Question

Explain why the boiling point of ethanol \(\left(78^{\circ} \mathrm{C}\right)\) is much higher than that of its isomer, dimethyl ether \(\left(-25^{\circ} \mathrm{C}\right)\) and why the boiling point of \(\mathrm{CH}_{2} \mathrm{~F}_{2}\left(-52{ }^{\circ} \mathrm{C}\right)\) is far above that of \(\mathrm{CH}_{4}\left(-128^{\circ} \mathrm{C}\right)\).

Short answer

The boiling point differences between ethanol and dimethyl ether, and between difluoromethane and methane are due to the different types and strengths of intermolecular forces present in these compounds. Ethanol has a higher boiling point than dimethyl ether because of the stronger hydrogen bonding present in ethanol. Similarly, difluoromethane has a higher boiling point than methane due to the presence (although weak) of hydrogen bonding in addition to van der Waals forces.

Step-by-step solution

Identify the molecular structure and intermolecular forces present in each compound.

To compare the boiling points, it's necessary to first understand the molecular structure and the types of intermolecular forces present in each molecule. Ethanol: \(C_2H_6O\), due to the presence of a hydroxyl group (-OH), hydrogen bonding occurs between ethanol molecules. Dimethyl ether: \(C_2H_6O\), due to the presence of oxygen atom and no hydroxyl group, it is only subject to van der Waals forces (London dispersion forces). Difluoromethane: \(CH_2F_2\), due to the presence of fluorine atoms, hydrogen bonding occurs between the molecules (though weaker than in ethanol, as the hydrogen is not bonded directly to fluorine). Methane: \(CH_4\), no hydrogen bonding or dipole-dipole interactions occur. It is only subject to van der Waals forces.

Compare the boiling points based on the intermolecular forces.

Now that we've identified the intermolecular forces present in each molecule, we can discuss how these forces are related to the boiling points. Boiling points are influenced by the strength of the intermolecular forces; stronger forces lead to higher boiling points. Ethanol vs. Dimethyl Ether: Ethanol has much stronger hydrogen bonding compared to the weaker van der Waals forces in dimethyl ether, leading to a higher boiling point for ethanol. Difluoromethane vs. Methane: Even though the hydrogen bonding present in difluoromethane is quite weak, it's sufficient to differentiate it from methane, which experiences only van der Waals forces. This leads to a higher boiling point for difluoromethane.

Conclude the reasons for boiling point differences.

In conclusion, the boiling point differences between ethanol and dimethyl ether, and between difluoromethane and methane, are due to the different types and strengths of intermolecular forces present in these compounds. Ethanol's higher boiling point compared to dimethyl ether is due to the stronger hydrogen bonding, while difluoromethane's higher boiling point compared to methane is due to the presence (although weak) of hydrogen bonding in addition to van der Waals forces.

Key concepts

Intermolecular Forces

An understanding of intermolecular forces (IMFs) is crucial for explaining the boiling points of various substances. IMFs are the forces of attraction or repulsion which act between neighboring particles (molecules, atoms or ions). They are distinct from the intramolecular forces that hold a molecule together, such as covalent bonds.

There are several types of IMFs, each varying in strength and impact on a substance's physical properties. The key types of IMFs include ionic bonds, hydrogen bonds, dipole-dipole interactions, and van der Waals forces which encompass London dispersion forces and dipole-induced dipole attractions. Among these, ionic bonds are generally the strongest, followed by hydrogen bonds, then dipole-dipole interactions, with van der Waals forces being the weakest.

• Ionic bonds form between positively and negatively charged ions.
• Hydrogen bonds typically form between a hydrogen atom of one molecule and a highly electronegative atom—such as oxygen, nitrogen, or fluorine—of another molecule.
• Dipole-dipole interactions occur between polar molecules, where positive ends are attracted to negative ends.
• Van der Waals forces are a catch-all term that includes several types of weaker attractions, such as instantaneous dipole-induced dipole attractions (London dispersion forces).

The boiling point of a substance is closely related to its IMFs; generally, the stronger the forces, the higher the boiling point because more energy is required to overcome these attractions and turn the liquid into a gas.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction and is significantly stronger than regular van der Waals forces. This form of bonding occurs when hydrogen is covalently bonded to a highly electronegative atom, and this hydrogen atom comes into close proximity with another electronegative atom with a lone pair of electrons.

Hydrogen bonds have a substantial impact on the physical properties of compounds, including their boiling points. Molecules with hydrogen bonds require more energy to separate from each other during boiling, resulting in higher boiling points.

For example, ethanol \(C_2H_6O\) with its hydroxyl group (\(-OH\)) can form hydrogen bonds between its molecules, which considerably increases its boiling point to \(78^\circ C\) compared to dimethyl ether (also \(C_2H_6O\)), which lacks this ability and boils at \( -25^\circ C\). It's the presence of the hydroxyl group in ethanol, allowing for hydrogen bonding, that is the key to its significantly higher boiling point.

Van der Waals Forces

Van der Waals forces, named after Dutch physicist Johannes Diderik van der Waals, are the weakest of the intermolecular forces. They include London dispersion forces, which arise instantaneously due to the random distribution of electrons creating temporary dipoles, and dipole-induced dipole forces, which occur when a dipole in one molecule induces a dipole in another.

London dispersion forces are present in all molecules, regardless of whether they are polar or nonpolar. These forces are particularly significant in nonpolar molecules, like methane \(CH_4\), which lacks other stronger forms of IMFs. Since these forces are relatively weak, substances that rely solely on London dispersion forces for cohesion tend to have low boiling points. This is exemplified in methane, which has a boiling point of \( -128^\circ C\), contrasting with difluoromethane \(CH_2F_2\), where even weak hydrogen bonds result in a comparatively higher boiling point of \( -52^\circ C\).

The strength of London dispersion forces typically increases with the number of electrons in the molecule, leading to higher boiling points in larger molecules due to these temporary, yet ubiquitous, attractive forces.