Question

What type of intermolecular force accounts for the following differences in each case? (a) \(\mathrm{CH}_{3} \mathrm{OH}\) boils at \(65^{\circ} \mathrm{C}, \mathrm{CH}_{3} \mathrm{SH}\) boils at \(6{ }^{\circ} \mathrm{C}\). (b) Xe is liquid at atmospheric pressure and \(120 \mathrm{~K}\), whereas Ar is a gas. (c) \(\mathrm{Kr}\), atomic weight 84 , boils at \(120.9 \mathrm{~K}\), whereas \(\mathrm{Cl}_{2}\), molecular weight about 71 , boils at \(238 \mathrm{~K}\). (d) Acetone boils at \(56{ }^{\circ} \mathrm{C}\), whereas 2 -methylpropane boils at \(-12{ }^{\circ} \mathrm{C}\).

Short answer

In summary, the primary intermolecular forces accounting for the differences in boiling points are: (a) Hydrogen bonding in CH3OH, (b) London dispersion forces in Xe and Ar, (c) A combination of London dispersion forces and dipole-dipole interactions in Kr and Cl2, and (d) Dipole-dipole interactions in acetone and London dispersion forces in 2-methylpropane.

Step-by-step solution

Case (a): CH3OH vs CH3SH

First, let's identify the chemical structure of the compounds: 1. CH3OH: Methanol, which is an alcohol and has a hydroxyl group (-OH). 2. CH3SH: Methanethiol, which is a thiol and has a sulfur-hydrogen group (-SH). The difference in boiling points can be attributed to the presence of a hydroxyl group in methanol. Methanol can form hydrogen bonds due to the high electronegativity of oxygen, while methanethiol forms weaker dipole-dipole interactions because sulfur is less electronegative. Thus, the primary intermolecular force responsible for the higher boiling point of CH3OH is hydrogen bonding.

Case (b): Xe vs Ar

Both substances are noble gases, which means they are monatomic elements and don't form any chemical compounds. The only type of intermolecular force acting between noble gases is the London dispersion force, which depends on the size of the atom or molecule. Xenon (Xe) has a larger atomic size and more electrons than Argon (Ar), which leads to stronger London dispersion forces between Xenon atoms and consequently, a higher boiling point. In this case, London dispersion forces are responsible for the difference in their boiling points.

Case (c): Kr vs Cl2

The given substances are: 1. Kr: Krypton, which is a noble gas and exists in a monatomic form. 2. Cl2: Chlorine, which exists as a diatomic molecule. Both substances have London dispersion forces as the primary intermolecular interaction. However, the molecular weight of Cl2 is around 71, which is less than Kr's atomic weight of 84. The larger number of electrons present in Kr creates stronger London dispersion forces compared to Cl2. Moreover, Cl2 molecules also experience dipole-dipole interactions. The stronger London dispersion forces in Kr are countered by the additional dipole-dipole interactions in Cl2, leading to a higher boiling point of Cl2. So, in this case, a combination of London dispersion forces and dipole-dipole interactions accounts for the difference in their boiling points.

Case (d): Acetone vs 2-methylpropane

The given substances are: 1. Acetone: A ketone with a carbonyl group (-C=O) in its structure. 2. 2-Methylpropane: An alkane, which is a hydrocarbon containing only single bonds. In acetone, the highly polar carbonyl group creates stronger dipole-dipole interactions compared to 2-methylpropane, which is a nonpolar molecule. The difference in boiling points is mainly due to the presence of stronger dipole-dipole interactions in acetone, whereas only London dispersion forces are present in 2-methylpropane.

Key concepts

Hydrogen Bonding

Hydrogen bonding is a special type of strong dipole-dipole attraction that occurs when a hydrogen atom is bonded to a highly electronegative element like nitrogen, oxygen, or fluorine. This results in hydrogen having a partial positive charge, making it highly attractive to nearby electronegative atoms.

This bonding is apparent in the case of methanol (\(\mathrm{CH}_{3}\mathrm{OH}\)). Methanol contains an -OH group where the hydrogen is bonded to oxygen, a highly electronegative atom. As a result, hydrogen bonding occurs, which dramatically increases the boiling point of methanol. On the other hand, methanethiol (\(\mathrm{CH}_{3}\mathrm{SH}\)) lacks this feature since sulfur is not as electronegative as oxygen, leading to weaker interactions. These hydrogen bonds are responsible for higher boiling points in substances where they occur because they require more energy to break.

• This type of bonding is crucial in biological molecules, helping stabilize structures like DNA and proteins.
• Hydrogen bonds can occur both between molecules (intermolecular) and within different parts of a single molecule (intramolecular).
• They are an essential factor in determining the solubility and boiling points of substances.

London Dispersion Forces

London dispersion forces are the weakest type of intermolecular forces but are present in all atoms and molecules. They arise due to temporary fluctuations in electron distribution within atoms or molecules, creating a temporary dipole that can induce a dipole in a neighboring atom or molecule.

This force is more pronounced in larger atoms or heavy molecules with more electrons. Take the comparison between xenon (Xe) and argon (Ar): xenon has a larger atomic radius and more electrons than argon, resulting in stronger London dispersion forces. This explains why xenon exists as a liquid at a temperature where argon remains gaseous. The increased size and electron cloud of xenon lead to greater temporary dipoles, enhancing these dispersion forces.

• These forces are the only intermolecular attractions in noble gases and nonpolar molecules.
• London dispersion forces increase with the size/mass of the molecules or atoms involved.
• They're also sometimes referred to as van der Waals forces.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules, where the positive end of one molecule attracts the negative end of another. This force is stronger than London dispersion forces but weaker than hydrogen bonds.

In the case of acetone (\(\text{C}_3\text{H}_6\text{O}\)), it has a polar carbonyl group (\(- \mathrm{C}=\mathrm{O}\)). This polarity results in dipole-dipole interactions, significantly increasing its boiling point compared to a non-polar molecule like 2-methylpropane, which only relies on weaker London dispersion forces.

Chlorine (\(\text{Cl}_2\)) and krypton (\(\text{Kr}\)) also exhibit interesting interactions. While krypton's larger electron cloud leads to stronger dispersion forces, chlorine's molecular structure allows for dipole-dipole interactions due to its molecular polarity, contributing to a higher boiling point than expected based solely on London dispersion forces.

• These interactions are significant in solutions, affecting how substances dissolve.
• They're crucial in determining the boiling and melting points of polar substances.
• Dipole-dipole interactions are directional, leading to specific orientations of molecules in a substance.