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?

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Answer 1 · 2016-11-28T07:46:13

Approx. $700*kJ$

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

$SiH_4(g) + 2O_2(g) rarr SiO_2(s) + 2H_2O(g)$ $DeltaH=-1429*kJ*mol^-1$ .

That is $-1429*kJ*mol^-1$ of energy are involved when $1*mol$ , $33*g$ of silane are combusted to give stoichiometric silicon oxide, and water.

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

$(15.7*g)/(32.12*g*mol^-1)=0.489*mol$

And $0.489*mol("of silane")xx-1429*kJ*mol^-1("of silane")$ $=$ $??kJ$ .

Given the combustion of silane... SiH_4(g)+2O_2(g)rarrSiO2(s)+2H_2O(g); DeltaH=−1429⋅kJ⋅mol^(−1)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*g15.7*g quantity?