Answer to Question #205207 in Molecular Physics | Thermodynamics for Ankush

Question

The energy from the combustion of Hydrazine, N2H4, is used to power rockets into space in the reaction:

N2H4(g)+O2(g)=N2(g)+2H2O(l) delta H degree= -627.6KJ

How many killograms of hydrazine would be necessary to produce 1.0 X 10^8 KJ of energy?

Hint: One mole of N2H4 produces 627.6KJ of energy. How many moles (and then grams) are required to produce 1.0 X 10^8 KJ of energy.

Answer 1

$n=\frac{1.0 \times 10^8}{627.6} = 1.5933 \times 10^5 \;mol$

MM(N₂H₄)= 32.04 g/mol

$m = 1.5933 \times 10^5 \times 32.04 = 51.05 \times 10^5 \;g = 51.05 \times 10^2 \;kg \\ = 5105 \;kg$

Answer: 5105 kg