Question 18

A researcher uses a table of average bond enthalpies to estimate the molar enthalpy change for the synthesis of hydrazine, a liquid rocket propellant, from nitrogen and hydrogen gas:

\text{N}_2(\text{g}) + 2\text{H}_2(\text{g}) \rightarrow \text{N}_2\text{H}_4(\text{g})

The average bond enthalpies are given in the table below:

Bond	Average bond enthalpy in $kJ/mol\text{kJ/mol}$
$N≡N\text{N}\equiv\text{N}$	944
$H−H\text{H}-\text{H}$	436
$N−N\text{N}-\text{N}$	163
$N−H\text{N}-\text{H}$	388

In this reaction:

Each nitrogen molecule ( $N2\text{N}_2$ ) contains one $N≡N\text{N}\equiv\text{N}$ triple bond.
Each hydrogen molecule ( $H2\text{H}_2$ ) contains one $H−H\text{H}-\text{H}$ single bond.
Each hydrazine molecule ( $N2H4\text{N}_2\text{H}_4$ ) contains one $N−N\text{N}-\text{N}$ single bond and four $N−H\text{N}-\text{H}$ single bonds.

Calculate the energy taken in (in $kJ\text{kJ}$ ) when the bonds in the reactants are broken.

Calculate the energy given out (in $kJ\text{kJ}$ ) when the bonds in the products are formed.

Use your answers to (i) and (ii) to calculate the molar enthalpy change, $ΔH\Delta H$ (in $kJ/mol\text{kJ/mol}$ ), for the synthesis of hydrazine.