Solids, cont

Types of Solids

• Metallic solids: these are different in that they have ions at the lattice points in a "sea" of shared electrons (Figures 10.18-.20, p 441-2 of Zumdahl). In metals that are very hard, such as chromium, the ions are also covalently bonded to each other.
• 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, p 444 of Zumdahl)
  • Diamond (covalent solid) - each atom in the bulk solid is covalently bonded to four others in a tetrahedral arrangement. This is the secret to diamonds great hardness: to break a piece off many strong covalent bonds must be broken.
  • 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 10.35-37, p 456-7 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 10.35, 10.37 p 456-7 of Zumdahl)

Phase Diagrams

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 10.49, Zumdahl p 467).

public domain image via Wikipedia Creative Commons

• 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 468).
  

  

  • 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.

Solutions

Solution Concentrations-a Review & Some New Stuff.

Solutions: a solution occurs when one chemical is completely dissolved or dispersed in another. We most commonly think of solutions as being liquid, but solid solutions also occur, such as the various metal alloys like steel, brass and bronze.

In a solution the substance present in highest concentration is considered to be the solvent, while components in lesser amounts are considered to be solutes. If you dissolve a sugar cube in water you get a sugar solution, where water is the solvent, and sugar is the solute.

FYI Example: What is the solvent in 80 proof rum: 80 proof = 40% alcohol in water, so water is the solvent. What is the solvent in 151 proof rum: 151 proof = 75.5% alcohol in water, so alcohol is the solvent.

 Concentration Measures

Concentration Terms:

Percent Concentration

Mass percent

ppt = parts/thousand (1mg/L of water); ppb = parts/billion (1 microgram/L of water)

Volume percent

Molarity: The most commonly used concentration term in chemistry = moles of solute dissolved in 1 L of solution.

Two types of situation arise giving two kinds of problems:

Making molar solutions.

FYI Example: Make up a 1.00000 L solution of 0.25 M NaCl (note that water is the "default" solvent). First weigh out o.25 moles of NaCl = (0.25 mole)(22.99 g + 35.45 g)/mole = 14.61 g Example: What is the concentration of a solution made by dissolving 10.00 g of KI in enough water to make 1.00000 L? First need to find the number of moles of KI: (10.00 g) / ({39.10 g+ 126.9 g}/mole) = 6.135 x 10-2 mole Thus the concentration will be 6.135 x 10-2 M

Dilution problems (see 17 February).

Molality: = moles of solute dissolved in 1 kg of solvent.

Mole fraction: = moles of solute dissolved in total moles of solution = n_a / S n

Example: What is the mole fraction of a solution of 10.0 moles of glycerol dissolved in 15.0 moles of water?

(10 mol) / (10 mol + 15 mol) = 10/25 = 0.400

 Solubility

All gases are completely soluble in each other.

Liquid solutions

Gases decrease in solubility with increasing temperature. [example of oxygen solubility]

Gases increase in solubility with increasing pressure.

Syllabus / Schedule

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Last modified 13 April 2009