Acetone, or propanone ($\text{CH}_3\text{COCH}_3$), is a simple organic compound commonly recognized as the fast-acting solvent in nail polish remover. The physical characteristics of any substance, such as its boiling point or solubility, are determined by the attractive forces that hold individual molecules together. These forces, known as intermolecular forces (IMFs), dictate the energy required to separate molecules and cause a change of state. Understanding acetone’s specific IMFs explains its distinct properties as a solvent.
Understanding Intermolecular Forces
Intermolecular forces are attractions between neighboring molecules, significantly weaker than the chemical bonds within a molecule. The weakest of these forces are the London Dispersion Forces (LDF), present in all substances regardless of polarity. LDF arise from the constant movement of electrons, which momentarily creates an uneven charge distribution, forming a temporary dipole. This brief charge imbalance induces a corresponding dipole in a nearby molecule, leading to a weak, transient attraction.
Dipole-Dipole interaction is a stronger attraction that occurs only between molecules possessing a permanent separation of charge, known as a net dipole moment. In these polar molecules, the slightly positive end of one molecule is consistently attracted to the slightly negative end of a neighbor. These forces are stronger than LDF because the charge separation is permanent, requiring more energy to overcome.
The strongest common IMF is a specialized form of attraction called Hydrogen Bonding. This highly specific interaction occurs only when a hydrogen atom is covalently bonded to nitrogen ($\text{N}$), oxygen ($\text{O}$), or fluorine ($\text{F}$). The high electronegativity of these atoms pulls electrons away from the hydrogen, leaving it with a large partial positive charge. This charged hydrogen then forms a strong bridge to a lone pair of electrons on an $\text{N}$, $\text{O}$, or $\text{F}$ atom of an adjacent molecule.
Acetone’s Molecular Structure and Polarity
Acetone is the simplest ketone, featuring a three-carbon chain with a central carbonyl group ($\text{C=O}$). The molecule consists of a carbon atom double-bonded to oxygen, with two methyl ($\text{CH}_3$) groups attached to the central carbon. This structure is responsible for the molecule’s overall polarity.
Oxygen is a highly electronegative atom, strongly attracting electrons in the chemical bond. In the carbonyl group, the oxygen atom pulls the double bond electrons toward itself, away from the carbon atom. This uneven sharing creates a significant, permanent bond dipole, leaving the oxygen atom partially negative and the carbon atom partially positive.
Because the two methyl groups are symmetrically positioned, they do not cancel out the strong polarity of the central carbonyl group. The resulting vector sum of all bond dipoles gives the entire acetone molecule a substantial net dipole moment, measured at approximately $2.8$ Debye. This permanent charge separation causes acetone to be classified as a polar molecule.
Identifying the Specific Forces in Acetone
The polarity of acetone dictates the two primary intermolecular forces present between its molecules: London Dispersion Forces and Dipole-Dipole interactions. LDF are universally present in all molecules, relying only on the movement of electrons. These forces contribute a baseline level of attraction that increases with molecular size.
The permanent net dipole moment of $2.8$ Debye ensures that neighboring acetone molecules align themselves so that the partially negative oxygen atom of one molecule attracts the partially positive carbon of another. This consistent, moderate Dipole-Dipole interaction is the dominant force influencing the liquid state of acetone.
Acetone molecules do not engage in hydrogen bonding with one another. Although acetone contains hydrogen atoms, they are bonded exclusively to carbon atoms, not to the highly electronegative oxygen. Therefore, the hydrogen atoms do not carry the high partial positive charge required to form a hydrogen bond.
How IMFs Influence Acetone’s Physical Properties
The combination of London Dispersion Forces and moderate Dipole-Dipole forces influences acetone’s physical properties. The Dipole-Dipole attractions hold the molecules together more strongly than non-polar molecules of similar size, such as pentane. However, the absence of hydrogen bonding results in weaker overall attractions compared to molecules that form these strong bridges, such as water or ethanol.
This intermediate strength of intermolecular attraction explains why acetone is a liquid at room temperature but is highly volatile. The moderate energy holding the molecules together is easily overcome by thermal energy, giving acetone a relatively low boiling point of approximately $56.3^\circ\text{C}$ at standard atmospheric pressure. This low boiling point causes it to evaporate very quickly, which is why it feels cool on the skin.
Acetone’s dual nature—having a polar section and non-polar methyl groups—makes it an excellent, versatile solvent that is completely miscible with water. The polar carbonyl group allows it to interact with and dissolve many polar compounds through dipole-dipole attractions and even form hydrogen bonds with water molecules. Simultaneously, the non-polar methyl groups on either side of the molecule enable it to engage in LDF with and dissolve a wide variety of non-polar substances.