Introduction

Lewis Structures illustrate the basic concepts of G.N. Lewis's representations of the valence shell electronic structures of atoms and molecules. The premise behind Lewis Structures is the octet rule: that all atoms would like to be surrounded with an octet of electrons. There are, of course, some exceptions: very small atoms (H and B) have less than an octet, and some main group atoms with low energy d orbitals (P, S, Cl, Br, and I) may have more than an octet as central atoms (especially when combined with highly electronegative atoms). Drawing Lewis Structures can be summarized into a series of steps:

Drawing Lewis Structures

Let's look at an example of how this works using a real molecule. Consider the molecule most responsible for the greenhouse effect, carbon dioxide (CO₂).

1. Count all the valence electrons for each atom:

Carbon 1 x 4 valence electrons = 4 electrons
Oxygen 2 x 6 valence electrons = 12 electrons
Total = 16 electrons

2. Determine which atoms are bonded to one another. Generally the least electronegative atom is the central atom. BUT if the only options are between a more electronegative atom and hydrogen, the more electronegative atom will be the central atom (E.g. Water). Hydrogen NEVER makes more than one bond and thus can NEVER be the central atom.

3. Connect each atom with a single pair of electrons or single bond: (This leaves 16 - 4 = 12 electrons)

4. Add electron pairs to outer atoms for octets:

No left over electrons, but central atom doesn’t have octet!

6. Form double bonds to form octet on central atom:

7. Minimize formal charges and look for resonance structures.

For some molecules, more than one structure can be drawn. Note that a Lewis structure for carbon dioxide can be written using a carbon-oxygen single bond on one side and carbon-oxygen triple bond on the other. How can these two possibilities be distinguished? How can the most important structure be chosen, or are they all equally likely? When several structures can be drawn, they are called resonance structures .

Resonance Structures

In resonance structures, all the atoms are in the same relative position to one another, but the distribution of electrons around them is different. To evaluate the importance of each structure, determine the formal charge on each atom.

Formal Charges

Electrons pairs in bonds between atoms are assumed to be split equally between the two atoms. Non-bonding electron pairs are counted as belonging to the atom on which they reside. This may be put into an equation:

or

The resonance structure which is most stable is the one in which:

For the structure with two double bonds, the formal charges can be calculated as follows:

Oxygens: Formal Charge = 6 - (4 + 1/2(4)) = 0

Carbon:Formal Charge = 4 - ( 0+ 1/2(8)) = 0

For the structure with a single and triple bond:

Oxygen (single): Formal Charge = 6 - (6 + 1/2(2)) = -1

Oxygen (triple): Formal Charge = 6 - (4 + 1/2(6)) = +1

Carbon: Formal Charge = 4 - ( 0+ 1/2(8)) = 0

So, while both structures work as Lewis Structures, the one which results in no formal charges for any of the atoms is more stable and thus more likely to exist in nature than one having charges on the two oxygen atoms.

Oxidation Numbers

Formal charges need to be distinguished from oxidation numbers (which can also be determined from Lewis Structures). Oxidation numbers are most often used when referring to compounds which are more ionic in character. When determining oxidation numbers, all the electrons in bonds between atoms are assigned to (“owned by”) the most electronegative element of the two atoms. The non-bonding electrons are assigned to the atom on which they reside. This equation is:

Now, as a contrast, how do oxidation numbers for this molecule differ from the formal charges given above for the most likely double bonded structure above?

Oxygen:  Oxidation Number = 6 - (4 + 4) = -2

Carbon:   Oxidation Number = 4 - (0 + 0) = +4

Once Lewis structures can be drawn successfully, they can be used to predict the electron group geometry, molecular shape and polarity of molecules and ions. For a thorough discussion, refer to Kotz 6th Edition, Chapter 9. In particular, look at the 3-dimensional representations for all the geometries and shapes, in Figures 9.8, 9.9 and 9.11. Polarity of molecules is discussed in Section 9.9. Bring your textbook with you to do this experiment in lab.

Electron Pair Geometry and Molecular Shape

Briefly, the geometry around a central atom is determined by the number of electron groups around it: 2, 3, 4, 5, or 6. Each set has a different name and arrangement in three dimensional space. Electron groups, all being negative, want to be as far away from one another as possible. This is called the valence shell electron pair repulsion theory (VSEPR). The shape of a complex is determined by the geometry, but distinguished from it since shape refers to the arrangement of the atoms in space. The number of non-bonding electron groups determines the shape. Consider the following table.

http://en.wikipedia.org/wiki/VSEPR_theory

For any given electron group arrangement around an atom, its geometry, and shape can be found.