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Chemical Bonds and the Shape of Molecules
Together with the strength and polarity of its bonds, the shape and size of a molecule largely determine the properties of a substance. Some of the most dramatic examples are seen in biochemical reactions and in substances produced by living species.
The sensations of smell and vision depend in part on molecular architecture. When we inhale, molecules in the air are carried past receptor sites in our nose. If the molecules are the right size and shape, they fit the sites, which transmit impulses to the brain. The brain then identifies these impulses as a particular aroma. Our nose can recognize differences in molecules as small as the different resonance forms of the molecule.
The overall shape of a molecule is determined by its bond angles, the angles made by the lines joining the nuclei of the atoms in the molecule. The bond angles, together with the bond lengths, accurately define the shape and size of the molecule.
The Valence Shell Electron-Pair Repulsion (VSEPR) Model
A bonding pair of electrons defines a region between two atoms in which the electrons will most likely be found. This region is called an electron domain. A nonbonding pair (lone pair) of electrons defines an electron domain located principally on one atom.
In general, each nonbonding pair, single bond, or multiple bond produces an electron domain around the central atom.
Because electron domains are negatively charged, they repel one another. The best arrangement of a given number of electron domains is the one that minimizes the repulsion between them. These arrangements are very predictable and are determined by the number of electron domains surrounding the central atom.

For molecules (and ions) with a central atom (A) surrounded by other atoms (B), ABn, there are five fundamental shapes:

The arrangement of electron domains around the center of an ABn molecule or ion is called its electron-domain geometry.

In contrast, the molecular geometry is the arrangement of only the atoms in a molecule or ion (non-bonding pairs are not part of molecular geometry).

The general steps used in the VSEPR model to predict shapes of molecules or ions are:

  1. Draw the Lewis structure of the molecule or ion, and count the total number of electron domains around the central atom. Each nonbonding electron pair, single bond, double bond, or triple bond counts as an electron domain.
     
  2. Determine the electron-domain geometry by arranging the electron domains around the central atom so the the repulsions among them are minimized.
     
  3. Use the arrangement of the bonded atoms to determine the molecular geometry.
In general, electron domains for nonbonded electron pairs exert greater repulsive forces on adjacent electron domains and thus tend to compress the bond angles.

In general, electron domains for multiple bonds exert a greater repulsive force on adjacent electron domains than do electron domains for single bonds.
These geometries are important because they include all the commonly
occurring shapes found for molecules and ions that obey the octet rule.
Bn
Electron-Domain
Geometry
Bonding
Domains
Nonbonding
Domains
Molecular Geometry
Lewis
Structure
2
2
0
3
3
0
2
1
4
4
0
3
1
2
2

Molecules with Expanded Valence Shells

  The most stable electron-domain geometry for five electron domains is the trigonal bipyramid (two pyramids sharing a base). Two of the five domains point toward what are called axial positions, and the remaining three domains point toward equatorial positions.
Each axial domain makes a 90o angle with any equatorial domain. Each equatorial domain makes a 120o angle with either of the other two equatorial domains.

  The most stable electron-domain geometry for six electron domains is the octahedron. An octahedron is a polyhedron with six vertices and eight faces, each of which is an equilateral triangle. If an atom has six electron domains around it, it can be visualized as being at the center of the octahedron with the electron domains pointing toward the six vertices.
All bond angles are 90o, and all six vertices are equivalent.

Valence-Bond Theory
The VSEPR model provides a simple means for predicting the shapes of molecules. But it does not explain WHY bonds exist between atoms. Combining Lewis' electron-pair bonds with quantum mechanical orbitals, leads to a model of chemical bonding called the valence-bond theory.

The valence-bond theory views the buildup of electron density between two nuclei occurring when a valence orbital of one atom merges with that of another atom. The orbitals then share a region of space, or overlap. This overlap of orbitals allows two electrons with opposite spin to share the common space between the nuclei, forming a covalent bond.

The theory does a good job of explaining covalent bonding, but must also explain the observed geometries of molecules. To do this, it is assumed that the atomic orbitals of an atom mix to form new orbitals called hybrid orbitals. The shape of any hybrid orbital is different from the shapes of the original atomic orbitals. The process of mixing atomic orbitals as atoms approach each other to form bonds is called hybridization.


According to the valence-bond model, a linear arrangement of electron domains implies sp hybridization.

Whenever a certain number of orbitals are mixed, the same number of hybrid orbitals are produced. Each of these hybrid orbitals is equivalent to the others but points in a different direction.

Multiple Bonds
Covalent bonds formed from the overlap of orbitals along a line connecting two atomic nuclei, (the internuclear axis), are called sigma (σ) bonds.

Multiple bonds are formed by a second type of bond, one that results from the overlap between two p orbitals oriented perpendicularly to the internuclear axis. This sideways overlap of p orbitals produces a pi (π) bond. A π bond is a covalent bond in which the overlap regions lie above and below the internuclear axis. Generally, π bonds are weaker than σ bonds.

Multiple (π) bonds require portions of a molecule to be planar, introducing rigidity into molecules.

Multiple bonds are more common in molecules made up of small atoms, especially C, N, and O. Larger atoms, such as S, P, and Si form π bonds less readily.

Delocalized π Bonding
When π or σ electrons are associated totally with the two atoms that form the bond, they are said to be localized. In molecules with π bonds that have two or more resonance structures, (such as benzene), electron pairs are associated with more than two bonding atoms. These electrons are said to be delocalized.

Delocalization of the electrons in its π bonds gives benzene a special stability. Delocalization of π bonds is also responsible for the color of many organic molecules.

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Jim Askew  

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