2.6 Resonance and Formal Charge

Resonance and formal charge help you pick the most accurate Lewis structure when more than one valid version can be drawn. Resonance means the real molecule is a hybrid of equivalent structures, which gives fractional bond orders, while formal charge helps you choose the best structure among nonequivalent options by keeping charges close to zero and placing negative charge on the more electronegative atom when needed. For AP Chemistry, use formal charge to defend the Lewis model you choose.

Resonance and Formal Charge Summary

Resonance is needed when more than one equivalent Lewis structure can be drawn for the same molecule or ion. The real structure is a resonance hybrid, so electrons are delocalized and a single Lewis diagram is not enough to describe the bonding accurately.

Formal charge helps you choose between nonequivalent Lewis diagrams. For AP Chemistry Topic 2.6, use the octet rule and formal charge to select the best model: keep formal charges as close to zero as possible, place negative formal charge on the more electronegative atom when needed, and remember that Lewis structures have limits.

Why This Matters for the AP Chemistry Exam

This topic builds directly on Lewis structures and pushes them further. You need to be able to construct Lewis diagrams and then refine them, either by adding resonance structures or by using formal charge to defend why one structure is better than another. That refinement is the key skill: supporting a claim about molecular structure with evidence from a particle-level model.

Resonance and formal charge also set up what comes next. Once you can draw an accurate structure, you can predict molecular shape, bond angles, and polarity in VSEPR. On the exam you may be asked to draw resonance structures, calculate formal charges, or explain which Lewis diagram is the better model and justify your choice.

Key Takeaways

• A molecule has resonance when you can draw more than one valid equivalent Lewis structure; the real molecule is an average of them.
• Resonance leads to fractional (partial) bond orders, found by dividing the number of bonds by the number of bonding positions.
• Formal charge = (valence electrons) - (lone pair electrons) - (number of bonds).
• The best Lewis structure keeps formal charges as close to zero as possible and places any negative formal charge on the more electronegative atom.
• The sum of all formal charges in a structure must equal the overall charge of the molecule or ion.
• The Lewis model has limits, especially for odd-electron species like NO.

Resonance

Sometimes there is more than one way to represent a molecule with a Lewis structure. When a structure can be drawn in multiple valid equivalent ways, it has resonance.

When you draw out the two or three resonance structures of a molecule, those structures together represent the entire molecule. The actual structure is an average of these structures. This can lead to bond orders that are fractions, such as 1/3 or 2/3.

A common misconception is that a molecule with resonance flips back and forth between different bond arrangements. It does not. Resonance is a way of representing one real molecule whose bonding is spread out, or delocalized, in a way that a single Lewis structure cannot show. The examples below make this clearer.

How do you know when a structure has resonance?

Try drawing the Lewis structure of the polyatomic ion NO3-.

Recall the steps from the Lewis diagrams guide:

• Count the total valence electrons. Nitrogen has 5 and each oxygen has 6: 5 + 6 + 6 + 6 = 23. NO3- carries a -1 charge, which adds one more electron, so the total is 24 valence electrons.
• Place the central atom. Nitrogen goes in the center since there is only one nitrogen atom.
• Surround nitrogen with the three oxygen atoms, connect them with single bonds, and fill in full octets.
• Count again. With all single bonds and full octets you get 26 electrons, which is 2 too many. When you have too many electrons, replace one single bond and one lone pair with a double bond. If you turned all three into double bonds you would not have enough electrons, so you only make one double bond. That single double bond is where resonance comes in.

When drawing the Lewis structure of a polyatomic ion, include the brackets and the charge.

Now the count works out to 24. This is one way to draw NO3-. So which oxygen gets the double bond?

Any of them. Each of the three possibilities exists at the same time in the real ion, so there are three ways to draw it on paper. All three are equally valid. On the AP Exam you would draw all three side by side with a double-headed arrow between them (↔).

In reality

The real ion does not have two single bonds and one double bond. It has a bond order of 4/3 for each nitrogen-oxygen bond. There is no way to draw 4/3 of a bond, which is exactly why you need resonance.

Bond order

How do you find that 4/3 bond order? Divide the number of bonds by the number of bonding positions: 4 lines (bonds) divided by 3 positions = 4/3. A 4/3 bond is between a single and a double bond, so it is stronger than a single bond but weaker than a double bond.

On the AP Exam, draw all valid resonance structures side by side with double-headed arrows between them. You can also write the word "resonance" on the page to make your reasoning clear.

Now try drawing the Lewis structure of nitrite (NO2-), another polyatomic ion. Check that it has resonance too.

Formal Charge

Formal charge is the charge assigned to an atom in a molecule when you assume the electrons in every bond are shared equally. It compares an atom's electron count in the structure to that of the isolated neutral atom, and it helps you decide where electrons should go.

How do you calculate formal charge?

Use this formula:

Formal charge = (valence electrons) - (lone pair electrons) - (number of bonds)

In words, take the atom's number of valence electrons, subtract the lone (nonbonding) electrons, then subtract the number of bonds attached to it.

When should you check formal charge?

Always glance at your formal charges, since it is easy to miscount electrons. Pay special attention when an element from period 3 or below (such as phosphorus or sulfur) is involved, because those atoms can hold more than an octet.

