Source: Quiz Section 1, Chapter 1 (Klein textbook)
Tags: Lewis structures, formal charge, octet rule, electron configuration, atomic orbitals, bonding patterns, expanded octet, organic chemistry, CHEM 237, University of Washington
Difficulty: Introductory | Prerequisites: General chemistry (atomic structure, periodic trends, electronegativity)
This material forms the foundation of everything in organic chemistry. Before you can understand reactions, mechanisms, or molecular behaviour, you need to be comfortable drawing Lewis structures, assigning formal charges, and recognising bonding patterns. These skills come from Chapter 1 of the Klein textbook and are tested early and often. If you are shaky on electron configurations or the octet rule from general chemistry, revisit those topics before going further.
Atoms bond by sharing or transferring electrons to achieve stable electron configurations, usually an octet. Lewis structures show how electrons are distributed in a molecule, and formal charge tells you whether an atom "owns" more or fewer electrons than it brought to the party. Knowing the standard bonding patterns for common atoms (C, N, O, H, halogens, B, P, Si) lets you draw and check structures quickly.
Atomic orbital
A region of space around the nucleus where there is a high probability of finding an electron. Orbitals are labelled by quantum numbers: the principal quantum number (n = 1, 2, 3...) and the subshell (s, p, d, f). In simple terms, think of it as the "address" an electron can have inside an atom.
Octet rule
Most atoms in the second row of the periodic table (C, N, O, F) are most stable when surrounded by eight electrons (four pairs). In simple terms, atoms want a full valence shell, and for second-row elements that means eight electrons.
Expanded octet
Atoms in the third row and below (e.g. S, P, Si) can accommodate more than eight electrons in their valence shell because they have access to empty d orbitals. Think of it as having extra room in the valence shell that second-row atoms do not have.
Incomplete octet (fewer than octet)
Some atoms, notably boron (B) and hydrogen (H), are stable with fewer than eight electrons. Boron commonly has six electrons around it; hydrogen is stable with two.
Lewis structure (electron-dot structure)
A diagram showing all valence electrons in a molecule as bonds (shared pairs) and lone pairs (non-bonding pairs). Think of it as a map of where every valence electron lives in a molecule.
Formal charge
A bookkeeping tool that assigns a hypothetical charge to each atom in a Lewis structure, calculated by comparing the number of valence electrons a free atom has with the number it "owns" in the structure. In simple terms, it tells you whether an atom is hogging electrons or giving them away relative to its neutral state.
Lone pair (non-bonding pair)
A pair of valence electrons on an atom that is not involved in bonding. Think of it as electrons an atom keeps to itself rather than sharing with a neighbour.
Bonding pair (shared pair)
A pair of electrons shared between two atoms, forming a covalent bond.
Valence electrons
The electrons in the outermost shell of an atom. These are the electrons involved in bonding and are the ones you count when drawing Lewis structures.
Electronic octet
When an atom has exactly eight electrons around it (counting both bonding and lone-pair electrons). The term is used to distinguish atoms that satisfy the octet rule from those with expanded or incomplete octets.
Orbitals are defined by the principal quantum number n (1, 2, 3...) and the angular momentum quantum number l (0 = s, 1 = p, 2 = d, 3 = f).
For a given principal quantum number n, the allowed subshells are l = 0 through l = n − 1.
n = 1: only 1s
n = 2: 2s, 2p
n = 3: 3s, 3p, 3d
n = 4: 4s, 4p, 4d, 4f
An orbital that does not exist: 2d is not possible. For n = 2, the maximum l is 1 (p), so there is no 2d orbital. This was tested directly in Quiz 1, Q1.40.
Energy ordering (approximate, for multi-electron atoms): 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p.
Among the orbitals listed in Q1.40 (4d, 5s, 2p, 3p, 2d, 6p): 2d does not exist. Of the remaining, 6p has the highest energy and 2p has the lowest energy.
Each orbital holds a maximum of 2 electrons (Pauli exclusion principle).
The octet rule: Second-row elements (C, N, O, F) tend to have eight electrons in their valence shell when bonded. This means four bonds for carbon, three bonds and one lone pair for nitrogen, two bonds and two lone pairs for oxygen, and one bond and three lone pairs for fluorine (all when neutral).
Expanded octets: Third-row elements and below (S, P, Si, Cl, Br, I) can exceed eight electrons because empty 3d (or higher) orbitals are available. Sulphur in SO₄²⁻, for example, can have 12 electrons around it.
Incomplete octets: Boron commonly has only six electrons (three bonds, no lone pairs). Hydrogen is always satisfied with two electrons (one bond).
When asked which atoms have a regular octet, an expanded octet, or fewer than an octet, count all electrons around each atom (bonding pairs + lone pairs):
8 electrons = regular octet (O, N, C with appropriate bonds)
More than 8 = expanded octet (typically S, P, or Si)
Fewer than 8 = incomplete octet (typically B or H)
In the quiz example, sulphur (S) had an expanded octet, boron (B) had fewer than an octet, and O, H, N, and C had regular octets.
