AP Chemistry Lewis Structures

How to draw Lewis structures for molecules and ions, with four worked examples and a diagram.

A Lewis structure (the course calls it a Lewis diagram) is bookkeeping for valence electrons: every electron is either in a bond or in a lone pair, and the total has to match what the atoms bring. Get the bookkeeping right and the geometry, polarity, and formal charge questions that follow in Unit 2 become much easier. This page gives a method that works every time, then runs it on four different cases.

The Six-Step Method for Drawing Any Lewis Structure

Count the total valence electrons. Add every atom's valence electrons; for an ion, add one electron per negative charge or subtract one per positive charge. The group number gives the valence electrons of a main-group element, as in electron configuration.
Arrange the atoms and join them with single bonds. Put the least electronegative atom that can form several bonds in the center (hydrogen never goes there) and link each outer atom to it with one single bond, which uses 2 electrons.
Complete the octets of the outer atoms with lone pairs. Hydrogen needs only 2 electrons.
Put any electrons that remain on the central atom.
If the central atom has fewer than 8 electrons, turn lone pairs from outer atoms into double or triple bonds with it until it does.
Check. The electrons drawn must equal the total from step 1. Then use formal charge to choose between alternatives or to spot resonance.

Every example below uses only octet-rule structures. The Unit 2 review notes that the Lewis model has limitations, such as molecules with fewer or more than 8 electrons around the central atom, so confirm with your teacher how far your course goes into those cases.

Worked Example 1: A Molecule With a Central Carbon (CH₂Cl₂)

Draw the Lewis structure for CH₂Cl₂.

Step 1: C 4 + 2 H (2 × 1) + 2 Cl (2 × 7) = 20 valence electrons.
Step 2: carbon is central, with four single bonds (to 2 H and 2 Cl): 4 × 2 = 8 electrons used, 12 remaining.
Step 3: each Cl already has 2 electrons from its bond and needs 6 more, three lone pairs. Two chlorines use 12 electrons, leaving none.
Steps 4–5: nothing is left, and carbon already has 8 electrons from its four bonds, so no multiple bonds are needed.
Check: 8 in bonds + 12 in lone pairs = 20.

Worked Example 2: Turning Lone Pairs Into a Triple Bond (HCN)

Draw the Lewis structure for HCN, with the atoms bonded in the order H–C–N.

Step 1: H 1 + C 4 + N 5 = 10 valence electrons.
Step 2: single bonds H–C and C–N use 4 electrons, leaving 6.
Step 3: hydrogen is complete with its bond. Nitrogen needs 6 more to reach an octet, so all 6 go on nitrogen as three lone pairs. Nothing is left for carbon.
Step 5: carbon has only 4 electrons (two bonds). Move one lone pair from nitrogen into the C–N bond to make a double bond, giving carbon 6; move a second to make a triple bond, giving carbon 8. Nitrogen keeps 6 + 2 = 8, now with one lone pair.
Check: H–C bond 2 + C≡N bond 6 + one lone pair 2 = 10.

Worked Example 3: A Polyatomic Ion With Resonance (Carbonate)

Draw the Lewis structure for the carbonate ion, CO₃²⁻.

Step 1: C 4 + 3 O (3 × 6) + 2 for the 2− charge = 24 valence electrons.
Step 2: carbon is central with three single bonds: 6 electrons used, 18 remaining.
Step 3: each oxygen needs 6 more electrons, so three lone pairs each, using all 18.
Step 5: carbon has only 6 electrons. Turn one lone pair of one oxygen into a C=O double bond; now carbon has 8.
Step 6: formal charges are 0 for carbon, 0 for the double-bonded oxygen, and −1 for each single-bonded oxygen, which adds up to the 2− charge on the ion.

The double bond could have gone to any of the three oxygens, so there are three equivalent structures. The real ion is a resonance hybrid of them: all three C–O bonds are the same length, between a single and a double bond, and each oxygen carries an equal share of the negative charge. The diagram below shows the finished structures from Examples 1 to 3, with carbonate drawn as one of its three resonance forms.

Worked Example 4: A Double Bond With No Lone Pairs (C₂H₄)

Draw the Lewis structure for ethene, C₂H₄, in which each carbon is bonded to two hydrogens.

Step 1: 2 C (2 × 4) + 4 H (4 × 1) = 12 valence electrons.
Step 2: the two carbons are joined to each other, and each carbon is bonded to two H: five single bonds use 10 electrons, leaving 2.
Steps 3–4: hydrogens are complete. If the last 2 electrons go on one carbon as a lone pair, that carbon reaches 8 electrons but the other has only 6, so the structure is unbalanced.
Step 5: instead, use those 2 electrons as a second shared pair between the carbons. Each carbon now has 4 + 4 = 8 electrons: a C=C double bond and no lone pairs.
Check: 4 C–H bonds (8) + C=C (4) = 12.

This case shows that multiple bonds do not always come from lone pairs sitting on outer atoms. Sometimes the leftover electrons themselves have to be shared. Once you have the structure, the VSEPR chart turns it into a molecular geometry.

Common Lewis Structure Mistakes

Forgetting the ion's charge in the electron count. Add electrons for a negative charge and subtract them for a positive one before drawing anything.
Giving hydrogen more than one bond. Hydrogen is complete with 2 electrons, so it is always on the outside.
Counting electrons per atom instead of in total. Lone pairs and bonds must add up to the single total from step 1, no more and no fewer.
Placing the leftover electrons on the outer atoms and stopping. Check the central atom's octet; if it is short, form a multiple bond.
Drawing only one resonance structure when several equivalent ones exist. Mention that the species is a hybrid and that the bond lengths are equal.
Leaving out brackets and the charge on an ion. A polyatomic ion's structure goes in square brackets with its charge at the upper right, as in the diagram.

Charged polyatomic species are worth practicing alongside this, since each has a name and charge to memorize, and the type of bond a Lewis structure shows is the subject of ionic versus covalent bonding. If you want to see how this unit's topics add up in a score, the AP Chem Score Calculator turns raw practice results into an estimated 1–5.

Frequently Asked Questions

What is the difference between a Lewis structure and a Lewis dot structure?

They are the same idea. AP Chemistry calls them Lewis diagrams. Bonds can be drawn as lines (each line is a shared pair of electrons) and lone pairs as dots, or every electron can be shown as a dot.

How do you count valence electrons for an ion?

Add up the valence electrons of every atom, then add one electron for each unit of negative charge or subtract one for each unit of positive charge. Carbonate, CO3 2−, has 4 + 3(6) + 2 = 24 valence electrons.

How do you know which atom goes in the center?

The central atom is usually the least electronegative atom that can form several bonds, often carbon. Hydrogen and the halogens normally form just one bond, so they sit on the outside. In a formula such as CH2Cl2 or HCN the first-listed non-hydrogen atom is usually central.

When do you use a double or triple bond?

Use one when, after placing all the electrons as lone pairs and single bonds, the central atom still lacks an octet. Move a lone pair from an outer atom into a shared position between the two atoms, and repeat until every atom has its octet (two electrons for hydrogen).

What does resonance mean in a Lewis structure?

When more than one valid Lewis structure can be drawn that differ only in where the double bond sits, the real species is a hybrid of them. In carbonate, all three carbon-oxygen bonds are identical in length, between a single and a double bond.

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