Chemistry, 15.04.2020 02:05, melissakm77
Given that Δ H ∘ f [ H ( g ) ] = 218.0 kJ ⋅ mol − 1 Δ H ∘ f [ C ( g ) ] = 716.7 kJ ⋅ mol − 1 Δ H ∘ f [ CH 4 ( g ) ] = − 74.6 kJ ⋅ mol − 1 calculate the average molar bond enthalpy of the carbon‑hydrogen bond in a CH 4 molecule
Answers: 2
is 415.825 kJ/mol.
![\Delta H^{0}=[1mol\times \Delta H_{f}^{0}(C)_{g}]+[4mol\times \Delta H_{f}^{0}(H)_{g}]-[1mol\times \Delta H_{f}^{0}(CH_{4})_{g}]](/tpl/images/0600/7413/0c0db.png)
![[1mol\times 716.7\frac{kJ}{mol}]+[4mol\times 218.0\frac{kJ}{mol}]-[1mol\times -74.6\frac{kJ}{mol}]=1663.3kJ](/tpl/images/0600/7413/961ee.png)
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Given that Δ H ∘ f [ H ( g ) ] = 218.0 kJ ⋅ mol − 1 Δ H ∘ f [ C ( g ) ] = 716.7 kJ ⋅ mol − 1 Δ H ∘ f...
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