Lewis Structures


Lewis structures are a simple way to keep track of bonding electrons, non-bonding electrons, and formal charges in molecules.

To make a Lewis structure, add all valence electrons from the component atoms of a molecule or ion. Arrange these electrons in 2-electron bonds and in non-bonding pairs so that each atom has a filled shell configuration.

You've seen how to draw the Lewis structures of some simple molecules. Let's look at a more difficult example, the compound potassium azide. This is a salt and it can be written KN3 or [K][N3]. By convention, the cation goes before the anion in a formula.

  1. Write out the Lewis structure of each atom or ion separately.

  2. Very electropositive metals usually lose electrons and form cations. (We'll worry about the charges later.)

  3. Sum the number of valence electrons from the atoms in the structure. Add 1 additional electron per negative charge if the structure is an anion. Subtract 1 electron per positive charge if the structure is a cation. This total number of electrons can be used to make 2-electron bonds and lone pairs.

  4. Use the formula of the molecule or ion to help you determine the arrangement of the atoms.

  5. Use valence electrons to make 2-electron bonds to connect the atoms in the structure. You can make single bonds (2 electrons), 2 bonds or a double bonds (4 electrons), 3 bonds or a triple bonds (6 electrons) in some cases. Be sure that you don't exceed the maximum electron count (filled valence shell number) for any atom.

  6. Use the remaining electrons to make lone pairs of electrons on atoms. When there is an odd number of electrons, there will be a single electron on some atom (radical).

  7. Does the number of electrons you have used equal the total valence count?

  8. There may be more than one possible Lewis structure. If so, the one with the fewest formal charges is usually the best. A correct Lewis structure always includes the formal charges.




Formal Charge

Formal charge can show us the electron rich and electron poor regions of a compound. It is an important part of the Lewis structure.

For each atom in the structure:
  1. Sum 1/2 of all electrons in bonds to that atom and add any other non-bonding electrons.

  2. Compare that number to the number of valence electrons of that atom.

    • total = number of valence electrons, formal charge is zero

    • total > number of valence electrons, the atom has a formal negative charge (-1 for every additional electron above the valence number)

    • total < number of valence electrons, the atom has a formal positive charge (+1 for every electron below the valence number)

  3. Use +/- signs and numbers to indicate the formal charge.



Electron Count

The electron count around an atom tells us about its reactivity. If the electron count is less than the number needed to fill the valence shell, the atom is reactive. All atoms will seek a filled valence shell.

For each atom:
  1. Sum the electrons from bonds to that atom (2 electrons per bond) and any other electrons on the atom.

  2. The electron count must be less than or equal to the number of electrons in the filled shell. For n=1, that is 2 electrons; for n=2, that is 8 electrons; and it is 18 for most other atoms.

Note that NO3 is a radical. There is an unpaired electron on oxygen and that oxygen atom has fewer electrons than it need to fill the valence shell. This is key to the chemistry of this molecule.




Professor Patricia Shapley, University of Illinois, 2012