Lewis structures are diagrams that show how valence electrons are arranged in molecules and polyatomic ions. They help chemists predict bonding, lone pairs, molecular shape, polarity, and reactivity. Learning a reliable step-by-step method prevents guessing and makes even unfamiliar formulas easier to draw.
These diagrams are especially useful in chemistry because many properties depend on where electrons are located.
The process starts by counting total valence electrons, choosing a central atom, and connecting atoms with single bonds. Then electrons are placed to complete outer atom octets before any remaining electrons go on the central atom. If the central atom lacks an octet, lone pairs can be converted into double or triple bonds.
For ions and molecules with multiple possible structures, formal charge helps identify the best Lewis structure.
Key Facts
- Total valence electrons = sum of valence electrons from all atoms, adjusted for charge.
- For anions, add electrons equal to the negative charge; for cations, subtract electrons equal to the positive charge.
- A single covalent bond contains 2 shared electrons.
- Most atoms follow the octet rule: stable main-group atoms often have 8 valence electrons.
- Hydrogen follows the duet rule and can hold only 2 electrons.
- Formal charge = valence electrons - nonbonding electrons - 1/2 bonding electrons.
Vocabulary
- Valence electron
- A valence electron is an outer-shell electron that can participate in chemical bonding.
- Lewis structure
- A Lewis structure is a diagram that represents atoms, bonds, and lone pairs using symbols and dots or lines.
- Lone pair
- A lone pair is a pair of valence electrons that belongs to one atom and is not shared in a bond.
- Octet rule
- The octet rule states that many main-group atoms are most stable when surrounded by 8 valence electrons.
- Formal charge
- Formal charge is a bookkeeping value used to compare possible Lewis structures by assigning electrons to atoms.
Common Mistakes to Avoid
- Forgetting to adjust for ion charge, which gives the wrong total electron count. Add electrons for negative charges and subtract electrons for positive charges before drawing bonds.
- Putting hydrogen in the center, which is wrong because hydrogen can form only one bond. Hydrogen should almost always be placed on the outside of a Lewis structure.
- Completing the central atom before outer atoms, which often produces incorrect octets. Fill the octets of outer atoms first, then place any remaining electrons on the central atom.
- Leaving the central atom with too few electrons when a multiple bond is needed, which makes the structure incomplete. Convert a lone pair from a neighboring atom into a double or triple bond when the central atom lacks an octet.
Practice Questions
- 1 Draw the Lewis structure for CO2. Count the total valence electrons, show all bonds and lone pairs, and identify whether single, double, or triple bonds are needed.
- 2 Draw the Lewis structure for NO3-. Count the total valence electrons, include brackets and the charge, and calculate the formal charge on each atom for one valid resonance form.
- 3 Explain why CH4 has carbon as the central atom and four single bonds, while H2O has oxygen as the central atom with two lone pairs.