Acetone can participate in hydrogen bonding, but only as an acceptor, not a donor. Its carbonyl oxygen (the C=O group) has lone pairs that readily form hydrogen bonds with molecules like water and alcohols, yet acetone lacks the O-H or N-H bond needed to donate a hydrogen bond in return. This one-sided capability creates a surprisingly rich set of behaviors that explain why acetone mixes so easily with water, why it evaporates faster than you might expect, and why its role in chemistry is sometimes misunderstood.
Why Pure Acetone Does Not Hydrogen Bond with Itself
A hydrogen bond needs two participants: a donor that supplies a hydrogen atom bonded to an electronegative atom like oxygen or nitrogen, and an acceptor that provides a lone pair of electrons to attract that hydrogen. Water checks both boxes because each molecule has two O-H bonds (donors) and lone pairs on the oxygen (acceptor). Acetone has the acceptor side covered through its carbonyl oxygen, but it has no O-H or N-H group to serve as a donor. Its hydrogens are bonded only to carbon atoms, and carbon is not electronegative enough to create the kind of strong, polarized bond that conventional hydrogen bonding requires.
This means that in a container of pure acetone, the molecules cannot form the classic hydrogen-bonded networks that define liquids like water or ethanol. Spectroscopic and computational studies confirm this: acetone does not self-associate through hydrogen bonding the way many polar molecules do. A study using NMR and near-infrared spectroscopy found that, unlike dimethyl sulfoxide (a structurally similar molecule that does form self-associated dimers), acetone shows no evidence of self-association in solution.1SpringerLink. Association/Hydrogen Bonding of Acetone in Polar and Non-polar Solvents: NMR and NIR Spectroscopic Investigations with Chemometrics
What Holds Liquid Acetone Together Instead
If acetone molecules cannot hydrogen bond with each other, what keeps acetone a liquid at room temperature rather than immediately flying off as a gas? The answer is a combination of dipole-dipole interactions and weaker dispersive forces. Acetone has a substantial permanent dipole moment because the C=O bond is highly polar, with significant negative charge concentrated on the oxygen. Neighboring acetone molecules align so that the partially positive carbon end of one molecule sits near the partially negative oxygen end of another. These electrostatic attractions are weaker than hydrogen bonds, but they are strong enough to keep acetone liquid up to about 56 °C.
A comprehensive computational study modeling acetone clusters confirmed this picture in detail. The analysis revealed that liquid acetone is held together by a cooperative network dominated by dipole-dipole interactions between oxygen and carbon atoms on neighboring molecules, supplemented by numerous weak dispersive contacts between hydrogen atoms. The researchers described liquid acetone as existing in a dynamic equilibrium of small, weakly bound clusters rather than the extended hydrogen-bonded networks seen in liquids like water or methanol.2PubMed Central. Comprehensive Modeling of Acetone Clusters: QTAIM Analysis and QCE Study
This distinction matters practically. It is one reason acetone’s boiling point (56 °C) is much lower than water’s (100 °C) and lower than ethanol’s (78 °C), despite acetone being a heavier molecule than either. Hydrogen-bonded liquids require extra energy to pull their molecules apart; acetone’s dipole-dipole forces, while meaningful, simply do not hold on as tightly.
Acetone as a Hydrogen Bond Acceptor
Where acetone really shows its hydrogen-bonding character is in mixtures with molecules that can donate hydrogen bonds. Water is the most familiar example. When you pour acetone into water, the water molecules use their O-H bonds to donate hydrogen bonds directly to acetone’s carbonyl oxygen. This is why acetone and water are fully miscible: the hydrogen bonds that form between the two species are energetically favorable enough to overcome any tendency for the liquids to separate.
The same principle applies to alcohols, carboxylic acids, and any other molecule with an O-H or N-H group. Acetone’s lone pairs on the carbonyl oxygen are excellent hydrogen bond acceptors. In fact, studies on acetone’s proton affinity show that the carbonyl oxygen is quite “hungry” for protons and can form one hydrogen bond with other polar molecules through that site.3Chemical Physics Letters. Structure, bonding and reactions within protonated acetone–methanol cluster ions The word “one” is important here: because the carbonyl group provides only one primary hydrogen-bonding site, acetone cannot build the branching, three-dimensional networks that water forms. It can anchor a single hydrogen bond per molecule to a donor, but that is where its capacity maxes out.
You can see indirect evidence of this interaction in infrared spectroscopy. When acetone is dissolved in a hydrogen-bond-donating solvent, the C=O stretching frequency shifts to a lower value compared to its frequency in a non-polar solvent. Different solvent environments push the carbonyl frequency around in predictable ways, with strongly interacting solvents pulling it furthest from its “free” position.4Optica Publishing Group. Chemical and Physical Effects Upon Carbonyl Stretching Frequencies: Ketones This frequency shift is a spectroscopic fingerprint of the hydrogen bond tugging on the carbonyl oxygen and weakening the C=O bond slightly.
