Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 110	General Chemistry	Fall 2003
Lecture Notes::Lec 17_6 October	© R. Paselk 2003
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Orbitals and Covalent Bonding

Bonding Review, cont.

Last time we began our review with some basics of Lewis Structures. Let's look at a couple of more examples illustrating what this classical model tells us:

• Covalent Compound Lewis Structure Examples:
  • Ammonium ion
  • Methane
  • Carbon dioxide
  • Carbon monoxide

Electronegativity. Electronegativity is a periodic measure of how electrons are shared by atoms. It enables us to guess the degree of polarity of a bond between two atoms (i.e. how the bonding electrons are shared), from non-polar covalent (equal sharing) to fully ionic bonds (no sharing). Recall that F has the highest electronegativity value for and Cs has the lowest. We have used two common ways of determining EN's:

• Look up values on table. Recall that a difference of 1.7 is a reasonable way to distinguish likely covalent (D EN< 1.7) from ionic (D EN> 1.7) bonds.
  • You should know values for Period 2, Li (1.0) to F (4.0) in steps of 0.5 and Hydrogen (2.1)
• Use the high, low, intermediate approximation. Recall that a combination of hi and lo gives ionic bonding, whereas other combinations are covalent.
  • all metals are low
  • the most electronegative of the non-metals are high (N, O, F, S, Cl, Br, I)
  • other elements are intermediate.

Lewis structures are of course quite limited - they work well only for the representative elements, and even then we have to stretch the concept to accommodate all covalent structures., Thus to follow the "Octet rule" we invented resonance for molecules which don't have enough electrons to give octets even with multiple bonding. Clark's rules can help you determine octet violations:

Clark's Method (abbreviated) for determining bonding in covalent Lewis Structures: Add up all of the valence electrons in the structure (remember to add one electron for each negative charge, or subtract one for each positive charge) If S e- = 6y + 2 where y = # atoms other than H, then octet rule is followed with single bonds only. If S e- < 6y + 2 then probably have multiple bonding with the number of multiple bonds = D/2 (remember a triple bond is 2 multiple bonds!). However, note the exceptions with small atoms (H, Li, Be, and B). If S e- > 6y + 2 then have an expanded valence shell. Note that if D = 2, then pentavalent (10 electrons in the valence shell) , and if D = 4, then hexavalent (12 electrons in the valence shell). If you can draw more than one structure, then chose the most symmetrical. If two or more structures are equally symmetrical, then you probably have resonance and should show all structures connected by double arrows.

• Resonance example (Using Clark's rules to help determine the structure):
  • Carbonate ion - CO₃^2-
    • valence electrons = 4 + 3 (6) + 2 = 24
    • 6y + 2 = 26, but S e^- = 24, therefore expect one multiple bond.
    • LS =
    • However, other equally symmetrical structures are possible, so:

And for p-block elements with available d-shells (Period 3 or greater) we had to invent expanded valence shells:

• Expanded Valence Shell Example:
  • SF₄
  • valence electrons = 6 + 4(7) = 34
  • 6y + 2 = 32, but S e^- = 34, therefore expect expanded valence shell with one extra electron pair.
  • LS =

© R A Paselk

Last modified 7 October 2003