2.6 Drawing Resonance Forms

Resonance Structures

Resonance structures show different ways to arrange electrons in a molecule without moving any atoms. The real molecule doesn't flip between these forms. Instead, it exists as a resonance hybrid, a blend of all valid resonance structures at once. Understanding resonance helps you predict stability, charge distribution, and reactivity across organic chemistry.

Resonance in Three-Atom Groupings

Three-atom systems are the simplest place to see resonance in action. In each case below, electrons (or charge) spread out across the terminal atoms through the central atom.

• Allyl cation ($CH_2=CH-CH_2^+$)
  • Two equivalent resonance forms exist. In one, the double bond sits between C1–C2 and the positive charge is on C3. In the other, the double bond is between C2–C3 and the positive charge is on C1.
  • The real structure is a hybrid: both terminal carbons carry a partial positive charge, and both C–C bonds have partial double bond character.
• Allyl anion ($CH_2=CH-CH_2^-$)
  • Two equivalent resonance forms, with the lone pair and negative charge alternating between the two terminal carbons.
  • The hybrid has partial negative charges on both terminal carbons and equal partial double bond character in both C–C bonds.
• Azide anion ($N_3^-$)
  • Works the same way as the allyl systems, but with nitrogen atoms. Two equivalent resonance forms place the negative charge on alternating terminal nitrogens.
  • The hybrid has partial negative charges on both terminal nitrogens and partial double bond character in both N–N bonds.

The pattern here: whenever you have three atoms in a row with a charge or lone pair on one end and a double bond on the other, you can "push" electrons across to draw a second resonance form.

Multiple Resonance Forms of Molecules

Some molecules have more than two resonance contributors because electrons can delocalize across a longer chain.

• 2,4-Pentanedione anion has three major resonance forms. The negative charge can sit on either terminal oxygen or on the central carbon (C3).
  • When the charge is on a terminal oxygen, that oxygen has a single bond to its adjacent carbon while the other oxygen forms a double bond to its carbon. This is the enolate form.
  • When the charge is on the central carbon, both oxygens form double bonds to their adjacent carbons and the central carbon bears the lone pair. This is the carbanion form.
  • The hybrid spreads the negative charge across all three positions (two oxygens and one carbon), which is why this anion is unusually stable for a carbon-based anion.

Resonance in Ions and Molecules

Resonance shows up constantly in common ions and functional groups. Here are the key examples you should know:

• Phosphate ion ($PO_4^{3-}$): Four equivalent resonance forms. Each form has one P=O double bond and three P–O single bonds (each single-bonded oxygen carries a $-1$ charge). The hybrid distributes the charge equally over all four oxygens.
• Nitrate ion ($NO_3^-$): Three equivalent resonance forms. Each form has one N=O double bond and two N–O single bonds with a negative charge on each single-bonded oxygen. The hybrid gives each N–O bond a bond order of 1.33.
• Carboxylate ion ($RCOO^-$): Two equivalent resonance forms. The negative charge alternates between the two oxygens. The hybrid places half a negative charge on each oxygen, and both C–O bonds are identical (bond order 1.5).
• Amide group ($RCONH_2$): Two major (but non-equivalent) resonance forms.
  • In the first, carbon has a double bond to oxygen (normal carbonyl) and a single bond to nitrogen.
  • In the second, the nitrogen lone pair delocalizes into the carbonyl, giving a C=N double bond and placing a positive charge on nitrogen and a negative charge on oxygen.
  • This partial C=N double bond character is why amide bonds are planar and why rotation around the C–N bond is restricted.
• Benzene ($C_6H_6$): Two equivalent resonance forms, each with three alternating double bonds around the ring. The only difference is which set of three C–C bonds are drawn as doubles. The hybrid has six identical C–C bonds (bond order 1.5) with the pi electrons fully delocalized around the ring.

Drawing and Evaluating Resonance Forms

Follow these steps when drawing resonance structures:

1. Use curved arrows to show where electron pairs move. Arrows always go from an electron source (lone pair or pi bond) toward an electron sink (an atom that can accept electrons).

2. Never move atoms. Only electrons change position between resonance forms. If you've moved an atom, you've drawn a different molecule, not a resonance structure.

3. Calculate formal charges on every atom in each form. Formal charge = (valence electrons) – (lone pair electrons) – (½ bonding electrons). This helps you verify your structure is valid.

4. Evaluate relative stability of non-equivalent resonance forms:

   • Structures with more covalent bonds are generally more stable.
   • Negative charges are more stable on more electronegative atoms (O > N > C).
   • Structures that maintain octets on all atoms (especially C, N, O) are preferred.
   • Minimizing formal charge separation makes a structure more favorable.

5. The resonance hybrid is the real molecule. For bonds that appear as single in some forms and double in others, the actual bond order is somewhere in between.