Vapor Pressure is the pressure created by the vapor, or gas, of a substance that forms above a liquid or solid of the same substance. All liquids, and even some solids, vaporize continuously. The term vapor pressure usually refers to equilibrium vapor pressure, or the pressure at which the rate that particles (atoms or molecules) leave the substance to form vapor equals the rate that particles reenter the substance from the vapor. Vapor pressure depends only on the temperature and the type of substance. Scientists can measure the temperature and vapor pressure of an unknown substance to help determine what it contains. Scientists usually measure vapor pressure in units of atmospheres (atm), millimeters of mercury (mm Hg), or torr. At 20° C (68° F), water has a vapor pressure of 0.023 atm (17.5 mm Hg). Isopropyl alcohol (rubbing alcohol) has a vapor pressure of 0.043 atm (33 mm Hg) at 20° C. Materials with a higher vapor pressure release more vapor particles and therefore evaporate more quickly than materials with low vapor pressure. A puddle of isopropyl alcohol will disappear more quickly than a water puddle of the same size because the alcohol has a higher vapor pressure than water.
Life is possible on Earth because the vapor pressure of water (at temperatures on Earth) is high enough to allow evaporation, but low enough that water can also exist as a liquid (and a solid). The evaporation of water and the existence of liquid water are both essential to the life processes of plants and animals.
A vapor exerts pressure because vapor particles fly about in random directions and at various speeds in the space above a liquid. If the particles are sealed in a container, they collide with other particles, the walls of the container, and the surface of the liquid. Each collision with the inside wall of the container exerts an outward force, and thousands of collisions per second occur on each square centimeter. The force of these collisions can be measured as pressure.
Most of the fast-moving vapor particles above a liquid simply bounce off the liquid’s surface, but the slower moving particles often rejoin, that is, condense back into, the liquid. As more liquid evaporates into vapor particles, more vapor particles become available to collide with the surface, and the condensation rate increases. After a while, the rate of condensation equals the rate of evaporation and the amounts of liquid and vapor remain constant, or at equilibrium. The pressure caused by the evaporated particles at this equilibrium point, in a sealed container, is the liquid’s equilibrium vapor pressure. Vapor pressure above a liquid or solid depends on how easily particles condense into and evaporate from the liquid or solid’s surface. Vapors, or gases, have high kinetic energy, or energy of motion, meaning their particles move around at high speeds. Liquids have lower kinetic energy than gases, and solids have even lower kinetic energy than liquids. When a liquid or solid particle on a material’s surface gains enough kinetic energy, it separates from the surface and joins the vapor. The amount of kinetic energy it needs to separate depends on the substance. The particles in different substances hold on to each other with different strengths in the liquid phase. This bond between liquid particles arises from forces between the particles. The amount of kinetic energy that the particles need to overcome this bond and evaporate, or the ease with which they evaporate, depends on the strength of their bonds. In most pure liquids and their vapors, these particles are molecules. (Some exceptions include mercury liquid and vapor, which are made of mercury atoms, and gallium liquid and vapor, which are made of gallium atoms.) All molecules in a pure liquid attract and repel each other because of electromagnetic forces, forces of attraction and repulsion between charged objects. Objects with opposite charges attract each other, while objects with like charges repel each other.
The electromagnetic forces between separate molecules are called intermolecular forces. They arise because molecules are made of atoms, which in turn consist of negatively charged electrons around a nucleus of positively charged protons (and in most cases, neutral neutrons). Even though molecules have no net electrical charge, they often have a region of positive charge and a region of negative charge. This imbalance occurs when one atom in a molecule pulls a shared electron more tightly and gains a slight negative charge, while the atom from which the electron is pulled gains a slight positive charge. The more tightly the electron is pulled, the more unbalanced the charges will be at different points on the molecule. Highly unbalanced molecules are called polar molecules. Polar molecules are strongly attracted to each other because the positive end of one molecule attracts the negative end of another molecule.