The Acetone-Water Bond Is Stronger Than You Might Think
One of the more counterintuitive findings about acetone’s hydrogen bonding is that the hydrogen bond between water and acetone is actually stronger than the hydrogen bond between two water molecules. Gas-phase measurements of the excess enthalpy of mixing water and acetone vapor showed that the interaction is exothermic, meaning energy is released when the two form a hydrogen bond. Researchers quantified this by determining the enthalpy of formation for both types of bonds: the water-water hydrogen bond was measured at roughly -16 kJ per mole, while the water-acetone hydrogen bond came in at about -20 kJ per mole.5The Journal of Chemical Thermodynamics. Water–acetone association. Second virial cross coefficients for water–acetone derived from gas phase excess enthalpy measurements
Why would acetone be a better hydrogen bond partner for water than water itself? The carbonyl oxygen on acetone is an unusually good acceptor. It carries a concentrated region of negative charge, and the two methyl groups flanking it are electron-donating, which makes the oxygen even more electron-rich. Water’s oxygen, by contrast, shares its electron density across two O-H bonds and two lone pairs while also serving as a donor. Acetone’s oxygen can devote more of its electron density to accepting a single hydrogen bond, which makes that bond stronger.
This also helps explain why acetone dissolves so readily in water. The energy gained by forming these favorable cross-species hydrogen bonds more than compensates for breaking some of the water-water hydrogen bonds that existed before mixing. In practical terms, this is why acetone is used as a solvent in so many laboratory and industrial settings where water compatibility matters.
The Gray Area of Weak C-H···O Contacts
The conventional answer that acetone “cannot donate” hydrogen bonds deserves an asterisk. While acetone’s C-H bonds are not polarized enough to form classical hydrogen bonds, there is growing recognition that very weak interactions between C-H groups and electronegative acceptors like oxygen do exist. These are sometimes called non-conventional or improper hydrogen bonds, and they occupy a gray zone between van der Waals forces and true hydrogen bonds.
In pure liquid acetone, the computational cluster study mentioned earlier identified numerous weak C-H···O contacts between neighboring molecules.2PubMed Central. Comprehensive Modeling of Acetone Clusters: QTAIM Analysis and QCE Study These are real interactions with measurable, if small, energetic contributions. They help stabilize the liquid structure, but their individual strength is far below that of a conventional O-H···O hydrogen bond, so they are classified under dispersive contacts rather than hydrogen bonding in most analyses.
When acetone interacts with a strong hydrogen bond donor like trifluoroacetic acid, the picture gets more interesting. Density functional theory calculations on complexes of trifluoroacetic acid with acetone revealed a cyclic structure: one strong O-H···O=C hydrogen bond (the acid donating to acetone’s carbonyl) paired with a weak C-H···O bond running in the opposite direction. The researchers found that this secondary weak bond contributes meaningfully to the total energy holding the complex together.6PubMed. IR Spectra of Hydrogen-Bonded Complexes of Trifluoroacetic Acid with Acetone and Diethyl Ether in the Gas Phase. Interaction between CH and OH Stretching Vibrations So while acetone’s C-H groups are not hydrogen bond donors in the textbook sense, they can participate in weak, stabilizing interactions that blur the boundary.
Whether these weak C-H···O contacts “count” as hydrogen bonds depends largely on how strict your definition is. The International Union of Pure and Applied Chemistry broadened its definition of hydrogen bonding in 2011 to potentially include C-H donors in certain contexts, but many chemists still treat these as a separate, lesser category. For most practical purposes in introductory chemistry, the answer remains that acetone does not donate hydrogen bonds. In research-level chemistry, the answer has more nuance.
How Acetone’s Hydrogen Bonding Shapes Its Behavior as a Solvent
Acetone’s lopsided hydrogen-bonding ability, strong acceptor but negligible donor, is exactly what makes it such a versatile solvent. It dissolves both polar and many nonpolar substances, a property that sets it apart from solvents that are more committed to one side of the polarity spectrum.
With polar and hydrogen-bonding solutes, acetone’s carbonyl oxygen accepts hydrogen bonds, stabilizing the dissolved molecules. With nonpolar solutes, acetone’s two methyl groups provide a hydrocarbon-like environment that accommodates nonpolar species. This dual nature is why acetone strips nail polish (dissolving nonpolar polymers), cleans laboratory glassware (breaking up a variety of residues), and mixes readily with water, all in one molecule.
