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 molecules? (d) 1-Propanol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH}\right)\) 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 Lewis structure of acetone consists of a trigonal planar arrangement around the carbonyl carbon (C=O) and tetrahedral geometries around the methyl carbon atoms. The acetone molecule is polar due to its asymmetric charge distribution and has a net dipole moment pointing towards the oxygen atom. Intermolecular attractive forces in acetone include dipole-dipole interactions and London dispersion forces. The difference in boiling points between acetone and 1-propanol can be attributed to 1-propanol's ability to form stronger hydrogen bonds, leading to higher boiling points.

Step-by-step solution

(a) Lewis Structure and Geometry

To draw the Lewis structure for acetone, first identify the central atom, which is the carbonyl carbon (C=O). - Surround the carbonyl carbon with three bonding pairs (one to the oxygen atom and two to the two methyl groups). - Place a double bond between the carbonyl carbon and oxygen atom, fulfilling the octet rule for both atoms. - Add the two methyl groups, attaching each to the available bonding sites on the carbonyl carbon with single bonds. - Complete the octets for the outer carbon atoms by adding six additional valence electrons to each (totaling three lone pairs on each). The geometry around the central atom (carbonyl carbon) is trigonal planar, while the geometry around the methyl carbon atoms is tetrahedral.

(b) Polarity of Acetone Molecule

Since the oxygen atom is more electronegative than the carbon atoms, it draws the electron density towards itself, resulting in a partially negative charge on oxygen and partial positive charges on the carbonyl carbon and methyl carbon atoms. However, due to the symmetry of the molecule and the orientation of the bond dipoles, the overall molecule is polar, as there is a net dipole moment pointing towards the oxygen atom.

(c) Intermolecular Attractive Forces

The intermolecular attractive forces that exist between acetone molecules are dipole-dipole interactions due to its polar nature, as well as London dispersion forces, which arise from the random fluctuations in electron distribution in all molecules.

(d) Boiling Points of Acetone and 1-Propanol

The difference in boiling points between acetone and 1-propanol, despite their similar molecular weights, can be explained by the intermolecular forces present in each molecule. 1-propanol has an -OH functional group, which allows for the formation of strong hydrogen bonds between its molecules. These hydrogen bonds result in stronger intermolecular forces, requiring more energy to break them and thus a higher boiling point. Acetone, on the other hand, only experiences dipole-dipole and London dispersion forces, which are weaker and result in a lower boiling point than 1-propanol.

Key concepts

Molecular Geometry

Understanding the molecular geometry of acetone involves looking at how the atoms are arranged around each central atom. In acetone, the molecule can be broken down primarily around two critical areas:

• The carbonyl carbon (C=O): This central carbon is connected to an oxygen atom by a double bond and to two methyl groups. The molecular geometry here is trigonal planar because the atom forms a shape with 120-degree angles between bonds.
• The methyl groups (CH₃): Around each methyl carbon, you have tetrahedral geometry. With single bonds to the carbonyl carbon and three hydrogen atoms, these carbon centers create a typical tetrahedral shape with bond angles of approximately 109.5 degrees.

Combining these geometries gives acetone its distinct shape, influencing its chemical behavior and interactions.

Polarity

When talking about polarity in acetone, it's all about how electrons are distributed. Oxygen in acetone is highly electronegative, drawing electron density towards itself. As a result, the oxygen atom becomes partially negative while other atoms like the carbonyl carbon and the methyl carbon attain partial positive charges.
Despite acetone having symmetrical sections, there’s an overall polarity because there's a net dipole moment. This is a vector sum of all individual dipoles, pointing towards the more electronegative oxygen. This imbalance gives acetone a polar character, making it have distinct interactions with other polar environments.

Intermolecular Forces

Intermolecular forces in acetone are primarily defined by its polar nature. Within a group of acetone molecules, you can expect:

• Dipole-dipole interactions: These occur because of the molecule's polarity. The partially negative oxygen of one molecule will attract the partially positive carbons of another, leading to this type of intermolecular attraction.
• London dispersion forces: Also known as van der Waals forces, these are present in all molecules, including non-polar ones. They arise due to momentary changes in electron density even if fleeting.

These forces dictate how acetone molecules stick together, though they are not the strongest when compared to some other molecular interactions like hydrogen bonds.

Boiling Points

Boiling points are greatly influenced by intermolecular forces. Acetone and 1-propanol provide an interesting comparison that highlights this concept. While both have similar molecular weights, their boiling points differ due to the types of intermolecular forces:

• Acetone has boiling interactions defined by dipole-dipole and dispersion forces. These forces are relatively weak, allowing acetone to have a lower boiling point (56.5°C).
• 1-Propanol, on the other hand, can form hydrogen bonds due to its -OH group. Hydrogen bonds are much stronger than dipole-dipole and dispersion forces, requiring more energy to break. Thus, 1-propanol has a higher boiling point (97.2°C).

These differences underscore the importance of hydrogen bonding in affecting the boiling points of substances even when they have similar molecular weights.

Hydrogen Bonding

Hydrogen bonding is a special type of attraction that occurs when a hydrogen atom, attached to an electronegative atom like oxygen, forms a weak bond with another electronegative atom. Though acetone does not partake in hydrogen bonding because it lacks an -OH or -NH group, understanding it is crucial when comparing boiling points with molecules like 1-propanol.
1-Propanol benefits from a hydroxyl (-OH) group that can form hydrogen bonds easily. This type of bonding is significantly stronger than dipole-dipole interactions found in acetone. It results in molecules being "stickier" to each other, hence requiring more energy (temperature) to be pulled apart, thus raising the boiling point. Hydrogen bonds are a key player in determining physical properties in similar-sized molecules.