Question

Look up and compare the normal boiling points and normal melting points of \(\mathrm{H}_{2} \mathrm{O}\) and \(\mathrm{H}_{2} \mathrm{~S}\). Based on these physical properties, which substance has stronger intermolecular forces? What kinds of intermolecular forces exist for each molecule?

Short answer

\(\mathrm{H}_{2}\mathrm{O}\) has stronger intermolecular forces due to hydrogen bonding, whereas \(\mathrm{H}_{2}\mathrm{S}\) has weaker forces.

Step-by-step solution

Gather Data on Boiling and Melting Points

First, we need to find the normal boiling points and melting points for both \(\mathrm{H}_{2}\mathrm{O}\) and \(\mathrm{H}_{2}\mathrm{S}\). For water (\(\mathrm{H}_{2}\mathrm{O}\)), the normal boiling point is approximately 100°C and the melting point is 0°C. For hydrogen sulfide (\(\mathrm{H}_{2}\mathrm{S}\)), the boiling point is approximately -60°C and the melting point is -82°C.

Analyze the Strength of Intermolecular Forces

Compare the boiling and melting points. \(\mathrm{H}_{2}\mathrm{O}\) has higher boiling and melting points than \(\mathrm{H}_{2}\mathrm{S}\), indicating that \(\mathrm{H}_{2}\mathrm{O}\) has stronger intermolecular forces. Higher boiling and melting points suggest stronger attractions between molecules, needing more energy to change phases.

Identify Intermolecular Forces

\(\mathrm{H}_{2}\mathrm{O}\) experiences hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Hydrogen bonding is exceptionally strong in \(\mathrm{H}_{2}\mathrm{O}\) due to the presence of highly electronegative oxygen. \(\mathrm{H}_{2}\mathrm{S}\) experiences weaker dipole-dipole interactions and London dispersion forces, as sulfur is less electronegative than oxygen and does not participate in hydrogen bonding as effectively.

Conclusion on Strength of Intermolecular Forces

Based on boiling and melting points and the types of forces present, \(\mathrm{H}_{2}\mathrm{O}\) has stronger intermolecular forces than \(\mathrm{H}_{2}\mathrm{S}\). This is primarily due to hydrogen bonding in water, which is not prevalent in hydrogen sulfide.

Key concepts

Boiling Points

Boiling points are crucial physical properties that provide insight into the strength of intermolecular forces within a substance. The boiling point is defined as the temperature at which a liquid turns into a vapor. It happens when a substance's vapor pressure equals the surrounding pressure. Essentially, higher boiling points indicate stronger molecular attractions.
For water (\(\mathrm{H}_{2}\mathrm{O}\)), the normal boiling point is about 100°C. In contrast, hydrogen sulfide (\(\mathrm{H}_{2}\mathrm{S}\)) boils at a much lower temperature, around -60°C. The significant difference in boiling points between these two substances signals that \(\mathrm{H}_{2}\mathrm{O}\) has much stronger intermolecular forces than \(\mathrm{H}_{2}\mathrm{S}\) where more energy is necessary for water molecules to escape their liquid state.
The boiling point is influenced by various factors such as:

• Intermolecular forces: Strong forces mean a higher boiling point.
• Molecular weight: Heavier molecules often have higher boiling points.
• Presence of functional groups: Such as hydroxyl groups that can form hydrogen bonds.

Melting Points

Melting points offer further insight into the strength of molecular forces within solids. It refers to the temperature at which solid turns to liquid. Similarly, to boiling points, higher melting points signify stronger intermolecular attractions.
In the case of water (\(\mathrm{H}_{2}\mathrm{O}\)), the melting point is 0°C, whereas hydrogen sulfide (\(\mathrm{H}_{2}\mathrm{S}\)) melts at -82°C. The higher melting point of water compared to hydrogen sulfide tells us that \(\mathrm{H}_{2}\mathrm{O}\) molecules are held together more securely than \(\mathrm{H}_{2}\mathrm{S}\) molecules.
The melting point can be affected by:

• Type and strength of intermolecular forces: Stronger forces need more energy to break.
• Symmetry and crystal structure: Highly organized structures often have higher melting points.
• Molecular size and mass: Larger molecules may require more energy to melt due to larger van der Waals forces.

Hydrogen Bonding

Hydrogen bonding is one of the strongest types of intermolecular forces. It occurs when hydrogen is covalently bonded to a highly electronegative atom such as oxygen, nitrogen, or fluorine, creating a significant dipole moment.
In \(\mathrm{H}_{2}\mathrm{O}\), hydrogen bonding heavily influences its physical properties. Each water molecule can form up to four hydrogen bonds, leading to a high boiling point due to the energy needed to break these bonds. This is absent or very weak in hydrogen sulfide (\(\mathrm{H}_{2}\mathrm{S}\)) because sulfur is less electronegative and does not support strong hydrogen bonding.
Key aspects of hydrogen bonding:

• Occurs in molecules where hydrogen is bonded to N, O, or F.
• Results in higher boiling and melting points.
• Creates unique attributes like water's high heat capacity and surface tension.

Dipole-Dipole Interactions

Dipole-dipole interactions are intermolecular forces between polar molecules, arising from attractions between partial positive and negative charges. They are generally weaker than hydrogen bonds but still significant in determining the physical properties of compounds.
In both \(\mathrm{H}_{2}\mathrm{O}\) and \(\mathrm{H}_{2}\mathrm{S}\), dipole-dipole interactions are present. However, the stronger dipole moment in water due to the highly electronegative oxygen results in more robust intermolecular attractions compared to the weaker interactions in hydrogen sulfide.
Importance of dipole-dipole interactions:

• Facilitate molecular alignment in a solid or liquid phase.
• Influence physical properties like boiling and melting points.
• More significant in smaller molecules where hydrogen bonding is absent.