Enthalpies of formation and combustion

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The enthalpy change when **1 mol** of a compound forms from its **elements in their standard states** (100 kPa, stated T).

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The enthalpy change when **1 mol** of a substance is **completely burned in oxygen** under standard conditions; always **negative**.

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**Zero** — e.g. O_{2}(g), N_{2}(g), C(graphite); there is nothing to form.

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$\Delta H^{\ominus} = \sum \Delta H_{f}^{\ominus}(\text{products}) - \sum \Delta H_{f}^{\ominus}(\text{reactants})$.

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$\Delta H^{\ominus} = \sum \Delta H_{c}^{\ominus}(\text{reactants}) - \sum \Delta H_{c}^{\ominus}(\text{products})$.

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Both reactants and products burn down to the **same products** (CO_{2} + H_{2}O), so the Hess cycle runs the other way → **reactants − products**.

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A pressure of **100 kPa** and a stated temperature (usually **298 K**), with all substances in their standard states.

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Enthalpy is a **state function** — ΔH depends only on the start and end states, so a 'paper' Hess route gives the same answer as experiment.

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Forgetting to multiply each value by its **stoichiometric coefficient** (e.g. 2 H_{2}O) or forgetting an **element is zero**.

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**Negative** (exothermic) — a quick check that you used the correct rule.

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**kJ mol⁻¹** (kilojoules per mole).