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Lewis Structures

Lewis Structures

8 min read Updated Mar 26, 2026

Imagine you are an architect sketching the floor plan of a building. You do not need to show every nail and wire - just the walls, doors, and rooms so someone can understand the layout at a glance. A Lewis structure does the same thing for a molecule. It is a simplified blueprint that shows which atoms are connected, where the bonding electrons sit, and where the lone pairs hide.

Learning to draw Lewis structures quickly and accurately is one of the highest-return skills for the MCAT. Once you have a correct Lewis structure, you can predict molecular geometry, polarity, formal charge, resonance, hybridization, and reactivity - all from a single diagram.

The Octet Rule

Most atoms “want” to have 8 electrons in their valence shell. This is because a filled valence shell (like the noble gases) represents maximum stability. Hydrogen is the exception - it is satisfied with just 2 electrons, matching the electron configuration of helium.

The octet rule is the driving force behind Lewis structures. Every step of the drawing process is designed to give as many atoms as possible a full octet.

Step-by-Step Procedure for Drawing Lewis Structures

Follow these five steps every time. With practice, you will be able to fly through them in under a minute.

Step 1: Count total valence electrons.

Add up the valence electrons for every atom in the molecule. For polyatomic ions, add one electron for each negative charge or subtract one for each positive charge.

Step 2: Draw the skeleton structure.

Place the least electronegative atom in the center (hydrogen and fluorine are always terminal). Connect each outer atom to the central atom with a single bond. Each single bond uses 2 electrons.

Step 3: Subtract bonding electrons.

Take the total from Step 1 and subtract the electrons used in your single bonds. The remaining electrons are what you have left to distribute.

Step 4: Distribute remaining electrons as lone pairs.

Place lone pairs on the outer atoms first, giving each a full octet (or duet for hydrogen). Then place any leftover electrons on the central atom.

Step 5: Check the central atom’s octet.

If the central atom has fewer than 8 electrons, convert one or more lone pairs from an adjacent atom into double or triple bonds until the central atom has a full octet.

Worked Example: Lewis Structure of CO₃²⁻

Let us walk through the carbonate ion step by step.

Step 1: Count valence electrons. Carbon has 4, each oxygen has 6, and the 2- charge adds 2 more: 4 + 3(6) + 2 = 24 electrons.

Step 2: Carbon is the central atom (least electronegative). Draw three C-O single bonds. This uses 6 electrons.

Step 3: Remaining electrons: 24 - 6 = 18 electrons.

Step 4: Distribute lone pairs on the three oxygens first. Each oxygen needs 6 more electrons (3 lone pairs) to complete its octet. That uses 3 x 6 = 18 electrons. All remaining electrons are now placed.

Step 5: Check carbon. Carbon currently has only 6 electrons (3 single bonds = 6 electrons). It needs 2 more. Convert one lone pair from an oxygen into a double bond. Now carbon has 8 electrons (2 single bonds + 1 double bond = 8 electrons), and the oxygen that donated the lone pair still has 8 electrons (1 double bond + 2 lone pairs = 8).

The result: carbon is double-bonded to one oxygen and single-bonded to the other two. Each single-bonded oxygen carries a formal charge of -1.

Formal Charge

Not all valid Lewis structures are equally good. Formal charge tells you how to pick the best one. It measures whether an atom “owns” more or fewer electrons than it brought to the molecule.

Rules for the Best Lewis Structure

When you can draw more than one valid Lewis structure, choose the one that:

  1. Minimizes formal charges. A structure where every atom has a formal charge of 0 is better than one with charges of +1 and -1.

  2. Places negative formal charges on more electronegative atoms. If formal charges are unavoidable, the negative charge should sit on the atom that attracts electrons more strongly (higher electronegativity).

  3. Avoids placing positive formal charges on electronegative atoms. Oxygen or fluorine with a +1 formal charge is a red flag.

Resonance Structures

Sometimes you can draw more than one equally valid Lewis structure for the same molecule. These are resonance structures. In our CO₃²⁻ example, the double bond could be placed on any of the three oxygens, giving three equivalent resonance structures.

The critical point: the real molecule is not any single resonance structure. It is a resonance hybrid - an average of all contributing structures.

Key Properties of Resonance

  • Resonance stabilizes a molecule by delocalizing electrons over a larger area. More resonance structures generally means greater stability.

