by $$4$$ yields the average energy of a single $$\text{C}-\text{H}$$ bond.