Question

Hydrazine (N2H4) is a fuel used by some spacecraft. It is normally oxidized by N2O4 according to the following equation:
N2H4(l)+N2O4(g)->2N2O(g)+2H2O(g)

Calculate delta Hrxn for this reaction using standard enthalpies of formation.

I am having trouble finding the standard enthalpies of all of the reactants and products

Answer 1

The equation for the reaction is:

\(\mathrm{N}_{2} \mathrm{H}_{4}(1)+\mathrm{N}_{2} \mathrm{O}_{4}(\mathrm{~g}) \rightarrow 2 \mathrm{~N}_{2} \mathrm{O}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\)

The enthalpy of the reaction \(\left(\Delta \mathrm{H}_{\mathrm{mx}}\right)\) is equal to the difference in \(\mathrm{s}\) tan dard enthalpies of products and reactants:

\(\begin{aligned} \Delta \mathrm{H}_{\mathrm{rm}} &=\Delta \mathrm{H}_{\mathrm{f}}(\text { products })-\Delta \mathrm{H}_{\mathrm{f}}(\text { reactan } \mathrm{ts}) \\ &=\left[2 \Delta \mathrm{H}_{\mathrm{f}}\left(\mathrm{N}_{2} \mathrm{O}(\mathrm{g})\right)+2 \Delta \mathrm{H}_{\mathrm{f}}\left(\mathrm{H}_{2} \mathrm{O}(\mathrm{g})\right)\right]-\left[\Delta \mathrm{H}_{\mathrm{f}}\left(\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{l})\right)+\Delta \mathrm{H}_{\mathrm{f}}\left(\mathrm{N}_{2} \mathrm{O}_{4}(\mathrm{~g})\right)\right] \\ &=[2(81.6 \mathrm{~kJ} / \mathrm{mol})+2(-241.8 \mathrm{~kJ} / \mathrm{mol})]-[50.6 \mathrm{~kJ} / \mathrm{mol}+11.1 \mathrm{~kJ} / \mathrm{mol}] \\ &=-382.1 \mathrm{~kJ} / \mathrm{mol} \end{aligned}\)