Problem 92 - Chemistry The Molecular Nature of Matter

Question

What kinds of intermolecular attractive forces are present in the following substances? (a) CC(=O)O (b) \(\mathrm{H}_{2} \mathrm{~S}\) (c) \(\mathrm{SO}_{3}\) (d) \(\mathrm{CH}_{3} \mathrm{NH}_{2}\)

Step-by-Step Solution

Answer

(a) Acetic acid has hydrogen bonding, dipole-dipole interactions, and London dispersion forces. (b) H2S exhibits dipole-dipole interactions and London dispersion forces. (c) SO3 has London dispersion forces. (d) Methylamine has hydrogen bonding, dipole-dipole interactions, and London dispersion forces.

Step 1: Analyze the Molecular Structure of CC(=O)O

The molecule CC(=O)O corresponds to acetic acid (CH3COOH). This molecule has a polar covalent bond due to the difference in electronegativity between hydrogen and oxygen, and carbon and oxygen. It also has hydrogen bonds due to the presence of a -OH group.

Step 2: Identify the Intermolecular Forces for H2S

Hydrogen sulfide (H2S) is a polar molecule due to the electronegativity difference between hydrogen and sulfur. However, it cannot form hydrogen bonds because sulfur is not as electronegative as nitrogen, oxygen, or fluorine. Therefore, the main intermolecular forces are dipole-dipole attractions and London dispersion forces.

Step 3: Determine the Intermolecular Forces Present in SO3

Sulfur trioxide (SO3) is a nonpolar molecule because it has a trigonal planar geometry which allows for the dipole moments to cancel out. The primary intermolecular force present in SO3 is London dispersion forces.

Step 4: Evaluate the Intermolecular Forces in CH3NH2

Methylamine (CH3NH2) is polar and has an -NH2 group that can form hydrogen bonds. Thus, the main types of intermolecular forces in CH3NH2 are hydrogen bonding, dipole-dipole interaction, and London dispersion forces.

Key Concepts

Polar Covalent Bonds

Polar covalent bonds are a type of chemical bond where two atoms share a pair of electrons, but they share them unequally. This happens because one atom has a greater electronegativity, or the ability to attract electrons, than the other. As a result, the electrons spend more time closer to the more electronegative atom, causing a partial negative charge on that atom, and a partial positive charge on the other.

For example, in water (H2O), oxygen is more electronegative than hydrogen, leading to polar covalent bonds. These polar bonds are essential because they can lead to the formation of dipoles in molecules, which are crucial for establishing intermolecular forces.

Hydrogen Bonding

Hydrogen bonding is a strong type of dipole-dipole interaction that occurs when a hydrogen atom bonded to a highly electronegative atom, such as nitrogen, oxygen, or fluorine (NOF), is attracted to a lone pair of electrons on another electronegative atom.

These bonds are much stronger than regular dipole-dipole interactions due to the high electronegativity of NOF atoms and the small size of hydrogen, allowing for closer intermolecular interactions. Substances with hydrogen bonds, like water, often have higher boiling points and unique properties due to these strong intermolecular forces.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules that have permanent dipoles, meaning there is an uneven distribution of electron density within the molecule. When the positive end of one polar molecule aligns with the negative end of another, a dipole-dipole attraction is formed.

This type of intermolecular force is significantly weaker than hydrogen bonding but is stronger than dispersion forces. The strength of dipole-dipole interactions is affected by the polarity of the molecules and their spatial arrangement relative to one another. Substances with strong dipole-dipole interactions typically have higher melting and boiling points than those with only London dispersion forces.

London Dispersion Forces

London dispersion forces, also known as van der Waals forces, are the weakest type of intermolecular force. They arise from the temporary fluctuations in electron distribution within atoms and nonpolar molecules, which create an instant, temporary dipole.

This induces a dipole in a neighboring atom or molecule, leading to an attraction between the two. London dispersion forces are present in all molecules, whether polar or nonpolar, but they are particularly significant in large, heavy atoms and molecules. The strength of these forces increases with the size of the molecules and the surface area in contact.

Molecular Geometry

Molecular geometry, the three-dimensional arrangement of atoms in a molecule, significantly influences the polarity of the molecule and the types of intermolecular forces it can participate in.

For instance, the linear shape of carbon dioxide (CO2) leads to symmetrical charge distribution, making it a nonpolar molecule even though it contains polar covalent bonds. In contrast, the bent shape of water (H2O) leads to an asymmetrical charge distribution, resulting in a polar molecule. A thorough understanding of molecular shapes is crucial to predict and explain the behavior of substances in various conditions.