Bonding
Lewis Structures
Diagrams showing valence electrons and bonds — the language of molecular structure
A Lewis structure is a 2D diagram showing the arrangement of valence electrons in a molecule — bonding pairs (lines or dots) and lone pairs (dots). Drawn following octet rule (8 electrons around each non-H atom; 2 for H). Predicts molecular shape, polarity, resonance. Introduced by Gilbert N. Lewis (1916). Steps: (1) Count total valence electrons. (2) Place least electronegative atom in center. (3) Connect with single bonds. (4) Add lone pairs to satisfy octets. (5) Use multiple bonds if needed. Foundation for VSEPR and other concepts.
- Introduced byGilbert N. Lewis, 1916
- ShowsValence electrons: bonding pairs + lone pairs
- Octet rule8 electrons around each non-H atom (typical)
- H exception2 electrons (duet)
- NotationLines for bond pairs; dots for lone pairs
- Formal chargeUsed to identify best resonance structure
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Why Lewis structures matter
- Predicting bonding. Single, double, triple bonds.
- Molecular geometry. Combined with VSEPR.
- Resonance. Identifies delocalized electrons.
- Reactivity. Lone pairs are reactive sites.
- Polarity. Dipole moment from structure.
- Education. Foundation chemistry concept.
- Drug design. Visualize functional groups.
Common misconceptions
- Octet rule is universal. Many exceptions.
- H needs 8 electrons. H needs only 2.
- Center atom is always C. Often, but can be N, P, S.
- Multiple bonds use more atoms. Same atoms; just more shared electrons.
- Best structure has all atoms in same plane. Lewis is 2D; geometry is 3D (VSEPR).
- Resonance means oscillating. Static average; not dynamic.
Frequently asked questions
How do you draw a Lewis structure?
Five steps. (1) Sum all valence electrons (group number for main-group; add 1 per negative charge, subtract 1 per positive). (2) Pick least electronegative atom as central (often C, sometimes N or P; H always terminal). (3) Connect surrounding atoms with single bonds (2 e- each). (4) Add remaining electrons as lone pairs to fill octets (start with outer atoms). (5) If central atom lacks octet, convert lone pairs to multiple bonds.
What's the octet rule?
Atoms try to achieve 8 valence electrons (like noble gas Ne, Ar, Kr...). Drives most bonding. H is exception — wants 2 (like He). Beyond row 2 (P, S, Cl...): can have expanded octet (>8) due to available d-orbitals. Be (4 electrons in BeF₂), B (6 electrons in BF₃) are common exceptions.
What's a formal charge?
Hypothetical charge on atom assuming equal bond electron distribution. Formula: FC = valence electrons - lone pair electrons - ½(bonding electrons). Used to evaluate Lewis structures: best structure has lowest formal charges (0 ideal) and any negative charges on most electronegative atoms. Doesn't always match physical charge.
How are multiple bonds drawn?
Double bond: two lines (=) between atoms. Triple bond: three lines (≡). Each line represents 2 shared electrons. CO₂: O=C=O. N₂: N≡N. Multiple bonds form when single bonds don't satisfy octet — share more electrons. Typically more reactive (π electrons exposed).
What's resonance?
When multiple Lewis structures equivalently describe a molecule. Real structure is "average" of resonance forms, with electrons delocalized. Examples: ozone (O₃) has two resonance forms with switching double bond; benzene with alternating double bonds. Real molecule has bonds intermediate in length and strength.
How do you handle ions?
Charged species. For cations: subtract electrons (e.g., NH₄⁺: N(5) + 4×H(1) - 1 = 8 valence e-). Anions: add electrons. Place brackets around the structure with charge outside: [structure]⁺ or [structure]ⁿ⁻. Formal charges on individual atoms shown if helpful.
When does the octet rule fail?
(1) H, Li (duet — 2 electrons). (2) Be (4 electrons in BeCl₂). (3) B, Al (6 electrons in BF₃). (4) Expanded octet for P, S, Cl, etc. (10+ electrons in PCl₅, SF₆). (5) Odd-electron molecules (NO, NO₂ have unpaired electrons). Don't blindly apply octet — recognize when other rules apply.