Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 107	Fundamentals of Chemistry	Fall 2008
Lecture Notes: 20 November	© R. Paselk 2005
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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 determined 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).

• Heating/Cooling Curves. If we monitor a substance such as water going from the solid to gaseous state we will get a heating curve, as seen below. (Figure 13.18. p 395 in text. Note that a cooling curve will be just the opposite.):

So what's going on? For each phase see a steady, linear increase in temperature with added energy (heat). But there are definite breaks where energy is added (lost on cooling) when a phase change takes place. Let's look in a bit more detail.

• Melting/Freezing
  • freezing/melting point: temperature at which solid and liquid are in equilibrium.
    • heat of fusion/crystallization
  • supercooling
    • Glasses: supercooled "solid" liquids
• Evaporation/Condensation
  • vapor pressure (equilibrium)
    • Heat of vaporization/condensation
      • cooling by evaporation
  • boiling
    • superheating

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© R A Paselk

Last modified 20 November 2008