Network (or covalent) solids - Carbon as example: carbon has allotropes (different physical forms of the same element) all of which have many atoms linked together in covalent networks, and two of which are covalent solids: (Figure 10.22, 24, 24, p 471–4 of Zumdahl)
Diamond (covalent solid) - looked at last time
Graphite (covalent solid) - as in diamond the carbon atoms are covalently linked to each other, but this time in layers, where each carbon is linked to three others with strong covalent bonds in a trigonal planar geometry. The layers are then very weakly held to each other, so they can readily slide over each other, making them "slippery." Note that the sheets are very strong, thus graphite fibers are used to reinforce other materials (graphite fishing poles etc.).
Fullerenes (molecular solids) - graphite like sheets are rolled up to make spheres (commonly 60 C's), and tubes.
Molecular solids: made up of a single type of covalent molecules, the crystal is held together by weak bonds which hold the molecules together. These solids tend to have low melting points.
Ionic solids: made up of two or more types of ions (e.g. Na+ and Cl-) held together by electrostatic attractions. (Figure 10.35, p 481 of Zumdahl). Note that cations (which tend to be small) often fit into the "holes of the unit cell made up the (larger) anions. (Figure 10.33, 10.34 p 480–81 of Zumdahl) Models.
Phase diagrams enable us to predict the behavior of a substance under different conditions of temperature and pressure. We will base our discussions on the behavior of two classic cases: water and carbon dioxide. (Note that phase diagrams strictly describe behavior only for pure substances, so they only hold exactly in closed laboratory systems. The main features do describe much of the contaminated world.)
Water (see Figure 10.47, Zumdahl p 491).
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First note that water is a bit unusual in that there is a negative slope for the transition from solid to liquid. This is a reflection of the fact that solid water is less dense than liquid water (a rare situation). (example: Ice skating)
Note the unique points:
Critical point. Defined by two values:
the Critical Temperature = the temperature above which liquid cannot exist,
the Critical Pressure = the pressure required to give a liquid at the critical temperature. Note that above the critical point liquids no longer have an equilibrium vapor pressure, nor do they boil, rather they undergo a fluid transition to a gas as the temperature is raised.
Triple Point. This is a unique set of values at which the gas, liquid, and solid phases can co-exist in equilibrium. The triple point of water is used as a standard and calibration point for temperature scales because it is precisely defined and reproducible anywhere.
Carbon Dioxide (see Figure 10.50, Zumdahl p 495).
Note that the solid-liquid transition has a positive slope in CO2, which is the common situation for most substances.
Note that the triple point for CO2 is above 1 atm, thus liquid CO2 is not seen in the open on Earth's surface. Rather the solid evaporates directly to a gas (sublimes) at -78 °C at 1 atm.
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© R A Paselk
Last modified 4 May 2015