Question

The burning of 80.3 g of SiH4 at constant pressure gives off 3790 kJ of heat. Calculate △H for this reaction. SiH4(g) + 2O2(g) ⟶ SiO2(s) + 2H2O(l).
A) -1520 kJ/mol rxn.
B) - 47.2 kJ/mol rxn.
C) - 4340 kJ/mol rxn.
D) -2430 kJ/mol rxn.
E) + 4340 kJ/mol rxn.

Answer 1

Answer  : The value of ΔH for this reaction is, -1516 kJ/mol

Explanation :

First we have to calculate the moles of $SiH_4$

$\text{Moles of }SiH_4=\frac{\text{Mass of }SiH_4}{\text{Molar mass of }SiH_4}$

$\text{Moles of }SiH_4=\frac{80.3g}{32.12g/mol}$

$\text{Moles of }SiH_4=2.5mol$

Now we have to calculate the ΔH for this reaction.

As, 2.5 mole of $SiH_4$ react to gives heat = -3790 kJ

So, 1 mole of $SiH_4$ react to gives heat = $-\frac{3790kJ}{2.5mol}$

                                                                                = -1516 kJ/mol

Therefore, the value of ΔH for this reaction is, -1516 kJ/mol