(VSEPR)

Valence Shell Electron Pair Repulsion theory (VSEPR) is a set of rules whereby the chemist may predict the shape of an isolated molecule. It is based on the premise that groups of electrons surrounding a central atom repel each other, and that to minimize the overall energy of the molecule, these groups of electrons try to get as far apart as possible. Groups of electrons can refer to electrons that participate in a bond (single, double, or triple) to another atom, or to non-bonding electrons (e.g. lone pair electrons).

The ideal electronic symmetry of a molecule consisting of a central atom surrounded by a number of substituents (bonded atoms and non-bonding electrons) is characteristic of the total number of substituents, and is determined solely by geometric considerations -- the substituents are arranged so as to maximize the distances amongst them. VSEPR is useful for predicting the shape of a molecule when there are between 2 and 6 substituents around the central atom (the case of one substituent is not discussed because it is trivial -- the only possible shape for such a molecule is linear). That means that there are only five unique electronic geometries to remember. For each electronic geometry, there may be a number of different molecular geometries (the shape of the molecule when only bonded atoms, not non-bonding electrons are considered). Molecular geometries are really just special cases of the parent electronic geometry -- this will hopefully be evident from the models shown on the pages linked to this one.

Since the molecular geometry is determined by how many bonding and non-bonding electron groups surround the central atom, the first thing one needs to do is count how many of each there are. There is a notation that simplifies this bookkeeping:

AB_{x}E_{y}

The **A** represents the central atom, **B** represents
the electron groups that form bonds to other atoms, and
**E** represents the non-bonding electron groups. The
subscripts, **x** and **y**, indicate how many of each
kind are present.

AB_{4}bonding groups: 4 non-bonding groups: 0 |
AB_{3}Ebonding groups: 3 non-bonding groups: 1 |
AB_{2}E_{2}bonding groups: 2 non-bonding groups: 2 |

CH _{4}methane |
NH _{3}ammonia |
H _{2}Owater |

AB4 AB3E AB2E2 ------- -------- --------- bonding groups: 4 3 2 non-bonding groups: 0 1 2 example compound: CH4 (methane) NH3 (ammonia) H2O (water)

Note that bonding "electron groups" does not necessarily imply single bonds; it can mean double or triple bonds as well:

AB_{2}bonding groups: 2 non-bonding groups: 0 |

CO _{2}Carbon Dioxide |

AB2 ------- bonding groups: 2 non-bonding groups: 0 example compound: C02 (carbon dioxide)

An incidental benefit of using the ABE notation is that it provides a convenient way of remembering the hybridization at the central atom. The total number of substituents (bonding plus non-bonding groups) is equal to the number of atomic orbitals that participate in the hybrid orbital.

Molecule | ABE representation | # of substituents | Hybridization |
---|---|---|---|

CH_{4} |
AB_{4} |
4 | sp^{3} |

NH_{3} |
AB_{3}E |
4 | sp^{3} |

H_{2}O |
AB_{2}E_{2} |
4 | sp^{3} |

CO_{2} |
AB_{2} |
2 | sp |

SF_{6} |
AB_{6} |
6 | sp^{3}d^{2} |

I_{3}^{-} ion |
AB_{2}E_{3} |
5 | sp^{3}d |

Molecule ABE Representation # of Substituents Hybridization CH4 AB4 4 sp3 NH3 AB3E 4 sp3 H2O AB2E2 4 sp3 CO2 AB2 2 sp SF6 AB6 6 sp3d2 (I3)- ion AB2E3 5 sp3d

Displayed in the following table are the five most important electronic symmetries. Each row in the table is linked to a page that shows the different molecular symmetries possible for that electronic symmetry.

Class | Hybridization | Electronic Symmetry |
---|---|---|

AB_{2} |
sp | linear |

AB_{3} |
sp^{2} |
trigonal planar |

AB_{4} |
sp^{3} |
tetrahedral |

AB_{5} |
sp^{3}d |
trigonal bipyramidal |

AB_{6} |
sp^{3}d^{2} |
octahedral |

Class Hybridization Electronic Symmetry ----- ------------- ------------------- AB2 sp linear AB3 sp2 trigonal planar AB4 sp3 tetrahedral AB5 sp3d trigonal bipyramidal AB6 sp3d2 octahedral