Drawing Complete Lewis Structures
A Lewis structure is a diagram that shows how the valence electrons of a molecule or polyatomic ion are arranged as bonding pairs between atoms and as lone pairs on atoms. In brief: for water, count 2 × 1 + 6 = 8 valence electrons, place oxygen in the center, draw two O–H single bonds (4 electrons) and put the other 4 electrons on oxygen as two lone pairs.
Lewis structures are the starting point for predicting the shape, polarity and reactivity of molecules. Once you know how to draw them, you can use them with VSEPR theory to find molecular geometry. This article gives a step-by-step procedure, shows how to use formal charge to choose between structures, covers resonance and the main exceptions to the octet rule, and works through several examples.
How do you draw a Lewis structure step by step?
- Count the total valence electrons. Add the valence electrons of all atoms. For an anion add one electron per negative charge, and for a cation subtract one electron per positive charge.
- Draw the skeleton. Put the least electronegative atom (never hydrogen) in the center, and connect the other atoms to it with single bonds. Carbon is nearly always a central atom.
- Complete the octets of the outer atoms by adding lone pairs. Hydrogen needs only 2 electrons.
- Place any remaining electrons on the central atom.
- Form multiple bonds if the central atom has fewer than 8 electrons: move a lone pair from an outer atom to make a double or triple bond.
- Check the formal charges and adjust the structure so that they are as close to zero as possible.
How many valence electrons does each atom have?
| Group | Valence electrons | Examples | Usual bonds in neutral molecules |
|---|---|---|---|
| 1 | 1 | H, Na | 1 |
| 2 | 2 | Be, Mg | 2 |
| 13 | 3 | B, Al | 3 |
| 14 | 4 | C, Si | 4 |
| 15 | 5 | N, P | 3 |
| 16 | 6 | O, S | 2 |
| 17 | 7 | F, Cl | 1 |
| 18 | 8 | Ne, Ar | 0 |
Worked examples
Example 1: carbon dioxide, CO2
- Valence electrons: 4 + 2 × 6 = 16.
- Skeleton O–C–O uses 4 electrons; 12 remain.
- Give each oxygen three lone pairs (12 electrons). Now carbon has only 4 electrons around it.
- Move one lone pair from each oxygen to form two C=O double bonds. Carbon now has 8 electrons, and each oxygen has 2 lone pairs plus a double bond.
The result is O=C=O, with every atom having a formal charge of zero.
Example 2: the ammonium ion, NH4+
Valence electrons: 5 + 4 × 1 − 1 = 8. Four N–H bonds use all 8 electrons, and nitrogen has no lone pair. The structure is written in square brackets with the charge +1 outside; the formal charge of +1 is on nitrogen.
Example 3: the nitrate ion, NO3−
- Valence electrons: 5 + 3 × 6 + 1 = 24.
- Three N–O bonds use 6 electrons; 18 remain and are placed as three lone pairs on each oxygen atom, which uses all of them.
- Nitrogen has only 6 electrons, so one oxygen lone pair is moved to form an N=O double bond.
The double bond can be on any of the three oxygens, so three equivalent resonance structures exist.
How does formal charge help?
The formal charge of an atom in a structure is:
FC = (valence electrons) − (nonbonding electrons) − ½ (bonding electrons)
The sum of the formal charges equals the charge on the molecule or ion. The best structure usually has the smallest formal charges, and negative charges on the more electronegative atoms. For example, in nitrate the oxygen with the double bond has FC = 6 − 4 − 2 = 0. Each oxygen with a single bond has FC = 6 − 6 − 1 = −1, and nitrogen has FC = 5 − 0 − 4 = +1. The sum is 0 − 1 − 1 + 1 = −1, which is the charge of the ion.
What is resonance?
When more than one valid Lewis structure can be drawn that differ only in the positions of electrons, the real molecule is a resonance hybrid, an average of the structures. In ozone, O3, both O–O bonds are identical, with a length of 128 pm, between a typical single bond (about 148 pm) and double bond (about 121 pm). In nitrate all three N–O bonds are equal, with a bond order of 4/3. The structures are not forms between which the molecule switches; they are limitations of the notation.
What are the exceptions to the octet rule?
| Type | Description | Examples |
|---|---|---|
| Incomplete octet | Central atom has fewer than 8 electrons | BF3 (6), BeCl2 (4) |
| Odd number of electrons | One unpaired electron (a radical) | NO (11 electrons), NO2 (17) |
| Expanded octet | Central atom of period 3 or higher has more than 8 | PCl5 (10), SF6 (12), XeF4 (12) |
Expanded octets are possible only for elements from the third period onwards. The older explanation, which used d orbitals for bonding, is now regarded as less important than the polarity of the bonds, but the Lewis structures with an expanded octet remain useful for prediction.
Who invented Lewis structures?
Gilbert N. Lewis introduced the idea of a shared electron pair in 1916 in a paper on “The Atom and the Molecule” and had sketched an early “cubical atom” version of the idea in his lecture notes in 1902. Irving Langmuir extended the ideas and in 1919 popularized the octet rule and introduced the term covalence. The structures are now drawn with lines for bonding pairs and dots for lone pairs, a notation that was in use by the 1920s.
Which mistakes are common?
- Miscounting electrons: forgetting to add or subtract electrons for the charge of an ion.
- Putting hydrogen in the center: hydrogen forms only one bond.
- Forgetting lone pairs on the central atom: they determine the molecular shape.
- Showing resonance as an equilibrium: a double-headed arrow is used, not a double arrow.
- Overusing expanded octets: do not expand the octet for elements in the second period.