Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109	General Chemistry	Summer 2002
Lecture Notes:: 3 July	© R. Paselk 2002
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Chemical Bonds, cont.

Covalent Bonds

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

Let's look at the formation of HCl as an example of the creation of a covalent bond:

In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:

• H₂ Æ 2 H^.
• Cl₂ Æ 2 Cl^. These radical then combine to form a bond with these two electrons shared between the two atoms.
• H^. + Cl^. Æ H:Cl (This is not a proper Lewis structure, I have only shown the bonding pair of electrons.)
• Lewis Structure:
• Covalent Compound Lewis Structure Examples:
  • Water
  • Ammonia
  • Ammonium ion
  • Methane
  • Hydrogen sulfide
  • Carbon dioxide
  • Carbon monoxide

Electronegativity

So how do we determine whether two atoms will form an ionic or a covalent bond? Use a new property - electronegativity (EN). Electronegativity is a periodic measure of how electrons are shared by atoms with the highest value for F and the lowest for Cs. There are a couple of ways of determining EN's:

• Look up values on table.
  • You should memorize 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:
  • all metals are low
  • the most electronegative of the non-metals are high (N, O, F, S, Cl, Br, I)
  • other elements are intermediate

Bond Type: So how do we use this to predict whether a bond is covalent or ionic?

• For numbers, use a difference of 1.7 to distinguish ionic (D EN> 1.7) and covalent (D EN< 1.7).
• Easier to use the hi, lo, intermediate system:
  • If combine hi + lo, then ionic
  • Otherwise, covalent (hi + hi, lo + lo, hi + inter., lo + inter., inter. + inter.)
  • Lewis Structure Examples:
    • Barium iodide - lo & hi, thus ionic:
    • Carbon disulfide - intermediate & hi, thus covalent:
      • In this case I knew it would have double bonds because it's "just like" carbon dioxide, which we saw above!
    • Arsenic triiodide - intermediate & hi, thus covalent:
    • Hydrogen selenide - intermediate & intermediate, thus covalent:
    • Oxygen difluoride - hi & hi, thus covalent:

Multiple Bonds & Resonance: Recall we must show an octet (or duet for Period I) in the outer-most shell (valence electrons). When this does not occur with single electron pairs (bonds) between atoms can sometimes make it happen with multiple bonds. You might find "Clark's Method" useful for determining the bonding patterns of various molecules:

• 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!).
    • If S e^- > 6y + 2 then have an expanded 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.
  • Multiple bond examples:
    • CO
      • valence electrons = 4 + 6 = 10
      • 6y + 2 = 14, thus 4 fewer electrons than required for all single bonds, 4/2 = 2 multi-bonds (2 double or 1 triple)
      • LS = :C:::O:
    • CO₂
      • valence electrons = 4 + 2x6 = 16
      • 6y + 2 = 20, thus 4 fewer electrons than required for all single bonds, 4/2 = 2 multi-bonds (2 double or 1 triple)
      • LS: from symmetry C will be central atom, therefore= :O::C::O:

• Resonance example:
  • 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:

• Expanded Valence Shell Example:
  • SF₄
  • valence electrons = 6 + 4(7) = 30
  • 6y + 2 = 32, but S e^- = 30, therefore expect expanded valence shell with one extra electron pair.
  • LS =

 

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Last modified 3 July 2002