Question

Describe the intermolecular forces that must be overcome to convert these substances from a liquid to a gas: (a) \(\mathrm{SO}_{2}\), (b) \(\mathrm{CH}_{3} \mathrm{COOH}_{\text {, }}\) (c) \(\mathrm{H}_{2} \mathrm{~S}\).

Short answer

To convert the given substances from a liquid to a gas, the following intermolecular forces must be overcome: (a) For SO2, dipole-dipole forces and London dispersion forces must be overcome. (b) For CH3COOH, hydrogen bonding forces, dipole-dipole forces, and London dispersion forces must be overcome. (c) For H2S, dipole-dipole forces and London dispersion forces must be overcome, but hydrogen bonding forces are absent due to the lower electronegativity of sulfur.

Step-by-step solution

Substance (a): SO2

Sulfur dioxide, SO2, is a bent-shaped polar molecule. The central sulfur has two oxygen atoms attached and one lone pair which cause it to have a bent shape. The S-O bonds are polar, creating a dipole moment. To convert SO2 from a liquid to a gas, dipole-dipole forces must be overcome, since these are the forces occurring between polar molecules. In addition, there are also London dispersion forces present in SO2, which are present in all molecules.

Substance (b): CH3COOH

Acetic acid, CH3COOH, is a carboxylic acid. The carboxyl group consists of a C=O bond and an O-H bond, which creates a polar environment due to the difference in electronegativity between the atoms involved. The presence of the O-H bond establishes a potential for hydrogen bonding. To convert acetic acid from a liquid to a gas, both hydrogen bonding forces, as well as dipole-dipole forces due to the polarity of the carboxyl group, must be overcome. And just like in any other substance, London dispersion forces are also present, which need to be overcome.

Substance (c): H2S

Hydrogen sulfide, H2S, is a bent-shaped polar molecule similar to H2O. However, the electronegativity difference between H and S is less than that between H and O, making H2S less polar than H2O. Nonetheless, the S-H bonds are polar, which creates a dipole moment. To convert H2S from a liquid to a gas, the dipole-dipole forces must be overcome. In addition, just like in the other two substances, London dispersion forces are present and must also be overcome. Keep in mind that despite the presence of S-H bonds, hydrogen bonding does not occur in H2S due to the lower electronegativity of sulfur compared to nitrogen, oxygen, and fluorine which are usually involved in hydrogen bonding.

Key concepts

Dipole-Dipole Interactions

Dipole-dipole interactions are an important type of intermolecular force that occur between polar molecules. These forces arise due to the presence of a permanent dipole moment in the molecule. A dipole moment occurs when there is a difference in electronegativity between atoms in a bond, causing an uneven distribution of electric charge. For example, in a molecule like sulfur dioxide (SO₂), which is bent and polar, dipole-dipole interactions play a crucial role.

These interactions effectively "stick" molecules together, making them difficult to separate when converting liquids into gases. This is because the negative end of one molecule is attracted to the positive end of a neighboring molecule.

In general, for molecules like hydrogen sulfide (H₂S) and acetic acid (CH₃COOH), the dipole-dipole forces must be overcome to change their state from liquid to gas. The stronger these dipole-dipole attractions, the higher the energy needed to break them apart in order for the molecules to enter the gas phase.

Hydrogen Bonding

Hydrogen bonding is a particularly strong type of dipole-dipole interaction. It occurs when hydrogen is directly bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. These bonds create a significant polarity, leading to strong attractions between molecules.

In acetic acid (CH₃COOH), hydrogen bonding occurs due to the presence of an O-H bond in its carboxyl group. This contributes further to the intermolecular forces that must be overcome to vaporize the substance.

It is important to note that hydrogen bonds are not present in all substances with hydrogen atoms. For instance, in hydrogen sulfide (H₂S), while there are S-H bonds, hydrogen bonding does not occur. This is because sulfur is not electronegative enough compared to nitrogen, oxygen, or fluorine to form these bonds.

• Hydrogen bonds significantly raise boiling points.
• They result in higher energy requirements to transition substances from liquids to gases compared to other dipole-dipole forces.

London Dispersion Forces

London dispersion forces, also known as van der Waals forces, are present in all molecules, regardless of their polarity. These forces stem from temporary shifts in electron density within an atom or molecule, creating an instantaneous dipole that can induce a similar effect in neighboring atoms or molecules.

Even though these forces are generally the weakest compared to dipole-dipole and hydrogen bonds, they become significant in larger molecules with more electrons. The more electrons a molecule has, the greater the potential for fluctuations in electron density, resulting in stronger London dispersion forces.

For example, substances like sulfur dioxide (SO₂), acetic acid (CH₃COOH), and hydrogen sulfide (H₂S) all experience some level of London dispersion forces. These forces must be considered alongside dipole-dipole attractions when changing the state of a substance from liquid to gas.

• Though weak individually, they cumulatively impact the physical properties of substances.
• They explain why even nonpolar substances can exist as liquids or solids under certain conditions.