Methane burns completely in oxygen according to the equation: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) The bond energies are: - C–H bond: 413 kJ/mol - O=O bond: 498 kJ/mol - C=O bond: 805 kJ/mol - O–H bond: 463 kJ/mol (a) One molecule of CH4 contains four C–H bonds, and each O2 molecu…

Answer: ((a)) 4 C–H bonds and 2 O=O bonds broken; energy absorbed = (4 x 413) + (2 x 498) = 1652 + 996 = 2648 kJ/mol ((b)) 2 C=O bonds and 4 O–H bonds formed; energy released = (2 x 805) + (4 x 463) = 1610 + 1852 = 3462 kJ/mol ((c)) ΔH = 2648 − 3462 = −814 kJ/mol; the reaction is exothermic.. Part (a): Bonds broken and energy absorbed One mole of CH4 contains four C–H bonds, so all four must be broken. The equation needs two moles of O2, and each O2 molecule has one O=O bond, so two O=O bonds must be broken. 4 × (C–H) = 4 × 413…

IGCSE Chemistry 0620: Bond Energy Calculation for the Combustion of Methane

Methane burns completely in oxygen according to the equation:

$\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g)$

The bond energies are:

C–H bond: $413\,\text{kJ/mol}$
O=O bond: $498\,\text{kJ/mol}$
C=O bond: $805\,\text{kJ/mol}$
O–H bond: $463\,\text{kJ/mol}$

(a) One molecule of $\text{CH}_4$ contains four C–H bonds, and each $\text{O}_2$ molecule contains one O=O bond. Determine the total number of each bond broken when one mole of $\text{CH}_4$ reacts completely, and calculate the total energy absorbed breaking these bonds. [3]

(b) One molecule of $\text{CO}_2$ contains two C=O bonds, and each $\text{H}_2\text{O}$ molecule contains two O–H bonds. Determine the total number of each bond formed, and calculate the total energy released forming these bonds. [3]

(c) Calculate the overall energy change, $\Delta H$ , for the complete combustion of one mole of methane, and state whether the reaction is exothermic or endothermic. [2]

Part (a): Bonds broken and energy absorbed

One mole of $\text{CH}_4$ contains four C–H bonds, so all four must be broken. The equation needs two moles of $\text{O}_2$ , and each $\text{O}_2$ molecule has one O=O bond, so two O=O bonds must be broken.

$4 \times (\text{C–H}) = 4 \times 413 = 1652\,\text{kJ/mol}$ $2 \times (\text{O=O}) = 2 \times 498 = 996\,\text{kJ/mol}$

Total energy absorbed breaking bonds: $1652 + 996 = 2648\,\text{kJ/mol}$

Part (b): Bonds formed and energy released

One mole of $\text{CO}_2$ is formed, containing two C=O bonds. Two moles of $\text{H}_2\text{O}$ are formed, and each $\text{H}_2\text{O}$ molecule has two O–H bonds, giving four O–H bonds in total.

$2 \times (\text{C=O}) = 2 \times 805 = 1610\,\text{kJ/mol}$ $4 \times (\text{O–H}) = 4 \times 463 = 1852\,\text{kJ/mol}$

Total energy released forming bonds: $1610 + 1852 = 3462\,\text{kJ/mol}$

Part (c): Overall energy change and classification

$\Delta H = (\text{energy absorbed breaking bonds}) - (\text{energy released forming bonds})$

$\Delta H = 2648 - 3462 = -814\,\text{kJ/mol}$

Since $\Delta H$ is negative, considerably more energy is released forming the new bonds in carbon dioxide and water than is absorbed breaking the bonds in methane and oxygen. The reaction is strongly exothermic, consistent with methane being burned as a fuel.

Final answers

(a) 4 C–H bonds and 2 O=O bonds broken; energy absorbed $=$ $2648\,\text{kJ/mol}$
(b) 2 C=O bonds and 4 O–H bonds formed; energy released $=$ $3462\,\text{kJ/mol}$
(c) $\Delta H = \boxed{-814}\,\text{kJ/mol}$ ; the reaction is exothermic.