Organic Chemistry II (University of Michigan) | Source: Topic 3 Lecture Notes
Difficulty: Introductory | Prerequisites: General chemistry (electron configuration, periodic table trends, electronegativity)
Lewis structures are the bread-and-butter notation for showing how atoms share electrons in covalent molecules. You will use them constantly in Organic Chemistry II to predict reactivity, understand charge distribution, and draw reaction mechanisms. This topic sits right at the start of the course because almost everything that follows (resonance, acid-base chemistry, nucleophilicity) depends on being able to draw a correct Lewis structure quickly and assign formal charges without hesitation. If you are rusty on valence electrons or the octet rule from general chemistry, revisit those before going further.
A Lewis structure is a diagram that shows every valence electron in a covalent molecule, with bonding pairs drawn as lines and lone pairs as dots. You draw one by working out the connectivity, counting total valence electrons, satisfying the octet (or duet) rule, and then assigning formal charges. Getting this right is the foundation for everything else in organic chemistry.
Lewis structure
A representation of a covalent molecule that shows all valence electrons: bonding pairs as lines between atoms and non-bonding (lone) pairs as dots. Think of it as the molecular blueprint that tells you where every electron lives.
Valence electrons
The electrons in an atom's outermost shell, which participate in bonding. In simple terms, these are the electrons that "do things" in chemistry.
Bonding pair (shared pair)
A pair of electrons shared between two atoms, forming a covalent bond. Drawn as a line connecting two atoms.
Non-bonding pair (lone pair)
A pair of valence electrons that belongs to one atom and is not shared. Drawn as two dots on the atom. These matter enormously in organic chemistry because lone pairs drive nucleophilic behaviour and affect molecular shape.
Octet rule
The principle that most atoms (C, N, O, F, Cl, etc.) are stable when surrounded by eight electrons. Think of it as the atom wanting a full outer shell, like the noble gases.
Duet rule
The hydrogen-specific version of the octet rule: H is stable with just two electrons (one bond).
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons (free atom) minus lone-pair electrons minus half the bonding electrons. It tells you which atoms are electron-rich or electron-poor in the structure.
Valency
The number of bonds an atom typically forms. Carbon is tetravalent (4 bonds), nitrogen is trivalent (3 bonds), oxygen is divalent (2 bonds), hydrogen and halogens are monovalent (1 bond).
Isomers
Compounds that share the same molecular formula but have different structural arrangements. In simple terms, same atoms, different wiring.
Step 1: Determine connectivity. Work out which atoms are bonded to which. Use typical valencies as your guide (see the table below). Hydrogen and halogens always go on the outside; carbon is almost always interior.
Step 2: Count total valence electrons. Add up all valence electrons from every atom. Adjust for charge: subtract one electron per positive charge, add one per negative charge.
Step 3: Satisfy the duet and octet rules. Place bonding pairs first (one per single bond), then distribute remaining electrons as lone pairs so that each atom reaches its target (2 for H, 8 for C/N/O/F/Cl). If electrons run short, form double or triple bonds.
Step 4: Assign formal charges. For each atom, calculate: formal charge = (valence electrons of free atom) − (lone pair electrons) − (½ bonding electrons). The best Lewis structure minimises formal charges and places any negative formal charge on the more electronegative atom.
Atom | Valence Electrons | Typical Bonds | Valency Name |
|---|---|---|---|
H | 1 | 1 | Monovalent |
F, Cl, Br, I | 7 | 1 | Monovalent |
O | 6 | 2 | Divalent |
N | 5 | 3 | Trivalent |
B | 3 | 3 | Trivalent |
C | 4 | 4 | Tetravalent |
Carbon is tetravalent. It always forms exactly four bonds (single, double, or triple in any combination that totals four). Carbon cannot make more than four bonds.
Oxygen is divalent. It forms two bonds and carries two lone pairs in a neutral structure. It can appear as −O− (two single bonds) or =O (one double bond).
Nitrogen is trivalent. It forms three bonds and carries one lone pair when neutral. Possible patterns: three single bonds, one single + one double, or one triple bond.
Halogens are monovalent. They form one bond and carry three lone pairs.
\text{Formal Charge} = (\text{valence } e^{-} \text{ of free atom}) - (\text{lone pair } e^{-}) - \tfrac{1}{2}(\text{bonding } e^{-})Equivalent shortcut: count lone-pair electrons + bonds (not bonding electrons, just the number of lines). If that total equals the atom's group number, formal charge is zero.
H has 1 valence electron, Cl has 7. Total = 8. One bonding pair connects them (uses 2 electrons). Remaining 6 electrons go as three lone pairs on Cl. Both atoms satisfied: H has its duet, Cl has its octet. All formal charges are zero.
