Readings for this section.

Petrucci: Section 10-7

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Experimental evidence clearly shows us that the Lewis Model of molecular
bonding, while having it's merits is far from complete. Take for example the
molecule Chlorofluoromethane (CH_{2}FCl) If we draw
the Lewis Dot Structure for this molecule, we get one of two possibilities:

These structures seem to show that there are two different
versions of this molecule, one in which the chlorine is adjacent to the
fluorine and one where it is across from it. Experimental evidence shows us
that there is only one molecule with the formula CH_{2}FCl, despite
there being two different ways to depict the molecule using only Lewis dot
theory.

Using similar logic, we see that the molecule CHFClBr has two distinct forms and experiment shows them to have different physical properties (optical properties). So, there are clearly examples where the Lewis dot theory breaks down. There must be further theories that can explain these observations.

It turns out that the flat representations produced in the Basic Lewis structures are the problem. Molecules are not generally flat but exist in three dimensions.

The Valence Shell Electron Pair Repulsion Theory (VSEPR), as
it is traditionally called helps us to understand the 3d structure of
molecules. Although we will speak often of electron pairs in this discussion,
the same logic will hold true for single electrons in orbitals, and for double
bonds, where one could think of the bond as consisting of two pairs of
electrons. In general, the region in space occupied by the pair of electrons
can be termed the *domain* of the electron pair. The domain is related
to the orbitals we have discussed earlier (and will elaborate on later) but
the two do not necessarily refer to the same thiing.

The general concept is that the pairs of electrons repel each other and wish to locate themselves as far as possible from each other about a given nucleus. Hence, for two pairs of electrons on a nucleus, the two pairs would locate themselves exactly opposite each other, forming a bond angle of exactly 180?. If three pairs exist, they will locate themselves in a plain about the nucleus at angles of 120? from each other. higher numbers of electrons form 3d arrangements as follows.

Table: Geometry and Electron Pair Arrangements. The angles given are the ideal angles for such an arrangements.

Electron Pairs |
e ^{-} pair
(domain) Geometry |
e ^{-}
pair diagram |
---|---|---|

2 | Linear | |

3 | Trigonal planar | |

4 | Tetrahedral | |

5 | Trigonal bipyramidal | |

6 | Octahedral |

We'll now go through a set of example molecules and/or ions and discuss their geometries. It is important to note that the shape of the molecules as we discuss them here is not always the same as the electron domain geometries described above.

We will consider the molecular shapes, starting with the simplest and working up to the more complicated examples. In addition, I'll mention a classification system which may be helpful in counting electron domains used herein.

- The classification system follows:
- A represents a central atom (any atom being considered. It need not
actually be the centre of the molecule)

X represents an atom bonded to A. It could be single, double or triple bonded. It makes no difference to the scheme.

E represents a non-bonded electron domain (lone pair).

For example, methane CH_{4} is an AX_{4} molecule while
ammonia NH_{3}
is an AX_{3}E molecule. Both of these molecules have four electron
domains and hence would have a tetrahedral domain geometry as listed above.
However, the shape of the molecules are not the same as we will see below.

BeCl_{2} |
AX_{2} |
LINEAR |

This molecule is linear. The Be does not fill its octet shell in this situation. To do so would put a large negative charge on it and a positive charge on the Chlorine atoms. This would simply not happen since Cl is so much more electronegative than Be.

BF_{3} |
AX_{3} |
Trigonal Planar shape |

4 electron pairs

CH_{4} |
AX_{4} |
tetrahedral | ||

NH_{3} |
AX_{3}E |
trigonal pyramidal | ||

H_{2}O |
AX_{2}E_{2} |
bent or angular | ||

HF | AX_{1}E_{3}or just AXE _{3} |
linear |

Note that in all these cases, the electron-domain geometry is
tetrahedral. However, the molecular shape is not always so. In the case of
CH_{4}, the molecule is actually tetrahedral in shape with a
perfect Tetrahedral angle of 109.5?. The next two examples have lone pairs
which occupy a larger domain volume (push more on the bonding pairs) and
reduce the bond angle to less than 109.5?. The last case, HF, is simply a
liner diatomic molecule. There is no bond angle.

PCl_{5} |
AX_{5} |
trigonal bipyramidal | ||

SF_{4} |
AX_{4}E |
See Saw or disphenoidal |
||

ClF_{3} |
AX_{3}E_{2} |
T-shaped | ||

XeF_{2} |
AX_{2}E_{3} |
Linear |

In these cases, the electron-domain geometry is always trigonal bipyramidal. However, only the first molecule is that shape with the ideal angles of 90 and 120 degrees for the axial and equatorial bonds, respectively.

In the case of SF_{4}, there is one lone pair and four
bonding pairs. The lone pair will preferentially locate itself in an
equatorial position since that position has only two other pairs of
electrons within 90 degrees while an axial position would have three.
Thus, the molecular would be see-saw shaped or the more technically
correct name, disphenoidal. The bond angles would be less than the ideal
angles of 90 and 120 degrees.

ClF3 has two lone pairs and they both locate themselves in equatorial positions for the same reasons as described in the previous case. This molecule is T-shaped with bond angles of less than 90 degrees.

SF_{6} |
AX_{6} |
Octahedral | ||

ClF_{5} |
AX_{5}E |
Square Pyramidal | ||

XeF_{4} |
AX_{4}E_{2} |
Square Planar |

In all these cases, the electron-domain geometry is octahedral and
in the case of SF_{6}, so is the shape. The molecule ClF_{5}
has one lone pair and five bonding pairs but since all positions in
the octahedral geometry are equivalent, it doesn't matter which
position the lone pair takes. I drew it on the bottom position here
for visual effect. In the case of XeF_{4}, the two lone pairs
will locate themselves on opposite sides of the square planar
molecule. In the case of the XeF_{4} molecule, the lone pairs
will orient themselves in a square plane and the molecule will be
linear in shape.

Copyright © 1997

Revised: August 29, 2013.