The low boiling point that comes from the absence of self-hydrogen-bonding is an advantage in many applications. Acetone evaporates quickly, which is useful when you want a solvent that does its job and then disappears. Compare this to water, which clings stubbornly to surfaces because its extensive hydrogen-bonded network resists evaporation, or to dimethyl sulfoxide, which has a similar molecular structure to acetone but self-associates through hydrogen bonding and consequently has a much higher boiling point of 189 °C.
The flip side is that acetone’s volatility and flammability require careful handling. Its vapors are heavier than air and can accumulate near the ground, creating a fire risk in poorly ventilated spaces. The same weak intermolecular forces that make it evaporate quickly mean it takes very little energy to ignite.
Acetone and Surface Hydrogen Bonding
Acetone’s acceptor capability also shows up in how it interacts with solid surfaces, particularly those bearing hydroxyl groups. Silica gel, one of the most common adsorbents in chemistry, has a surface covered in Si-OH groups. When acetone vapor contacts silica gel, the surface hydroxyl groups donate hydrogen bonds to acetone’s carbonyl oxygen. This interaction has been directly observed through infrared spectroscopy: the O-H stretching frequency of the silica surface shifts downward when acetone is present, a clear sign of hydrogen bond formation. At the same time, acetone’s own C=O stretching frequency also shifts slightly lower, confirming that both partners feel the bond.7J-STAGE. FTIR Study of Adsorption on Silica Gel for Organic Solvents Diluted in Supercritical Carbon Dioxide
These surface interactions are relevant in practical settings like chromatography, where silica gel is used to separate mixtures. Acetone’s ability to hydrogen bond with the silica surface affects how strongly it and other molecules stick to the column, influencing separation efficiency. It also matters in industrial gas sensing and adsorption processes, where the strength of the acetone-surface interaction determines how effectively acetone vapor can be captured or detected.
Interestingly, when the same silica-acetone system was studied under increasing pressure with supercritical carbon dioxide as a carrier, the strength of the hydrogen bonding between acetone and the silica surface changed. At lower pressures, the amount of hydrogen-bonded acetone on the surface increased, reaching a maximum below the critical pressure of CO₂, and then gradually declined as pressure continued to rise.7J-STAGE. FTIR Study of Adsorption on Silica Gel for Organic Solvents Diluted in Supercritical Carbon Dioxide The dense supercritical fluid essentially competed with acetone for space near the surface, peeling it away. This kind of pressure-tunable adsorption is one reason supercritical CO₂ is used in extraction and cleaning processes: it can selectively remove molecules that are hydrogen-bonded to surfaces.
Why Textbooks Sometimes Get This Wrong
If you search for whether acetone has hydrogen bonding, you will find conflicting answers, and the confusion is understandable. Some sources say “no” because acetone cannot donate hydrogen bonds, and in introductory chemistry the phrase “hydrogen bonding” often implicitly means self-hydrogen-bonding in the pure liquid. Under that framing, acetone does not hydrogen bond, period. Other sources say “yes” because acetone’s carbonyl oxygen readily accepts hydrogen bonds from donors like water. Both answers are correct within their own framing, but neither tells the full story alone.
The most common misconception is that acetone is a nonpolar molecule. Its symmetrical shape can fool people: the two methyl groups flanking the carbonyl might suggest a balanced, nonpolar structure. But the C=O bond is strongly polar, and the overall molecule has a dipole moment of about 2.9 debye, which is quite substantial. Acetone is definitively a polar molecule. It just happens to be a polar molecule that cannot donate hydrogen bonds in the classical sense.
Another source of confusion involves boiling points. Students sometimes learn that molecules with hydrogen bonding have higher boiling points, observe that acetone boils at a relatively low temperature, and conclude it must have no hydrogen-bonding capability whatsoever. The flaw in this reasoning is that boiling point reflects self-interaction in the pure liquid, and acetone’s inability to hydrogen bond with itself is what keeps its boiling point low. This says nothing about its ability to accept hydrogen bonds from other molecules. A ketone that cannot self-associate can still be an excellent hydrogen bond acceptor in a mixture.
For anyone studying chemistry or working with solvents professionally, the practical takeaway is that acetone’s behavior always depends on what it is mixed with. In a pure bottle, it behaves like a polar but non-hydrogen-bonded liquid, volatile and easy to evaporate. In water or an alcohol, it integrates into the hydrogen-bonded network as an acceptor, becoming far more interactive. And on a hydroxylated surface, it latches on through hydrogen bonds that are strong enough to influence adsorption and separation processes. The molecule itself does not change. What changes is whether a hydrogen bond donor happens to be nearby.