Question

Acetone, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO},\) 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, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH},\) 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 Lewis structure of acetone shows trigonal planar geometry around each carbon atom. The molecule is polar due to C=O bond polarity and exhibits dipole-dipole interactions and London dispersion forces as intermolecular forces. The higher boiling point of 1-propanol compared to acetone can be attributed to the presence of hydrogen bonding in 1-propanol, a stronger intermolecular force.

Step-by-step solution

Drawing the Lewis Structure for Acetone

To draw the Lewis structure for acetone, first, count the total number of valence electrons. Carbon (C) has 4 valence electrons, Oxygen (O) has 6 valence electrons, and Hydrogen (H) has 1 valence electron. In acetone, there are 2 CH3 groups (or 2 sets of Carbon and 3 Hydrogens), 1 Carbon, and 1 Oxygen. Total valence electrons = (2 x (1C + 3H)) + 1C + 1O = 2 x (4 + 3) + 4 + 6 = 24 Now, draw the skeletal structure of the molecule, with carbon in the center, one oxygen bonded to the central carbon, and two CH3 groups bonded to the central carbon. Next, place the remaining electrons around the atoms to complete their octets, keeping in mind the molecule has 24 valence electrons in total.

Predicting the Geometry around Carbon Atoms

To determine the geometry around each carbon atom, we will use the VSEPR (Valence Shell Electron Pair Repulsion) theory. The central carbon atom is bonded to two other carbon atoms and one oxygen atom, making a total of 3 bonding domains around the central carbon. There are no lone pairs in this case. According to the VSEPR theory, with three bonding domains and no lone pairs, the geometry around the carbon atom is trigonal planar. For the other two carbon atoms (in CH3 groups), each carbon atom is bonded to three hydrogen atoms, with a total of 3 bonding domains, and no lone pairs. This also results in a trigonal planar geometry around these carbon atoms.

Determining Polarity of Acetone Molecule

To check the polarity of the molecule, we need to identify the polar bonds and the molecular geometry. There is a difference in electronegativity between Carbon and Oxygen (C-O bond) and Carbon and Hydrogen (C-H bond). The C=O bond is considered polar, while the C-H bonds are less polar. Due to the trigonal planar geometry around the oxygen atom, the C=O bond vectors do not cancel each other, leading to a net molecular dipole moment. Therefore, the acetone molecule is polar.

Identifying Intermolecular Attractive Forces

The types of intermolecular attractive forces that exist between acetone molecules are: 1. Dipole-dipole interactions: These occur between the polar molecules of acetone due to a difference in electronegativity between the carbon and oxygen atoms. 2. London dispersion forces: These are induced-dipole/induced-dipole interactions that occur in all molecules, including polar molecules like acetone.

Comparing Boiling Points of Acetone and 1-Propanol

We are given that acetone boils at \(56.5^{\circ} \mathrm{C}\), and 1-propanol boils at \(97.2^{\circ} \mathrm{C}\). The higher boiling point of 1-propanol can be attributed to the presence of hydrogen bonding between the molecules, due to the presence of an OH group in its structure. In acetone, dipole-dipole interactions and London dispersion forces exist, but hydrogen bonding is absent. Hydrogen bonding is a stronger intermolecular force compared to dipole-dipole interactions and London dispersion forces, which explains the higher boiling point of 1-propanol compared to acetone despite similar molecular weights.

Key concepts

Lewis Structure

A Lewis structure is a visual way to represent the valence electrons in a molecule. It shows how atoms are bonded together and can indicate the presence of lone pairs. When drawing the Lewis structure for acetone, \((\text{CH}_3)_2\text{CO}\), follow these steps:

• Count the total number of valence electrons. Acetone has a total of 24 valence electrons: 16 from the carbon atoms (2 carbons with 4 electrons each) plus 6 from the oxygen atom, and 2 from the hydrogen atoms.
• Arrange the atoms in a skeletal structure, with two methyl groups (\(\text{CH}_3\)) attached to a central carbon atom, which is also bonded to an oxygen atom.
• Distribute the electrons to fulfill the octet rule, ensuring each atom (except hydrogen) has 8 electrons to achieve a stable configuration.

The goal is a structure where the central carbon atom is double-bonded to the oxygen and single-bonded to two other carbon atoms, each of which is bonded to three hydrogens.

Intermolecular Forces

Intermolecular forces are the attractions between molecules, affecting physical properties like boiling point and solubility. Acetone exhibits two primary types of intermolecular forces:

• Dipole-Dipole Interactions: Because acetone molecules are polar, they have positive and negative sides. The C=O (carbonyl) group creates a significant dipole moment due to the difference in electronegativity between carbon and oxygen. This means that acetone molecules tend to align so that opposite charges are near each other, creating dipole-dipole attractions.
• London Dispersion Forces: These are present in all molecules, caused by temporary shifts in electron density that create instantaneous dipoles. Even though they are the weakest intermolecular forces, they still contribute to acetone's overall physical properties.

Intermolecular forces play a crucial role in determining how often molecules interact, influencing properties like volatility and boiling point.

Polarity

The concept of polarity involves the distribution of electrical charge over the atoms in a molecule. A molecule is polar if it has regions of positive and negative charge leading to a net dipole moment. In acetone, the presence of the carbon-oxygen double bond (C=O) creates polarity. Oxygen is more electronegative than carbon, pulling electron density towards itself and away from the carbon.
Even though C-H bonds are considered nonpolar due to a small difference in electronegativity, the overall shape of the molecule matters. Acetone has a trigonal planar geometry around the central carbon, meaning the polar C=O bond creates an imbalance in charge distribution, rendering the whole molecule polar.
This polarity is important as it affects solubility and the types of reactions acetone can participate in.

Boiling Point

The boiling point of a substance is the temperature at which it changes from liquid to gas. It is influenced by the strength of intermolecular forces between its molecules. Acetone has a boiling point of 56.5°C, mainly due to dipole-dipole interactions and London dispersion forces. In contrast, 1-propanol, even with a similar molecular weight, boils at 97.2°C.
The key factor here is hydrogen bonding in 1-propanol. It has an O-H bond capable of forming hydrogen bonds, which are much stronger than dipole-dipole interactions. Hydrogen bonds require more energy to break, thus raising the boiling point of 1-propanol.
Factors such as molecular weight, shape, and polarizability also contribute, but the presence of hydrogen bonding in 1-propanol is the primary reason it has a significantly higher boiling point than acetone.