Lewis Structures
High-Yield Summary
- Bond order = number of shared electron pairs: single = 1, double = 2, triple = 3. Higher bond order → shorter bond length, higher bond energy.
- Polarity comes from an electronegativity difference; dipole moment p = qd (Debye units), pointing from + to − charge.
- A coordinate covalent bond forms when one atom donates both bonding electrons (Lewis base → Lewis acid), e.g. NH₃ + H⁺ → NH₄⁺.
- Formal charge = (valence e⁻) − (lone pair e⁻) − ½(bonding e⁻); helps judge which Lewis structure is most reasonable.
- Resonance: multiple valid Lewis structures for one molecule; the real structure is a hybrid, ranked by full octets > minimized formal charge > negative charge on more electronegative atom > equal contribution.
Steps to Draw a Lewis Structure
- 1Sum valence electrons of all atoms (add 1 e⁻ per negative charge, subtract 1 per positive charge for a polyatomic ion).
- 2Determine connectivity — central atom is usually the least electronegative (never H); in oxyacids, H attaches to O, O attaches to the central atom.
- 3Distribute electrons: single bonds first (2 e⁻ each), then complete outer atoms' octets, then place leftovers on the central atom in pairs.
- 4Check the central atom's octet — if short, convert outer atoms' lone pairs into double/triple bonds with the central atom.
- 5Final check: confirm the exact total electron count from step 1 was used.
Dipole Moment & Formal Charge
p = qd | Formal charge = (valence e⁻) − (lone pair e⁻) − ½(bonding e⁻)
- p = Dipole moment (vector, measured in Debye)
- q = Magnitude of the partial charge
- d = Displacement vector between the partial charges
- Worked check: NH₄⁺ nitrogen — 4 bonds, 0 lone pairs → 5 − 0 − ½(8) = +1, matching the ion's overall charge.
Resonance Structure Priority (Most → Least Important)
| Rule | What It Favors |
|---|---|
| 1. Satisfy the octet rule | More atoms with complete octets |
| 2. Minimize formal charges | Small/zero formal charges over large ones |
| 3. Place charges correctly | Negative formal charge on the more electronegative atom |
| 4. Equal contribution | If 1–3 don't distinguish, structures contribute equally (e.g., NO₃⁻'s 3 forms) |
Common MCAT Trap
- Hydrogen can never be the central atom — it can only form one bond.
- In a resonance hybrid, real bonds are NOT literally alternating single/double — every equivalent bond (e.g., each N–O in NO₃⁻) has an identical, intermediate bond order.
- A coordinate covalent bond is chemically identical to a regular covalent bond once formed — the 'coordinate' label only matters for electron bookkeeping (formal charge), not bond strength or behavior.
Quick Recall
How does bond order affect bond length and bond energy?
What are the two most common examples of coordinate covalent bond formation?
In NO₃⁻, why do all three resonance structures contribute equally?