Question

In the presence of excess oxygen, methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) ΔH = -890.0 kJ Calculate the value of q (kJ) in this exothermic reaction when 1.80 g of methane is combusted at constant pressure.

Answer 1

Answer: The value of q for the reaction will be -100.1 kJ

Explanation:

To calculate the number of moles, we use the equation:

$\text{Number of moles}=\frac{\text{Given mass}}{\text{Molar mass}}$

Given mass of methane = 1.80 g

Molar mass of methane = 16 g/mol

Putting values in above equation, we get:

$\text{Moles of }CH_4=\frac{1.80g}{16g/mol}=0.1125mol$

For the given chemical reaction:

$CH_4(g)+2O_2(g)\rightarrow CO_2(g)+2H_2O(l);\Delta H^o_{rxn}=-890.0kJ$

By Stoichiometry of the reaction:

When 1 mole of methane reacts, the heat released is 890.0 kJ

So, when 0.1125 moles of methane will react, the heat released will be $\frac{890.0kJ}{1mol}\times 0.1125mol=100.1kJ$

Sign convention of heat (q):

When heat is absorbed, the sign of heat is taken to be positive and when heat is released, the sign of heat is taken to be negative.

Hence, the value of q for the reaction will be -100.1 kJ