How to draw Lewis structures

Valence electrons and the octet rule

Valence electrons are the electrons in an atom's outermost shell, the ones that take part in bonding. For main-group elements, you can read the number from the periodic table: it matches the last digit of the group number in the 1 to 18 numbering.

Atoms tend to bond so that each one ends up with eight electrons around it, counting both shared and unshared pairs, the same as a noble gas. This is the octet rule. Hydrogen is the main exception: it needs only two electrons, a duet, so it always forms exactly one bond and never sits in the middle of a structure.

• Hydrogen: 1 valence electron.
• Carbon, group 14: 4.
• Nitrogen, group 15: 5.
• Oxygen, group 16: 6.
• Fluorine, chlorine and the other halogens, group 17: 7.

The step method

Follow these steps in order for any small molecule or ion:

• Count the total valence electrons of all the atoms. For a negative ion, add one electron for each negative charge; for a positive ion, subtract one for each positive charge.
• Choose the central atom: usually the one that can form the most bonds and is least electronegative, often the single atom written first, such as C in CO₂. Never hydrogen, and rarely fluorine.
• Draw the skeleton, connecting each outer atom to the centre with a single bond. Each bond uses 2 electrons; subtract them from the total.
• Complete the octets of the outer atoms with lone pairs, except hydrogen, which is already full with one bond.
• Place any electrons still left as lone pairs on the central atom.
• If the central atom has fewer than eight electrons, move a lone pair from an outer atom into a shared position to make a double bond, and repeat if needed to make a triple bond.
• Check: the electrons drawn must equal the total from step 1, and every atom should have an octet, or a duet for hydrogen.

Worked example: water, H₂O

Water is the simplest case with lone pairs on the central atom. Hydrogen cannot be central, so oxygen goes in the middle.

Worked example: ammonia, NH₃

Nitrogen has 5 valence electrons and needs 3 more for an octet, so it typically forms three bonds and keeps one lone pair.

Worked example: carbon dioxide, CO₂

Carbon dioxide shows why step 6 matters. After the first pass, carbon is short of an octet, so lone pairs have to become double bonds.

Formal charge: choosing the best structure

Sometimes more than one structure satisfies the octet rule. Formal charge helps you choose. For each atom, formal charge equals its valence electrons, minus the electrons in its lone pairs, minus the number of bonds it forms (each bond counts once, as half of its two shared electrons).

The best structure has formal charges as close to zero as possible, and any negative formal charge on the more electronegative atom. The formal charges in a molecule add up to zero, and in an ion they add up to the ion's charge.

For CO₂ drawn as O=C=O, each oxygen has 6 − 4 − 2 = 0 and carbon has 4 − 0 − 4 = 0. An alternative with one triple and one single bond also gives every atom an octet, but the triple-bonded oxygen has 6 − 2 − 3 = +1 and the single-bonded oxygen has 6 − 6 − 1 = −1. All zeros beats +1 and −1, so O=C=O is the better structure.

Ions and exceptions to the octet rule

For a polyatomic ion, adjust the electron count for the charge and draw the finished structure in square brackets with the charge outside. The ammonium ion, NH₄⁺, has 5 + 4 − 1 = 8 electrons: four N-H bonds and no lone pairs. The carbonate ion, CO₃²⁻, has 4 + 3 × 6 + 2 = 24 electrons, and its structure has one C=O double bond and two C-O single bonds. The double bond could sit on any of the three oxygens, so carbonate has three equivalent resonance structures; the real ion is a blend of all three, with every C-O bond the same length.

Some molecules break the octet rule:

• Fewer than eight: boron in BF₃ has only 6 electrons around it, and beryllium in BeCl₂ has 4.
• More than eight: atoms in period 3 and below, such as phosphorus in PCl₅ (10 electrons) or sulfur in SF₆ (12), can hold extra electrons around the centre.
• Odd numbers of electrons: nitrogen monoxide, NO, has 11 valence electrons, so one atom must be left with an unpaired electron.

Common mistakes

Check your structure against this list:

• Counting all of an atom's electrons instead of only its valence electrons. Oxygen has 8 electrons but 6 valence electrons.
• Forgetting to add or subtract electrons for an ion's charge.
• Putting hydrogen in the centre or giving it more than one bond.
• Drawing more or fewer electrons than you counted in step 1. Count every bond as 2 electrons and every lone pair as 2.
• Adding extra lone pairs to make an octet instead of forming a double bond, which changes the total electron count.
• Assuming the Lewis structure shows shape. H₂O looks straight on paper but is bent, and NH₃ is a pyramid, because lone pairs push the bonds closer together.

How to practise and check your work

Work through a sequence that adds one new idea at a time: CH₄ (four single bonds), H₂O and NH₃ (lone pairs), CO₂ (double bonds), N₂ (a triple bond, 10 electrons), then ions such as NH₄⁺ and OH⁻.

After each one, run three checks. Count the electrons drawn and compare with your step 1 total. Count the electrons around every atom: 8, or 2 for hydrogen. Work out the formal charges and make sure they add up to the overall charge. If all three pass, your structure is almost certainly right.