Worked example: phosphate (PO4 3-)

• Count valence electrons: 5 (P) + 6 + 6 + 6 + 6 (four O) + 3 (from the -3 charge) = 32 total.
• Place phosphorus in the center with four oxygen atoms around it, all single bonds, full octets.
• This uses all 32 valence electrons, so it looks complete. But phosphorus is in period 3, so check formal charges before trusting it.
• Formal charges in the all-single-bond version:
  • P: 5 - 0 - 4 bonds = +1
  • Each O: 6 - 6 - 1 bond = -1
• You want the central atom near a formal charge of 0, and any negative formal charge on the most electronegative atom. A formal charge of 0 means the electrons sit in their most stable arrangement.
• Lower phosphorus to 0 by adding one double bond:
  • P: 5 - 0 - 5 bonds = 0
  • The three single-bonded oxygens: 6 - 6 - 1 bond = -1 each
  • The double-bonded oxygen: 6 - 4 - 2 bonds = 0
• Now two atoms sit at a formal charge of 0, which is a more stable representation.

Because the double bond can go to any of the four oxygens, this structure also has resonance, so you would draw all four versions. The phosphorus-oxygen bonds have a bond order of 5/4.

Quick check: the formal charges should add up to the overall charge. Three oxygens at -1 sum to -3, which matches the -3 charge of phosphate.

How to Use This on the AP Chemistry Exam

Free Response

• When a question says some bonds are equal in length, that is a signal for resonance. Draw every valid structure side by side with a double-headed arrow (↔) between them.
• Include brackets and the overall charge for polyatomic ions. Forgetting the charge can cost a point.
• Do not draw an equilibrium arrow (two single half-arrows). Use one full double-headed arrow for resonance.

Problem Solving

• To pick between nonequivalent Lewis structures, calculate formal charges for both and choose the one with charges closest to zero and the negative charge on the more electronegative atom.
• Show your formal charge math when a question asks you to justify a choice. State the formal charge on each relevant atom, then explain which structure is better and why.
• Use the sum check: formal charges must add to the ion's overall charge. If they do not, recount.

Common Trap

• Bond order in a resonance structure is fractional. Find it with (number of bonds) / (number of bonding positions), then describe the bonds as stronger than a single bond and weaker than a double bond, or vice versa.

Common Misconceptions

• Resonance does not mean the molecule switches between forms. The real molecule is a single, averaged structure with delocalized electrons. The separate drawings are just a limitation of the Lewis model.
• Formal charge is not the same as real charge. It is a bookkeeping tool based on equal electron sharing, used to compare possible structures, not a measured charge on the atom.
• A formal charge of zero is the goal, but not the only rule. You also want any negative formal charge on the more electronegative atom. Both criteria matter when choosing the best structure.
• Not every atom must have a full octet. Elements in period 3 and beyond can exceed an octet, and odd-electron species like NO cannot satisfy the octet rule at all. These are real limits of the Lewis model.
• Brackets and charge are part of the answer for ions. Leaving them off an FRQ Lewis structure can lose credit.

AP FRQ Practice Questions

These questions come from previous AP Chemistry exams. You can find full exams on the College Board website.

AP Chemistry Exam 2016 - #2e

The HCO3- ion has three carbon-to-oxygen bonds. Two of the carbon-to-oxygen bonds have the same length and the third carbon-to-oxygen bond is longer than the other two. The hydrogen atom is bonded to one of the oxygen atoms. In the box below, draw a Lewis electron-dot diagram (or diagrams) for the HCO3- ion that is (are) consistent with the given information.

Try it on your own before checking the guidelines below.

AP Chemistry Exam 2017 - #1c

S2Cl2 is a product of the reaction. In the box below, complete the Lewis electron-dot diagram for the S2Cl2 molecule by drawing in all of the electron pairs.

AP Chemistry Exam 2017 - #2a

Two possible Lewis electron-dot diagrams for fulminic acid, HCNO, are shown. Explain why one diagram is the better representation for the bonding in fulminic acid. Justify your choice based on formal charges.

2016 #2e Scoring Guidelines

You are given the ion and its charge and asked to draw the Lewis structure. For HCO3-, count 23 valence electrons plus one from the negative charge: 24 total.

Place carbon in the center as the central atom, attach the oxygen atoms to it, and bond the hydrogen to one of the oxygens.

After drawing the atoms and filling full octets, you count 26 electrons, which is 2 too many. Replace one single bond and lone pair with a double bond. Since two single bonds could become the double bond, this structure has resonance.

Two points are available:

• One point for a correct Lewis structure with brackets and a charge.
• One point for showing resonance by drawing both structures with a double-headed arrow between them.

The hint about two bonds being equal in length and one being longer is the clue that resonance is involved. Note: make sure you use a resonance arrow, not an equilibrium arrow.

2017 #1c Scoring Guidelines

A correct, completed Lewis diagram of S2Cl2 with all electron pairs drawn earns the one available point.

2017 #2a Scoring Guidelines

This question is about calculating formal charges, and the prompt tells you so. The best structure puts the negative formal charge on the more electronegative atom.

This part was worth two points: one for the correct formal charge calculation and one for the correct explanation of which structure is better.

A sample response:

• Point 1: In one diagram the C atom has a formal charge of 0 and the O atom has a formal charge of -1. In the other diagram the C atom has a formal charge of -1 and the O atom has a formal charge of 0.
• Point 2: The better representation puts the negative formal charge on oxygen, which is more electronegative than carbon.

Related AP Chemistry Guides

• 2.1 Types of Chemical Bonds
• Unit 2 Overview: Compound Structure and Properties
• 2.2 Intramolecular Force and Potential Energy
• 2.3 Structure of Ionic Solids
• 2.4 Structure of Metals and Alloys
• 2.7 VSEPR and Bond Hybridization