Answer :
Final Answer:
69 kJ/mol is the average C−F bond energy, thus the correct answer is: a) 69 kJ/mol
Explanation:
The average C₋F bond energy can be calculated using the heat of atomization data for CH₄ and CH₂F₂. The difference in the heats of atomization provides the energy required to break the C₋H bonds in CH₄ and replace them with C₋F bonds in CH₂F₂.
Firstly, find the energy change for breaking C₋H bonds in CH₄:
ΔH(CH₄) = Heat of atomization of CH₄
ΔH(CH₄) = 1660 kJ/mol
Then, find the energy change for breaking C₋H bonds and forming C₋F bonds in CH₂F₂:
ΔH(CH₂F₂) = Heat of atomization of CH₂F₂
ΔH(CH₂F₂) = 1798 kJ/mol
Now, subtract the energy change for CH₄ from that of CH₂F₂ to get the energy change for replacing C₋H with C₋F bonds:
ΔH(C₋F) = ΔH(CH₂F₂) - ΔH(CH₄)
ΔH(C₋F) = 1798 kJ/mol - 1660 kJ/mol
ΔH(C₋F) = 138 kJ/mol
Since there are two C₋F bonds formed in the process, calculate the average C₋F bond energy:
Average C₋F bond energy = ΔH(C₋F) / Number of bonds
Average C₋F bond energy = 138 kJ/mol / 2
Average C₋F bond energy = 69 kJ/mol
Considering the given options, the closest answer is: a) 69 kJ/mol