Question

In dichloromethane, \(\mathrm{CH}_{2} \mathrm{Cl}_{2}(\mu=1.60 \mathrm{D})\), the disper sion 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 \((\mu=1.43 \mathrm{D}),(\mathrm{b})\) in difluo romethane \((\mu=1.93 \mathrm{D}) ?\) Explain.

Short answer

The relative importance of dispersion forces and dipole-dipole forces in dibromomethane is expected to be similar to that in dichloromethane, due to its slightly smaller dipole moment and stronger dispersion forces. In difluoromethane, the relative importance of dipole-dipole forces is expected to be greater compared to dichloromethane, due to its larger dipole moment and weaker dispersion forces.

Step-by-step solution

Understanding Dispersion and Dipole-Dipole Forces

Dispersion forces, also known as London dispersion forces, are weak intermolecular forces that arise from the temporary fluctuations in electron distribution around molecules. These forces are present in all molecules, regardless of polarity. Dipole-dipole forces are intermolecular forces that occur between polar molecules. These forces result from the interaction between the positive pole of one polar molecule and the negative pole of another polar molecule. These forces depend on the magnitude of the dipole moment.

Comparing Dibromomethane to Dichloromethane

Dibromomethane has a dipole moment of \(1.43 D\), which is slightly smaller than dichloromethane's dipole moment of \(1.60 D\). This would suggest that the dipole-dipole forces could be somewhat weaker in dibromomethane compared to dichloromethane. However, the size and polarizability of the bromine atoms are larger than those of chlorine atoms. This means that dispersion forces are expected to be stronger in dibromomethane than in dichloromethane. As a result, the relative importance of the two kinds of intermolecular attractive forces will not differ significantly in dibromomethane compared to dichloromethane.

Comparing Difluoromethane to Dichloromethane

Difluoromethane has a dipole moment of \(1.93 D\), which is significantly greater than the dipole moment of dichloromethane. This would suggest that dipole-dipole forces are stronger in difluoromethane compared to dichloromethane. However, the size and polarizability of the fluorine atoms are smaller than those of chlorine atoms, resulting in weaker dispersion forces in difluoromethane than in dichloromethane. Thus, we can expect the relative importance of dipole-dipole forces to be greater in difluoromethane compared to dichloromethane.

Conclusion

In summary, the relative importance of dispersion forces and dipole-dipole forces in dibromomethane is expected to be similar to that in dichloromethane. However, in difluoromethane, the relative importance of dipole-dipole forces is expected to be greater compared to dichloromethane due to its larger dipole moment and weaker dispersion forces.

Key concepts

Dipole-Dipole Interactions

Dipole-dipole interactions are a type of intermolecular force that occur in polar molecules, which have regions of positive and negative charge due to an unequal distribution of electrons. These interactions arise because the positive end of one polar molecule is attracted to the negative end of a neighboring molecule.

The force of attraction increases with the molecule's dipole moment, a measure of the strength of its polarity. The dipole moment is determined by the amount of charge separation and the distance between the charges. In substances where dipole-dipole interactions dominate, such as water (H₂O), the molecules tend to have strong attractions, leading to higher boiling points and greater viscosity.

Key Takeaways

• Dipole-dipole interactions are stronger with higher dipole moments.
• These forces are only significant in polar molecules.
• They can greatly affect the physical properties of a substance.

London Dispersion Forces

London dispersion forces, also known as instantaneous dipole-induced dipole forces, are a type of intermolecular force that are present in all molecules, whether polar or nonpolar. These forces are the result of the momentary fluctuations in electron distribution within molecules, which results in a temporary dipole. When this occurs, the temporary dipole can induce a dipole in an adjacent molecule.

The strength of London dispersion forces increases with the size and polarizability of the molecule. Larger atoms or molecules with more electrons are more easily polarizable, which means that the electron cloud can distort in response to temporary dipoles more readily. Molecules like noble gases and nonpolar molecules, such as methane (CH₄), rely entirely on these forces for their intermolecular attractions.

Key Takeaways

• Present in all molecules, including nonpolar ones.
• Strength depends on the size and polarizability of the molecule.
• Crucial for determining the properties of nonpolar substances.

Dipole Moment

The dipole moment is a vector quantity that measures the separation of charge in a molecule and the polarity of the bond. It is defined mathematically as the product of the charge difference between the atoms and the distance separating them. The unit of measurement for dipole moment is the Debye (D).

A higher dipole moment indicates a more significant separation of charge within the molecule, meaning the molecule is more polar. Substances with significant dipole moments will experience stronger dipole-dipole interactions, as seen in the example of difluoromethane (CH₂F₂), which has a higher dipole moment compared to dichloromethane (CH₂Cl₂).

Key Takeaways

• Measures the polarity of a molecule.
• Expressed in Debye (D).
• Higher values indicate stronger dipole-dipole interactions.