Difficulty: Introductory | Prerequisites: General chemistry (atomic structure, periodic trends)
Lewis structures are diagrams that show every valence electron in a molecule, both shared (bonding) and unshared (lone pairs). Getting them right is the foundation for understanding molecular shape, polarity, and reactivity in organic chemistry. Once you can draw a correct Lewis structure, you can assign formal charges, predict which bonds are polar, and work out where electron density sits, all of which feeds directly into resonance, hybridisation, and reaction mechanisms later in the course.
Lewis structure (electron-dot structure)
A diagram showing all valence electrons in a molecule as dots or lines (bonds). Lone pairs are drawn as paired dots on the atom; bonding pairs are drawn as lines between atoms.
In simple terms: it is the blueprint that tells you where every outer-shell electron lives.
Valence electrons
The electrons in an atom's outermost shell, which participate in bonding. Carbon has 4, nitrogen has 5, oxygen has 6, each halogen has 7, and hydrogen has 1.
Think of these as the currency atoms spend to form bonds.
Lone pair (non-bonding pair)
A pair of valence electrons on an atom that is not shared with another atom. Oxygen in water has two lone pairs; each one occupies space and affects molecular geometry.
In simple terms: electrons that belong to one atom and stay home.
Bonding pair
A pair of electrons shared between two atoms, forming a covalent bond. A single bond is one bonding pair; a double bond is two; a triple bond is three.
Octet rule
Most main-group atoms (C, N, O, halogens) are stable when surrounded by eight electrons (four pairs). Hydrogen is the exception: it needs only two electrons (a duet).
Think of it as each atom wanting a full outer shell, and eight is the magic number for row-2 elements.
Formal charge
A bookkeeping charge assigned to an atom in a Lewis structure. Calculated as: valence electrons minus lone-pair electrons minus half the bonding electrons. A formal charge of zero on every atom is ideal; structures that minimise formal charges are generally more stable.
In simple terms: it tells you whether an atom "owns" more or fewer electrons than it brought to the party.
Electronegativity
A measure of how strongly an atom attracts shared electrons towards itself. Fluorine is the most electronegative element. Electronegativity increases across a period (left to right) and up a group on the periodic table.
Bond polarity
When two bonded atoms differ in electronegativity, the shared electrons are pulled closer to the more electronegative atom, creating a partial negative charge (δ–) on that atom and a partial positive charge (δ+) on the other. The bond is said to be polar.
Think of it as a tug-of-war for electrons, where the stronger puller gets more than its share.
Dipole
A separation of charge within a bond or molecule. A bond dipole points from the δ+ atom towards the δ– atom. The overall molecular dipole is the vector sum of all bond dipoles.
Count total valence electrons. Add up the valence electrons for every atom. For anions, add one electron per negative charge; for cations, subtract one per positive charge.
Example: NO₃⁻ has 5 (N) + 3×6 (O) + 1 (charge) = 24 valence electrons.
Draw the skeleton. Place the least electronegative atom in the centre (never hydrogen). Connect each outer atom to the centre with a single bond.
Distribute remaining electrons as lone pairs. Start with the outer atoms; give each an octet (or duet for H). Then place leftover electrons on the central atom.
Check octets. If the central atom lacks an octet, convert one or more lone pairs on an adjacent atom into bonding pairs (forming double or triple bonds).
Assign formal charges. Use the formula (see Formulas section below). Minimise formal charges where possible; place any unavoidable negative formal charges on the more electronegative atom.
Dichloromethane, CH₂Cl₂
Total valence electrons: 4 (C) + 2×1 (H) + 2×7 (Cl) = 20.
Carbon is central, bonded to two H and two Cl atoms by single bonds.
Each Cl carries three lone pairs (six electrons). H has no lone pairs.
All formal charges are zero. Carbon's octet is satisfied by four single bonds (eight shared electrons).
Hydrazine, N₂H₄
Total valence electrons: 2×5 (N) + 4×1 (H) = 14.
The two nitrogen atoms are bonded to each other by a single bond. Each N also bonds to two H atoms.
Each nitrogen carries one lone pair.
All formal charges are zero.
Note: the worksheet answer key shows a triple bond between the nitrogens, but the correct Lewis structure for hydrazine (N₂H₄) has a single N–N bond with a lone pair on each nitrogen.
Nitrate ion, NO₃⁻
Total valence electrons: 5 (N) + 3×6 (O) + 1 (charge) = 24.
Nitrogen is central, bonded to three oxygen atoms.
One N–O bond is a double bond; the other two are single bonds (in any given resonance structure).
Formal charges: the double-bonded O has formal charge 0; each single-bonded O has formal charge –1; nitrogen has formal charge +1.
Three equivalent resonance structures exist (the double bond can sit on any of the three oxygens).
Oxygen atoms: the hydroxyl (–OH) oxygen has two lone pairs; the carbonyl (C=O) oxygen has two lone pairs.
Nitrogen atom: the nitrogen carries one lone pair (it forms three bonds: one to the aromatic ring carbon, one to the carbonyl carbon, and one to hydrogen).
Polarity of highlighted bonds:
C–O bond: oxygen is more electronegative, so δ– sits on O and δ+ on C.
O–H bond: oxygen is more electronegative, so δ– on O and δ+ on H.
N–H bond: nitrogen is more electronegative than hydrogen, so δ– on N and δ+ on H.
