Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 110	General Chemistry	Summer 2006
Lecture Notes::Lec 11_13 June	© R. Paselk 2006
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Orbitals and Covalent Bonding

Bonding Review

Chemical bonds are the strongest forces that exist between atoms. They are the forces that hold atoms together in molecules and atoms or ions together in solids. We will look at other weak bonds and forces later.

The two most important and common strong bond types in chemistry are ionic bonds and covalent bonds, a third bond type, found in metallic solids, will be discussed later.

Ionic Bonds. An ionic bond is the result of the electrostatic force of attraction between ions that carry opposite electrical charges, as described by Coulomb's Law:

where r is the distance between ion centers in nm.

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:

H₂ + Cl₂ 2 HCl

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 2 H^.
• Cl₂ 2 Cl^. These radical then combine to form a bond with these two electrons shared between the two atoms.
• H^. + Cl^. 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:
  • 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 element 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 (delta EN< 1.7) from ionic (delta 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 = delta/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 delta= 2, then pentavalent (10 electrons in the valence shell) , and if delta= 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 =

VSEPR Theory and Molecular Geometry

Another great limitation of Lewis structures is that they tell us nothing about molecular shape. So to determine shape we added another tool, VSEPR Theory, to our chemical toolbox.

VSEPR (Valence Shell Electron Pair Repulsion) Theory is based on three assumptions:

• Electron pairs will orient around a central point to minimize repulsion.
• Lone-pairs of electrons will have greater repulsion than bonded pairs of electrons (note that the atoms are ignored in terms of repulsion).
• Repulsion is strong at 90° and weaker at 120° (weakest at 180°).

© R A Paselk

Last modified 13 June 2006