Question

Explain why the boiling point of ethanol $\left({78}^{\circ }\mathrm{C}\right)$ is much higher than that of its isomer, dimethyl ether $(-{25}^{\circ }\mathrm{C})$, and why the boiling point of ${\mathrm{CH}}_2{\mathrm{F}}_1(-{52}^{\circ }\mathrm{C})$ is far above that of ${\mathrm{CH}}_4(-{128}^{\circ }\mathrm{C})$.

Short answer

The boiling point of ethanol $({78}^{\circ }\mathrm{C})$ is much higher than that of its isomer, dimethyl ether $(-{25}^{\circ }\mathrm{C})$, due to the stronger hydrogen bonding present in ethanol. These hydrogen bonds require more energy to break, leading to a higher boiling point. On the other hand, the boiling point of ${\mathrm{CH}}_2{\mathrm{F}}_1(-{52}^{\circ }\mathrm{C}$) is far above that of ${\mathrm{CH}}_4(-{128}^{\circ }\mathrm{C}$ because of stronger dipole-dipole interactions in ${\mathrm{CH}}_2{\mathrm{F}}_1$ compared to the weaker London dispersion forces in ${\mathrm{CH}}_4$. More energy is needed to separate ${\mathrm{CH}}_2{\mathrm{F}}_1$ molecules, resulting in a higher boiling point.

Step-by-step solution

Identify the Type of Intermolecular Forces

First, let's identify the type of intermolecular forces present in each of the four molecules. 1. Ethanol: Ethanol has an OH group, which allows it to form hydrogen bonds with other molecules. 2. Dimethyl Ether: This molecule cannot form hydrogen bonds, but it has polar C-O bonds, which can create dipole-dipole interactions. 3. CH2F1: This molecule also has polar C-F bonds, leading to dipole-dipole interactions. 4. CH4: This molecule is nonpolar and has only London dispersion forces acting between its molecules.

Compare the Strength of Intermolecular Forces

Now that we have identified the intermolecular forces present in each molecule, let's compare their strengths: 1. Hydrogen bonding (ethanol) > dipole-dipole interactions (dimethyl ether): Hydrogen bonds are typically stronger than other dipole-dipole interactions because the positively charged hydrogen atom is directly attracted to the highly electronegative elements like oxygen, nitrogen, or fluorine. 2. Dipole-dipole interactions (CH2F1) > London dispersion forces (CH4): Dipole-dipole interactions are stronger than London dispersion forces since they involve the attraction between charged molecular dipoles, whereas London dispersion forces are induced by temporary fluctuations in electron densities.

Explain Differences in Boiling Points

Finally, let's explain why the boiling points of ethanol and CH2F1 are higher than those of their respective isomers: 1. Ethanol vs. Dimethyl Ether: Ethanol has a higher boiling point than dimethyl ether because it can form stronger hydrogen bonds with neighboring ethanol molecules. This means more energy is required to break these bonds, leading to a higher boiling point. 2. CH2F1 vs. CH4: CH2F1 has a higher boiling point than CH4 due to the stronger dipole-dipole interactions present between CH2F1 molecules compared to the weaker London dispersion forces between CH4 molecules. As a result, more energy is needed to separate CH2F1 molecules, leading to a higher boiling point.

Key concepts

Intermolecular Forces

Understanding the influence of intermolecular forces is crucial when analyzing the boiling points of various substances. Intermolecular forces are the attractions between molecules that hold them together in the liquid phase. These forces must be overcome to transition from a liquid to a gas at the boiling point.

The three primary types of intermolecular forces are hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Each type has a different strength and impact on boiling points. Stronger intermolecular forces require more energy—in the form of heat—to be overcome, which leads to higher boiling points. As you delve into this concept, remember that these forces explain why substances with similar molecular weights can have dramatically different boiling points.

Hydrogen Bonding

Hydrogen bonding arises when a hydrogen atom is covalently bonded to a highly electronegative atom, such as oxygen, nitrogen, or fluorine, and is in close proximity to another electronegative atom.

This results in an interaction that is partial ionic and partial covalent in nature. Among the intermolecular forces, hydrogen bonds are one of the strongest, leading to significantly higher boiling points for compounds that have them. This is why ethanol, with its ability to form hydrogen bonds due to its OH group, has a much higher boiling point than many other compounds of similar size that do not engage in hydrogen bonding.

Dipole-Dipole Interactions

Atoms with different electronegativities can create polar molecules with partial positive and negative charges or 'dipoles.' When polar molecules are near each other, the positive end of one molecule is attracted to the negative end of another, creating a dipole-dipole interaction.

These are the second strongest type of intermolecular forces after hydrogen bonds. For example, molecules like CH2F2, with polar C-F bonds, experience these dipole-dipole attractions that elevate their boiling points above nonpolar molecules. While not as potent as hydrogen bonds, dipole-dipole interactions still require a significant amount of energy to break, which means substances with such interactions will have higher boiling points compared to those with only weaker forces.

London Dispersion Forces

London dispersion forces are the weakest type of intermolecular force, yet they are universal as they occur in all molecules, polar and nonpolar alike. They result from temporary shifts in the electron cloud of molecules, causing a fleeting dipole that induces another dipole in a nearby molecule. This transient attraction increases with the size of the electron cloud, which is why larger, nonpolar molecules can still have relatively high boiling points.

However, for smaller molecules like CH4 (methane), London dispersion forces are very weak, leading to much lower boiling points compared to larger molecules or those with stronger intermolecular forces, such as hydrogen bonds or dipole-dipole interactions.