Question

Acetone \(\left[\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO}\right]\) is widely used as an industrial solvent. (a) Draw the Lewis structure for the acetone molecule and predict the geometry around each carbon atom. (b) Is the acetone molecule polar or nonpolar? (c) What kinds of intermolecular attractive forces exist between acetone mol-ecules? (\boldsymbol{d} 1 Propanol ~ ( C H ~ \(_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH}\) ) has a molecular weight that is very similar to that of acetone, yet acetone boils at \(56.5^{\circ} \mathrm{C}\) and 1 -propanol boils at \(97.2^{\circ} \mathrm{C}\) . Explain the difference.

Short answer

The acetone molecule, \((\mathrm{CH}_{3})_{2} \mathrm{CO}\), has a trigonal planar geometry around each carbon atom. It is a nonpolar molecule due to the cancelation of dipole moments, and its main intermolecular attractive forces are London dispersion forces. In contrast, 1-propanol has a similar molecular weight but a higher boiling point, as it forms stronger hydrogen bonds due to the -OH group.

Step-by-step solution

(Drawing the Lewis structure for acetone)

First and foremost, we will draw the Lewis structure for acetone (\((\mathrm{CH}_{3})_{2} \mathrm{CO}\)). The carbon atom in the middle is the central atom, double bonded to an oxygen and single bonded to two methyl groups. The hydrogen atoms are bonded to each of the carbon atoms in the methyl group. Here's the Lewis structure: H H H | | | H-C-C-C=O | | | H H H Now let's predict the geometry around each carbon atom in acetone.

(Predicting the geometry around each carbon atom in acetone)

Since the carbon atoms in the two methyl groups have three neighboring atoms and no lone pairs, they exhibit a trigonal planar geometry (120° bond angles). The central carbon atom is double bonded to an oxygen atom and single bonded to two methyl groups, so it also has three neighboring atoms with no lone pairs, and therefore, also has trigonal planar geometry (120° bond angles).

(Determining if the acetone molecule is polar or nonpolar)

To determine if the acetone molecule is polar or nonpolar, we need to consider the electronegativities of the constituent atoms and the geometry of the molecule. Oxygen is more electronegative than carbon, so the carbon-oxygen double bond is polar. However, due to the symmetry of the acetone molecule resulting from its trigonal planar geometry, the dipole moments around the central carbon atom cancel out each other. Therefore, the overall acetone molecule is nonpolar.

(Identifying the types of intermolecular attractive forces in acetone)

Between acetone (\((\mathrm{CH}_{3})_{2} \mathrm{CO}\)) molecules, the main intermolecular attractive forces are London dispersion forces. These are relatively weak forces and arise due to temporary fluctuations in electron distribution which induce temporary dipoles.

(Explaining the difference in boiling points between acetone and 1-propanol)

Acetone and 1-propanol have similar molecular weights, but their boiling points are significantly different (acetone boils at \(56.5^{\circ} \mathrm{C}\), while 1-propanol boils at \(97.2^{\circ} \mathrm{C}\)). The main reason for this difference is the type of intermolecular attractive forces present in each molecule. 1-propanol (\(CH_{3} CH_{2} CH_{2} OH\)) has an -OH group, which can form hydrogen bonds with other molecules. Hydrogen bonding is a much stronger intermolecular force than London dispersion forces, which are present in acetone. As a result, 1-propanol has a higher boiling point since more energy is required to break these stronger hydrogen bonds. In summary, the difference in boiling points between acetone and 1-propanol is due to the difference in the types and strength of intermolecular attractive forces present in each molecule—with hydrogen bonding in 1-propanol being stronger than London dispersion forces in acetone.

Key concepts

Lewis structure

One of the first steps in understanding the structure of acetone is drawing its Lewis structure, which visually represents the arrangement of atoms and electrons within the molecule. Acetone, represented as \((\text{CH}_3)_2\text{CO}\), consists of two methyl groups \((\text{CH}_3)\) and a central carbon-oxygen double bond \((\text{C=O})\). The central carbon atom sits at the heart of the molecule, creating a backbone with oxygen and the two methyl groups extending from it. Each carbon in the methyl group is bonded to three hydrogen atoms. This gives acetone a straightforward skeleton, which can be easily visualized as:

• Central carbon \((\text{C})\) double bonded to oxygen \((\text{O})\)
• The central carbon single bonded to two methyl groups \((\text{CH}_3)\)
• Each carbon in the methyl group bonded to three hydrogen atoms

This structure provides a foundation for predicting the geometry and analyzing other properties of acetone.

molecular geometry

In acetone, the molecular geometry is critical to understanding how it behaves chemically and physically. The geometry around each carbon atom is determined by counting the number of atoms bonded to it and any lone pairs of electrons. For acetone:

• The carbon atoms in the methyl groups are each bonded to three atoms (three hydrogens) and have no lone pairs, resulting in a trigonal planar geometry. This means each bond angle is about 120°.
• The central carbon, which links two methyl groups and an oxygen atom, adopts the same trigonal planar arrangement. It is bonded to three other atoms (two from the methyl groups and one from the oxygen) and has no lone pairs.

This consistent geometry throughout the molecule contributes to acetone's symmetrical shape, influencing its physical behavior and interactions.

intermolecular forces

The behavior of acetone at the molecular level is influenced significantly by the types of intermolecular forces present. In acetone, the primary force at play is the London dispersion force. These are temporary forces that occur due to the momentary distribution of electrons around atoms, leading to a temporary dipole.

In acetone:

• London dispersion forces arise because of these temporary dipoles, which attract neighboring molecules.
• These forces are weak compared to other types of intermolecular forces, such as hydrogen bonds, but they are crucial in determining the liquid properties of acetone.

Despite these forces being weaker, they are enough to keep acetone in liquid form until its boiling point is reached.

polarity

The polarity of a molecule like acetone helps us predict how it will interact with other substances. Polarity arises when there is a difference in electronegativity between bonded atoms, causing a dipole moment.

For acetone:

• The carbon-oxygen double bond is polar because oxygen is more electronegative than carbon, creating a dipole.
• However, the overall acetone molecule is considered nonpolar. This is due to its symmetrical trigonal planar shape, which cancels out the individual dipole moments.

Even though it has polar bonds internally, the molecular symmetry means that the exterior shows a nonpolar characteristic. This influences how acetone dissolves in solvents and engages in chemical reactions.

boiling point comparison

Understanding why acetone has a lower boiling point compared to a similar molecular weight compound such as 1-propanol involves examining intermolecular forces.

Acetone's boiling point is lower at 56.5°C compared to 97.2°C for 1-propanol due to:

• The nature of intermolecular forces. Acetone has London dispersion forces, which are relatively weak.
• Contrastingly, 1-propanol has an -OH group, which enables hydrogen bonding — a much stronger force than London dispersion forces.

The stronger the intermolecular forces, the more energy (thus a higher temperature) is required to break these interactions. Hence, 1-propanol's higher boiling point is attributed to the stronger hydrogen bonds that hold its molecules together, compared to the weaker forces in acetone.