Question

Use average bond energies to estimate the value of Delta H for the reaction below. CH_4(g) + 2O_2(g) rightarrow CO_2(g) + 2H_2O(g)

Answer 1

To estimate the ΔH for the reaction, the total energy for breaking the CH4 and O2 bonds is calculated and then subtracted from the energy released by forming the CO2 and H2O bonds. The result, approximately 417.4 kJ, suggests the reaction is exothermic.

Step-by-step explanation:

To estimate the value of Delta H (ΔH) for the reaction CH4(g) + 2O2(g) → CO2(g) + 2H2O(g), we'll use average bond energies. First, we need to calculate the total energy required to break the bonds in the reactants (bond dissociation energy).

CH4: 4 C-H bonds x 415 kJ/mol each = 1660 kJ

2 O2: 2 O=O bonds x 498.7 kJ/mol each = 997.4 kJ

Then, we add up the energy released by forming the new bonds in the products:

CO2: 2 C=O bonds x -192.0 kJ/mol each = -384 kJ

2 H2O: 4 O-H bonds (since there are 2 H2O molecules, each with 2 O-H bonds) x -464 kJ/mol each = -1856 kJ

We calculate the approximate enthalpy change (ΔH) by subtracting the energy of bond formation (the energy released) from the energy of bond dissociation (the energy required):

ΔH = (Energy required to break bonds) - (Energy released by forming bonds)

ΔH = (1660 kJ + 997.4 kJ) - (-384 kJ - 1856 kJ)

ΔH = 2657.4 kJ - (-2240 kJ)

ΔH = 2657.4 kJ + 2240 kJ

ΔH ≈ 417.4 kJ

Therefore, the estimated value of ΔH for the given reaction is approximately 417.4 kJ, indicating that the reaction is exothermic.