Problem 81 - Chemistry The Molecular Science

Question

Explain why the boiling point \(\left(5.9^{\circ} \mathrm{C}\right)\) of methanethiol, \(\mathrm{CH}_{3} \mathrm{SH},\) is much lower than the boiling point \(\left(64.7{ }^{\circ} \mathrm{C}\right)\) of methanol, \(\mathrm{CH}_{3} \mathrm{OH}\).

Step-by-Step Solution

Answer

Methanol's hydrogen bonding leads to a higher boiling point than methanethiol's weaker forces.

Step 1: Understand Boiling Point

The boiling point is the temperature at which a substance transitions from a liquid to a gas. It is influenced by intermolecular forces; stronger intermolecular forces result in higher boiling points.

Step 2: Compare Molecular Structure

Methanol (\(CH_3OH\)) has a hydroxyl group (OH) that can form hydrogen bonds, which are, compared to other molecular forces, very strong due to the presence of the highly electronegative oxygen atom bonded to hydrogen. Methanethiol (\(CH_3SH\)), on the other hand, contains a thiol group (SH) which cannot form hydrogen bonds because sulfur is less electronegative and larger than oxygen.

Step 3: Analyze Hydrogen Bonding

Hydrogen bonding is a significant intermolecular force that greatly affects a substance's boiling point. In methanol, hydrogen bonding occurs between the hydrogen attached to the oxygen and the lone pair on the oxygen of another molecule. This additional bonding requires more energy to break, thus raising the boiling point.

Step 4: Evaluate Weak Forces in Methanethiol

Methanethiol's main intermolecular forces are van der Waals forces and dipole-dipole interactions, which are weaker than hydrogen bonding. Consequently, the amount of energy required to break these weaker bonds is less, resulting in a lower boiling point.

Step 5: Conclusion of Comparison

Methanol, with its ability to form hydrogen bonds, has a much higher boiling point than methanethiol, which relies only on weaker van der Waals forces and dipole-dipole interactions. Therefore, the greater the intermolecular forces, the higher the boiling point.

Key Concepts

Intermolecular Forces

Intermolecular forces are crucial in determining the physical properties of substances, such as their boiling points. Think of these forces as the glue that holds molecules together in the liquid state. The strength of these forces determines how much energy is needed to separate the molecules during the transition from liquid to gas, affecting the boiling point.
There are several types of intermolecular forces:

• London Dispersion Forces: These are the weakest and occur in all molecular structures due to temporary shifts in electron density.
• Dipole-Dipole Interactions: Occur in polar molecules where permanent dipoles lead to attraction between opposite charges.
• Hydrogen Bonding: A stronger type of dipole-dipole interaction, specifically involving hydrogen bonded to electronegative atoms like oxygen or nitrogen.
• Ion-Dipole Forces: These are particularly strong, but occur in solutions with ions.

Methanol ( CH_3OH ) , with its strong hydrogen bonds, showcases how these forces impact boiling points. In contrast, methanethiol ( CH_3SH ) relies on weaker dipole-dipole interactions and dispersion forces, leading to its lower boiling point.

Hydrogen Bonding

Hydrogen bonding is a special type of strong dipole-dipole interaction that occurs when a hydrogen atom is bonded to a highly electronegative atom, namely fluorine, oxygen, or nitrogen. This strong bond forms because the hydrogen atom's partial positive charge is attracted to the lone pair of electrons on a nearby electronegative atom.
In methanol (\(CH_3OH\)), hydrogen bonding greatly contributes to its relatively high boiling point. Here’s why:

• The oxygen atom is highly electronegative, meaning it pulls electrons towards itself, creating a polarity.
• The hydrogen next to oxygen becomes slightly positive, allowing it to strongly attract an oxygen from another molecule.
• This results in a network of hydrogen bonds which require considerable energy to break.

Without hydrogen bonds, methanol would have a much lower boiling point. Methanethiol (\(CH_3SH\)), lacks this strong bonding due to sulfur being less electronegative than oxygen, which prevents effective hydrogen bonding. Thus, the energy needed to change methanethiol from liquid to gas is much lower.

Molecular Structure Comparison

The molecular structures of methanol (\(CH_3OH\)) and methanethiol (\(CH_3SH\)) are quite similar at first glance, as both have a single carbon atom bonded to three hydrogens and one other atom. However, the key difference lies in the group bonded to this carbon.
In methanol, the hydroxyl group (OH) is bonded to the carbon, while in methanethiol, the thiol group (SH) is bonded instead.

• The (OH) group in methanol allows for hydrogen bonding due to the small size and high electronegativity of oxygen.
• The (SH) group in methanethiol, on the other hand, contains sulfur which is larger and less electronegative than oxygen, thus unable to establish strong hydrogen bonds.

Because methanol's structure allows for significant hydrogen bonding, more energy is needed to separate its molecules compared to methanethiol, which predominantly relies on weaker van der Waals forces. This accounts for the distinct difference in boiling points, illustrating how molecular structure shapes intermolecular forces and subsequently physical properties.