Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109	General Chemistry	Summer 2002
Lecture Notes::10 July	© R. Paselk 2002
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Liquids & Solids

Gases vs. Liquids vs. Solids:

• Gases are either monatomic (i.e. inert gases) or covalently bonded molecules.
• Pure liquids at "comfortable" temperatures (20 - 40 °C) all consist of molecules (uncharged) or are metals (Hg, Fr, Cs, Ga, and Rb).
• At higher temperatures (around 300 - 1200 °C) ionic compounds become liquids.
• Melting points for pure metals range up to 3680 K, while the highest melting non-metal is carbon at 3820 K.

These melting points are detemined by the types of forces involved (van der Waals, ionic, metallic, or covalent), and, to a lesser extent by the sizes of the particles.

Liquids: The particles of a liquid are in continuous motion, but the distances between collisions are very short compared to those of gases. Thus liquids are largely incompressible - need to increase pressure about a million-fold to halve volume. Diffusion though liquids is much slower than in gases (hours to days vs. seconds to minutes).

• Evaporation
  • cooling by evaporation
    • number of molecules vs. KE diagram (see Fig 10.39, p476 of Zumdahl 5th ed)
  • vapor pressure (equilibrium)
  • boiling (occurs when the v.p. = atmospheric pressure)
  • superheating (occurs because its hard to initiate bubbles)
• Freezing
  • freezing/melting point: temperature at which solid and liquid are in equilibrium.
  • heat of fusion/crystallization
    • Heating/cooling curves (see Fig 10.42, p490 of Zumdahl 5th ed)
  • supercooling (occurs because its hard to "seed" crystals - will see later when solids are discussed)
    • Glasses: supercooled "solid" liquids
• Viscosity
  • due to van der Waals forces and long molecules
  • due to strong H-bonding (ethylene glycol, HOCH₂CH₂OH has a viscosity much like high MW syrup)
• Surface Tension: due to attraction of molecules of liquid for each other (diagram).
  • meniscus
• detergents/surfactants and wetting (water striders, ducks etc.)

Weak Bonds

Weak bonds range from about 10% as strong as a covalent or ionic bond to <1% as strong. Note the examples in the table below:

Interaction Type Example Average Strength, kcal/mol (kJ/mol) Range** Charge-dipole     1/r2 Dipole-dipole      1/r3 Dipole-induced dipole*   0.1-0.2 (0.4-4)  1/r6 Dispersion* (induced dipole-induced dipole)   0.1-0.2 (0.4-4) 1/r6 Hydrogen bond   3-8 (12-30)    van der Waals repulsion     1/r12 *van der Waals interactions, **from Zubay Biochemistry 3rd. Table 4.3, pg. 89.

The van der Waals bonds are strictly the dipole-induced dipole and dispersion types, but is also often used to refer to other weak bonds other than hydrogen bonds. Notice that the bonds are not only very weak (about 0.1 - 0.3% as strong as a covalent bond), they also do not act at a distance. Essentially they are contact bonds - they sort of act like a weak tape. The corollary is that they increase in importance with increases in molecular size (and thus contact surface).

Thus for hydrocarbons, which are essentially completely non-polar, we see a very low boiling point for methane (CH₄) of - 161°C and a fairly regular increase in boiling point as carbons are added (ethane, C₂H₆ - 88°C; butane, C₄H₁₀ - 0.5°C; hexane, C₆H₁₄ 69°C; octane, C₈H₁₈ 126°C; etc.) until very large molecules such as paraffin (about 100 C's) and polyethylene (>1,000 C's) are essentially non-volatile. Note also though, in these very large molecules the forces holding the substance together have become significant due to the very large contact areas.

Hydrogen bonds are a special case of weak bonds. Note that they are significantly stronger (>100 fold) than the other weak bonds at about 4-10% as strong as a covalent bond. Hydrogen bonds only occur when a hydrogen bound to a small, very electronegative atom is brought close to another small, very electronegative atom. Essentially this means that we only see hydrogen bonds between hydrogens bound to N, O, or F (second Period electronegative elements) and N, O, or F. So we can have O-H O, O-H N, O-H F, N-H O, N-H N etc. hydrogen bonds. This is because hydrogen bonds involve dipole-dipole interactions, but they also have covalent character (about 10% of the sharing we see in true covalent bonds) which requires that the participating atoms be small enough to get close enough to allow such partial sharing. (Grp IVA-VIIA bp plot, Fig 10.4

Hydrogen bonding accounts for much of the special properties of water, such as its very high boiling point (261°C higher than methane with only a 10% increase in MW), high viscosity, high heat capacity etc. which in turn are due to the strong bonds between the individual molecules so they stick together.

Examples of water excluding non-polar substances to force the formation of biomembranes, separate out oils etc.

Equilibrium Vapor Pressure

Occurs when rate of evaporation = rate of condensation. Must have some liquid (or solid for sublimation) present. (Figure 10.37, p 485)

Recall that:

• Equilibrium vapor pressure depends only on T and substance.
• Boiling occurs when vapor pressure = atmospheric pressure.
• Heat of Vaporization is the energy required to convert a mole of liquid to a gas.

Quantitative variation of vapor pressure with temperature:

For the linear plot can find the equation (y = ax + b):

This expression is known as the Clausius-Clapeyron Equation.We can use this equation to find useful information such as the boiling points of liquids at different elevations (and thus pressures).

Example: Find the boiling point of water at 10,000 ft elevation if the atmospheric pressure is 508.4 mmHg. DH_vap = 4.39 x 10⁴ J/mol.

How do we solve this? If we take the difference between the two situations we get:

ln P₁ - ln P₂ = -DH_vap/R (1/T₁ - 1/T₂) + b - b

reaarranging and recalling that log a - log b = loga/b

ln P₁/ P₂ = DH_vap/R (1/T₂ - 1/T₁)

and

1/T₂ - 1/T₁ = (R/DH_vap) (ln P₁/ P₂)

putting in numbers

1/T₂ = (8.315 JK^-1mol^-1/ 4.39 x 10⁴ J/mol) ln (760 mmHg / 508.4 mmHg) + 1/373.15 K

1/T₂ = 7.62 x 10^-5 + 2.68 x10^-3 = 2.76 x 10^-3

T₂ = 362.7 K = 89.7 °C

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