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

Explanation :

First we have to calculate the moles of

Now we have to calculate the ΔH for this reaction.

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

So, 1 mole of react to gives heat =

                                                                                = -1516 kJ/mol

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

b) mutualism

hope this : )

answer: enthalpy

explanation:

for a exothermic reaction, the energy is released i.e. the energy of the products is lower than the energy of the reactants.

= enthaply

= energy of products

= energy of reactants

as , thus comes out to be negative. negative value of enthalpy shows that the reaction is exothermic.