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, as well as between CH₄ and CH₂F₂, can be attributed to the differences in intermolecular forces present within these substances. Ethanol has a higher boiling point than dimethyl ether due to the presence of hydrogen bonding. In contrast, CH₂F₂ has a higher boiling point than CH₄ because of the existence of dipole-dipole interactions.

Step-by-step solution

Identifying the intermolecular forces in each substance

It is important to identify the types of intermolecular forces present in ethanol and dimethyl ether as well as in CH₄ and CH₂F₂. In ethanol (C₂H₅OH), there is a presence of an O-H bond, which leads to hydrogen bonding. In dimethyl ether (CH₃OCH₃), there is a C-O bond, but no O-H bond, so hydrogen bonding is absent and only dipole-dipole and London dispersion forces are present. In CH₄, only London dispersion forces are present. In CH₂F₂, there is a dipole-dipole interaction, however, hydrogen bonding is not present as there is no hydrogen atom connected to a highly electronegative atom (F, O or N).

Comparing the intermolecular forces and their effects on boiling points

Generally, the stronger the intermolecular forces, the higher the boiling point. For ethanol and dimethyl ether, the main difference in intermolecular forces is the presence of hydrogen bonding in ethanol and the absence in dimethyl ether. Hydrogen bonding is the strongest intermolecular force among the three types of forces. Since ethanol has hydrogen bonding while dimethyl ether does not, the boiling point of ethanol is much higher than that of dimethyl ether. For CH₂F₂ and CH₄, the key difference in intermolecular forces between these two substances is the presence of dipole-dipole interactions in CH₂F₂ and the absence in CH₄. Dipole-dipole interactions are stronger than London dispersion forces, which result in a higher boiling point for CH₂F₂ compared to CH₄.

Conclusion

The main difference in boiling points of ethanol and dimethyl ether as well as CH₄ and CH₂F₂ can be explained by the difference in intermolecular forces present within these substances. Ethanol has a higher boiling point than dimethyl ether due to the presence of hydrogen bonding, while the higher boiling point of CH₂F₂ compared to CH₄ can be attributed to the presence of dipole-dipole interactions.

Key concepts

Hydrogen Bonding

Hydrogen bonding is a type of strong intermolecular interaction that significantly influences the physical properties of substances, especially their boiling points. It occurs when a hydrogen atom is covalently bonded to a highly electronegative element like oxygen, nitrogen, or fluorine, and this hydrogen is also attracted to an electronegative atom in another molecule.

The result is a particularly strong dipole-dipole attraction, often visualized as a dotted line between the partially positive hydrogen and the partially negative electronegative atom. For example, in ethanol (C₂H₅OH), hydrogen bonding occurs because of the O-H bond, leading to a high boiling point. This type of force is generally stronger than other types of dipole-dipole and London dispersion forces, resulting in substances with hydrogen bonding often having much higher boiling points than their counterparts without it.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules, where the positive end of one molecule is attracted to the negative end of another molecule. These interactions are caused by the presence of permanent dipoles within the molecules, which happen due to an uneven distribution of electrons. Substances with dipole moments, such as CH₂F₂, exhibit stronger intermolecular forces than nonpolar molecules like CH₄.

Imagine two bar magnets with a north and a south pole. Just as the opposite poles of these magnets are attracted to each other, the opposite charges in polar molecules gravitate towards one another. This interaction is intermediate in strength, being stronger than London dispersion forces but weaker than hydrogen bonding. Consequently, molecules with dipole-dipole interactions tend to have higher boiling points than those with just London dispersion forces.

London Dispersion Forces

London dispersion forces, also known as Van der Waals forces, are the weakest form of intermolecular forces and are present in all molecules, whether polar or nonpolar. They originate from the temporary fluctuations in the electron density of molecules, which creates a momentary dipole. Even in nonpolar molecules like CH₄ (methane), these temporary dipoles can induce dipoles in adjacent molecules, bringing about a slight attraction between them.

The strength of London dispersion forces depends on the size of the molecules and the number of electrons contained within them. Generally, larger molecules with more electrons exhibit stronger London dispersion forces. However, despite their ubiquitous nature, these forces are the weakest link when it comes to the boiling points of substances. Thus, molecules relying solely on London dispersion forces, such as nonpolar molecules with symmetrical shapes, will typically have lower boiling points compared to those with hydrogen bonds or dipole-dipole interactions.