Question

How much heat will be released when 12g of ${H_2}$ reacts with 76 g of ${O_2}$ according to the following equation?
$2{H_2} + {O_2} \to 2{H_2}{{O \Delta H = - 571}}{\text{.6kJ}}$

Answer 1

This could be simply solved by applying the concepts of mole calculations. Also, we know that limiting reagent in a chemical reaction is a reactant that is totally consumed when the chemical reaction is completed, so here we need these concepts.

Formula used:
Here, we will use the number of moles formula:
${\text{no of moles = }}\dfrac{{{\text{given mass}}}}{{{\text{molecular mass}}}}$

Complete step by step answer
We are given that 12g of ${H_2}$ reacts with 76 g of ${O_2}$ .
Given equation:
$2{H_2} + {O_2} \to 2{H_2}{\text{O }}$
We are going to identify the moles of respective reactants.
Moles of dihydrogen= $\dfrac{{12{\text{g}}}}{{2{\text{g}}{\text{.mo}}{{\text{l}}^{{\text{ - 1}}}}}} = 6{\text{mol}}$
Moles of dioxygen= $\dfrac{{76{\text{g}}}}{{32{\text{g}}{\text{.mo}}{{\text{l}}^{{\text{ - 1}}}}}} = 2.38{\text{mol}}$
Given the stoichiometry, clearly, there is an insufficient molar quantity of dioxygen for complete combustion.
According to the given equation.
At most $4.75$ mol hydrogen can react which will be = $2 \times 2.38{\text{mol}}$
And so, energy released can based on the molar quantity of dioxygen gas = $2.38 \times - 576.6{\text{kJ}}{\text{.mo}}{{\text{l}}^{{\text{ - 1}}}} = 1372.308{\text{kJ}}$
Dioxygen gas is the limiting reagent.
Thus, energy released can be based on the molar quantity of dioxygen gas = $1.3 \times {10^3}{\text{kJ}}$ .

Note:
It should always be kept in mind that limiting reagent is defined as the reactant in a chemical reaction that limits the amount of product that can be formed. The reaction will stop when all of the limiting reactant is consumed. Excess Reactant - The reactant in a chemical reaction that remains when a reaction stops when the limiting reactant is completely consumed. In much the same way, a reactant in a chemical reaction can limit the amounts of products formed by the reaction. When this happens, we refer to the reactant as the limiting reactant (or limiting reagent).