We are asked to calculate ΔH for the following reaction: CH_{4}(g) + 4 F_{2}(g) → CF_{4}(g) + 4 HF(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.

You may want to reference (Pages 191 - 192) Section 5.8 and Appendix C (Pages 1088 - 1091) while completing this problem.

It is interesting to compare the "fuel value" of a hydrocarbon in a hypothetical world where oxygen is not the combustion agent. The enthalpy of formation of CF_{4}(g) is -679.9 kJ/mol

**Calculate ΔH for the following reaction: CH _{4}(g) + 4 F_{2}(g) → CF_{4}(g) + 4 HF(g)**

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