Lewis structures are the starting point for predicting molecular shape (via VSEPR), which in turn determines physical properties like boiling point and solubility. Pharmaceutical chemists draw Lewis structures to understand how a drug molecule will interact with an enzyme's active site, because the position of lone pairs and formal charges dictates where the molecule can donate or accept electrons.
Students often forget lone pairs on terminal atoms (especially halogens and oxygen). Every electron must be accounted for; if your total count does not match, you have a missing lone pair somewhere.
Students sometimes give carbon five bonds. Carbon can never exceed four bonds. If your structure shows five lines to a carbon, redraw it.
Students confuse formal charge with oxidation state. Formal charge splits bonding electrons equally between the two atoms; oxidation state gives them all to the more electronegative atom. They answer different questions.
Students assume that a Lewis structure with formal charges is always wrong. Formal charges are sometimes unavoidable (e.g. in nitromethane, ozone). The goal is to minimise them, not eliminate them at any cost.
⚠️ Drawing lone pairs on every heteroatom (O, N, halogens) is tested constantly. Leaving them off costs marks and leads to wrong formal charges.
⚠️ Formal charge calculation is a staple exam question. You will be asked to assign formal charges to atoms in a given structure, or to choose the "best" Lewis structure from several options.
⚠️ Knowing typical valencies cold (C = 4, N = 3, O = 2, H/halogens = 1) lets you spot errors instantly and speeds up drawing under time pressure.
⚠️ Practice problems in Chem 210 frequently ask you to draw all lone pairs on a given line-angle structure. If you cannot do this quickly, mechanisms will be painful.
True or False: Oxygen typically forms three bonds in a neutral Lewis structure.
False. Oxygen is divalent (two bonds).
Fill in the blank: The formal charge on an atom is calculated as valence electrons minus ______ minus half the bonding electrons.
lone-pair electrons
True or False: Carbon can form a maximum of four bonds.
True.
Fill in the blank: Halogens (F, Cl, Br, I) are ______ and form one bond.
monovalent
True or False: A Lewis structure with non-zero formal charges is always an incorrect structure.
False. Formal charges are sometimes unavoidable; the goal is to minimise them.
Q: What is the formal charge on the nitrogen atom in the ammonium ion (NH₄⁺)?
A: Nitrogen has 5 valence electrons. In NH₄⁺ it has 0 lone-pair electrons and 8 bonding electrons (4 bonds). Formal charge = 5 − 0 − ½(8) = +1.
Q: Draw the Lewis structure of CH₂O (formaldehyde). How many lone pairs does the oxygen carry?
A: Carbon forms a double bond to oxygen and two single bonds to hydrogen. Oxygen carries two lone pairs. Total valence electrons = 4 + 2(1) + 6 = 12. The C=O double bond uses 4 electrons, the two C−H bonds use 4, and 4 remain as two lone pairs on O.
Q: A student draws a Lewis structure in which carbon has five bonds. What is wrong?
A: Carbon is tetravalent. It can form a maximum of four bonds (eight electrons in its valence shell). Five bonds would require ten electrons around carbon, which violates the octet rule for second-row elements.
Q: In the molecule HCl, how many non-bonding electron pairs does chlorine have?
A: Three. Chlorine has 7 valence electrons; one is used in the bond with H, leaving 6 as three lone pairs.
Q: Why does hydrogen follow the duet rule rather than the octet rule?
A: Hydrogen's valence shell is the 1s orbital, which can hold a maximum of 2 electrons. Filling that orbital gives hydrogen a noble-gas (helium) configuration.
This connects directly to resonance structures (Topic 4 in most Chem 210 sequences), because resonance is just the idea that some molecules can be represented by more than one valid Lewis structure. You cannot draw resonance structures without being comfortable with lone pairs and formal charges first.
Formal charge also feeds into acid-base chemistry: the stability of a conjugate base depends on where the negative charge sits, which you assess through formal charge and electronegativity arguments.
Units of unsaturation (covered in the companion notes) use the connectivity rules from this topic to work backwards from a molecular formula to the number of rings and pi bonds a molecule must contain.
Lewis dot structure, electron dot diagram, Lewis dot diagram, bonding pair, shared pair, lone pair, non-bonding electrons, octet rule, duet rule, formal charge, formal charge calculation, valence electrons, valency, monovalent, divalent, trivalent, tetravalent, covalent bonding, covalent molecule, connectivity, skeletal structure, line-angle structure, isomers, structural isomers, Chem 210, organic chemistry II, University of Michigan