# Given the combustion of silane... SiH_4(g)+2O_2(g)rarrSiO2(s)+2H_2O(g); DeltaH=−1429⋅kJ⋅mol^(−1).. How much heat will evolve given combustion of 15.7*g quantity?

Nov 28, 2016

Approx. $700 \cdot k J$

#### Explanation:

Thermochemical reactions ALWAYS quote enthalpy input/output PER MOLE OF REACTION AS WRITTEN:

$S i {H}_{4} \left(g\right) + 2 {O}_{2} \left(g\right) \rightarrow S i {O}_{2} \left(s\right) + 2 {H}_{2} O \left(g\right)$ $\Delta H = - 1429 \cdot k J \cdot m o {l}^{-} 1$.

That is $- 1429 \cdot k J \cdot m o {l}^{-} 1$ of energy are involved when $1 \cdot m o l$, $33 \cdot g$ of silane are combusted to give stoichiometric silicon oxide, and water.

And thus we simply work out the molar quantity of silane:

$\frac{15.7 \cdot g}{32.12 \cdot g \cdot m o {l}^{-} 1} = 0.489 \cdot m o l$

And $0.489 \cdot m o l \left(\text{of silane")xx-1429*kJ*mol^-1("of silane}\right)$ $=$ ??kJ.