We are asked to calculate the enthalpy change of formation of CH_{4}(g)

$\overline{){\mathbf{\Delta H}}{{\mathbf{\xb0}}}_{{\mathbf{rxn}}}{\mathbf{=}}{\mathbf{\Delta H}}{{\mathbf{\xb0}}}_{\mathbf{f}\mathbf{,}\mathbf{}\mathbf{product}}{\mathbf{-}}{\mathbf{\Delta H}}{{\mathbf{\xb0}}}_{\mathbf{f}\mathbf{,}\mathbf{reactant}}}$

Note that we need to *multiply each ΔH˚ by the stoichiometric coefficient* since ΔH˚ is in kJ/mol.

Balanced reaction : C(s) + 2 H_{2}(g) →CH_{4}(g) .

The methane molecule, CH_{4}, has the geometry shown in following figure. Imagine a hypothetical process in which the methane molecule is "expanded," by simultaneously extending all four C–H bonds to infinity. We then have the process CH_{4}(g) → C(g) + 4 H(g).

Calculate the enthalpy change of formation of CH_{4}(g)

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