Q.What is meant by the term average bond enthalpy? Why is there difference in bond enthalpy of O—H bond in ethanol (C2H5OH) and water?

Concept: Average bond enthalpy and its dependence on molecular environment.

Average bond enthalpy is the mean energy required to break one mole of a particular type of bond in the gaseous state, averaged over a range of different compounds containing that bond. Since the same bond (say C—H) has slightly different strengths in different molecules due to varying electronic environments, we use an average value for thermochemical calculations.

The O—H bond enthalpy differs between ethanol and water because bond strength depends on the molecular environment. In water, the oxygen atom is bonded to two hydrogen atoms and carries two lone pairs; in ethanol, the oxygen is bonded to one hydrogen and one carbon (the ethyl group). The electron-withdrawing or electron-donating effect of the neighboring carbon chain alters the electron density around the O—H bond, changing the force constant and hence the bond dissoci …

What is average bond enthalpy?

Bond enthalpy measures the strength of a chemical bond: specifically, the energy required to break one mole of that bond in the gaseous state, separating the atoms completely. But here's the subtlety—the same type of bond (say, C—H or O—H) does not have exactly the same strength in every molecule. The electron cloud around a bond, and hence its strength, depends on what else is attached to those atoms.

Average bond enthalpy is the arithmetic mean of the bond dissociation enthalpies for a particular type of bond measured across many different molecules. For instance, the C—H bond appears in methane, ethane, propane, and countless other compounds, each time with a slightly different strength. We average all these values to get a single representative number—the average C—H bond enthalpy—that we can use in thermochemical calculations.

This averaging is practical: it lets us estimate enthalpy changes for reactions without needing the exact bond strength in every specific molecule. The trade-off is that our estimates are approximate, not exact.

Why does O—H bond enthalpy differ in ethanol and water?

The O—H bond does not exist in isolation. Its strength is influenced by the molecular environment—the atoms and groups bonded to the oxygen.

1. In water ($\ce{H2O}$): Oxygen is bonded to two hydrogen atoms. The molecule is small, highly polar, and the oxygen atom experiences a particular electron density distribution. When we break the first O—H bond in water, we get:

Breaking the second O—H bond in the hydroxyl radical $\ce{OH}$ requires a different energy:

These two dissociation energies are not equal—the first is about 498 kJ/mol, the second about 428 kJ/mol. The average O—H bond enthalpy in water is the mean of these two: roughly 463 kJ/mol.

2. In ethanol ($\ce{C2H5OH}$): Oxygen is bonded to one hydrogen and to an ethyl group ($\ce{C2H5}$). The ethyl group is electron-donating compared to hydrogen, which increases electron density on the oxygen atom. This alters the polarity and strength of the O—H bond. The dissociation:

requires a different energy—typically around 436 kJ/mol—because the $\ce{C2H5O}$ radical left behind is stabilized differently than $\ce{OH}$.

3. The root cause: …