Prepare for AP Chemistry Topic 2.5 with this free practice quiz. Test your ability to draw accurate Lewis electron-dot structures and obey the octet rule to maximize your AP score.
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The Pre-Quiz Review
Lewis diagrams show how valence electrons are distributed around atoms in covalent molecules and polyatomic ions. For AP Chemistry, the key skill is being able to quickly determine the correct number of valence electrons, distribute them according to the octet rule, and recognize exceptions such as incomplete and expanded octets.
Valence Electrons
Valence electrons are the electrons in the outermost occupied shell of an atom. For main-group elements, the periodic table group generally tells you the number of valence electrons:
- Group 1 → 1 valence electron
- Group 2 → 2
- Group 13 → 3
- Group 14 → 4
- Group 15 → 5
- Group 16 → 6
- Group 17 → 7
- Group 18 → 8, except He, which has 2
Lewis diagrams use only valence electrons.
For a molecule, add the valence electrons from every atom. For a polyatomic ion, adjust the total for the charge:
- Negative charge → add electrons.
- Positive charge → subtract electrons.
Example: SO₄²⁻
S contributes 6 electrons, four O atoms contribute 24, and the 2− charge adds 2 more:
Total = 32 valence electrons.
The total electron count is one of the fastest ways to check whether a Lewis diagram is possible.
Basic Lewis Diagram Procedure
Use this sequence:
- Count all valence electrons.
- Choose the central atom.
- Connect atoms with single bonds.
- Complete the octets of terminal atoms.
- Place remaining electrons on the central atom.
- If the central atom lacks an octet, form multiple bonds when appropriate.
- Check the final electron count.
Hydrogen is never the central atom because it can form only one bond and follows a duet rule rather than the octet rule.
The central atom is often the least electronegative atom, excluding hydrogen. Carbon is commonly central in carbon-containing compounds. Halogens are usually terminal atoms because they normally form one bond.
Bonding Pairs
A bonding pair is a pair of electrons shared between two atoms in a covalent bond.
In a Lewis diagram:
- A single bond contains 2 electrons.
- A double bond contains 4 electrons.
- A triple bond contains 6 electrons.
Each line represents one bonding pair.
For example:
H–Cl
The H–Cl bond is one bonding pair. Chlorine also has three lone pairs.
Bonding electrons count toward the octet of both atoms involved in the bond. This is why a single bond contributes two electrons around each bonded atom when determining whether its octet is satisfied.
Lone Pairs
A lone pair is a pair of valence electrons that is not shared in a covalent bond.
For example, in H₂O, oxygen has:
- 2 bonding pairs
- 2 lone pairs
The four electron pairs around oxygen give it an octet.
Lone pairs are especially important because they affect molecular geometry and bond angles. In VSEPR theory, lone pairs occupy electron domains and generally repel more strongly than bonding pairs.
For Lewis diagram questions, distinguish carefully between:
- bonding pairs: electrons shared between atoms
- lone pairs: electrons localized on one atom
The Octet Rule
The octet rule states that many main-group atoms tend to form bonds so that they have 8 electrons in their valence shell.
Atoms can achieve an octet by:
- gaining electrons
- losing electrons
- sharing electrons through covalent bonds
For Lewis structures of covalent compounds, sharing electrons is the primary mechanism.
Carbon, nitrogen, oxygen, and fluorine commonly obey the octet rule.
Examples:
CH₄: carbon has four bonds and therefore 8 electrons around it.
NH₃: nitrogen has three bonds plus one lone pair, giving 8 electrons.
H₂O: oxygen has two bonds plus two lone pairs, giving 8 electrons.
Atoms in the second period are particularly important because they cannot exceed an octet in ordinary Lewis structures. These include B, C, N, O, and F.
Hydrogen and the Duet Rule
Hydrogen is an important exception to the octet rule.
Hydrogen is stable with 2 electrons because its only occupied shell is the first energy level, which contains only the 1s orbital.
Therefore:
- H forms one single bond.
- H can have only 2 electrons around it.
- H does not form double or triple bonds in ordinary introductory Lewis structures.
For example, in H₂, each hydrogen shares one pair of electrons and achieves a duet.
Incomplete Octets
Some atoms can be stable with fewer than 8 electrons around them. This is called an incomplete octet.
The most important AP Chemistry examples involve boron and beryllium.
Boron commonly forms three bonds and can have only 6 electrons around it.
Example:
BF₃
Boron forms three B–F bonds. The three bonding pairs give boron 6 electrons, so boron has an incomplete octet.
Beryllium can form two bonds and have only 4 electrons around it.
Example:
BeCl₂
Beryllium forms two Be–Cl bonds, giving it 4 electrons around the central atom.
Do not automatically add a double bond simply because the central atom lacks an octet. Incomplete octets are legitimate exceptions, particularly for Be and B.
Expanded Octets
Elements in the third period and beyond can sometimes have more than 8 electrons around the central atom. These are called expanded octets.
Common examples include:
- P
- S
- Cl
- Br
- I
- Xe
For example, phosphorus can form PCl₅, in which phosphorus has five bonding pairs, or 10 electrons around the central atom.
Sulfur can appear with expanded octets in compounds such as SF₆, where six S–F bonds give sulfur 12 electrons around it.
Second-period elements cannot have expanded octets in standard AP Chemistry Lewis structures.
