Question

Look up and compare the normal boiling points and normal melting points of \(\mathrm{H}_{2} \mathrm{O}\) and \(\mathrm{H}_{2} \mathrm{~S}\). (a) Based on these physical properties, which substance has stronger intermolecular forces? What kind of intermolecular forces exist for each molecule? (b) Predict whether solid \(\mathrm{H}_{2} \mathrm{~S}\) is more or less dense than liquid \(\mathrm{H}_{2} \mathrm{~S}\). How does this compare to \(\mathrm{H}_{2} \mathrm{O}\) ? Explain. (c) Water has an unusually high specific heat. Is this related to its intermolecular forces? Explain.

Short answer

(a) H2O has stronger intermolecular forces due to its higher boiling and melting points compared to H2S. H2O has hydrogen bonding, while H2S has dipole-dipole interactions. (b) Solid H2S is likely to have a similar or slightly greater density than liquid H2S, while H2O has lower density in the solid phase (ice) compared to the liquid phase due to the structure of ice created by hydrogen bonding. (c) Water's high specific heat is related to its strong intermolecular forces (hydrogen bonding), as more energy is required to disrupt these bonds, resulting in more heat being absorbed for a given temperature change.

Step-by-step solution

Comparing normal boiling and melting points

Lookup the boiling and melting points of H2O and H2S. We have: - H2O: melting point = 0°C (273 K), boiling point = 100°C (373 K) - H2S: melting point = -85.5°C (187.6 K), boiling point = -60.7°C (212.5 K)

Identifying the stronger intermolecular forces and types of forces

Since H2O has higher boiling and melting points compared to H2S, it has stronger intermolecular forces. The types of intermolecular forces for each molecule are: - H2O: hydrogen bonding (a type of dipole-dipole interaction) - H2S: dipole-dipole interaction (lacks hydrogen bonding due to the absence of a highly electronegative atom like O, N, or F)

Predicting the density comparison

As H2S does not experience hydrogen bonding, it is likely that solid H2S has a similar or slightly greater density than liquid H2S (molecules would be packed more closely in solid phase). On the other hand, H2O has a lower density in its solid phase (ice) than its liquid phase.

Explaining the density difference in H2O

The density difference between solid and liquid H2O arises due to the structure of ice being hexagonal and having open spaces due to the hydrogen bonding. This organized structure results in lower density compared to the liquid phase, where molecules are more randomly arranged and occupy less space.

Explaining the relationship between specific heat and intermolecular forces

The high specific heat of water is linked to its strong intermolecular forces (hydrogen bonding). As more energy is needed to disrupt these strong hydrogen bonds between water molecules, more heat is absorbed. This results in a high specific heat value as it takes more energy to raise the temperature of water compared to substances with weaker intermolecular forces.

Key concepts

Boiling Point

The boiling point of a substance is the temperature at which its liquid phase transitions to a gas phase. This change requires enough energy to overcome the forces holding the molecules in the liquid state, meaning the boiling point is closely related to intermolecular forces. In our comparison, water (\(\mathrm{H}_{2} \mathrm{O}\)) has a significantly higher boiling point at 100°C (373 K) than hydrogen sulfide (\(\mathrm{H}_{2} \mathrm{~S}\)) at -60.7°C (212.5 K).

• This high boiling point in water suggests stronger intermolecular forces are at play, namely hydrogen bonding, which is a subset of dipole-dipole interactions.
• Hydrogen sulfide, with its lower boiling point, does not have hydrogen bonds and relies on weaker dipole-dipole interactions due to its less electronegative sulfur atom.

Thus, the substantial difference in boiling points indicates water has much stronger intermolecular forces than hydrogen sulfide.

Melting Point

Melting point defines the temperature at which a solid turns into a liquid. Like boiling point, it reflects how tightly molecules are held together by intermolecular forces. \(\mathrm{H}_{2} \mathrm{O}\) has a melting point of 0°C (273 K), whereas \(\mathrm{H}_{2} \mathrm{~S}\) has a melting point of -85.5°C (187.6 K).

• A higher melting point in water indicates stronger intermolecular forces at work, specifically hydrogen bonding, which stabilizes the solid structure by forming a network of bonds.
• In contrast, \(\mathrm{H}_{2} \mathrm{~S}\) melts at a much lower temperature since its dipole-dipole interactions are less effective at maintaining a solid structure.

Again, the difference in melting points aligns with stronger hydrogen bonds in water versus mere dipole-dipole forces in hydrogen sulfide.

Density Comparison

Density is the mass per unit volume of a substance. It changes between phases due to the arrangement of molecules. For \(\mathrm{H}_{2} \mathrm{~S}\), without hydrogen bonding, its solid phase is likely similar or denser than its liquid phase. This means molecules are close in the solid form as expected.

• In water (\(\mathrm{H}_{2} \mathrm{O}\)), the situation is unique. Ice (solid water) is less dense than liquid water due to the open hexagonal structure created by hydrogen bonding.
• This hexagonal arrangement in ice means molecules are further apart, hence less dense.

While typical substances become denser upon solidifying, the unique hydrogen-bonded structure of ice results in a lower density compared to liquid water, creating the well-known phenomenon that ice floats.

Specific Heat

Specific heat is the amount of heat energy required to raise the temperature of a substance by a certain amount. Water has a notably high specific heat, which can be attributed to its hydrogen bonding.

• These strong bonds mean considerable energy is needed to increase the kinematic energy of the water molecules.
• Thus, water can absorb more heat without a substantial rise in temperature, which is why it exhibits such a high specific heat compared to other liquids.

The ability to hold heat makes water an excellent thermal buffer, important in both natural ecosystems and technological applications, where thermal regulation is crucial.

Hydrogen Bonding

Hydrogen bonding is a type of dipole-dipole attraction between molecules, where hydrogen is bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine. \(\mathrm{H}_{2} \mathrm{O}\) showcases prominent hydrogen bonding, responsible for many of its unique physical properties.

• The presence of oxygen allows water molecules to form extensive hydrogen-bond lattice networks, explaining why water has such high boiling and melting points.
• This bonding also influences water’s density and specific heat, contributing to phenomena such as ice's lower density and water’s ability to store heat energy.

In the case of \(\mathrm{H}_{2} \mathrm{~S}\), absence of a highly electronegative atom limits hydrogen bonding, leading to weaker intermolecular connections and subsequently lower boiling and melting points.