Question

Which of the following liquids has the higher vapor pressure at \(25^{\circ} \mathrm{C}:\) water or mercury?

Short answer

Water has a higher vapor pressure at 25°C than mercury due to weaker intermolecular forces.

Step-by-step solution

Understanding Vapor Pressure

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid phase at a given temperature. A high vapor pressure means the liquid evaporates more easily at that temperature.

Identify Key Properties

We need to compare the vapor pressures of water and mercury at the same temperature, which is important for understanding which liquid has more molecules in the vapor phase in equilibrium.

Consider Molecular Interactions

Water has strong hydrogen bonding between molecules, leading to lower vapor pressure as more energy is needed to escape the liquid phase. Mercury, being a metal with minimal intermolecular forces in liquid form, has a higher vapor pressure compared to substances like water with strong intermolecular forces.

Consider Liquid Types and Data

Water is a polar molecule with significant intermolecular forces, thus having a lower vapor pressure due to hydrogen bonding. Mercury is a heavy, non-polar element with weaker intermolecular forces in the liquid phase, resulting in higher vapor pressure.

Conclude with Known Values

Water is known to have a vapor pressure of about 23.8 mm Hg at 25°C, while the vapor pressure of mercury is much lower, around 0.0012 mm Hg at the same temperature due to its high molecular weight and metallic bond.

Key concepts

Intermolecular Forces

Intermolecular forces are the attractions between molecules, which play a key role in determining a substance's physical properties, such as boiling and melting points. These forces are essentially what determine how a molecule in a liquid interacts with its neighbors and thus its tendency to remain a liquid or convert into gas.

There are several types of intermolecular forces, including:

• Van der Waals forces, which are weak and include dispersion forces and dipole-dipole attractions.
• Hydrogen bonds, which are a special, stronger type of dipole-dipole interaction.
• Metallic bonds, observed in metals, where electrons are shared over many nuclei, contributing to unique properties like high electrical conductivity.

Liquids with strong intermolecular forces have molecules that cling together tightly, requiring more energy to separate into a gaseous state. This results in lower vapor pressures compared to liquids with weaker intermolecular forces. In the example of water versus mercury, water has stronger intermolecular forces due to hydrogen bonding, making it less prone to evaporation.

Hydrogen Bonding

Hydrogen bonding occurs when a hydrogen atom is covalently bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine. This creates a strong type of dipole-dipole interaction because the hydrogen atom becomes slightly positive, attracting nearby electronegative atoms.

In water, hydrogen bonds are particularly significant. Each water molecule can form hydrogen bonds with up to four neighboring water molecules. This aggregation of hydrogen bonds results in water having unique physical properties, such as a high surface tension and a low vapor pressure relative to its molecular weight. The energy needed to break these extensive hydrogen bonds is substantial; thus, it requires more energy to evaporate, explaining water's relatively low vapor pressure.

These bonds also contribute to the high boiling point of water compared to other molecules with similar molar masses. Understanding hydrogen bonding is crucial when studying liquid properties, because such interactions significantly alter thermal and physical characteristics.

Molecular Polarity

Molecular polarity refers to the distribution of electrical charge over the atoms in a molecule. If a molecule has areas where there is a buildup of charge (either positive or negative), it is considered polar.

Water is a classic example of a polar molecule. Its bent structure gives rise to a partial negative charge near the oxygen atom and partial positive charges near the hydrogen atoms. This asymmetrical charge distribution allows for strong hydrogen bonding.

• Polar molecules tend to have higher boiling points and lower vapor pressures due to their strong intermolecular forces.
• Non-polar molecules, in contrast, have a more even charge distribution, leading to weaker intermolecular forces like London's dispersion forces.

Mercury, however, is a metal and does not form polar molecules. Its molecules are bound by metallic bonds, which are different in nature from the dipole interactions seen in water. These differences mean that mercury has weaker cohesion forces in its liquid state than water, despite its heavier atomic structure, accounting for a generally higher vapor pressure compared to water.