A.P. CHEMISTRY CHAPTER NINE. COVALENT BONDING: ORBITALS#74. In Exercise 75 in Chapter 8, the Lewis structures for benzene (C6H6) were drawn. Using one of the Lewis structures, estimate ΔH of f for...

A.P. CHEMISTRY CHAPTER NINE. COVALENT BONDING: ORBITALS

#74. In Exercise 75 in Chapter 8, the Lewis structures for benzene (C6H6) were drawn. Using one of the Lewis structures, estimate ΔH of f for C6H6(g) using bond energies and given that the standard enthalpy of formation of C(g) is 717 kJ/mol. The experimented ΔH of f value of C6H6(g) is 83 kJ/mol. Explain the discrepancy between the experimental value and the calculated ΔH of f value for C6H6(g).

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llltkl | College Teacher | (Level 3) Valedictorian

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Relevant bond energy values are (in kJ/mol) H-H 432, H-C 411, C-C 346, C=C 602. This gives, ΔHf for C6H6(g) as =6*411+3*346+3*602 = 5310 kJ/mol.

6*717 + 3*432 = 5598 kJ/mol energy required to produce separate atoms from reactant molecules. 5310 kJ/mol is the ∆Hf of benzene. That means this amount of energy is given off when benzene is formed. So, theoretically (5598-5310) = 288 kJ/mol energy should have been absorbed when benzene is formed. But the observed (experimental) value of ∆Hf of benzene is only 83 kJ/mol. So the real benzene molecule is (288-83) = 205 kJ/mol more stable than any single structure would predict. This is due to resonance stabilization and this difference in energy is known as resonance energy of benzene.

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