Vapor pressure is a fundamental physical property that describes a liquid or solid’s tendency to transition into a gaseous state, or vapor. It is the pressure that the vapor exerts when it is in a state of dynamic equilibrium with its condensed phase in a closed container. Dynamic equilibrium is a balanced condition where the rate at which molecules escape the liquid phase through evaporation equals the rate at which vapor molecules return to the liquid phase through condensation. Although molecules are constantly moving back and forth between the two states, the total amount of liquid and vapor remains constant. Understanding this equilibrium pressure is important because it indicates how easily a substance will evaporate.
The Influence of Temperature
Temperature is one of the most powerful factors determining a substance’s vapor pressure, and the two are directly related. As the temperature of a liquid increases, the kinetic energy of its molecules also increases. This energy dictates how fast the individual molecules move and vibrate within the liquid. When molecules gain enough kinetic energy, they can overcome the attractive forces holding them within the liquid phase and escape into the space above as vapor. At any given temperature, increasing the temperature dramatically increases the fraction of molecules that possess the necessary energy to escape. This higher rate of escape leads to a greater number of vapor molecules gathering above the liquid. Since pressure is caused by these vapor molecules colliding with the container walls, the increased molecular activity results in a higher vapor pressure. The relationship is not linear; a small rise in temperature can cause an exponential jump in vapor pressure.
The Role of Intermolecular Forces
The identity of the substance itself is the second major factor influencing vapor pressure, which is determined by the strength of its intermolecular forces (IMFs). Intermolecular forces are the attractive forces that exist between neighboring molecules in a liquid. These forces must be overcome for a molecule to escape the liquid and become a vapor. The relationship between the strength of these forces and vapor pressure is inverse: stronger IMFs result in lower vapor pressure, while weaker IMFs lead to higher vapor pressure.
When molecules are held tightly together by strong attractions, they require significantly more kinetic energy to break free and enter the vapor phase. Consequently, fewer molecules can escape at a given temperature, resulting in a low vapor pressure. A good example of a strong force is hydrogen bonding, which is present in water; this force holds water molecules together tightly, giving water a relatively low vapor pressure. In contrast, liquids like diethyl ether, which primarily rely on much weaker London dispersion forces, have high vapor pressures because their molecules escape into the vapor phase easily.
Variables That Do Not Affect Vapor Pressure
There are several physical properties that might seem to affect vapor pressure but do not, once the system has reached dynamic equilibrium. Vapor pressure is an intensive property, meaning it is independent of the amount of substance present. Specifically, neither the surface area of the liquid nor the total volume of the liquid affects the vapor pressure.
While increasing the surface area of a liquid does increase the rate at which evaporation occurs, it also increases the surface area for condensation. In a closed system, a larger surface area simply allows the liquid and vapor to reach the equilibrium state faster, but the final equilibrium pressure itself remains unchanged. Similarly, changing the total volume of the liquid does not alter the pressure that the vapor exerts once the balance between evaporation and condensation has been established.