Page 20 - Mechanism and Theory in Organic Chemistry

Molecular Orbitals 9

If the groups are not all equivalent, the angles will deviate from the ideal
values. Thus in NH, (four electron <groups, three in N-H bonds, one an unshared
pair), the unshared pair, being attracted only by the nitrogen nucleus, will be
closer to the nitrogen on the average than will the bonding pairs, which are also
attracted by a hydrogen nucleus. Therefore the repulsion between the unshared
pair and a bonding pair is greater than between two bonding pairs, and the
bonding pairs will be pushed closer to each other. The H-N-H angle should
therefore be less than 109.5". It is found experimentally to be 107". Similarly, in
H,O (four electron groups, two unshared pairs, and two 0-H bonds), the angle
is 104.5".
Ambiguity may arise when more than one structure contributes. Then un-
shared pairs in one structure may become multiple bonds in another, so that the
number of electron groups around a given atom is not the same in both structures.
An example is methyl azide (19). The central nitrogen is clearly linear (two
electron groups), but the nitrogen bonded to CH, has three electron groups in

19a and four in 19b. In such a situation, the number of electron groups is deter-
mined from the structure with the larger number of honds. Thus the nitrogen in
question in 19 is trigonal, not tetrahedral.

Conventions for Structural Formulas
This book contains large numbers of Lewis structural formulas. Frequently we
shall not write out the full Lewis structure; unshared pairs of electrons not shown
explicitly are implied. When there are two or more contributing structures, we
shall show them all only if that is essential to the point being illustrated; again, it
will be assumed that the reader will understand that the missing structures are
implied.

1.2 MOLECULAR ORBITALS

Lewis structures serve admirably for many aspects of mechanistic organic
chemistry. Frequently, however, we need a more accurate bonding model.

Models Based on the Quantum Theory
The description of chemical bonding must ultimately be based on an understand-
ing of the motions of electrons. In order to improve our model, we need to appeal
to the quantum theory, which summarizes the current understanding of the be-
havior of particles of atomic and subatomic size.
The quantum theory provides the mathematical framework for describing
the motions of electrons in molecules. When several electrons are present, all
interacting strongly with each other through their mutual electrostatic repulsion,
the complexity is so great that exact solutions cannot be found. Therefore
approximate methods must be used even for simple molecules. These methods