Question

Which of the following statements about intermolecular forces is(are) true? a. London dispersion forces are the only type of intermolecular force that nonpolar molecules exhibit. b. Molecules that have only London dispersion forces will always be gases at room temperature \((25^{\circ} \mathrm{C})\). c. The hydrogen-bonding forces in \(\mathrm{NH}_{3}\) are stronger than those in \(\mathrm{H}_{2} \mathrm{O}\). d. The molecules in \(\mathrm{SO}_{2}(\mathrm{~g})\) exhibit dipole-dipole intermolecular interactions. e. \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{3}\) has stronger London dispersion forces than does \(\mathrm{CH}_{4}\).

Short answer

Statements A, D, and E are true. Statement A is true because London dispersion forces are the only type of intermolecular force that nonpolar molecules exhibit. Statement D is true because SO₂ has a bent molecular geometry and exhibits dipole-dipole intermolecular interactions. Statement E is true because propane, with a larger molecular mass and more electrons than methane, has stronger London dispersion forces.

Step-by-step solution

Evaluate Statement A

London dispersion forces are the weakest type of intermolecular force, and they arise due to temporary dipoles in nonpolar molecules. Although these forces can be relatively weak, they are indeed the only type of intermolecular force that nonpolar molecules exhibit. Thus, statement A is true.

Evaluate Statement B

London dispersion forces are weak and can generally be overcome with relatively low amounts of energy. However, it is not always true that molecules with only London dispersion forces will be gases at room temperature. For example, larger nonpolar molecules with greater surface areas can have significant London dispersion forces, making them liquids or even solids at room temperature. Thus, statement B is false.

Evaluate Statement C

Hydrogen bonding occurs when hydrogen is bonded to a highly electronegative atom, such as nitrogen, oxygen, or fluorine. The electronegativity difference between hydrogen and the highly electronegative atom leads to a strong dipole-dipole interaction. Comparing H₂O and NH₃, both of which have hydrogen bonding, the hydrogen bonding in water is stronger due to the higher electronegativity of oxygen compared to nitrogen. Therefore, statement C is false.

Evaluate Statement D

SO₂ has a bent molecular geometry and is a polar molecule due to the presence of the highly electronegative oxygen atoms causing a net dipole moment. Polar molecules exhibit dipole-dipole intermolecular interactions, in addition to London dispersion forces. Therefore, statement D is true.

Evaluate Statement E

CH₃CH₂CH₃ (propane) and CH₄ (methane) are both nonpolar hydrocarbon molecules. They both exhibit London dispersion forces due to the temporary dipoles in their electron clouds. Since propane has a larger molecular mass and more electrons than methane, its London dispersion forces are stronger. Therefore, statement E is true. In conclusion, statements A, D, and E are true, while statements B and C are false.

Key concepts

London Dispersion Forces

London dispersion forces are the weakest type of intermolecular forces. They occur due to temporary, short-lived shifts in electron density, which create temporary dipoles within nonpolar molecules.
These forces are universally present in all atoms and molecules, but they are the only intermolecular forces available to nonpolar molecules.
Strength of London Dispersion Forces: - The strength of these forces increases as the size of the molecule increases.
- Larger molecules have more electrons, thus creating a larger instantaneous dipole. - Greater surface area in the molecules also enhances these forces because it permits more points of contact between the molecules.
An important point to keep in mind is that while London dispersion forces are generally weak, they can become quite significant in larger molecules. For example, they can cause certain nonpolar substances to be liquids or even solids at room temperature.

Dipole-Dipole Interactions

Dipole-dipole interactions occur in polar molecules, which have a permanent net dipole moment.
These forces occur when the positive end of a permanent dipole in one molecule is attracted to the negative end of a dipole in another molecule.
Understanding Dipole Moment: - A dipole moment is a measure of the separation of positive and negative charges in a molecule.
- Dipole-dipole interactions are stronger than London dispersion forces because they depend on permanent polarization rather than temporary fluctuations in the electron cloud.
Factors Influencing Strength: - The greater the dipole moment, the stronger the dipole-dipole interaction.
- Geometry of the molecule also affects these interactions, as seen in bent or asymmetrical molecules like sulfur dioxide (SO₂), which demonstrates strong dipole-dipole forces due to its polar nature.
In summary, dipole-dipole interactions are crucial in determining the physical properties of polar molecules, such as boiling and melting points.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction and one of the strongest among intermolecular forces.
A hydrogen bond occurs when a hydrogen atom is covalently bonded to a highly electronegative atom, such as nitrogen, oxygen, or fluorine.
Characteristics of Hydrogen Bonds: - The electronegative atom attracts the electron cloud towards itself, leaving the hydrogen atom with a significant positive charge.
- This positive hydrogen is then attracted to lone pairs on the electronegative atoms of adjacent molecules.
Why Water's Hydrogen Bonding is Stronger: - Water ( H₂O) exhibits stronger hydrogen bonds compared to ammonia ( NH₃) because oxygen is more electronegative than nitrogen.
- Stronger electronegativity results in a more prominent partial positive charge on hydrogen, thus enhancing the strength of hydrogen bonding. Hydrogen bonding profoundly affects the properties of compounds, such as the high boiling and melting points of water, enabling life to thrive.