  • All resonance structures must be valid Lewis structures (correct number of electrons, no exceeded octets for second-period elements).

  • Equivalent resonance structures (like the three structures of CO₃²⁻) contribute equally to the hybrid. Non-equivalent structures contribute unequally, with the more stable structure contributing more.

  • Resonance structures are connected by a double-headed arrow (↔), not an equilibrium arrow (⇌). The double-headed arrow means “these are representations of the same molecule,” not “these interconvert.”

Common MCAT Resonance Examples

  • CO₃²⁻ (carbonate): three equivalent structures, each C-O bond order is 43\frac{4}{3}
  • NO₃⁻ (nitrate): three equivalent structures, identical to carbonate’s pattern
  • O₃ (ozone): two equivalent structures, each O-O bond order is 1.5
  • Benzene (C₆H₆): two equivalent structures with alternating single and double bonds; the hybrid has six identical bonds of order 1.5

Lewis structures: count first, then draw

Bonding
The procedure 1 Count valence electrons add up the group numbers; add one per negative charge, subtract one per positive 2 Work out what a full set needs 8 for every atom except hydrogen, which needs 2 3 Subtract, then halve (needed − available) ÷ 2 is the number of bonds 4 Place the rest as lone pairs most electronegative atom outside, least electronegative in the centre 5 Check formal charges valence − dots − lines. The best structure keeps them closest to zero CO₃²⁻: 24 available, 32 needed, (32 − 24) ÷ 2 = 4 bonds so three C–O single bonds and one C=O double bond Resonance: three ways to place that double bond O O − O − C O − O O − C O − O − O C ↔↔ The ion is the average, not the alternation All three C–O bonds are identical in length, between a single and a double, and each oxygen carries two thirds of a negative charge. When the octet rule genuinely does not apply Incomplete octet BF₃ 6 around boron boron and beryllium are stable with fewer than eight; boron routinely stops at six Expanded octet SF₆, PCl₅ 12 and 10 period 3 and below have empty d orbitals to pu t the extra pairs in, so they can exceed eight Odd electron count NO, NO₂ an unpaired electron an odd total cannot pair up completely; these are radicals and are correspondingly reactive F F F B S N O unpaired
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Formal charge, in one lineFormal charge = valence electrons − lone-pair electrons − number of bonds. Work it out for every atom; the sum must equal the overall charge on the species. When two structures both obey the octet rule, the better one is the one with formal charges nearest zero, and with any negative charge sitting on the most electronegative atom.

Resonance is not flippingThe three carbonate structures are not states the ion moves between. The real ion is a single average of all three, which is why all three C–O bonds are identical in length — longer than a double bond, shorter than a single — and why the charge is spread evenly over all three oxygens.

Which elements may expandOnly period 3 and beyond. Nitrogen cannot have five bonds and oxygen cannot have three in a neutral structure, however tempting it looks; phosphorus and sulfur, one row down, can. That single restriction resolves most disputed Lewis structures on the exam.

The count comes first, the picture second. Available minus needed, halved, gives the bonds; everything left over is lone pairs. Draw it that way and the octet rule is arithmetic rather than a guess.
What is the formal charge on the double-bonded oxygen in one resonance structure of CO₃²⁻?
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0. Using FC = Valence - Dots - Sticks: the double-bonded oxygen has 6 valence electrons, 4 lone pair electrons, and 2 bonds. FC = 6 - 4 - 2 = 0. The two single-bonded oxygens each have FC = 6 - 6 - 1 = -1, and the carbon has FC = 4 - 0 - 4 = 0. The total is 0 + (-1) + (-1) + 0 = -2, matching the ion’s charge.

How do resonance structures differ from one another?
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Only in the placement of electrons (bonds and lone pairs). The atoms remain in the same positions. If atom positions change, the structures are constitutional isomers, not resonance structures. The real molecule is a hybrid (average) of all valid resonance structures.

You draw two Lewis structures for a molecule. In Structure A, all atoms have formal charges of 0. In Structure B, one atom has +1 and another has -1. Which is the better (more stable) Lewis structure?
Click to reveal answer

Structure A. The best Lewis structure minimizes formal charges. A structure with all formal charges of zero is more stable than one with separated charges (+1 and -1), even if both structures are technically valid.

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