Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 107	Fundamentals of Chemistry	Fall 2008
Lecture Notes: 14 October	© R. Paselk 2005
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• Average (mean) = 51.5
• Standard deviation = 10.6
• Range = 8-95 [5 ≥ 90; 11 = 80-89; 14 = 70-79; 23 = 60-69; 22 = 50-59; 13 = 40-49; 14 = 30-39; 5 = 20-29; 1= 10-20; 1 = 0-9]

  85.5	A-
  76	B-
  61.75	C-
  47.5	D

Molecular Geometry & Polarity

Five molecular geometries

Last time we left off with the five molecular geometries possible with an octet of valaence electrons.

Linear

• Carbon monoxide, CO =
• Carbon dioxide, CO₂ = or

Trigonal planar

• Formaldehyde, CH₂O =

Tetrahedral

• 
• Methane, CH₄ = or, another view =

Trigonal pyramidal

• Ammonia, NH₃ = = view from beneath N=

Bent

• Water, H₂O = =

Other geometries occur when there are more than eight electrons (as frequently occurs with transition metals), but we won't worry about them here.

Polarity: So now we can predict bonding and shape in representative group molecules (and thus most biomolecules). How can we predict electron density and thus charge distribution? Need two bits of information:

1. Shape (based on VSEPR Theory)
2. Electron distribution within a bond (based on electronegativity).
   • So what are the considerations for determining charge distribution in a molecule?
     • A bond between two atoms of different electronegativities is polar.
       • The more electronegative atom has a partial negative charge (δ-) while the less electronegative atom has a partial positive charge (δ-+)
       • The two atoms together comprise a dipole.
       • The dipole is symbolized by an arrow with the positive sign (tail) at the electropositive atom (lower EN), and the arrow pointing toward the electronegative atom (higher EN).

Examples:

• Carbon monoxide
• Carbon dioxide
• Water
• Ammonia

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Last modified 14 October 2008