12/11/09
Chemical Bonding
Join hands and sing along…..
Review
•  Polar Bonds v. Non-Polar Bonds v. Ionic
Bonds
–  What’s the difference??
•  Valence Electrons
–  What are they?
–  C: 1s22s22p2 Which electrons in carbon are
its valence electrons?
•  Dot Diagramming Valence Electrons
Review
•  What in the world is electronegativity??
–  Describe the trend.
–  Increases across periods, decreases down
groups.
•  Differences in electronegativity can help
us predict how chemicals react and the
type of bond present.
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12/11/09
Review
•  ΔE = 0, no polar character, considered non-polar
covalent
•  0 < ΔE < 1.2, considered non-polar covalent
•  1.2 < ΔE < 1.6, considered polar covalent
•  1.6 < ΔE , considered ionic
•  All are less than or equal to.
Bonding
•  This is all models to describe what we see.
•  The atoms do not know we are talking
about them.
•  As a consequence there are exceptions to
the rules.
•  Learn to understand why exceptions exist
to understand the models better.
Lewis Structures!!
•  This is a method, but not the only way, for
drawing the structure of molecules.
–  Add up the total number of valence electrons.
–  Draw a simple skeleton structure (the least
EN element usually goes in the center).
–  Use two valence electrons to form bonds
–  Add the left over electrons as lone pairs
–  Make sure the octet rule is followed (mind the
exceptions).
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12/11/09
Practice
•  CH4
•  PCl3
Exceptions Time!
•  Polyatomic Ions (NO3-)
–  Don’t forget charges.
–  Put in brackets when finished.
•  Resonance: Same atoms, different
arrangement of electrons.
•  Double and Triple Bonds
Formal Charge
•  This is a tool to help figure out if our
structure is correct.
•  Formal Charge = (Valence e-) – (1/2
bonding electrons) – (lone pair electrons)
•  Should be zero
–  If not, put negative charges on most
electronegative elements.
–  Sum of formal charge = charge on molecule
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12/11/09
Exceptions
•  Remember::
–  Sulfur can bond up to six times (12 electrons)
–  Phosphorus can bond five times (10
electrons)
–  Nitrogen can bond four times (rare)
–  Oxygen can triple bond (only in CO), double
bond, and single bond (uncommon).
Coordinate Compounds
•  Coordinate Compounds contain
coordinate bonds.
–  In a coordinate bond both electrons in the
bond are from only one of the elements.
–  Example, NH4+
•  This happens in polyatomics.
•  Rare in neutral compounds.
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