Question

Which type of intermolecular force accounts for each of these differences? (a) Acetone, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO},\) boils at \(56^{\circ} \mathrm{C} ;\) dimethyl sulfoxide or DMSO, (CH \(\left._{3}\right)_{2}\) SO, boils at \(189^{\circ}\) C. (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 \(82.6^{\circ} \mathrm{C}\).

Short answer

(a) The difference in boiling points between acetone and DMSO is due to dipole-dipole interactions, which are stronger in DMSO due to its polar S=O bond. (b) The difference in physical states between CCl4 and CH4 is due to London dispersion forces, which are stronger in CCl4. (c) The difference in boiling points between H2O and H2S is mainly due to hydrogen bonding, which is stronger in H2O. (d) 1-propanol has a higher boiling point than 2-propanol due to stronger hydrogen bonding interactions resulting from the position of the O–H group in the molecular structure.

Step-by-step solution

(a) Acetone and DMSO: Intermolecular forces in play

Acetone and DMSO have different boiling points due to the dipole-dipole interactions. Both molecules have a polar bond, which means that there is an unequal distribution of electrons between the participating atoms. Acetone has a polar C=O bond, and DMSO has a polar S=O bond. The dipole-dipole interaction in DMSO is stronger than in acetone because sulfur is more polarizable and has a larger size than the carbon atom in acetone. This leads to a stronger intermolecular attraction in DMSO, hence its higher boiling point.

(b) CCl4 and CH4: Intermolecular forces in play

CCl4 and CH4 have different states at room temperature: liquid and gas, respectively. The main intermolecular force responsible for this difference is London dispersion forces. These forces are caused by the temporary fluctuations in electron distribution around the molecules. CCl4 has a larger molar mass and more electrons than CH4, leading to a greater dispersion force, which causes the CCl4 to be a liquid at room temperature, while CH4 remains a gas.

(c) H2O and H2S: Intermolecular forces in play

H2O has a much higher boiling point than H2S due to the presence of hydrogen bonding. In H2O, the highly electronegative oxygen atoms attract the hydrogen electrons, leaving the hydrogen atoms electron-deficient. This induces a partial positive charge on the hydrogen atoms and partial negative charge on the oxygen atoms. When H2O molecules come close to each other, these charges lead to the formation of hydrogen bonds, which are much stronger than the dipole-dipole interactions present in H2S. This results in H2O having a significantly higher boiling point than H2S.

(d) 1-propanol and 2-propanol: Intermolecular forces in play

The difference in boiling points between 1-propanol and 2-propanol can be attributed to hydrogen bonding. Both molecules have O–H groups that can form hydrogen bonds, but the positioning of this group in the molecular structure affects the extent of hydrogen bonding the molecules can form. In 1-propanol, the OH group is located at an end of the molecule, which allows it to form more hydrogen bonds with adjacent molecules. In 2-propanol, the OH group is located at the middle of the molecule, which results in weaker hydrogen bonding interactions due to steric hindrance from surrounding methyl groups. This causes 1-propanol to have a higher boiling point than 2-propanol due to the stronger hydrogen bonding interactions.

Key concepts

Dipole-Dipole Interactions

Dipole-dipole interactions are one form of intermolecular force that occur between polar molecules. A polar molecule has an uneven distribution of electrons, creating a positive and a negative end, or "dipole." These oppositely charged ends attract each other, similar to magnets. For example, in acetone - Molecules experience dipole-dipole interactions because the oxygen atom is more electronegative than carbon, - This creates a partial negative charge on the oxygen and a partial positive charge on the carbon. In comparison, DMSO has sulfur bonded to oxygen. The sulfur atom is larger and more polarizable than carbon, - Making the dipole-dipole interactions stronger than in acetone. - Stronger dipole-dipole interactions lead to higher boiling points as the molecules stick together more firmly, requiring more energy, or heat, to separate them. This is why DMSO boils at a much higher temperature than acetone. Dipole-dipole interactions play a critical role in determining the physical properties of polar molecules.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction. It involves a hydrogen atom that is covalently bonded to a very electronegative atom such as oxygen, nitrogen, or fluorine. This setup creates significant differences in electron density, with the hydrogen acquiring a partial positive charge and the electronegative atom acquiring a partial negative charge. For example, water (

• Each hydrogen atom is bonded to oxygen, forming an especially strong dipole because of the great difference in electronegativity.
• These strong dipoles allow for the formation of hydrogen bonds between water molecules, leading to a considerably higher boiling point in comparison to hydrogen sulfide (H₂S), where sulfur replaces oxygen.

Hydrogen bonds

• Are particularly influential in determining water's unique physical properties, such as its high boiling point, surface tension, and specific heat capacity.
• These properties are pivotal for life as they direct water’s behavior in biological systems.

In propanol, hydrogen bonding also affects boiling points because

• The O–H group can form hydrogen bonds with neighboring molecules, influencing 1-propanol’s higher boiling point compared to 2-propanol.

The position of the OH group affects how effectively hydrogen bonds can form.

London Dispersion Forces

London dispersion forces are weak intermolecular forces that arise from temporary fluctuations in electron distribution within atoms or molecules. These are the only type of intermolecular force that can act between nonpolar molecules, but they are present in all molecules. These forces are significant when bigger atoms or molecules with - More electrons permit greater temporary shifts in electron density. In the case of CCl₄ (carbon tetrachloride) - The molecule is nonpolar despite its symmetric tetrahedral shape. - With a larger electron cloud than methane (CH₄), - CCl₄ has stronger London dispersion forces, allowing it to be a liquid at room temperature, while CH₄ remains a gas. London forces become more influential as the size and mass of the molecule increase, making them quite essential in substances made of heavy atoms or large molecules. Though weaker individually than dipole-dipole or hydrogen bonds, - These forces can add up to create significant attractions in substances consisting of large numbers of particles.