Question

Consider two different organic compounds, each with the formula \(\mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O}\) . One of these compounds is a liquid at room conditions and the other is a gas. Write Lewis structures consistent with this observation, and explain your answer. (Hint: The oxygen atom in both structures satisfies the octet rule with two bonds and two lone pairs.)

Short answer

The C2H6O molecular formula can form two different structural formulas: one with a hydroxyl group (O-H bond) and the other with an ether group (O in between two carbons). The hydroxyl group compound forms hydrogen bonds, leading to a higher boiling point and a liquid state at room temperature. The ether group compound has weaker van der Waals forces and a lower boiling point, making it a gas at room temperature. Hydroxyl group structure (liquid): H H O H | | | | C-C-O-H C-C | | | | H H H H Ether group structure (gas): H H H H | | | | C-C-O-C-C | | | | H H H H

Step-by-step solution

1. Identify the possible structural formulas

: To begin, we need to identify the possible structural formulas for the two compounds. We know that they must have the molecular formula C2H6O and the oxygen atom should have two bonds and two lone pairs, according to the octet rule: 1. Hydroxyl group: The oxygen atom forms a single bond to one carbon and a single bond to a hydrogen atom. 2. Ether group: The oxygen atom forms a single bond to each of the carbon atoms. Now we have two possible structures to work with: one with a hydroxyl group and the other with an ether group.

2. Draw Lewis structures for each compound

: Using the structural formulas, we can now draw Lewis structures for each compound: 1. Hydroxyl group compound: In this structure, the oxygen atom forms a single bond with one carbon and a single bond with a hydrogen atom, with two lone pairs of electrons on the oxygen atom. H H O H | | | | C-C-O-H C-C | | | | H H H H 2. Ether group compound: In this structure, the oxygen atom forms a single bond with each of the carbon atoms, with two lone pairs of electrons on the oxygen atom. H H H H | | | | C-C-O-C-C | | | | H H H H

3. Determine which compound is a liquid and which is a gas at room temperature

: Now that we have both the Lewis structures, let's analyze them to determine which one would be a liquid and which one would be a gas at room temperature. 1. Hydroxyl group compound: This compound has an O-H bond, which leads to the formation of hydrogen bonding between the oxygen atom of one molecule and the hydrogen atom of another molecule. Hydrogen bonding is a strong intermolecular force, which allows this compound to maintain a relatively higher boiling point and be in liquid form at room temperature. 2. Ether group compound: In this compound, there is no hydrogen bonding as both the hydrogen atoms are connected to carbon atoms. The intermolecular forces present in this compound are van der Waals forces (London dispersion forces). These forces are weaker when compared to hydrogen bonding, so this compound would have a lower boiling point, making it a gas at room temperature. In conclusion, the compound with the hydroxyl group is a liquid at room temperature and the compound with the ether group is a gas at room temperature.

Key concepts

Lewis Structure

In organic chemistry, a Lewis structure is a helpful tool for representing the arrangement of atoms in a molecule. It uses dots to symbolize lone pairs of electrons and lines to depict chemical bonds. Lewis structures are particularly useful for visualizing the valence electrons, and they help in predicting the molecule's geometry and reactivity.
To fulfill the octet rule, each atom strives for eight electrons in its valence shell. For the molecular formula \(\mathrm{C_{2}H_{6}O}\), we need to draw Lewis structures for both possible compounds—one featuring a hydroxyl group and another with an ether group. Oxygen, a key player here, aims for an octet through two bonds and two lone pairs.
Keep in mind:

• Bond pairs are represented as lines (e.g., \(\mathrm{O-H}\) or \(\mathrm{C-O}\)).
• Lone pairs around atoms like oxygen are depicted by dots beside the element symbol.

Drawing these accurately ensures that we correctly predict how these two organics differ in properties such as boiling points.

Hydroxyl Group

A hydroxyl group is a distinctive functional group in organic chemistry, consisting of an oxygen and hydrogen atom connected by a covalent bond (\(\mathrm{O-H}\)). It is a key component of alcohols—a common type of organic compound. As seen in one variant of \(\mathrm{C_{2}H_{6}O}\), a hydroxyl group can attach to a carbon chain, creating an alcohol.
The presence of the hydroxyl group is significant because it plays a role in hydrogen bonding, a strong type of intermolecular force. Because the oxygen atom is more electronegative than hydrogen, it attracts the shared electrons more strongly. This creates a partial negative charge on the oxygen and a partial positive charge on the hydrogen. These partial charges enable attraction between molecules, leading to hydrogen bonds.
This interaction significantly impacts physical properties:

• Substantial increase in boiling point
• Solubility in water due to hydrogen bonds with \(\mathrm{H_{2}O}\)

Hence, molecules with hydroxyl groups, such as the discussed alcohol, often remain liquid at room temperature.

Ether Group

An ether group is another functional group type in which an oxygen atom is bonded to two alkyl or aryl groups (\(\mathrm{R-O-R'}\)). In one structural form of \(\mathrm{C_{2}H_{6}O}\), oxygen is sandwiched between the carbon atoms, forming an ether.
Ethers are quite different in their chemical behavior compared to alcohols with hydroxyl groups. Specifically, ethers lack hydrogen atoms bonded to the oxygen that could participate in hydrogen bonding with other molecules.
For this reason:

• Ethers have lower boiling points compared to alcohols.
• They exhibit weaker intermolecular forces known as van der Waals forces or London dispersion forces.

Consequently, ethers are often gases at room temperature. This distinction helps in understanding why two molecules with the same formula, \(\mathrm{C_{2}H_{6}O}\), can differ so much in their physical state based on their structure.

Intermolecular Forces

Intermolecular forces are the forces of attraction or repulsion between neighboring molecules, influencing the physical properties of the material. There are various types of intermolecular forces, including hydrogen bonds, dipole-dipole interactions, and London dispersion forces.
With regard to \(\mathrm{C_{2}H_{6}O}\) compounds:

• The hydroxyl-containing compound relies on hydrogen bonds. These are strong interactions due to the large difference in electronegativity between oxygen and hydrogen, allowing for a network of cohesive forces that keep the molecules together, thus retaining a liquid state at room temperature.
• The ether-containing compound exhibits primarily London dispersion forces. These are much weaker and stem from temporary fluctuating charges in electron clouds. Hence, compounds primarily held together by these forces tend to be gases under similar conditions.

Understanding these forces provides insight into why some compounds boil or freeze at different temperatures, which is vital for predicting and explaining their behavior in practical applications.