Calculate valence electrons, bonding pairs, and formal charges
Step 1: Add valence electrons from each atom
Step 2: Add electrons for negative charge (subtract for positive)
Example: H₂O = 2(1) + 1(6) = 8 electrons
Example: NH₄⁺ = 1(5) + 4(1) - 1 = 8 electrons
1. Octet Rule: Most atoms want 8 valence electrons (H wants 2)
2. Bonding: Each bond = 2 shared electrons (1 electron from each atom)
3. Lone Pairs: Non-bonding electron pairs on an atom
4. Formal Charge: FC = V - (L + B/2), where V = valence, L = lone pair e⁻, B = bonding e⁻
5. Best Structure: Minimize formal charges, place negative charge on most electronegative atom
Lewis structures (also called Lewis dot structures or electron dot structures) are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist.
These structures help us:
Key Concept:
In Lewis structures, dots represent valence electrons, lines represent bonds (2 electrons each), and the arrangement follows the octet rule (atoms tend to gain, lose, or share electrons to have 8 valence electrons).
Add up valence electrons from all atoms:
For ions: Add electrons for (-) charge, subtract for (+) charge
Place least electronegative atom in center (usually):
Connect each terminal atom to central atom with single bond (2 electrons):
Complete octets for terminal atoms first, then central atom:
If central atom has fewer than 8 electrons:
Calculate formal charge for each atom:
FC = V - (L + B/2)
V = valence electrons, L = lone pair electrons, B = bonding electrons
Draw the Lewis structure for carbon dioxide (CO₂).
Step 1: Count Valence Electrons
C: 4 electrons | O: 6 electrons each
Total = 4 + 2(6) = 16 electrons
Step 2: Arrange Atoms
Carbon is less electronegative, so it's central: O-C-O
Step 3: Draw Single Bonds
Two C-O bonds: 2 × 2 = 4 electrons used
Remaining: 16 - 4 = 12 electrons
Step 4: Complete Octets
Each oxygen gets 6 more electrons (3 lone pairs): 2 × 6 = 12 electrons
But now carbon only has 4 electrons! Need to form double bonds.
Step 5: Form Multiple Bonds
Convert 2 lone pairs from each oxygen into bonding pairs with carbon.
Final structure: O=C=O (two double bonds)
Each oxygen has 2 lone pairs, carbon has 8 electrons in bonds.
Step 6: Check Formal Charges
Carbon: FC = 4 - (0 + 8/2) = 0
Each Oxygen: FC = 6 - (4 + 4/2) = 0
All formal charges are zero ✓
Final Answer:
O=C=O with each oxygen having 2 lone pairs. This structure satisfies the octet rule for all atoms and has zero formal charges on all atoms.
Formal charge is a way to estimate the distribution of electrical charge within a molecule. It helps determine the most stable Lewis structure when multiple possibilities exist.
FC = V - (L + B/2)
V = Number of valence electrons in free atom
L = Number of lone pair electrons
B = Number of bonding electrons
Rules for Best Structure:
| Atom | Valence (V) | Bonds | Lone Pairs | Formal Charge |
|---|---|---|---|---|
| N in NH₄⁺ | 5 | 4 (8 e⁻) | 0 (0 e⁻) | +1 |
| N in NH₃ | 5 | 3 (6 e⁻) | 1 (2 e⁻) | 0 |
| O in H₂O | 6 | 2 (4 e⁻) | 2 (4 e⁻) | 0 |
| C in CO₂ | 4 | 4 (8 e⁻) | 0 (0 e⁻) | 0 |
Some atoms are stable with fewer than 8 electrons:
Molecules with odd number of electrons (radicals):
Period 3+ can have more than 8 electrons (use d orbitals):
Understanding Lewis structures helps pharmaceutical chemists predict how drug molecules will interact with biological targets, design more effective medications, and minimize side effects.
Chemists use Lewis structures to predict reaction mechanisms, identify reactive sites, and design synthesis routes for complex organic molecules.
Lewis structures help predict material properties like conductivity, hardness, and reactivity, guiding the development of polymers, semiconductors, and nanomaterials.
Understanding molecular structure helps predict pollutant behavior, design remediation strategies, and develop green chemistry alternatives.
Lewis structures explain enzyme mechanisms, protein-ligand interactions, and the behavior of biological molecules like ATP, DNA, and amino acids.
Lewis structures are fundamental in chemistry education, helping students visualize bonding and understand molecular properties from first principles.
Ions have different electron counts than neutral molecules.
Correct: NH₄⁺ has 8 electrons (not 9), OH⁻ has 8 electrons (not 7)
Most atoms follow the octet rule. Exceptions are specific cases.
Correct: Form multiple bonds if needed to satisfy octets (e.g., O=C=O)
Hydrogen can only form one bond, so it's always terminal.
Correct: H is always on the outside of the structure
Remember: FC = V - (L + B/2), where B is total bonding electrons, not bonds.
Correct: A double bond = 4 bonding electrons, not 2
Group 1: 1 electron
Group 14: 4 electrons
Group 15: 5 electrons
Group 16: 6 electrons
Group 17: 7 electrons
Group 18: 8 electrons
FC = V - (L + B/2)
V = valence, L = lone pair e⁻, B = bonding e⁻
Single bond: 2 electrons
Double bond: 4 electrons
Triple bond: 6 electrons
Lone pair: 2 electrons
Most atoms: 8 electrons
Hydrogen: 2 electrons
Period 3+: Can exceed 8
Be, B: Often < 8