Grade12 AP Chemistry

Tuesday Feb. 21, 2007

We have been drawing Lewis structures of molecules on a 2-dimensional medium (paper).  In these
structures, molecules with 4 bonds are shown in a square configuration with 90o angles
between them.  There are several reasons why we think that this is not a realistic configuration.
  1) From the side, you can see that in a flat molecule there is a lot of wasted space above and 
     below the molecule (left-hand 2 structures).  A 3-dimensional arrangement of bonds around 
     the central atom results in a better use of space and a greater angle of separation (right-hand
     3 structures).
     

  2) Molecules such as CH2F2, if they were flat squares would have two
     possible configurations.  Since C-F bonds are polar, the upper right-hand structure would be 
     polar since the C-F bonds are not opposite each other and cannot cancel each other out.  The 
     upper left-hand structure would not be polar, however, since the C-F bonds are opposite each 
     other and can cancel each other out.
     

     However, when we examine Carbon dihydride difluoride, we see that all molecules are polar.
     So, the flat configuration cannot explain the properties of the molecule.  In a 3-dimensional
     molecule, such as is shown in the lower part of the diagram above, there are no two bonds 
     which are opposite each other, so no bonds can cancel each other out if the bonds are not 
     all the same.

VSEPR (Valence Shell Electron Pair Repulsion Theory
The shape of a molecule (the angles between bonds) is determined by the mutual repulsion of
electron groups (usually pairs, but more in the case of multiple bonds) for each other.  The
VSEPR theory is based on the assumption that electron groups will be most stable when they have 
the maximum possible separation from each other.  The bond angles are described with respect to a 
central atom (sometimes more than one, in which case there will be several sets of bond angles).
 
The arrangement of independent electron groups (bonds or lone pairs) around the central atom 
in a molecule depends on the number of these groups: the more groups, the smaller the angles
of separation.  

The theory can be summarized as follows:
   - Only valence electrons influence the bond angles.
   - Valence electrons are paired.
   - Valence electron groups (lone pairs or bonds) repel each other.
   - The repulsion of electron groups is more or less equal, but there are differences:
     repulsion of 2 lone pairs > repulsion between a lone pair and a bond > repulsion of 2 bonds
   - The molecule shape is defined by the angular separation of electron groups and the 
     number of them that are bonds rather than lone pairs.

There are several families of simple molecule shapes (simple means 1 central atom) depending
on the number of electron groups.

    # e- Groups    Name                 Angles between e- groups (see below)
         2         Linear                  180o
         3         Trigonal Planar         120o
         4         Tetrahedral             109.5o
         5         Trigonal bipyramidal    a mix of 120o and 90o
         6         Octahedral              90o



The actual shape of the molecule depends on how many of the electron groups are bonds
as opposed to lone pairs.  We use a code to specify how many bonds and lone pairs surround the
central atom.

AXnEm is a molecule with n bonds and m lone pairs.

Family             Code   Example     Shape Name
Linear             AX2E0   linear
Trigonal Planar    AX3E0   trigonal planar
                   AX2E1   bent
Tetrahedral        AX4E0   tetrahedral
                   AX3E1   trigonal pyramidal
                   AX2E2   bent
Trigonal           AX5E0   trigonal bipyramidal
Bipyramidal        AX4E1   see saw
                   AX3E2   T-shaped or planar triangular
                   AX2E3   linear
Octahedral         AX6E0   octahedral
                   AX5E1   square pyramid
                   AX4E2   square planar

Note: all AX1En molecules were omitted.  They are diatomic.

To see pictures of each of the above shapes, follow this link

To determine the shape of a molecule, you need to draw the correct Lewis structure, count the 
number of bonds and lone pairs around the central atom, generate the shape code, and match that
code to the correct shape.

The precise shape that results from each shape code results not only from the code, but from the 
relative repulsion of the different electron groups.  For example, the molecule ICl41- could either
be a see-saw shaped molecule or a square planar molecule.  However, the lone pairs are more repulsive,
so the molecule is more stable if the two lone pairs are on opposite sides of the molecule, rather 
than only 90o apart.

VSEPR is a better predictor of molecule shape when the central atom is from the top part of the 
periodic table.  For example, H2O has bond angles of about 109o as predicted
by VSEPR.  H2S, which has the same shape code as water has bond angles of about 92o,
much smaller than predicted.

VSEPR also offers good support for multiple bonds.  C2H4 has a double bond between
the two Carbons.  This gives each Carbon atom three bonds and no lone pairs.  Whe C2H4 is
examined, the bond angles around each Carbon are found to be approximately 120o as predicted
by VSEPR.