Compare the value you obtained from average bond energies: (260Kj) to the actual standard enthalpy of formation of gaseous benzene, 82.9 kj. What does the difference between these two values tell you about the stability of benzene?
To compare the value obtained from average bond energies (260 kJ) to the actual standard enthalpy of formation of gaseous benzene (82.9 kJ), we can calculate the difference between these two values.
The difference is calculated by subtracting the actual standard enthalpy of formation from the value obtained from average bond energies:
Difference = Value from average bond energies - Actual standard enthalpy of formation
Difference = 260 kJ - 82.9 kJ
Difference = 177.1 kJ
The positive value indicates that the value obtained from average bond energies is higher than the actual standard enthalpy of formation of gaseous benzene. This suggests that the actual enthalpy of formation requires less energy input, making it a more stable configuration.
In other words, the difference between these two values suggests that benzene is more stable than what would be expected based on the average bond energies. The stability of benzene can be attributed to its delocalized π-electron system which provides additional stabilization compared to the individual bond strengths present in the molecule.
To compare the two values and understand what they suggest about the stability of benzene, we need to recognize the relationship between bond energies and the standard enthalpy of formation.
The average bond energy represents the average amount of energy required to break a particular type of bond, on average, in a variety of different molecules. In this case, the average bond energy of benzene is given as 260 kJ.
On the other hand, the standard enthalpy of formation of a compound represents the change in enthalpy when one mole of that compound is formed from its constituent elements in their standard states. In this case, the standard enthalpy of formation of gaseous benzene is 82.9 kJ.
Comparing the two values, we can make the following observations:
1. The average bond energy of benzene (260 kJ) is higher than the standard enthalpy of formation (82.9 kJ).
2. The bond energy represents the energy required to break the existing bonds in benzene, while the enthalpy of formation represents the energy released when the new bonds are formed.
3. The higher bond energy suggests that it requires more energy to break the bonds in benzene than is released when the compound is formed.
4. Thus, the higher bond energy indicates that benzene is relatively stable because it requires a significant amount of energy to break its bonds.
In summary, the difference between the average bond energy and the standard enthalpy of formation of benzene suggests that benzene is a stable compound since it requires more energy to break its bonds than is released when the compound is formed.