Question

hydrazine (N2H4) decomposes according to the following reaction:
3N2H4 yields 4NH3 + N2

(a) Given that the standard enthalpy of formation of Hydrazine is 50.42 kJ/mol, calculate delta H for its composition.

(b) Both hydrazine and ammonia burn in oxygen to produce H2O(l) and N2(g). Write a balanced equation for each of these processes, and calculate delta H for each ofthem.

(c) On a mass basis (per kg), which would be the better fuel: hydrazine or ammonia?

Answer 1

3 N2H4(l) -> 4NH3(g) + N2(g)

A) There should be a table you can use to find the actual ΔH°for each substance, then multiply by the coefficient to get your answer.
ΔH°=[4(NH3 kj/mol) + 1(0 kj/mol)]-[3(50.42 kj/mol)]

B) Comparing ΔH°rxn of ammonia gas to liquid hydrazine.
N2H4 + O2 -> N2(g) +2H2O(l)
ΔH°rxn= [1(0 kJ/mol) + 2(H2O kJ/mol)] - [1*(N2H4 kJ/mol) +1(0 kJ/mol)]

4NH3(g) + 3 O2(g) -> 2N2(g) + 6H2O(g) (complete balanced eqn)

2NH3(g) + 3/2 O2(g) -> N2(g) + 3H2O(g) (complete balanced eqn-I divided the eqn by 2)
ΔH°rxn=[1*(0 kj/mol) +3*(H2O kJ/mol)] - [1*(NH3 kj/mol) +(3/2)(0 kj/mol)]

C) This part I'm a little unsure about... here's what I would do:

Take the deltaH was for both liquid hydrazine and ammonia found in part B and multiply by the molar mass of each to find the mass/kJ:

[1 mol/(ΔH°N2H4) kJ][ 18g N2H4/mol)](.001 kg)

Then for ammonia

[1 mol/(ΔH°NH3) kJ][ 10g NH3/mol)](.001 kg)