Question

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

Short answer

To convert the given substances from a liquid to a gas, the intermolecular forces present in each substance must be overcome. For (a) SO2, dipole-dipole interactions and London dispersion forces need to be overcome. For (b) CH3COOH, strong hydrogen bonds and weaker London dispersion forces must be broken. For (c) H2S, the main intermolecular forces are dipole-dipole interactions and London dispersion forces. In all cases, increasing the temperature provides the necessary energy to break these forces and allow the molecules to transition to a gaseous state.

Step-by-step solution

(a) Identifying the intermolecular forces in SO2

SO2, or sulfur dioxide, is a polar molecule. This means that the electrons in the molecule are unevenly distributed, creating a polar or uneven charge distribution across the molecule. In SO2, the main intermolecular forces are dipole-dipole interactions and London dispersion forces.

(a) Overcoming intermolecular forces in SO2

In order to convert SO2 from a liquid to a gas, the intermolecular forces holding the molecules together in the liquid state, i.e., dipole-dipole interactions and London dispersion forces, have to be overcome. Increasing the temperature provides the energy required to break these forces and allows the molecules to separate from each other, thus transitioning to a gaseous state.

(b) Identifying the intermolecular forces in CH3COOH

CH3COOH, or acetic acid, is a polar molecule containing oxygen and hydrogen atoms bonded together, which allows for the formation of hydrogen bonds, a strong type of dipole-dipole interaction. Additionally, London dispersion forces are present, albeit weaker compared to hydrogen bonds.

(b) Overcoming intermolecular forces in CH3COOH

To convert CH3COOH from its liquid state to a gaseous state, the strong hydrogen bonds holding the molecules together must be overcome, along with the weaker London dispersion forces. By increasing the temperature, sufficient energy is provided to break these intermolecular forces, allowing the molecules to transition to a gaseous state.

(c) Identifying the intermolecular forces in H2S

H2S, or hydrogen sulfide, is a polar molecule due to the presence of sulfur, which is more electronegative than hydrogen. In H2S, the significant intermolecular forces are dipole-dipole interactions and London dispersion forces. It is worth mentioning that H2S does not form hydrogen bonds as sulfur is not as electronegative as oxygen, nitrogen, or fluorine.

(c) Overcoming intermolecular forces in H2S

To convert H2S from a liquid to a gas, the intermolecular forces acting on it must be overcome, mainly the dipole-dipole interactions and London dispersion forces. By applying heat to raise the temperature, the energy required to break these intermolecular forces is provided, thus allowing the H2S molecules to separate and transition to a gaseous state.

Key concepts

Phase Transition

Understanding the concept of a phase transition is crucial when studying the behavior of substances in response to temperature changes. In simple terms, a phase transition is the transformation from one state of matter to another, such as from liquid to gas, which is called vaporization. During this process, specific intermolecular forces that hold the particles together must be overcome. For instance, as the temperature of a liquid rises, the particles gain kinetic energy. When this energy becomes high enough to disrupt the attractive forces between them, they can break free from the liquid state and become a gas. Each substance requires a different amount of energy for this transition, influenced by the strength and type of intermolecular forces involved.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules, which have regions of positive and negative charge due to an uneven distribution of electrons. These molecules have what's known as a dipole moment. One end of a polar molecule carries a slight positive charge, while the other end carries a slight negative charge. When such molecules are in close proximity, the positive region of one molecule is attracted to the negative region of another, causing them to align and form an intermolecular attraction. In the case of SO₂ and H₂S, the presence of dipole-dipole interactions explains why energy, typically in the form of heat, must be added to these substances to disrupt the attractive forces and facilitate a phase transition from liquid to gas.

Hydrogen Bonding

Hydrogen bonding is a more powerful form of dipole-dipole interaction. It occurs when a hydrogen atom is covalently bonded to a highly electronegative atom such as oxygen, nitrogen, or fluorine, and is attracted to another electronegative atom in a nearby molecule. This specific attraction is what gives hydrogen bonds their significant strength. Acetic acid (CH₃COOH) is an example of a substance that forms hydrogen bonds due to the -OH (hydroxyl) group it contains. These hydrogen bonds are so strong that raising the temperature to induce a phase transition from liquid to gas requires more energy than would be needed if only weaker intermolecular forces were present.

London Dispersion Forces

London dispersion forces, also known as induced dipole-induced dipole interactions, are the weakest of all intermolecular forces. They stem from the momentary fluctuations in electron density within atoms and nonpolar molecules, which create temporary dipoles. These transient dipoles can then induce similar dipoles in neighboring molecules, leading to an attraction between them. Even though these forces are weak compared to dipole-dipole interactions and hydrogen bonding, they still require a certain amount of energy to be overcome during a phase transition. London dispersion forces are present in all molecules, but they are most notable in nonpolar compounds and noble gases, and they also contribute to the intermolecular forces in polar molecules like SO₂, CH₃COOH, and H₂S.