Intermolecular attractions are generally stronger in liquids with polar molecules, such as water, than in nonpolar liquids, such as ether. In polar liquids, surface molecules are strongly attracted to their neighbors and resist being bumped into the vapor phase. Therefore, liquids with strong intermolecular attractions evaporate more slowly than do those with weak attractive forces. They reach equilibrium with few vapor molecules above the surface, so their equilibrium vapor pressures are lower. Liquids that evaporate easily and rapidly at normal temperatures are called volatile liquids. They have weak intermolecular attractions. Volatile liquids release more particles above the liquid surface than nonvolatile liquids. Therefore, the vapor pressure of a volatile liquid is higher than that for a nonvolatile liquid when both liquids are at the same temperature.
Scientists can measure vapor pressure of a liquid sample by using an open-tube manometer, a U-shaped tube partially filled with liquid mercury. If a liquid is sealed in a container with air above it, it can develop a measurable vapor pressure. To measure this pressure, the scientist connects the manometer to a flask filled with air, such that the pressure inside the apparatus equals the outside air pressure. When these pressures are equal, the mercury reaches the same height on both sides of the tube. The scientist then adds a small amount of the liquid to be measured to the flask. As the liquid evaporates, the pressure inside the sealed flask increases, forcing the mercury downward on the flask’s side of the manometer and upward on the side open to the atmosphere. At equilibrium, the difference in the levels of mercury (in millimeters) measures the increase in pressure produced by the liquid’s vapor, which is the liquid’s equilibrium vapor pressure. One torr of pressure is equal to 1 mm Hg, which means that one torr of pressure increases the level of mercury in one of these devices by one millimeter. Vapor pressure depends only on the substance that is vaporized and its temperature. Higher temperature increases vapor pressure because heat causes particles to move faster and gain more kinetic energy. When they have higher kinetic energy, more of the particles on the liquid’s surface can evaporate. The vapor contains more particles, and those particles move faster and are less likely to condense. This increases the number of collisions the particles have with the walls of the container, and, on average, the collisions exert more force than before.
Therefore, the vapor pressure increases with temperature. For example, the equilibrium vapor pressure of water at 20o C is 17.5 mm Hg, and at 50o C it is 92.5 mm Hg. Increasing the volume of liquid in a container has no effect on vapor pressure. For example, a sealed rectangular can that is taller than it is wide and half-filled with a liquid will develop a vapor pressure. If the cap is removed and more liquid is added to the can, the liquid volume increases and forces vapor from the can into the air. Once the cap is replaced, the amount of vapor above the liquid decreases such that the vapor concentration, or number of particles per unit volume, is the same as before. The number of collisions per unit area does not change. Therefore, the equilibrium vapor pressure remains unchanged. Increasing the volume of vapor above a liquid also has no effect on vapor pressure. If liquid is poured out of a container and the cap is immediately replaced, the vapor in the container will occupy a greater volume than before. As a result, the vapor particles spread farther apart and the rate of collisions decreases, which initially lowers the vapor’s pressure. However, evaporation increases the number of vapor particles until equilibrium is restored, and the liquid’s vapor exerts the same pressure as before. Finally, increasing the liquid’s surface area has no effect on vapor pressure either. For example, if the tall, rectangular can is turned on its side, the surface area of the liquid becomes larger. The increased surface area increases the total rate of evaporation. However, it also increases the rate at which particles condense and return to the liquid state. The larger surface area affects the evaporation and condensation rates equally, and the vapor pressure remains the same. The vapor pressure of a liquid is related to its boiling point. If a liquid is heated to a high enough temperature, particles start vaporizing throughout the liquid and bubbles of vapor form and rise to the surface. This process is called boiling, and the temperature at which it occurs is defined as the liquid’s boiling point. For example, in an open container at sea level, the normal boiling point of water is 100° C (212° F).
For liquids in open containers, scientists define the boiling point as the temperature at which the vapor pressure of the liquid equals the atmospheric pressure. These pressures must be equal because liquid particles below the surface need to overcome both their intermolecular attractions and the external pressure above the liquid, which pushes the particles close together and prevents them from expanding into a vapor bubble. Heat increases the kinetic energy of the liquid particles and therefore the vapor pressure. When the vapor pressure of the liquid equals the external pressure, the particles beneath the surface have enough kinetic energy to separate into a vapor. At this point, the liquid begins to boil. At higher altitudes, such as on the top of mountains, the atmospheric pressure is lower than at sea level. In places such as these, the vapor pressure will equal the atmospheric pressure when the liquid is at a lower temperature. This is why the boiling point for a particular liquid is a lower temperature at higher altitudes.
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