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Methane cannot be made cleanly from carbon and hydrogen, so its Δf H° is found with Hess's law.

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Methane cannot be made cleanly from carbon and hydrogen, so its $\Delta_f H^\ominus$ is found with Hess's law. The standard enthalpies of combustion are $-393.5\ \text{kJ/mol}$ for graphite, $-285.8\ \text{kJ/mol}$ for hydrogen and $-890.3\ \text{kJ/mol}$ for methane. Find $\Delta_f H^\ominus$ of $\text{CH}_4\text{(g)}$ .

Need a reminder? The key ideaShow

Hess's law: the enthalpy change of a reaction is the sum of the enthalpy changes of any set of steps that lead from the same start to the same end. Equations add like algebra: reverse one and its $\Delta H$ changes sign, multiply one and its $\Delta H$ is multiplied. For gases, $\Delta_r H \approx \sum(\text{bonds broken}) - \sum(\text{bonds formed})$ , using mean bond enthalpies.

The worked solution

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Methane cannot be made cleanly from carbon and hydrogen, so its Δf H° is found with Hess's law.

The worked solution

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