Line (skeletal) structures hide carbon atoms and most hydrogen atoms. To convert:
Every vertex and endpoint in a line structure is a carbon atom.
Add hydrogens to each carbon until it has four bonds total.
Write the molecular formula by counting all C, H, and heteroatoms.
Worksheet examples:
(a) A four-carbon alcohol: molecular formula C₄H₈O (think: butanol-like).
(b) A six-carbon amine: molecular formula C₆H₁₅N.
(c) A cyclic ester structure: molecular formula C₆H₁₀O₂.
\text{Formal Charge} = (\text{Valence electrons}) - (\text{Lone pair electrons}) - \tfrac{1}{2}(\text{Bonding electrons})Common valence electron counts (memorise these):
Element | Group | Valence Electrons | Typical Bonds | Typical Lone Pairs |
|---|---|---|---|---|
H | 1A | 1 | 1 | 0 |
C | 4A | 4 | 4 | 0 |
N | 5A | 5 | 3 | 1 |
O | 6A | 6 | 2 | 2 |
F, Cl, Br, I | 7A | 7 | 1 | 3 |
Bond polarity determines how drug molecules interact with proteins. Tylenol (acetaminophen) works partly because its polar N–H and O–H bonds can form hydrogen bonds with the COX enzyme active site. Without understanding where partial charges sit, you cannot predict binding, solubility, or reactivity in pharmaceutical design.
Lewis structures also matter in materials science: the lone pairs on nitrogen in nylon are what allow hydrogen bonding between polymer chains, giving the material its strength.
Forgetting lone pairs on inner atoms. Students often fill octets on outer atoms and forget the central atom can also carry lone pairs (nitrogen in amines, for example).
Putting hydrogen in the centre. Hydrogen can only form one bond. It is always a terminal atom, never the central atom in a Lewis structure.
Confusing 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.
Drawing polarity arrows backwards. The δ– goes on the more electronegative atom. A common slip is labelling carbon as δ– in a C–O bond.
⚠️ Lewis structures are the basis for nearly every topic that follows in organic chemistry. If you cannot draw them quickly and correctly, resonance, hybridisation, and reaction mechanisms will be significantly harder.
⚠️ Formal charge calculation appears on virtually every introductory organic exam. Know the formula cold.
⚠️ Bond polarity (δ+/δ– notation) is tested both on its own and as part of predicting nucleophilic and electrophilic sites. Expect to be asked to mark polarity on a structure like Tylenol.
⚠️ Converting line structures to full structures (showing all C and H atoms) and writing molecular formulas is a bread-and-butter exam skill. Speed matters here.
True or False: Hydrogen can have a lone pair in a valid Lewis structure.
Answer: False. Hydrogen follows the duet rule and can hold at most two electrons (one bond, zero lone pairs).
Fill in the blank: The formal charge on an atom equals its valence electrons minus its lone-pair electrons minus ______ its bonding electrons.
Answer: half.
True or False: In a C–O bond, the δ– sits on carbon because carbon has more electrons.
Answer: False. The δ– sits on oxygen because oxygen is more electronegative.
Fill in the blank: Nitrogen has ______ valence electrons and typically forms ______ bonds with one lone pair.
Answer: 5; 3.
True or False: The nitrate ion (NO₃⁻) has three equivalent resonance structures.
Answer: True.
Q: Draw the Lewis structure for dichloromethane (CH₂Cl₂). How many lone pairs does each chlorine carry?
A: Carbon is central with single bonds to two H atoms and two Cl atoms. Each chlorine carries three lone pairs (six non-bonding electrons). All formal charges are zero.
Q: Calculate the formal charge on nitrogen in the nitrate ion (NO₃⁻) where nitrogen is double-bonded to one oxygen and single-bonded to the other two.
A: Nitrogen has 5 valence electrons, 0 lone-pair electrons, and 8 bonding electrons (one double bond = 4e, two single bonds = 4e). Formal charge = 5 – 0 – 4 = +1.
Q: In acetaminophen (Tylenol), which atom in the C=O bond carries the δ– and why?
A: Oxygen carries the δ– because it is more electronegative than carbon. It pulls the shared electrons closer to itself.
Q: A line structure shows a chain with four vertices, an OH group at one end, and no other heteroatoms. What is the molecular formula?
A: Four vertices plus the endpoint carbons can give different counts depending on the drawing, but for a simple four-carbon alcohol: C₄H₈O (or similar, depending on branching and saturation). The key skill is counting every implicit carbon at each vertex and endpoint, then adding hydrogens until each carbon has four bonds.
Q: Why does nitrogen in hydrazine (N₂H₄) carry a lone pair even though it could form another bond?
A: Each nitrogen already forms three bonds (one N–N and two N–H), giving it six bonding electrons plus one lone pair (two electrons) for a total of eight. The octet is satisfied, so the lone pair stays.
This connects to resonance because once you can draw one valid Lewis structure, you can push electrons to generate alternative resonance contributors (covered in Part 2 of these notes). The formal charges you assign here determine which resonance structure is the major contributor.
This also connects to VSEPR and molecular geometry: lone pairs occupy space around an atom and change bond angles. You need to know where every lone pair sits (from the Lewis structure) before you can predict shape.
Looking ahead, hybridisation (sp³, sp², sp) is assigned based on the number of electron groups around an atom, and you read those groups straight off the Lewis structure.
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