SAT Chemistry Liquids, Solids, and Phase Changes - Liquids

SAT Chemistry Liquids, Solids, and Phase Changes - Liquids

Importance of Intermolecular Interaction
A liquid can be described as a form of matter that has a definite volume and takes the shape of its container. In a liquid, the volume of the molecules and the intermolecular forces between them are much more important than in a gas. When you conside r that in a gas the molecules constitute far less than 1% of the total volume, while in the liquid state the mol­ecules constitute 70% of the total volume, it is clear that in a liquid the forces between mol­ecules are more important. Because of this decreased volume and increased intermolecular intera „don, a liquid expands and contracts only very slightly with a change in temperature and lacks the compressibility typical of gases.

Kinetics of Liquids
Even though the volume of space between molecules has decreased in a liquid and the mutual attraction forces between neighboring molecules can have great effects on the mol¬ecules, they are still in motion. This motion can be verified under a microscope when colloidal particles are suspended in a liquid. The particles’ zigzag path, called Brownian movement, indicates molecular motion and supports the Kinetic-Molecular Theory.

Increases in temperature increase the average kinetic energy of molecules and the rapidity of their movement. This is shown graphically in Figure 22. The molecules in the sample of cold liquid have, on the average, less kinetic energy than those in the warmer sample. Hence, the temperature reading T1 will be less than the temperature reading T2. If a particular mol¬ecule gains enough kinetic energy when it is near the surface of a liquid, it can overcome the attractive forces of the liquid phase and escape into the gaseous phase. This is called a change of phase. When fast-moving molecules with high kinetic energy escape, the average energy of the remaining molecules is lower; hence, the temperature is lowered.

Viscosity is the friction or resistance to motion that exists between the molecules of a liquid when they move past each other. It is logical that the stronger the attraction between the molecules of a liquid, the greater its resistance to flow—and thus the greater its viscosity. The viscosity of a liquid depends on its intermolecular forces. Because hydrogen bonds are such strong intermolecular forces, liquids with hydrogen bonds tend to have high viscosities. Water, for example, is strongly hydrogen bonded and has a relatively high viscosity. You may have noticed how fast liquids with low viscosity, such as alcohol and gasoline, flow.

Surface Tension
Molecules at the surface of a liquid experience attractive forces downward, toward the inside of the liquid, and sideways, along the surface of the liquid. On the other hand, molecules in the center of the liquid experience uniformly distributed attractive forces. This imbalance of forces at the surface of a liquid results in a property called surface tension. The uneven forces make the surface behave as if it had a tight film stretched across it. Depending on the mag¬nitude of the surface tension of the liquid, the film is able to support the weight of a small object, such as a razor blade or a needle. Surface tension also explains the beading of rain¬drops on the shiny surface of a car.

Capillary Action
Capillary action, the attraction of the surface of a liquid to the surface of a solid, is a property closely related to surface tension. A liquid will rise quite high in a very narrow tube if a strong attraction exists between the liquid molecules and the molecules that make up the surface of the tube. This attraction tends to pull the liquid molecules upward along the surface against the pull of gravity. This process continues until the weight of the liquid balances the gravita¬tional force. Capillary action can occur between water molecules and paper fiber, causing the water molecules to rise up the paper. When a water soluble ink is placed on the paper, the ink moves up the paper and separates into its various colored components. This separation occurs because the water and the paper attract the molecules of the ink components differ¬ently. These phenomena are used in the separation process of paper chromatography, as shown in the paper chromatography experiment on page 307. Capillary action is at least partly responsible for the transportation of water from the roots of a plant to its leaves. The same process is responsible for the concave liquid surface, called a meniscus, that forms in a test tube or graduated cylinder. (See the drawing on page 49.)

Figure 23 shows water in a container enclosed by a bell jar. Observation of this closed system would show an initial small drop in the water level, but after some time the level would become constant. The explanation is that, at first, more energetic molecules near the surface are escaping into the gaseous phase faster than some of the gaseous water molecules are returning to the surface and possibly being caught by the attractive forces that will retain them in the liquid phase. After some time the rates of evaporation and condensation equalize. This is known as phase equilibrium.

