Given the following chemical reaction: 2H2(g) + O2(g) → 2H2O(l), which of the following best describes how to calculate the enthalpy change (ΔH) for the overall reaction using standard enthalpies of formation?

• A. ΔH = ΔH_f(H2O(l)) - ΔH_f(H2(g))
• B. ΔH = ΔH_f(H2O(l)) + ΔH_f(H2(g)) + ΔH_f(O2(g))
• C. ΔH = [2 × ΔH_f(H2O(l))] - [2 × ΔH_f(H2(g)) + ΔH_f(O2(g))]
• D. ΔH = [2 × ΔH_f(H2(g)) + ΔH_f(O2(g))] - [2 × ΔH_f(H2O(l))]

Answer: (C)

Verified step by step guidance

1. Identify the balanced chemical equation: \$2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$, noting the coefficients for each species.
2. Recall that the enthalpy change of the reaction, \(\Delta H\), can be calculated using standard enthalpies of formation (\(\Delta H_f^\circ\)) with the formula: \(\Delta H = \sum \Delta H_f^\circ (\text{products}) - \sum \Delta H_f^\circ (\text{reactants})\).
3. Apply the formula by multiplying the standard enthalpy of formation of each compound by its coefficient in the balanced equation: for products, multiply \(\Delta H_f^\circ(H_2O(l))\) by 2; for reactants, multiply \(\Delta H_f^\circ(H_2(g))\) by 2 and add \(\Delta H_f^\circ(O_2(g))\) multiplied by 1.
4. Write the expression explicitly as: \(\Delta H = [2 \times \Delta H_f^\circ(H_2O(l))] - [2 \times \Delta H_f^\circ(H_2(g)) + 1 \times \Delta H_f^\circ(O_2(g))]\).
5. Remember that the standard enthalpy of formation for elemental forms in their standard state (like \(H_2(g)\) and \(O_2(g)\)) is zero, which simplifies the calculation but should still be included in the formula for completeness.