This distinction is frequently tested:
- Period 2 → maximum of 8 electrons.
- Period 3 and beyond → expanded octets may occur.
Single, Double, and Triple Bonds
If a central atom does not have an octet after all terminal atoms have been completed, a lone pair on a neighboring atom may be used to form an additional bond.
For example, CO₂ initially gives carbon two single bonds:
O–C–O
After completing the oxygen octets, carbon has only 4 electrons around it. One lone pair from each oxygen can be used to form additional bonds, producing:
O=C=O
Carbon now has four bonding pairs and an octet.
Bond order affects both electron distribution and bond properties:
- Single bond → 1 shared pair
- Double bond → 2 shared pairs
- Triple bond → 3 shared pairs
Higher bond order generally means a shorter and stronger bond.
Formal Charge and Lewis Structures
Formal charge is useful for deciding between possible Lewis structures.
For an atom in a Lewis structure:
Formal charge = valence electrons − nonbonding electrons − number of bonds
When comparing structures, prefer structures that:
- give atoms appropriate octets when possible
- minimize the magnitude of formal charges
- place negative formal charge on more electronegative atoms when possible
The sum of all formal charges must equal the overall charge of the species.
For a neutral molecule, the formal charges must add to 0.
For a 2− ion, they must add to −2.
Formal charge is particularly useful when more than one Lewis structure appears possible. AP Chemistry questions may provide several structures and ask which one is most reasonable.
Lewis Structures of Polyatomic Ions
For polyatomic ions, the charge must be included when counting electrons.
For a negative ion, add electrons. For a positive ion, subtract electrons.
Example: NO₃⁻
N contributes 5 electrons.
Three O atoms contribute 18.
The −1 charge adds 1 electron.
Total = 24 valence electrons.
After constructing the structure, the entire ion is normally enclosed in brackets with the charge shown outside:
[NO₃]⁻
The bracket indicates that the displayed atoms and electrons together represent an ion.
Resonance
Some species cannot be represented adequately by a single Lewis structure because equivalent electron arrangements are possible. These are resonance structures.
For example, NO₃⁻ has three equivalent arrangements in which the N=O double bond can be placed between nitrogen and any one of the three oxygen atoms.
The actual species is not rapidly switching between three structures. Instead, the real electron distribution is a resonance hybrid.
In equivalent resonance structures, the atoms remain in the same positions while the placement of electrons and multiple bonds changes.
AP questions may ask you to identify resonance structures or determine whether multiple Lewis structures represent resonance.
Electron Counting as a Verification Tool
Electron counting is one of the fastest ways to catch errors.
After drawing a Lewis structure, count:
- electrons in all bonds
- electrons in all lone pairs
The total must equal the number of valence electrons calculated at the beginning.
For a neutral molecule, there should be no missing or extra electrons.
For an ion, the total must reflect the charge.
A structure can look chemically reasonable but still be incorrect if the electron count does not match.
Octet Counting
When checking an atom’s octet, count all electrons surrounding it:
- single bond → 2 electrons
- double bond → 4 electrons
- triple bond → 6 electrons
- each lone pair → 2 electrons
For example, an atom with two single bonds and two lone pairs has:
2 + 2 + 2 + 2 = 8 electrons
An atom with one double bond, one single bond, and one lone pair also has:
4 + 2 + 2 = 8 electrons.
Remember that the octet rule concerns electrons around an atom, not simply the number of lines attached to it.
Central Atom and Bonding Patterns
Common bonding patterns are useful for rapid AP Chemistry questions:
- H → usually 1 bond, 0 lone pairs
- F, Cl, Br, I → usually 1 bond, 3 lone pairs
- O → commonly 2 bonds, 2 lone pairs
- N → commonly 3 bonds, 1 lone pair
- C → commonly 4 bonds, 0 lone pairs
These are patterns rather than absolute rules because formal charge, resonance, and expanded or incomplete octets can alter them.
For example, nitrogen can have four bonds in NH₄⁺, giving nitrogen a positive formal charge.
Lewis Structures and Molecular Geometry
Lewis structures also provide the electron-domain information needed for VSEPR analysis.
Count regions of electron density around the central atom:
- single bond = 1 electron domain
- double bond = 1 electron domain
- triple bond = 1 electron domain
- lone pair = 1 electron domain
Thus, a double bond does not count as two electron domains for VSEPR purposes.
For example, CO₂ has two electron domains around carbon because carbon has two double bonds. Its electron-domain geometry is linear.
H₂O has four electron domains around oxygen: two bonding pairs and two lone pairs. Its electron-domain geometry is tetrahedral, while its molecular geometry is bent.
This distinction is commonly tested: lone pairs count as electron domains even though they are not shown as atoms in the molecular geometry.
Fast Recognition of Common Exceptions
For rapid Lewis diagram questions, recognize these patterns immediately:
Hydrogen:
- duet, not octet
Beryllium:
- may have 4 electrons
Boron:
- may have 6 electrons
Second-period elements:
- cannot exceed 8 electrons
Third-period and heavier main-group elements:
- may have expanded octets
Negative ions:
- add electrons
Positive ions:
- subtract electrons
Polyatomic ions:
- use brackets and show the overall charge
Equivalent electron arrangements:
- may indicate resonance
A correct Lewis structure must simultaneously satisfy the electron count, appropriate bonding pattern, charge, and applicable octet-rule exceptions.