Lewis Structures
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Core Concept
Lewis structures are electron dot diagrams that show:
How valence electrons are arranged in molecules and ions
Which atoms are bonded together
Where lone pairs (nonbonding electrons) are located
Key Purpose: Bridge between individual atoms and molecular behavior.
Practice Tips
Miscounting valence electrons: Always double-check your total electron count.
Forgetting formal charges: Use formal charges to find the most stable structure, especially for ions.
Ignoring octet rule exceptions: Remember that hydrogen, boron, and elements in period 3 or higher may not follow the octet rule.
Not drawing all resonance structures: If there are multiple ways to arrange double bonds or lone pairs, draw all resonance structures.
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LABORATORY
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Step-by-Step Guide
Step 1: Count Total Valence Electrons
Determine the number of valence electrons for each atom in the molecule. This information is based on the group number for each element in the periodic table.
Add up the valence electrons for all atoms in the molecule. For polyatomic ions:
Add electrons if the ion is negatively charged.
Subtract electrons if the ion is positively charged.
Step 2: Arrange the Atoms
Place the least electronegative atom (usually the atom that can form the most bonds) in the center. Hydrogen is always a terminal (outer) atom, as it can only form one bond.
Arrange other atoms around the central atom, symmetrically if possible.
Step 3: Form Single Bonds
Connect each outer atom to the central atom using a single line, which represents a single bond (2 electrons).
Subtract the bonding electrons (2 electrons per bond) from the total valence electron count.
Step 4: Distribute Remaining Electrons as Lone Pairs
Place lone pairs of electrons around each outer atom to fulfill the octet rule (8 electrons), starting with the most electronegative atoms.
Any remaining electrons should go to the central atom.
Step 5: Check the Octet Rule
Ensure each atom (except hydrogen) has a complete octet (8 electrons around it). Hydrogen should have only 2 electrons.
If the central atom has fewer than 8 electrons, consider forming double or triple bonds by sharing lone pairs from surrounding atoms.
Step 6: Verify and Adjust
Count the total electrons in the structure to ensure it matches the original total valence electrons.
For polyatomic ions, add brackets around the structure and indicate the charge outside the brackets (e.g., $[ \text{NO}_3^- ]$).
Example:
Drawing the Lewis Structure for CO₂
Count Valence Electrons: Carbon has 4, and each oxygen has 6.
Total = 4+6+6=16 electrons.
Arrange the Atoms: Place carbon (the least electronegative) in the center with oxygens on each side.
Form Single Bonds: Connect each oxygen to carbon with a single bond. This uses 2×2=4 electrons, leaving 16 − 4 = 12 electrons.
Distribute Remaining Electrons: Place 6 lone pairs (12 electrons) around the oxygens to fulfill their octets.
Check the Octet Rule for Carbon: Carbon has only 4 electrons, so form double bonds by converting lone pairs from each oxygen into bonding pairs.
Verify: The structure should have two double bonds, with each atom following the octet rule.
The final structure is: O=C=O
Study Tips
Practice Strategy:
Start simple - Master H₂O, NH₃, CH₄ first
Work systematically - Follow all 6 steps every time
Draw by hand - Don't just look at structures, draw them
Check your work - Count electrons and verify octets
Memory Aids:
"Hydrogen is never central" - H can only form 1 bond
"Outer atoms first" - Complete outer octets before central
"Count twice, draw once" - Verify electron count before starting
