11.8
Vapor Pressure
When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally c…
In a closed system, at vapor–liquid equilibrium, condensation and vaporization occur at the same rate, with no net change in the masses of the two phases.
The partial pressure exerted by the gas phase in dynamic equilibrium with its liquid is called its vapor pressure. The more molecules there are in the vapor phase, the higher the vapor pressure will be.
Thus, vapor pressure is a reflection of the tendency of liquid molecules to escape into the vapor phase at a given temperature. It is a measurable quantity, governed by intermolecular forces.
Volatility qualitatively describes this tendency, based on the vapor pressures of liquids held under the same conditions.
For example, compare hexane and water held at the same temperature. Because hexane displays weak dispersion forces and water exhibits strong hydrogen bonds, hexane vaporizes more readily than water does.
In a closed system at equilibrium, hexane has a higher vapor pressure than water does: hexane is volatile, while water is nonvolatile.
The distribution of thermal energies in the liquid phase is a function of temperature. Heating a liquid raises its temperature, indicating that the molecules have higher thermal energies, which leads to a higher vaporization rate and higher vapor pressure.
When the vapor pressure equals the external pressure, the liquid begins to boil, and the temperature at which this happens is called the boiling point of the liquid.
The normal boiling point of a liquid is the temperature at which the liquid’s vapor pressure equals 1 atm. However, at a different external pressure, the liquid will boil at a temperature different from its normal boiling point.
For instance, at standard sea level, where the atmospheric pressure is 1 atm, water boils at 100 °C. At a higher altitude — where the atmospheric pressure is less than 1 atm — the vapor phase requires fewer molecules to equal the lower external pressure. This explains why water will boil at a lower temperature.
In a pressure cooker, the higher external pressure demands more vapor-phase molecules, thereby requiring the water to be at a higher temperature to boil.
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Q1: What is vapor pressure and how does it relate to dynamic equilibrium?
Vapor pressure is the partial pressure exerted by gas molecules in equilibrium with a liquid in a closed system. At dynamic equilibrium, vaporization and condensation occur at equal rates, with molecules continuously exchanging between phases but no net change in liquid or vapor mass. This measurable quantity reflects the tendency of liquid molecules to escape into the vapor phase at a given temperature.
Q2: How do intermolecular forces affect a liquid's vapor pressure?
Strong intermolecular forces impede vaporization and favor recapture of gas molecules, resulting in low vapor pressure. Weak intermolecular attractions present less barrier to vaporization, yielding high vapor pressures. For example, diethyl ether has weak dispersion forces and high vapor pressure, while water exhibits strong hydrogen bonding and lower vapor pressure.
Q3: Why does heating a liquid increase its vapor pressure?
Heating raises the liquid's temperature, increasing molecular thermal energies. At higher temperatures, a greater fraction of molecules possess sufficient energy to overcome intermolecular forces and escape into the vapor phase. Both the increased escape rate and greater average speed of escaping molecules contribute to higher vapor pressure.
Q4: What is the relationship between vapor pressure and boiling point?
Boiling occurs when a liquid's vapor pressure equals the external atmospheric pressure. The normal boiling point is the temperature at which vapor pressure reaches 1 atm. At higher external pressures, more vapor molecules are needed to balance the pressure, so boiling occurs at higher temperatures. Conversely, at lower pressures, boiling happens at lower temperatures.
Q5: How does volatility differ from vapor pressure?
Volatility qualitatively describes the tendency of liquid molecules to escape into the vapor phase, while vapor pressure is the quantitative measure of this tendency under specific conditions. Volatility is a comparative term based on vapor pressures of liquids held under identical conditions. For instance, hexane is volatile relative to water because hexane has weaker intermolecular forces and higher vapor pressure.
Q6: Why does water boil at different temperatures at different altitudes?
At higher altitudes, atmospheric pressure is lower than 1 atm. Water's vapor pressure requires fewer molecules to equal this reduced external pressure, so boiling occurs at lower temperatures. Conversely, in a pressure cooker with higher external pressure, water must reach higher temperatures before its vapor pressure equals the elevated external pressure.
Q7: Does the size of a container or liquid surface area affect vapor pressure?
Neither the surface area of the liquid in contact with vapor nor the size of the vessel affects the vapor pressure at a given temperature. However, these factors do influence the time required for vapor-liquid equilibrium to be established. Vapor pressure depends only on the chemical identity of the substance and temperature.