Drawing a correct Lewis structure is a step-by-step process:
Count total valence electrons. Add up valence electrons for every atom. For anions, add one electron per negative charge. For cations, subtract one electron per positive charge.
Draw the skeleton. Connect atoms with single bonds. The least electronegative atom usually goes in the centre (hydrogen and fluorine are always terminal).
Distribute remaining electrons. Place lone pairs on outer atoms first to satisfy their octets, then on the central atom.
Check octets. If the central atom lacks an octet, convert lone pairs on adjacent atoms into bonding pairs (forming double or triple bonds).
Assign formal charges (see next section).
Azide ion, N₃⁻
Total valence electrons: 3 × 5 + 1 (for the charge) = 16. The structure is linear: N=N=N with formal charges of −1, +1, −1 on the three nitrogens (left to right), or the resonance form N≡N−N with charges 0, +1, −2. The symmetric structure with two double bonds is the more important resonance contributor.
Isocyanic acid, HNCO (atom connectivity H−N=C=O)
Total valence electrons: 1 + 5 + 4 + 6 = 16. Two valid structures exist:
Without formal charges: H−N=C=O (nitrogen has one lone pair, oxygen has two lone pairs)
With formal charges: H−N≡C−O⁻ (nitrogen carries no formal charge, carbon carries +1, oxygen carries −1)
Butane, C₄H₁₀ (two possible structures)
n-Butane: CH₃−CH₂−CH₂−CH₃ (straight chain)
Isobutane: (CH₃)₃CH (branched, with a central carbon bonded to three CH₃ groups and one H)
Both have the same molecular formula but different connectivity. This is an early preview of constitutional isomers.
Ketene, C₂H₂O (no ring)
Structure: H₂C=C=O. The left carbon has two C−H bonds and a C=C double bond. The right carbon has a C=C double bond and a C=O double bond. All formal charges are zero. Oxygen has two lone pairs.
The formal charge formula is central to this chapter and will follow you through the entire course.
Formal charge = (valence electrons) − (lone pair electrons) − (shared electrons ÷ 2)
Or equivalently:
Formal charge = (valence electrons) − (lone pair electrons) − (number of bonds)
Since each bond represents two shared electrons, dividing shared electrons by 2 gives you the number of bonds.
Carbon (4 valence e⁻): 4 bonds, 0 lone pairs → FC = 0. Three bonds + 1 lone pair → FC = −1. Three bonds + 0 lone pairs → FC = +1.
Nitrogen (5 valence e⁻): 3 bonds, 1 lone pair → FC = 0. Four bonds, 0 lone pairs → FC = +1. Two bonds, 2 lone pairs → FC = −1.
Oxygen (6 valence e⁻): 2 bonds, 2 lone pairs → FC = 0. Three bonds, 1 lone pair → FC = +1. One bond, 3 lone pairs → FC = −1.
Hydrogen (1 valence e⁻): 1 bond → FC = 0. No bonds, but paired → FC = −1 (hydride).
The best Lewis structure minimises formal charges and places any negative formal charges on the more electronegative atom.
Memorise these. They let you spot errors in Lewis structures at a glance.
H: 1 bond, 0 lone pairs
C: 4 bonds, 0 lone pairs
N: 3 bonds, 1 lone pair
O: 2 bonds, 2 lone pairs
F / Cl / Br / I (halogens): 1 bond, 3 lone pairs
B: 3 bonds, 0 lone pairs (incomplete octet, only 6 electrons)
P: 3 bonds, 1 lone pair (same as nitrogen, but can also expand)
Si: 4 bonds, 0 lone pairs (same pattern as carbon, but can expand)
N⁺: 4 bonds, 0 lone pairs (like carbon's neutral pattern, e.g. NH₄⁺)
O⁺: 3 bonds, 1 lone pair (like nitrogen's neutral pattern, e.g. H₃O⁺)
O⁻: 1 bond, 3 lone pairs
C⁻: 3 bonds, 1 lone pair (carbanion)
B⁻: 4 bonds, 0 lone pairs (like carbon's neutral pattern, e.g. BH₄⁻)
The pattern is consistent: gaining a bond costs a lone pair and adds a positive formal charge; losing a bond gains a lone pair and adds a negative formal charge.
Formal Charge Formula
FC = V − L − B
Where V = number of valence electrons for the free atom, L = number of lone-pair electrons on the atom, B = number of bonds to the atom.
Allowed Orbitals Rule
For principal quantum number n, the allowed subshell values are l = 0, 1, 2, ... , (n − 1). An orbital is "not possible" if l ≥ n. For example, 2d requires l = 2 but n = 2 only allows l up to 1.
Electron Count for Octet Check
Total electrons around an atom = (2 × number of bonds) + (lone-pair electrons). If this equals 8, the atom has a regular octet.
Formal charge and Lewis structures are not just exam exercises. Chemists use them to predict which atoms in a molecule are most reactive, which is the foundation for understanding reaction mechanisms later in the course. Pharmaceutical chemists draw Lewis structures to reason about drug-receptor interactions, and materials scientists use bonding patterns to design polymers with specific properties.
