Water, cont.

We left off last time compring some of water's properties to other substances:

• extremely high mp and bp (0 °C & 100 °C K, compare to methane, -183 °C & -161 °C, with a MW of 16 vs. water's 18, see figure below or Figure 10.20, p 392 in text.)

• a high heat capacity (18 cal/°C mol vs. 8 cal/°C mol for methane)
• a high viscosity
• its solid form is less dense than the liquid form at the same temperature (ice floats on water - very rare)
• it has a high dielectric constant (78.5 vs. 1.9 for hexane)
• it has hi surface tension
• it has high heats of fusion (melting)/freezing and vaporization/condensation.

The high mp, bp, and heat capacity all predict relatively strong bonding between water molecules, H-bonding. Note environmental consequences - Earth's weather is much more pleasant because it is moderated by water, especially along coasts. Ice floating prevents "solid" seas, definitely a downer in environmental terms.

Water of course is a covalent structure: H-O-H. But what gives it its special properties is the polarity of its O-H bonds and the resultant dipole moments of the bonds and the molecule itself.

The water molecule itself is bent, with an angle of 104.5° between the hydrogens (compare to 109.5° for sp³ tetrahedron) [slide]. Because of the very strong dipole moments of these bonds and the very small size of the hydrogen substituents on water, a slight degree of orbital overlap occurs between adjacent water oxygens and hydrogens to give partial covalent bonds known a H-bonds (effectively, can only form with O, N, & F). Note that the partial covalent character means that they are directional!

Within solid bulk water (ice) every water molecule is bonded to 4 others, as in the ice structure seen in Figure 10.12 on p 381 [note solid vs. liquid images] In liquid water the molecules are still bonded to a large degree (the heat of fusion for ice is only 13% of the heat of vaporization for ice, thus most of the H-bonds must survive melting). Of course in liquid water the bonds are very unstable (average lifetime about 10 psec = 10^-11 sec), exchanging constantly to give a "flickering cluster" structure. The various properties of water arise from this structure. (Note hi bp & mp, heat cap., viscosity, and, less obviously, that ice floats. This is because the molecules are in an open lattice rather than close-packed. G&G note that close-packed molecules would only occupy about 57% of volume. This would lead one to expect that ice would float "high." It doesn't because most of the structure remains in the liquid phase at 0° C.)

Water is also an excellent solvent for polar substances since its dipolar structure enables it to insulate them from each other and it can make good dipole-dipole and dipole-charge bonds. Figure 11.10 on pg. 423 shows the hexavalent liganding of water to sodium and chloride ions to form hydration shells (For sodium ions, the waters in the inner hydration-shell exchange every 2-4 nsec.). Anything which can H-bond will also of course be quite soluble.

Solids and Crystals

Solids: Recall earlier definition - solids have fixed or definite shapes and volumes. By this definition solids are strictly limited to the crystalline solids. (The amorphous (noncrystalline) solids discussed in our text are what we have discussed as supercooled liquids.)

• Properties:
  • Hard and rigid - they have virtually no tendency to flow or diffuse.
  • Nearly incompressible - need to increase pressure about 1,000,000 times to decrease volume by half.
  • Very low thermal coefficients of expansion.
  • Crystal lattice
    • 3 kinds of cubic lattice: simple cubic, body-centered cubic, and face-centered cubic. [slide]
    • For identical spheres can get two kinds of close packing: cubic close packing, which turns out to be a face-centered cubic lattice, and hexagonal close packing.
  • Melting and freezing points are sharp - all units in the interior of a perfect crystal have the same relationships, and therefore the same bonds. Thus when enough energy is added to break the bonds for one unit, there is enough to break bonds with all, so melting is sudden as all the particles break bonds with each other at same temperature and thus same energy.
    • heat of fusion/crystallization (saw last time with liquids)
  • Nearly all solids expand when they melt (after all the particles are moving faster). As a consequence, nearly all solids will sink in their liquid forms (water is of course a notable exception - we'll look at why later).

Types of Solids - Read Text chapter 10.4

• Ionic solids: made up of two or more types of ions (e.g. Na⁺ and Cl^-) held together by electrostatic attractions. [text figure 10.37 p405]
• 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.
• Covalent (or covalent network) 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:
  • 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. [text figure 10.40 p407]
  • Graphite (covalent solid) - as in diamond the carbon atoms are covalently linked to each other, but this time in layers, where each carbon in 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.). [text figure 10.43 p408]
  • Fullerenes (molecular solids) - graphite like sheets are rolled up to make spheres (commonly 60 C's), and tubes.

Solutions

Definitions:

• Solution - a homogeneous mixture of substances.
  • The major component is known as the solvent.
  • The substances dissolved in it (by definition minor components) are solutes.
• Solubility - a measure of how much solute will dissolve in a given solvent under given conditions (default = 20 °C, 1 atm).
  • Liquids which dissolve in each other are called miscible.
  • Liquids which do not dissolve in each other are considered immiscible.

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 22 September).

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