Question

Which type of intermolecular force accounts for each of these differences? (a) Acetone, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO},\) boils at \(56^{\circ} \mathrm{C}_{i}\) dimethyl sulfoxide or \(\mathrm{DMSO},\left(\mathrm{CH}_{3}\right)_{2} \mathrm{SO},\) boils at \(189^{\circ} \mathrm{C} .(\mathbf{b})\) \(\mathrm{CCl}_{4}\) is a liquid at atmospheric pressure and room temperature, whereas \(\mathrm{CH}_{4}\) is a gas under the same conditions. \((\mathbf{c})\) \(\mathrm{H}_{2} \mathrm{O}\) boils at \(100{ }^{\circ} \mathrm{C}\) but \(\mathrm{H}_{2} \mathrm{~S}\) boils at \(-60{ }^{\circ} \mathrm{C}\). (d) 1 -propanol boils at \(97^{\circ} \mathrm{C}\), whereas 2 -propanol boils at \(82.6^{\circ} \mathrm{C}\). CC(C)=O DMSO

Short answer

(a) Dipole-dipole interactions in DMSO are stronger. (b) CCl4 exhibits stronger London forces. (c) H2O forms hydrogen bonds. (d) 1-Propanol's structure supports stronger hydrogen bonding.

Step-by-step solution

Analyze Acetone and DMSO Boiling Points

Acetone \((\text{CH}_3)_2 \text{CO}\) has a boiling point of 56°C, whereas dimethyl sulfoxide (DMSO) \((\text{CH}_3)_2 \text{SO}\) boils at 189°C. The large difference in boiling points suggests stronger intermolecular forces in DMSO. DMSO exhibits dipole-dipole interactions due to the highly polar S=O bond, while acetone has weaker dipole-dipole interactions because of the less polar C=O bond.

Compare CCl4 and CH4 Phases at Room Temperature

CCl4 is a liquid at room temperature, while CH4 is a gas. The key difference is due to London dispersion forces. CCl4, being a larger and heavier molecule, has stronger London dispersion forces compared to the smaller CH4 molecule, enabling CCl4 to remain a liquid under the same conditions.

Contrast Boiling Points of H2O and H2S

Water \(\text{H}_2\text{O}\) boils at 100°C, whereas hydrogen sulfide \(\text{H}_2\text{S}\) boils at -60°C. Water molecules form hydrogen bonds, which are much stronger intermolecular forces compared to the weaker dipole-dipole forces and dispersion forces present in H2S molecules. This results in a higher boiling point for water.

Analyze Boiling Points of 1-Propanol and 2-Propanol

1-Propanol boils at 97°C while 2-propanol boils at a lower temperature of 82.6°C. Both have hydrogen bonding; however, the molecular structure affects the strength of these interactions. 1-Propanol has more linear molecular geometry compared to the branched structure of 2-propanol, leading to stronger hydrogen bonding interactions and a higher boiling point.

Key concepts

Boiling Points Comparison

Boiling points are fundamental in understanding the strength of intermolecular forces within different substances. It essentially tells us how much energy is needed to convert a liquid into a gas. Generally, the higher the boiling point, the stronger the intermolecular forces at play. Comparing boiling points between substances, like acetone and DMSO or 1-propanol and 2-propanol, highlights how various intermolecular forces such as dipole-dipole interactions and hydrogen bonding influence physical properties of compounds such as volatility and phase. Knowing the boiling points of substances helps us make sense of their everyday uses and behavior. For instance, DMSO's high boiling point allows it to be used as a solvent in pharmaceutical applications, while the differences in boiling points between isomers like 1-propanol and 2-propanol reflect their structural variations which influence intermolecular interactions.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules that have permanent dipoles. These are the forces that arise from the attraction between the positively charged end of one polar molecule and the negatively charged end of another. This type of intermolecular force is relatively strong compared to London dispersion forces but weaker than hydrogen bonding.An illustrative example can be seen when comparing acetone \((\text{CH}_3)_2\text{CO}\) and DMSO \((\text{CH}_3)_2\text{SO}\). DMSO, due to its highly polar S=O bond, shows strong dipole-dipole interactions, resulting in a significantly higher boiling point than acetone, where the C=O bond is less polar. Thus, stronger dipole-dipole interactions increase the boiling point, as seen with DMSO.

London Dispersion Forces

London dispersion forces are a type of intermolecular force that can exist between all molecules, whether they are polar or nonpolar. These forces arise from temporary shifts in electron density, creating instantaneous dipoles that can attract one another. Despite being the weakest form of intermolecular force, they play a crucial role in determining the state of smaller or nonpolar molecules at room temperature.Take, for example, \(\text{CCl}_4\) and \(\text{CH}_4\). At room temperature, \(\text{CCl}_4\) is a liquid because it experiences stronger London dispersion forces due to its larger size and higher molar mass compared to the gas \(\text{CH}_4\). These dispersion forces are crucial for keeping \(\text{CCl}_4\) in a liquid state under standard conditions.

Hydrogen Bonding

Hydrogen bonds are a special type of strong dipole-dipole interaction that occur when hydrogen is bound to highly electronegative elements like nitrogen, oxygen, or fluorine. These bonds are significantly stronger than other dipole-dipole interactions and are one of the main reasons for the high boiling points of certain substances.Water \(\text{H}_2\text{O}\) is a prime example of a substance with extensive hydrogen bonding. This bonding results in a much higher boiling point \(100^{\circ}\text{C}\) compared to hydrogen sulfide \(\text{H}_2\text{S}\), which relies on weaker dipole-dipole interactions and London dispersion forces and boils at \(-60^{\circ}\text{C}\). Additionally, within the same class of alcohols, the presence of more hydrogen bonds in 1-propanol compared to 2-propanol contributes to its higher boiling point.