Question

In dichloromethane, \(\mathrm{CH}_{2} \mathrm{Cl}_{2}(\mu=1.60 \mathrm{D})\), the dispersion force contribution to the intermolecular attractive forces is about five times larger than the dipole-dipole contribution. Would you expect the relative importance of the two kinds of intermolecular attractive forces to differ (a) in dibromomethane \((\mu=1.43 \mathrm{D})\), (b) in difluoromethane \((\mu=1.93 \mathrm{D}) ?\) Explain.

Short answer

In dibromomethane, the dispersion forces will be more significant, and the dipole-dipole interactions will be less significant compared to dichloromethane, leading to a larger ratio between the two kinds of attractive forces. In difluoromethane, the dispersion forces will be less significant, and the dipole-dipole interactions will be more significant compared to dichloromethane, leading to a smaller ratio between the two kinds of attractive forces.

Step-by-step solution

Compare polarizability and size

We'll start by comparing the polarizability and size of three given compounds. More polarizable and larger molecules will have stronger London dispersion forces. 1. Difluoromethane \((\mathrm{CH}_{2} \mathrm{F}_{2})\): Fluorine is the smallest of the halogens, leading to a smaller and less polarizable molecule compared to dichloromethane and dibromomethane. 2. Dichloromethane \((\mathrm{CH}_{2} \mathrm{Cl}_{2})\): Chlorine is larger and more polarizable than fluorine, leading to stronger dispersion forces. 3. Dibromomethane \((\mathrm{CH}_{2} \mathrm{Br}_{2})\): Bromine is the largest and most polarizable among the three halogens, leading to the strongest dispersion forces.

Compare dipole moments

Now, we will compare the dipole moments of the given molecules. Higher dipole moments indicate stronger dipole-dipole interactions. 1. Difluoromethane \((\mu=1.93 \mathrm{D})\): This molecule has the highest dipole moment among the three molecules, indicating stronger dipole-dipole interactions. 2. Dichloromethane \((\mu=1.60 \mathrm{D})\): The exercise provides this molecule's dipole-dipole interaction as a reference point. 3. Dibromomethane \((\mu=1.43 \mathrm{D})\): This molecule has the lowest dipole moment among the three molecules, indicating weaker dipole-dipole interactions.

Analyze the relative importance of intermolecular forces

Let's analyze the relative importance of dispersion forces and dipole-dipole interactions in the given molecules. (a) In dibromomethane \((\mathrm{CH}_{2} \mathrm{Br}_{2})\): - Dispersion forces: Dibromomethane has the most polarizable and largest molecule, leading to stronger dispersion forces than in dichloromethane. - Dipole-dipole interactions: Dibromomethane has the lowest dipole moment among the three molecules, leading to weaker dipole-dipole interactions than in dichloromethane. Overall, the dispersion forces will be more significant, and the dipole-dipole interactions will be less significant in dibromomethane. Therefore, the ratio between the two kinds of attractive forces will be even larger than it is in dichloromethane. (b) In difluoromethane \((\mathrm{CH}_{2} \mathrm{F}_{2})\): - Dispersion forces: Difluoromethane has the least polarizable and smallest molecule, leading to weaker dispersion forces than in dichloromethane. - Dipole-dipole interactions: Difluoromethane has the highest dipole moment among the three molecules, leading to stronger dipole-dipole interactions than in dichloromethane. Overall, the dispersion forces will be less significant, and the dipole-dipole interactions will be more significant in difluoromethane. Therefore, the ratio between the two kinds of attractive forces will be smaller than it is in dichloromethane.

Key concepts

Dispersion Forces

Dispersion forces, also known as London dispersion forces, are a type of intermolecular force that occurs in all molecules, whether they are polar or nonpolar. These forces are the result of temporary dipoles that occur when the electron distribution within a molecule becomes momentarily asymmetric, creating a temporary positive charge on one side and a negative charge on the other side.

These temporary dipoles can induce dipoles in adjacent molecules, leading to a weak attraction between them. The strength of these forces is dependent on the size of the molecules and their electron cloud's polarizability. Larger molecules with more electrons have more significant dispersion forces because their electron clouds are more easily distorted, creating stronger temporary dipoles. In essence, dispersion forces are stronger in molecules that are larger and more polarizable.

Regarding the exercise, in dibromomethane, the bromine atoms are larger and have more diffuse electron clouds than chlorine atoms in dichloromethane, resulting in stronger dispersion forces. On the other hand, difluoromethane, having smaller fluorine atoms, will have weaker dispersion forces compared with dichloromethane.

Dipole-Dipole Interactions

Dipole-dipole interactions are intermolecular forces that occur between polar molecules, where each molecule has a permanent dipole moment. A dipole moment occurs due to the uneven distribution of electrons in a molecule, leading to a separation of charges with a positive end and a negative end.

These interactions can be illustrated by considering the positive end of one polar molecule attracting the negative end of another, resulting in a force that holds the molecules close together. The strength of dipole-dipole interactions depends on the magnitude of the dipole moments and the distance between the molecules. Higher dipole moments, indicating a greater separation of charge within the molecule, lead to stronger dipole-dipole interactions.

In the context of the given exercise, difluoromethane, which has the highest dipole moment, will exhibit stronger dipole-dipole interactions than dichloromethane or dibromomethane. Thus, despite the weaker dispersion forces, the dipole-dipole interactions play a more significant role in the overall intermolecular attractions within difluoromethane.

Molecular Polarizability

Molecular polarizability refers to the extent to which the electron cloud of a molecule can be distorted by an external electric field, resulting in an induced dipole. Highly polarizable molecules have loosely held valence electrons, which means that the electron cloud can easily shift and produce a significant dipole moment when influenced by nearby electrical charges.

Polarizability is a crucial factor in determining the strength of dispersion forces—the more polarizable a molecule, the stronger the dispersion forces. It is affected by the number of electrons in a molecule and the volume of the electron cloud, which is why larger atoms or molecules with more electrons are typically more polarizable.

• For instance, dibromomethane is more polarizable than dichloromethane, contributing to its stronger dispersion forces.
• Conversely, difluoromethane, with less polarizable fluorine atoms, will have weaker dispersion forces.

Understanding molecular polarizability helps to predict and explain the relative strength of intermolecular attractions in different substances, as seen in the solutions provided in the exercise.