Acetone has two types of intermolecular forces: dipole-dipole interactions and London dispersion forces. It does not form hydrogen bonds with itself, despite being fully miscible with water. This combination of moderate intermolecular forces explains why acetone evaporates quickly, feels cool on the skin, and works as an effective solvent for so many substances.
Why Acetone Is a Polar Molecule
Acetone’s chemical formula is (CH₃)₂CO. The key feature is its carbonyl group: a carbon atom double-bonded to an oxygen atom right in the middle of the molecule. Oxygen pulls electrons much more strongly than carbon does, so the electrons in that double bond spend more time near the oxygen end. This creates a region of partial negative charge on the oxygen and partial positive charge on the carbon, giving the molecule a permanent dipole.
Acetone’s dipole moment measures about 2.91 Debye, which is substantial for a small organic molecule. That strong polarity is the reason acetone dissolves so many different compounds and has a relatively high dielectric constant for an organic solvent.
Dipole-Dipole Interactions
Because every acetone molecule has a permanent positive end and a permanent negative end, neighboring molecules orient themselves so that opposite charges face each other. In liquid acetone, pairs of molecules tend to line up in an antiparallel arrangement, with their carbonyl groups pointing in opposite directions. This lets the partially negative oxygen on one molecule sit near the partially positive carbon on its neighbor, lowering the overall energy and holding the liquid together.
These dipole-dipole forces are the dominant intermolecular attraction in pure acetone. They’re significantly stronger than London dispersion forces for a molecule this size, and they’re the main reason acetone is a liquid at room temperature rather than a gas.
London Dispersion Forces
Every molecule experiences London dispersion forces, including acetone. These arise from the constant, random motion of electrons. At any given instant, the electrons in a molecule may be distributed unevenly, creating a fleeting dipole that induces a complementary dipole in a neighboring molecule. The two molecules briefly attract each other before the electron cloud shifts again.
London dispersion forces are typically weak, but they scale with the size of the electron cloud. Acetone has a molar mass of about 58 g/mol, which is relatively small, so its dispersion forces are modest. They contribute to the total attraction between acetone molecules but play a secondary role compared to the dipole-dipole interactions.
Why Acetone Doesn’t Hydrogen Bond With Itself
Hydrogen bonding requires two things: a hydrogen atom bonded to a highly electronegative atom (like oxygen or nitrogen) and a lone pair on a nearby electronegative atom to accept that hydrogen bond. Acetone has the second part covered. Its carbonyl oxygen carries lone pairs and is a strong hydrogen bond acceptor, with an acceptor strength rated at 5.7 on standard solvent scales.
What acetone lacks is a hydrogen bond donor. Its hydrogen atoms are all bonded to carbon, which isn’t electronegative enough to create the kind of strong partial positive charge on hydrogen that drives hydrogen bonding. So in a container of pure acetone, no hydrogen bonds form between molecules. This is why acetone is classified as a polar aprotic solvent: polar because of the dipole, aprotic because it can’t donate hydrogen bonds.
Acetone Can Hydrogen Bond With Water
When you mix acetone with water, something different happens. Water molecules have O-H bonds, making them excellent hydrogen bond donors. The carbonyl oxygen on acetone readily accepts hydrogen bonds from water, forming a bridge between the two types of molecules. Studies using molecular dynamics simulations show that the hydrogen atoms of water approach within about 2.4 angstroms of acetone’s carbonyl oxygen, a distance consistent with a true hydrogen bond.
This is why acetone and water mix in any proportion. The hydrogen bonds between them are strong enough to overcome the energy cost of disrupting water’s own hydrogen bonding network. If acetone couldn’t accept hydrogen bonds, it would behave much more like an oil in water.
How These Forces Shape Acetone’s Properties
Acetone’s intermolecular forces are strong enough to keep it liquid at room temperature but weak enough that it evaporates easily. Its boiling point is about 56°C (133°F), well below water’s 100°C. That gap reflects the absence of hydrogen bonding in pure acetone. Water molecules cling to each other through hydrogen bonds, dipole-dipole interactions, and dispersion forces all at once, requiring much more energy to pull apart.
Acetone’s vapor pressure at 20°C is roughly 246 millibar, meaning it produces a lot of vapor even at room temperature. If you’ve ever spilled nail polish remover and noticed how fast it disappears, that’s the practical result of having only dipole-dipole and dispersion forces holding the liquid together.
The surface tension of acetone at 20°C is 23.7 mN/m, less than a third of water’s 72.8 mN/m. Its viscosity is just 0.32 centipoise, making it feel noticeably thinner and more “watery” than water itself. Both of these properties reflect weaker intermolecular attractions. Fewer and weaker forces between molecules mean less resistance to flow and less cohesion at the surface.
Comparing Acetone to Similar Molecules
A useful way to see how intermolecular forces work is to compare molecules of similar size but different polarity. Propane has nearly the same molar mass as acetone (about 44 vs. 58 g/mol) but is completely nonpolar, relying only on London dispersion forces. It boils at a frigid -42°C and is a gas at room temperature.
Ethanol has a molar mass of 46 g/mol, close to acetone’s, but it can both donate and accept hydrogen bonds through its O-H group. It boils at 78°C, more than 20 degrees higher than acetone. The progression from propane to acetone to ethanol neatly illustrates how each additional type of intermolecular force (dispersion alone, then dipole-dipole, then hydrogen bonding) raises the boiling point by requiring more energy to separate molecules from the liquid.