Question

Which type of intermolecular force accounts for each of these differences: (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 \(\mathrm{Ar}\) is a gas under the same conditions. (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

(a) The difference in boiling points between CH₃OH and CH₃SH can be attributed to hydrogen bonding in methanol. (b) The difference between Xe and Ar can be explained by the stronger London dispersion forces in Xenon. (c) The higher boiling point of Cl₂ compared to Kr is due to stronger dipole-dipole interactions in Cl₂. (d) The difference in boiling points between acetone and 2-methylpropane can be attributed to the stronger dipole-dipole interactions in acetone.

Step-by-step solution

(a) Comparing CH3OH and CH3SH boiling points

We need to compare methanol (CH3OH) with methanethiol (CH3SH). Methanol is an alcohol with an OH group that can participate in hydrogen bonding, a strong intermolecular force. Methanethiol, on the other hand, has an SH group which lacks the hydrogen bonding capability. Both molecules have similar sizes, so the difference in boiling points (65°C for CH3OH and 6°C for CH3SH) can be attributed to the presence of hydrogen bonding in methanol.

(b) Comparing Xe and Ar properties at 120 K

Xenon (Xe) and Argon (Ar) are both noble gases, which means that they do not form molecules and have complete electron shells. The only intermolecular force present in these gases is the London dispersion force. As Xe has a larger atomic mass and more electrons than Ar, the London dispersion forces are stronger in Xenon, causing it to be in a liquid state at 120 K and atmospheric pressure. In contrast, Ar is still in the gaseous state under the same conditions.

(c) Comparing boiling points of Kr and Cl2

Krypton (Kr) is a noble gas with an atomic weight of 84, and Cl2 (chlorine) is a diatomic molecule with a molecular weight of about 71. In the case of Kr, the only intermolecular force present is the London dispersion force. For Cl2, there is a dipole-dipole interaction between the molecules due to the presence of polar covalent bonds between the two chlorine atoms. The dipole-dipole interaction in Cl2 is stronger than the dispersion forces in Kr, which explains the higher boiling point of Cl2 (238 K) compared to Kr (120.9 K).

(d) Comparing boiling points of acetone and 2-methylpropane

Acetone is a polar molecule with a ketone functional group (C=O), whereas 2-methylpropane (also known as isobutane) is a nonpolar hydrocarbon. The primary intermolecular forces in acetone are dipole-dipole interactions between the polar C=O groups, while the primary intermolecular forces in 2-methylpropane are the weaker London dispersion forces. This difference in intermolecular forces accounts for the difference in boiling points, with acetone boiling at 56°C and 2-methylpropane boiling at -12°C.

Key concepts

Hydrogen Bonding

Hydrogen bonding is a type of intermolecular force that occurs when a hydrogen atom is bonded to an electronegative atom like oxygen, nitrogen, or fluorine. This creates a strong attraction between molecules because hydrogen, being small, can get very close to a lone pair of electrons on other molecules, creating a "bridge" or bond that is particularly strong.

In the context of boiling points, hydrogen bonding significantly increases the amount of energy needed to separate molecules from one another. For instance, methanol ( CH_3OH) features an -OH group capable of hydrogen bonding, which accounts for its relatively high boiling point of 65°C compared to methanethiol ( CH_3SH), which lacks this ability. This strong hydrogen bonding in methanol means it requires more energy (or heat) to break these interconnections before it can transition from liquid to gas.

When studying molecular interactions, always look for hydrogen bonding as they can explain why certain substances appear more 'stuck together' at a given temperature.

London Dispersion Forces

London Dispersion Forces (LDFs) are the weakest form of van der Waals forces, found in all molecules, whether polar or nonpolar. They are temporary attractions that occur when electrons in adjacent molecules create instantaneous dipole-induced dipole attractions. Though individually weak, their collective strength can increase with larger atoms or molecules, which possess more electrons that contribute to greater polarizability.

This is evident in comparing noble gases like xenon (Xe) and argon (Ar). Despite being nonpolar and having no formal permanent dipoles, Xe remains liquid at 120 K, while Ar is gaseous. This difference is due to Xe's larger number of electrons creating stronger dispersion forces, thus requiring more energy for its molecules to escape into a gas.

Understanding LDFs helps explain many observations about the physical states of different substances, especially those without significant molecular polarity.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between molecules that have permanent dipoles – that is, molecules with unevenly distributed charges resulting in a partial positive and partial negative end. These interactions are stronger than London dispersion forces. The positive end of one molecule is attracted to the negative end of another, creating an intermolecular attraction that affects physical properties like boiling points.

In comparing krypton (Kr) and chlorine ( Cl_2), chlorine experiences dipole-dipole interactions due to its polar covalent bonds, which are greater than the London dispersion forces present in Kr. This results in a higher boiling point for Cl_2 despite its lower molecular weight.

Recognizing the presence of dipole-dipole interactions in polar molecules can help predict and explain their behavioral characteristics in various chemical reactions and states of matter.

Boiling Points

Boiling points are highly influenced by the type and strength of intermolecular forces present within a substance. A higher boiling point indicates stronger intermolecular forces, meaning more energy is required to separate the molecules and transition from liquid to gas.

In the case of acetone compared to 2-methylpropane, acetone molecules engage in dipole-dipole interactions because of their polar nature, resulting in a higher boiling point of 56°C. This contrasts with 2-methylpropane, which primarily offers London dispersion forces, resulting in a significantly lower boiling point of -12°C.

Understanding the relationship between intermolecular forces and boiling points allows us to predict how substances will behave under different thermal conditions. By analyzing the molecular structure and type of bonding, one can ascertain the amount of energy needed for phase transitions.