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 1.70 g of methane is combusted at constant pressure.
(A) -94.6 kJ
(B) 0.0306 kJ
(C) -0.0106 kJ
(D) 32.7 kJ
(E) -9.46 × 10^4 kJ

Answer 1

Answer:

The energy released will be -94.56 kJ or -94.6 kJ.

Explanation:

The molar mass of methane is 16g/mol

The given reaction is:

$CH_{4}(g) + 2O_{2} (g) --> CO_{2} (g)+ 2H_{2}O(l)$

the enthalpy of reaction is given as  ΔH = -890.0 kJ

This means that when one mole of methane undergoes combustion it gives this much of energy.

Now as given that the amount of methane combusted = 1.70g

The energy released will be:

$=\frac{energy released by one moleXgiven mass}{molarmass} =\frac{-890X1.7}{16}= -94.56 kJ$

Answer 2

-94.6 kJ of heat is evolved at constant pressure.

The equation of the reaction is;

CH4 (g) 2O2 (g) → CO2 (g) 2H2O(l) ΔH = -890.0 kJ

Number of moles of methane = mass/molar mass

Molar mass of methane = 16 g/mol

Mass of methane reacted =  1.70 g

Number of moles of methane = 1.70 g /16 g/mol = 0.106 moles

If 1 mole of CH4 produces  -890.0 kJ of heat

0.1 moles of CH4 produces  0.106 moles × -890.0 kJ /1 mole

= -94.6 kJ of heat

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