Students often think that all atoms follow the octet rule. They do not. Hydrogen is stable with 2 electrons, boron with 6, and third-row elements can exceed 8.
Students confuse the number of bonds with the number of electrons in a bond. A single bond is 2 electrons, a double bond is 4 electrons, a triple bond is 6 electrons. When counting electrons around an atom for octet purposes, count all the electrons in every bond to that atom.
Students sometimes forget to include the extra electron when drawing Lewis structures for anions, or forget to subtract an electron for cations. This throws off the entire structure.
Students assume that the structure with formal charges is always wrong. Both structures (with and without formal charges) can be valid Lewis structures. The one that minimises formal charges is generally preferred, but the other is still a legitimate resonance contributor.
⚠️ The formal charge formula (FC = V − L − B) appears on nearly every exam in some form. Be able to calculate it quickly for any atom.
⚠️ Expect questions asking you to identify which orbital designation is "not possible" (e.g. 2d, 1p, 3f). Know that the subshell number must be less than the principal quantum number.
⚠️ Drawing Lewis structures with correct formal charges is a foundational skill tested repeatedly. You will be asked to draw structures both with and without formal charges and compare them.
⚠️ Bonding patterns for neutral and charged atoms (Section 1.3C) are assumed knowledge for the rest of the course. If carbon has five bonds in your structure, something is wrong.
⚠️ Constitutional isomers (like n-butane vs. isobutane for C₄H₁₀) will return in later chapters. Get comfortable with the idea that the same molecular formula can have different structures.
True or false: A 3f orbital exists.
False. For n = 3, maximum l = 2 (d subshell). f requires l = 3.
Fill in the blank: Neutral oxygen forms ___ bonds and has ___ lone pairs.
2 bonds, 2 lone pairs.
True or false: Boron typically satisfies the octet rule.
False. Boron commonly has only 6 electrons around it.
Fill in the blank: The formal charge on nitrogen in NH₄⁺ is ___.
+1. (5 valence electrons − 0 lone pair electrons − 4 bonds = +1.)
True or false: Sulphur can have more than 8 electrons in its valence shell.
True. Sulphur is in the third row and has access to empty d orbitals.
Q: Which of the following orbitals is not possible: 4d, 5s, 2p, 3p, 2d, 6p?
A: 2d is not possible. For n = 2, the maximum angular momentum quantum number is l = 1 (p subshell). A d subshell requires l = 2, which requires n ≥ 3.
Q: In a molecule containing B, C, N, O, H, and S, how do you determine which atoms have expanded octets, regular octets, or fewer than an octet?
A: Count the total electrons around each atom (bonding + lone pair). H will have 2 (fewer). B will typically have 6 (fewer). C, N, O will have 8 (regular). S can have 10 or 12 (expanded) if it has more than four bonds or additional lone pairs beyond the octet.
Q: Draw a Lewis structure for the azide ion (N₃⁻) and assign formal charges to each atom.
A: The symmetric structure is :N=N=N:⁻ with formal charges −1, +1, −1 from left to right. Total valence electrons: 3(5) + 1 = 16.
Q: What is the formal charge on oxygen in H₃O⁺?
A: +1. Oxygen has 6 valence electrons, 1 lone pair (2 electrons), and 3 bonds. FC = 6 − 2 − 3 = +1.
Q: Draw two constitutional isomers of C₄H₁₀.
A: n-Butane (CH₃CH₂CH₂CH₃, a straight chain of four carbons) and isobutane ((CH₃)₃CH, a central carbon bonded to three methyl groups and one hydrogen).
Q: For isocyanic acid (HNCO), draw one Lewis structure with no formal charges and one with two formal charges.
A: No formal charges: H−N=C=O (N has one lone pair, O has two). With formal charges: H−N≡C−O⁻ (N has one lone pair and FC = 0, C has FC = +1, O has three lone pairs and FC = −1).
This material connects directly to resonance structures (Chapter 2), where you will draw multiple valid Lewis structures for the same molecule and evaluate their relative contributions. The formal charge skills you build here are essential for that.
It also feeds into acid-base chemistry (Chapter 3), where understanding lone pairs and electronegativity helps predict which atoms can donate or accept protons.
Later, when you study reaction mechanisms, every curved arrow you draw represents the movement of electron pairs, so a solid grasp of where electrons sit in a Lewis structure is the starting point for every mechanism.
Lewis dot structure, electron dot diagram, formal charge calculation, octet rule, expanded octet, incomplete octet, valence electrons, bonding patterns, lone pairs, non-bonding electrons, covalent bonds, quantum numbers, atomic orbitals, 2d orbital, impossible orbitals, azide ion, isocyanic acid, ketene, butane isomers, constitutional isomers, CHEM 237, organic chemistry, University of Washington, Klein textbook Chapter 1, resonance structures, electron configuration, valence shell, shared electrons