In a closed system like this, when opposing changes are taking place at equal rates, the system is said to have dynamic equilibrium. At higher temperatures, since the number of molecules at higher energies increases, the number of molecules in the liquid phase will be reduced and the number of molecules in the gaseous phase will be increased. The rates of evaporation and condensation, however, will again become equal.
The behavior of the system described above illustrates what is known as Le Chatelier’s Principle. It is stated as follows: When a system at equilibrium is disturbed by the application of a stress (a change in temperature, pressure, or concentration), it reacts so as to minimize the stress and attain a new equilibrium position.
In the discussion above, if the 20°C system is heated to 30°C, the number of gas molecules will be increased while the number of liquid molecules will be decreased:
Heat + H2O (l) — H2O (g)

The equation shifts to the right (any similar system that is endothermic shifts to the right when temperature is increased) until equilibrium is reestablished at the new temperature.
The molecules in the vapor that are in equilibrium with the liquid at a given temperature exert a constant pressure. This is called the equilibrium vapor pressure at that temperature.

The vapor pressure-temperature relation can be plotted on a graph for a closed system. (See Figure 24.) When a liquid is heated in an open container, the liquid and vapor are not in equilibrium, and the vapor pressure increases until it becomes equal to the pressure above the liquid. At this point the average kinetic energy of the molecules is such that they are rapidly converted from the liquid to the vapor phase within the liquid as well as at the surface. The temperature at which this occurs is known as the boiling point. Notice that in this graph, water’s normal boiling point is at 760 mm Hg pressure and 100°C temperature.
Figure 24. Vapor Pressure-Temperature Relationship for Carbon Tetrachloride and Water

There are conditions for particular substances when it is impossible for the liquid or gaseous phase to exist. Since the kinetic energy of a molecular system is directly proportional to the Kelvin temperature, it is logical to assume that there is a temperature at which the kinetic energy of the molecules is so great that the attractive forces between molecules are insufficient for the liquid phase to remain. The temperature above which the liquid phase of a substance cannot exist is called its critical temperature. Above its critical temperature, no gas can be liquefied regardless of the pressure applied. The minimum pressure required to liquefy a gas at its critical temperature is called its critical pressure.

Whereas particles in gases have the highest degree of disorder, the solid state has the most ordered system. Particles are fixed in rather definite positions and maintain definite shapes. Because of their variation in packing, solids can be divided into three categories: Crystalline
solids have a three-dimensional representation much like a brick wall. They have a regular structure, in which the particles pack in a repeating pattern from one edge of the solid to the other. Amorphous solids (literally, “solids without form”) have a random structure, with little if any long-range order. Polycrystalline solids are an aggregate of a large number of small crystals or grains in which the structure is regular, but the crystals or grains are arranged in a random fashion.

Particles in solids do vibrate in position, however, and may even diffuse through the solid. (Example: Gold clamped to lead shows diffusion of some gold atoms into the lead over long periods of time.) Other solids do not show diffusion because of strong ionic or covalent bonds in network solids. (Examples: NaCl and diamond, respectively.)
When heated at certain pressures, some solids vaporize directly without passing through the liquid phase. This is called sublimation. Solids like solid carbon dioxide and solid iodine exhibit this property because of unusually high vapor pressure.
The temperature at which atomic or molecular vibrations of a solid become so great that the particles break free from fixed positions and begin to slide freely over each other in a liquid state is called the melting point. The amount of energy required at the melting point temperature to cause the change of phase to occur is called the heat of fusion. The amount of this energy depends on the nature of the solid and the type of bonds present.

The simplest way to discuss a phase diagram is by an example, such as Figure 25.
A phase diagram ties together the effects of temperature and also pressure on the phase changes of a substance. In Figure 25 the line BD is essentially the vapor-pressure curve for the liquid phase. Notice that at a pressure of 760 millimeters of mercury (1 atmosphere) the water will boil (change to the vapor phase) at 100°C (point F). However, if the pressure is raised, the boiling point temperature increases; and, if the pressure is less than 760 millimeters, the boiling point decreases along the BD curve down to point B.
Figure 25. Partial Phase Diagram for Water (distorted somewhat to distinguish the triple point from the freezing point)

At 0°C the freezing point of water is found along the line BC at point E for pressure at 1 atmosphere or 760 millimeters. Again, this point is affected by pressure along the line BC so that, if pressure is decreased, the freezing point is slightly higher up to point B or 0.01°C.
Point B represents the point at which the solid, liquid, and vapor phases may all exist at equilibrium. This point is known as the triple point. It is the only temperature and pressure at which three phases of a pure substance can exist in equilibrium with one another in a system containing only the pure substance.



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