- #1
zenterix
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- Homework Statement
- For a diatomic molecule the bond energy is equal to the change in internal energy for the reaction
$$\mathrm{X_2(g)=2X(g)}$$
at 0K.
Of course, the change in internal energy and the change in enthalpy are the same at 0K.
Calculate the enthalpy of dissociation of ##\mathrm{O_2(g)}## at 0K.
The enthalpy of formation of ##\mathrm{O(g)}## at 298.15K is ##249.173\mathrm{kJ\ mol^{-1}}##.
In the range ##0-298.15\text{K}##, the average value of the heat capacity of ##\mathrm{O_2(g)}## is ##\mathrm{29.1 J\ K^{-1}mol^{-1}}## and the average heat capacity of ##\mathrm{O(g)}## is ##22.7\mathrm{J\ K^{-1}mol^{-1}}##.
- Relevant Equations
- What is the value of the bond energy in electron volts?
We can compute the enthalpy of reaction of ##\mathrm{O_2(g)\rightarrow 2O(g)}## at 0K by heating reactant to 298.15K, doing the reaction and obtaining the product at this temperature, and then cooling the product down to 0K.
The result is
$$\Delta H_r=\mathrm{493.48kJ\ mol^{-1}}$$
According to the problem statement, this equals ##\Delta U## for the reaction.
How do we calculate the bond energy? The problem statement says that the bond energy equals the change in ##U## (equivalently the change in ##H##).
When I convert ##\Delta H## to ##\text{eV}## I get ##3.08\times 10^{24}\text{eV}##.
The back of the book says the answer is ##5.115\text{eV}##.
What am I missing here?
The result is
$$\Delta H_r=\mathrm{493.48kJ\ mol^{-1}}$$
According to the problem statement, this equals ##\Delta U## for the reaction.
How do we calculate the bond energy? The problem statement says that the bond energy equals the change in ##U## (equivalently the change in ##H##).
When I convert ##\Delta H## to ##\text{eV}## I get ##3.08\times 10^{24}\text{eV}##.
The back of the book says the answer is ##5.115\text{eV}##.
What am I missing here?
