Molecular Geometry and Bonding Theories

Multiple Bonds


Multiple Bonds

The "internuclear axis" is the imaginary axis that passes through the two nuclei in a bond:

The covalent bonds we have been considering so far exhibit bonding orbitals which are symmetrical about the internuclear axis (either an s orbital - which is symmetric in all directions, or a p orbital that is pointing along the bond towards the other atom, or a hybrid orbital that is pointing along the axis towards the other atom)

Bonds in which the electron density is symmetrical about the internuclear axis are termed "sigma" or "s" bonds

In multiple bonds, the bonding orbitals arise from a different type arrangement:

This type of overlap of two p orbitals is called a "pi" or "p" bond. Note that this is a single p bond (which is made up of the overlap of two p orbitals)

In p bonds:


Generally speaking:


C2H4 (ethylene; see structure above)

What about the electron configuration?

Carbon: 1s2 2s2 2p2

Experimentally:

C2H2 (acetylene)

Delocalized Bonding

localized electrons are electrons which are associated completely with the atoms forming the bond in question

In some molecules, particularly with resonance structures, we cannot associate bonding electrons with specific atoms

C6H6 (Benzene)

Benzene has two resonance forms

The apparent hybridization orbital consistent with the geometry would be sp2 (trigonal planar arrangement)

With six p electrons we could form three discrete p bonds

The best model is one in which the p electrons are "smeared" around the ring, and not localized to a particular atom

Benzene is typically drawn in two different ways:


Structure of NO3-

The Lewis structure of NO3- ion suggests that three resonance structures describe the molecular structure

How will this arrangement look as far as the orbital diagrams?

What might we expect for the electron configuration if we just started with the N atom?

Thus, the correct way to determine electron configurations appears to be:


1996 Michael Blaber