Lewis or Kekulé Structures.

Various models have been introduced to visualize the formation of molecules. One of the first successful ones was introduced by Lewis and application of this model gives Lewis structures (or, if lines instead of dots are used to indicate electrons, Kekulé structures). Although the ideas of how electrons are shared have changed since Lewis put forward the idea, these structures are simple to remember and to draw and so remain popular, especially in organic chemistry.

Basically, Lewis suggested that an atom having less than eight electrons in its outer shell (or layer - in wave theory, outer shell means largest n value) could share electrons in pairs (or accept or donate electrons) up to a total of eight in the outer shell.

Using the octet theory to draw molecular structures

To draw a Lewis (actually a Kekulé) structure I recommend learning the basic forms for the columns in the periodic table as illustrated below.

In this table

a line pointing away from an atomic symbol represents an electron pair bond; two electrons shared by the atom and another atom.

a line next to the atomic symbol represents a pair of electrons which are associated with that atom but which are not involved in bonding, a so-called lone pair of electrons.

When you draw a structure, each atom in the structure will 'look like' the corresponding atom in this table.

Number of outer electrons	3	4	5	6	7
Bonds / lone pairs	3 / 0	4 / 0	3 / 1	2 / 2	1 / 3
Neutral molecule
Positive ion one less electron.
Negative ion one more electron.

This table illustrates the typical single bonding arrangements in the compounds containing boron, carbon, nitrogen, oxygen, and fluorine. Lone pairs are illustrated by lines rather than two dots. Note especially the neutral species, and then what happens when an electron is added to give the negative ions or subtracted to give the positive ions.

For other elements*, note:

hydrogen forms only one bond and has no lone pairs.

phosphorus(III), P, is like nitrogen.

sulfur(II), S, is like oxygen

halogens(I) are like fluorine.

* The (III), or (II), or (I) indicates the combining power (valence) of the element.

To find the structure of a compound:

a) put together atoms that can form more than one bond first,

b) then add those (such as hydrogen, halogen) which form only one bond

making sure that each atom "looks like" the one in the table.

Examples: sketch the structures of a molecules with the formulae CH₅N, C₂H₆, CH₄O.

Answer: first, connect together atoms which can form more than one bond (C-N; C-C; C-O respectively) - not forgetting the lone pairs; second add the hydrogens as appropriate.

In each of these, check with the table to see that each atom conforms to the listed view.

Double (4-electron) and triple (6-electron) bonds.

Multiple bonded compounds are formed in much the same way, remembering that an atom that can form more than one bond may form more than one bond to an adjacent atom leading to the following possibilities (amongst others):

Note again how, for example, carbon has four electron pair bonds, even if there are two or three bonds to the same neighbouring atom.

More examples:

A page of examples of Kekulé structures of neutral molecules

Kekulé structures for electron rich and electron poor species