AP Chemistry Unit 2 Study Notes

AP Chemistry 2.5: Lewis diagrams

Represent valence electrons, bonding pairs, lone pairs, and octets.

Aligned to Compound Structure and Properties from the current College Board AP Chemistry course outline. Exam weighting for this unit: 7%-9% of the multiple-choice score range listed by College Board.

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Lewis Diagrams Overview 5.1 What Is a Lewis Diagram? 5.2 Bonding Pair 5.3 Lone Pair 5.4 The Octet Rule 5.5 Hydrogen Exception 5.6 How to Draw Lewis Structures 5.7 Example: H₂O 5.8 Example: CO₂ 5.9 Example: N₂ 5.10 Lewis Structures of Ions 5.11 Example: NH₄⁺ 5.12 Example: CN⁻ 5.13 Octet Rule Exceptions 5.14 Incomplete Octets 5.15 Expanded Octets 5.16 Do Not Expand Period-2 Atoms

Source note

These detailed Unit 2 notes were organized from the provided study document. For further study, visit Khan Academy. All Khan Academy content is available for free at www.khanacademy.org.

Lewis Diagrams Overview Open

Khan Academy's Lewis-diagram section covers drawing structures, molecular and polyatomic-ion examples, and exceptions to the octet rule.
- This is one of the biggest Unit 2 topics.

5.1 What Is a Lewis Diagram? Open

A Lewis diagram shows how valence electrons are distributed in a molecule or polyatomic ion.
- It represents:
- bonding pairs
- with lines
- and:
- lone pairs
- with dots.

5.2 Bonding Pair Open

Two electrons shared between two atoms form a covalent bond.
- Example:
- H:H
- is usually written:
- H—H
One line means:
- 2 bonding electrons

5.3 Lone Pair Open

A lone pair is a pair of valence electrons not used in bonding.
- Example in water:
- oxygen has two lone pairs in addition to two O—H bonds.
- These lone pairs strongly influence molecular shape.

5.4 The Octet Rule Open

Many main-group atoms tend toward:
- 8 valence electrons
- through bonding.
- This resembles the stable electron configuration of a noble gas.
- Important:
The octet rule is a useful model, not a universal law .
There are exceptions.

5.5 Hydrogen Exception Open

Hydrogen only needs:
- 2 electrons
- because the first shell contains only the 1s orbital.
- So H normally forms one bond.
Do not give hydrogen an octet.

5.6 How to Draw Lewis Structures Open

Use this general procedure.
Step 1: Count total valence electrons
Add valence electrons from every atom.
For ions:
- negative charge → add electrons
- positive charge → subtract electrons
Step 2: Choose the central atom
Usually:
- the least electronegative atom
- goes in the center.
Hydrogen is almost never the central atom.
Halogens are usually terminal.
Step 3: Connect atoms using single bonds
Each bond uses:
- 2 electrons
Step 4: Complete terminal-atom octets
Place remaining electrons around outer atoms first.
Step 5: Place remaining electrons on the central atom
Step 6: Check the central atom's octet
If it lacks an octet, consider creating:
double bonds
triple bonds
- using lone pairs from neighboring atoms.

5.7 Example: H₂O Open

Valence electrons:
- O = 6
- 2 H = 2
- Total:
- 8 electrons
Skeleton:
- H—O—H
Two bonds use:
4 electrons.
Remaining:
- 4 electrons
Place them as two lone pairs on oxygen.
Oxygen now has:
2 bonding pairs + 2 lone pairs
- = 8 electrons around it.

5.8 Example: CO₂ Open

Total valence electrons:
- C = 4
- 2 O = 12
- Total:
- 16
Start:
- O—C—O
After completing oxygen octets, carbon does not have a complete octet.
Move lone-pair electrons into bonding positions.
Final structure:
- O=C=O
Carbon now has an octet.

5.9 Example: N₂ Open

Each nitrogen:
- 5 valence electrons
- Total:
- 10
A single bond does not satisfy both octets.
Eventually we obtain:
- N≡N
- with one lone pair on each N.

5.10 Lewis Structures of Ions Open

For a negative ion:
- add electrons equal to the negative charge.
For a positive ion:
- subtract electrons.
- Then place the entire Lewis structure inside:
- brackets
- and show the charge.

5.11 Example: NH₄⁺ Open

Valence electrons:
- N = 5
- 4 H = 4
Total before charge:
- 9
- Because charge is +1:
- subtract 1.
- Total:
- 8 electrons
Four N—H bonds use all 8 electrons.
NH₄⁺ has four bonds and no lone pairs on nitrogen.

5.12 Example: CN⁻ Open

C:
- 4
N:
- 5
−1 charge:
- +1 electron
- Total:
- 10 electrons
The structure is essentially:
- [C≡N]⁻
- with a lone pair associated with each atom in the complete Lewis representation.
Khan Academy includes CN⁻ as one of its current worked Lewis-diagram examples.

5.13 Octet Rule Exceptions Open

Three main categories matter.
- 1. Incomplete octet
Some atoms are stable with fewer than eight electrons.
- 2. Odd-electron species
Some molecules contain an odd total number of electrons.
- 3. Expanded valence shell
Some atoms from period 3 or beyond can be represented with more than eight electrons around the central atom in common Lewis models.

5.14 Incomplete Octets Open

Boron and beryllium are common examples.
BF₃ often has:
- three B—F bonds.
Boron has only:
- 6 electrons around it
- yet this is a valid structure.

5.15 Expanded Octets Open

Elements in period 3 or below on the periodic table can sometimes be represented with more than eight electrons around them.
Examples include compounds with:
- P
- S
- Xe
Khan Academy includes XeF₂ as an expanded-octet Lewis-structure example.

5.16 Do Not Expand Period-2 Atoms Open

Second-period elements such as:
- C
- N
- O
- F
- cannot exceed an octet in the usual AP Lewis-structure model.
- This is a major test trap.