The molar enthalpy of vapourisation of acetone is less than that of water. Why?

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Q. The molar enthalpy of vapourisation of acetone is less than that of water. Why?

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Concept understanding — Enthalpy and Types of Enthalpy Change

Enthalpy, $H$ , is a state function defined as $H = U + PV$ , introduced because the overwhelming majority of real chemical reactions take place in open vessels at essentially constant atmospheric pressure, rather than at constant volume. At constant pressure, the change in enthalpy equals the heat exchanged with the surroundings, $\Delta H = q_p$ — precisely the quantity a calorimeter measures for a reaction carried out in an open flask, container, or solution. Enthalpy and internal energy are related, for gas-phase reactions at constant temperature, through $\Delta H = \Delta U + \Delta n_g RT$ , where $\Delta n_g$ is the change in moles of gas across the reaction.

Chemists have defined a family of precisely standardized enthalpy changes, each describing one specific, carefully defined process, denoted by the standard-state superscript $^\circ$ (usually $298\ \text{K}$ , $1\ \text{bar}$ ):

Standard enthalpy of formation ( $\Delta_fH^\circ$ ) — one mole of a compound formed from its elements in their standard states; by convention, an element in its standard state has $\Delta_fH^\circ = 0$ .
Standard enthalpy of combustion ( $\Delta_cH^\circ$ ) — one mole of a substance completely burnt in excess oxygen; always exothermic (negative), and readily measured by bomb calorimetry.
Standard enthalpy of atomization ( $\Delta_aH^\circ$ ) — one mole of a substance completely converted into gaseous atoms; always endothermic (positive).
Bond dissociation enthalpy — the energy to break one mole of a specific covalent bond in a gaseous molecule; dividing an atomization enthalpy by the number of equivalent bonds gives an average bond enthalpy.
Enthalpy of ionization (first, second, …) — the energy to remove one mole of electrons from one mole of gaseous atoms (first) or from the resulting gaseous cations (second, always larger than the first, since removing an electron from an already-positive species is harder). …

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The difference arises from the strength of intermolecular forces: water has strong hydrogen bonding, while acetone has only dipole-dipole interactions. Hence, water requires more energy to vaporise.

The key is to think about what molar enthalpy of vapourisation ( $\Delta H_{\text{vap}}$ ) actually measures. It is the energy needed to overcome all intermolecular attractions holding molecules together in the liquid phase, so that they can escape into the gas phase. The stronger the forces between molecules, the more energy you must supply per mole to break them.

Water molecules are held together by hydrogen bonds — the strongest type of dipole-dipole interaction, formed between a hydrogen atom bonded to a highly electronegative atom (oxygen) and a lone pair on another electronegative atom. Each water molecule can form up to four hydrogen bonds, creating an extensive network.

Acetone molecules, on the other hand, interact primarily through dipole-dipole forces (the C=O bond is polar) and weaker London dispersion forces. There is no O–H or N–H bond, so no hydrogen bonding between acetone molecules.

Let’s compare the numbers to make it concrete.

Identify the intermolecular forces in each liquid.

Water: hydrogen bonding (dominant), plus dipole-dipole and dispersion.

Acetone: dipole-dipole (moderate) and dispersion (weak to moderate).
Relate force strength to energy required.

Hydrogen bonds are roughly 10–40 kJ/mol each, while typical dipole-dipole interactions are about 2–10 kJ/mol. To vaporise water, you must break a significant fraction of its hydrogen bonds — a large energy cost. For acetone, you only need to overcome weaker dipole-dipole attractions.
Check the actual values (for context).

At their boiling points:
- $\Delta H_{\text{vap}}$ of water ≈ 40.7 kJ/mol
- $\Delta H_{\text{vap}}$ of acetone ≈ 29.1 kJ/mol
The difference (~11.6 kJ/mol) is substantial and directly reflects the extra energy needed to break water’s hydrogen-bond network. …

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Q.The molar enthalpy of vapourisation of acetone is less than that of water. Why?

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