Bonding in Organic Compounds

There are several fundamental ways to describe bonding. We will talk about the Lewis formulation, with its various patches, and molecular orbital theory on a very qualitative basis. We will also discuss, briefly, molecular mechanics, which is a parameterized theoretical approach, based on the characteristics of molecules that we can measure.

We will use as much as we need, but usually not more. There is an increased "expense" in using higher levels of theory (usually more time). We'll work a lot at the Lewis, or valence bond level, because it will help us understand reactions, which involve rearrangements of bonding.

Properties of Carbon Atoms

We will start with a discussion of the measured properties of carbons in typical organic molecules. While this may seem odd in a discussion of theory, you should remember that the theories are developed to explain reality, not the other way around. When a theory conflicts with a well-designed and executed experiment, the theory loses.

Below is a table of properties which are generally predicted for carbons bonded to 4, 3, and 2 substituents. One should note that significant strain in the molecule can alter these values quite a bit, but that these are pretty accurate for unstrained molecules.

Property for Carbon

4 substituents

3 substituents

2 substituents

Units

Number of Bonds

4

4

4

C-C bond length

154

133

120

pm (0.01Å)

C-H bond length

110

108

106

pm

C-C bond strength

88

152

200

kcal/mol

C-H bond strength

98

103

125

kcal/mol

Geometry

Tetrahedral

Trigonal planar

Linear

Bond Angles

109

120

180

°

Hybridization of Carbon

sp3

sp2

sp


Valence Bond Theory

This is also loosely refered to as Lewis bonding analysis. The advantages are that it is very quick, and helps us focus our attention properly on electrons. It can be viewed as a bookkeeping device that accounts quickly but approximately for where all of the electrons are. The Lewis formulation allows the calculation of overall charges on molecules, assigning full "formal charges" to each atom.

The disadvantages are that it does not tell us (without help) about the polarity of bonds, shapes (aside from simple VSEPR guesses), flexibilities, energies, etc.

These disadvantages led to several "patches" to the theory. We will discuss these in the order: Pauling electronegativity, resonance, hybridization.

Formal Charge

Formal Charge = (# e- required) - (# e- actually there)

Formal charge = (# valence electrons) &emdash; Electron Count

or = (# e- in free atom) - (# e- in bonds)/2 - (# non-bonded e-)

 

Polar Bonding

Pauling added to Lewis bonding the concept of unequal sharing of electrons, made specific by his scale of electronegativities. The degree of polarity can be determined by examining the numerical difference in electronegativities between the two bonded elements. While the choice of cutoff points is somewhat arbitrary, the following works reasonably well for assigning polarity to bonds.

The dipole moment of the bond (its polarity) is by convention drawn with an arrow from atom of lower electronegativity to higher. Physically, the atom of higher electronegativity will have an overabundance of electron density, leading to a small negative charge. The other atom will have a deficiency of electron density, leading to a positive charge. While one can calculate the exact charge difference for simple two-atom systems, most organic molecules have a more complex interrelationship that defies easy calculation simply from electronegativities.

In practice, you will want to be able to look at a bond and, with a table of electronegativities, tell whether or not it is polar. You should also be able to compare two bonds and predict which is more polar.

By the way, Pauling was also responsible for the concept of "covalent radius" which helps us know how big atoms are in molecules. These radii are used in the VSEPR analysis of molecules.

Resonance

Resonance helps us to address the sharing of pairs of electrons among more than two atoms. This is really a kluge, which we retain because it has surprisingly good predictive effects, if used in simple molecular systems.

The general idea of resonance came from the realization that the valence bond representations we write are often wrong to varying degrees. The part that is wrong is the requirement that each bond involves only two atoms and two electrons. Resonance is a formalized way to show the interaction of a pair of electrons with more than two atoms. (It also works with single electrons.)

Resonance works by starting with a valid Lewis structure of a molecule. Electrons are moved, usually a pair or two at a time, to generate new structures. Reasonable structures are mentally combined to give a composite picture which is usually more accurate than the starting structure. This structure is called the "resonance hybrid."

When working resonance examples, remember some simple rules: don't allow atoms to move (the molecule is being viewed as a stationary structure); move electrons one or two at a time, refiguring the formal charge each time; and evaluate each structure you generate for stability. A page of rules with examples will help you evaluate the structures.

Molecular Orbital Approaches

Molecular orbital approaches to describing bonding are discussed on another page. Topics include: