Drawing Lewis Symbols and Lewis Dot Structures
A Lewis symbol shows an element’s valence electrons as dots around its chemical symbol, and a Lewis structure shows how atoms share or transfer those electrons in a molecule or ion, with lines for bonding pairs and dots for lone pairs. In brief: to draw a Lewis structure, count the total valence electrons, place the least electronegative atom in the center (except hydrogen), connect atoms with single bonds, complete the octets of the outer atoms, and put any extra electrons on the central atom, making double or triple bonds if the central atom lacks an octet.
Lewis structures are the simplest way of picturing how atoms are bonded in a molecule and where the lone pairs are. They are the starting point for predicting the shape, polarity and reactivity of molecules. This article explains how to write Lewis symbols, gives a step-by-step method for Lewis structures, applies it to molecules, polyatomic ions and compounds with multiple bonds, includes tables of electron counts, works through examples and gives a short history.
What are Lewis symbols?
A Lewis symbol (or Lewis dot symbol) consists of the element’s symbol, which represents the nucleus and inner electrons, surrounded by dots for the valence electrons. The dots are placed one at a time on the four sides of the symbol (top, right, bottom, left), and then paired. Unpaired dots show how many bonds the atom can form.
| Element | Valence electrons | Unpaired dots | Lone pairs | Usual number of bonds |
|---|---|---|---|---|
| Li | 1 | 1 | 0 | 1 |
| Be | 2 | 2 (or 0 with 1 pair) | 0 | 2 |
| B | 3 | 3 | 0 | 3 |
| C | 4 | 4 | 0 | 4 |
| N | 5 | 3 | 1 | 3 |
| O | 6 | 2 | 2 | 2 |
| F | 7 | 1 | 3 | 1 |
| Ne | 8 | 0 | 4 | 0 |
How do you draw a Lewis structure?
- Count the total valence electrons. Add the valence electrons of all atoms. For an anion, add one electron for each negative charge. For a cation, subtract one for each positive charge.
- Choose the skeleton. Put the least electronegative atom in the center (hydrogen and fluorine are never central). Arrange the other atoms around it. The central atom is often the one that forms the most bonds, such as carbon, nitrogen, sulfur or phosphorus.
- Connect with single bonds. Draw a single line (2 electrons) between each outer atom and the central atom.
- Complete the octets of the outer atoms (the duet for hydrogen) with lone pairs, using the remaining electrons.
- Put any left-over electrons on the central atom.
- If the central atom has fewer than 8 electrons, convert lone pairs from outer atoms into double or triple bonds until it has an octet.
- Check that the total number of electrons is correct and that formal charges are as small as possible (see the next article).
Valence electron counts for common species
| Species | Calculation | Total electrons | Bonding pairs | Lone pairs |
|---|---|---|---|---|
| H2O | 2(1) + 6 | 8 | 2 | 2 |
| NH3 | 5 + 3(1) | 8 | 3 | 1 |
| CH4 | 4 + 4(1) | 8 | 4 | 0 |
| CO2 | 4 + 2(6) | 16 | 4 (two double bonds) | 4 |
| N2 | 2(5) | 10 | 3 (triple bond) | 2 |
| NH4+ | 5 + 4(1) − 1 | 8 | 4 | 0 |
| OH− | 6 + 1 + 1 | 8 | 1 | 3 |
| NO3− | 5 + 3(6) + 1 | 24 | 4 (one double bond) | 8 |
Worked examples
- Water, H2O. Total = 2 × 1 + 6 = 8 electrons. Oxygen is the center, bonded to two hydrogen atoms (4 electrons used). The remaining 4 electrons form two lone pairs on oxygen. Oxygen has 8 electrons (2 bonds + 2 lone pairs), and each hydrogen has 2.
- Carbon dioxide, CO2. Total = 4 + 12 = 16. Skeleton O–C–O uses 4 electrons. Completing the oxygen octets uses 12 more, and the carbon has only 4 electrons. Convert one lone pair from each oxygen into a bonding pair: O=C=O. Each atom now has 8 electrons: carbon has 4 shared pairs, each oxygen has 2 shared pairs and 2 lone pairs. Total electrons: 8 (bonds) + 8 (lone pairs) = 16 ✓.
- Hydrogen cyanide, HCN. Total = 1 + 4 + 5 = 10. Carbon is in the center: H–C–N uses 4 electrons. The nitrogen takes 6 to complete its octet, and carbon has only 4. Two pairs from nitrogen become bonding pairs, giving H–C≡N with one lone pair on N. Check: 2 + 6 + 2 = 10 ✓.
- Ammonium ion, NH4+. Total = 5 + 4 − 1 = 8. Four N–H bonds use all 8 electrons, so nitrogen has no lone pair. The ion is placed in brackets with the charge 1+.
- Nitrate ion, NO3−. Total = 5 + 18 + 1 = 24. Three N–O single bonds use 6, and the oxygen octets use 18 more, which totals 24. But nitrogen has only 6 electrons. Convert one lone pair from one oxygen into an N=O double bond. Then N has 8 electrons. There are three equivalent places for the double bond, which gives resonance structures (see the article on resonance).
- Ethene, C2H4. Total = 2 × 4 + 4 = 12. Skeleton: H2C–CH2 uses 5 bonds (10 electrons), leaving 2 electrons. Each carbon has only 6 electrons, so the 2 remaining electrons go between the carbon atoms as a second bond: H2C=CH2. Each carbon has 4 bonds, and 12 electrons are used.
How do you draw ionic compounds?
For an ionic compound, show the transfer of electrons. Draw the Lewis symbol of each ion in brackets, with the charge. For sodium chloride, Na+ has no dots (its outer shell is empty after losing the 3s electron), and Cl− has 8 dots in brackets: [Na]+ [Cl with 8 dots]−. For magnesium chloride, one Mg2+ with no dots and two chloride ions with 8 dots each: [Mg]2+ 2[Cl with 8 dots]−.
How do you check the structure?
- Count the electrons in the drawing: each line counts as 2, and each dot counts as 1. The sum must equal the total valence electrons.
- Check the octet (duet) for every atom, apart from known exceptions.
- Calculate formal charges and prefer structures with the smallest charges, with negative charges on the more electronegative atoms.
- Check that common atoms form their usual numbers of bonds: H 1, C 4, N 3, O 2, halogens 1.
Who invented the Lewis structure?
Gilbert N. Lewis introduced dot diagrams in 1902 in his notes and published his theory of the shared electron pair in a paper of 1916 titled “The Atom and the Molecule”. The idea was developed in his 1923 book Valence and the Structure of Atoms and Molecules. Lewis was nominated many times for the Nobel Prize but never received it. Lewis diagrams remain a basic tool of teaching because they are easy to draw, even though more complex theories exist.
Which mistakes are common?
Mistake 1: Miscounting the total electrons
Always add or subtract electrons for the charge of an ion.
Mistake 2: Putting hydrogen in the center
Hydrogen forms only one bond and is always on the outside.
Mistake 3: Forgetting lone pairs
Draw all the lone pairs, because they determine the shape of the molecule.
Mistake 4: Giving an atom more than its allowed electrons without a reason
Period 2 atoms never exceed an octet. Only the elements of period 3 and below may have an expanded octet.