Source: Chapter 1, Dr Uyeda lecture slides
Tags: Lewis structure, formal charge, resonance, VSEPR, molecular geometry, bond angle, dipole moment, polar molecule, line drawing
Difficulty: Introductory – Intermediate Prerequisites: Atomic Structure and Bonding Types notes (Part 1 of this chapter). You need to be comfortable with valence electrons, electronegativity, and the octet rule.
Once you know where electrons sit in isolated atoms, the next step is to draw how they arrange themselves in molecules. Lewis structures are the single most-used tool for this in organic chemistry, and formal charge tells you whether your drawing is plausible. Some molecules need more than one Lewis structure to capture their real electron distribution, which is where resonance comes in. Finally, VSEPR lets you predict the 3D shape of a molecule from its Lewis structure, which in turn determines whether the molecule is polar. All of this feeds directly into reaction mechanisms from Chapter 2 onward.
Lewis structures depict valence electrons as dots or lines. Formal charge checks whether each atom "owns" the right number of electrons. Resonance structures are needed when a single Lewis drawing cannot represent the real electron distribution. VSEPR predicts 3D geometry by maximising the distance between electron groups around a central atom.
Lewis dot structure (Lewis structure)
A diagram showing only the valence electrons of each atom in a molecule, drawn as dots (lone pairs) or lines (shared pairs/bonds). Introduced by G. N. Lewis in 1916. Think of it as the molecular blueprint that tells you where every valence electron lives.
Lone pair
A pair of valence electrons on an atom that is not shared with another atom. Drawn as two dots. In simple terms, these are electrons that belong entirely to one atom and are not involved in bonding.
Bonding pair
A pair of electrons shared between two atoms, forming a covalent bond. Drawn as a line (–) or two dots between the atoms.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: (number of valence electrons in the free atom) – (number of lone-pair electrons) – (half the number of bonding electrons). Think of it as an accounting check: does each atom have the electron count it started with?
Resonance structures
Two or more valid Lewis structures for the same molecule that differ only in the placement of electrons (not atoms). The real molecule is a weighted average (resonance hybrid) of all contributors. In simple terms, the true electron distribution is a blend of these drawings.
Resonance hybrid
The actual structure of a molecule for which resonance structures can be drawn. Bond lengths and charges are intermediate between those shown in individual resonance contributors.
VSEPR (valence-shell electron-pair repulsion)
A model for predicting the 3D shape of a molecule. Electron groups (bonds and lone pairs) around a central atom repel each other and arrange themselves to be as far apart as possible.
Bond angle
The angle formed between two bonds that share a common atom. Predicted by VSEPR based on the number of electron groups.
Dipole moment
A measure of the overall polarity of a molecule. It is the vector sum of all individual bond dipoles. If bond dipoles cancel (due to molecular symmetry), the molecule has no net dipole moment and is non-polar.
Line drawing (skeletal structure)
A simplified way to draw organic molecules. Lines represent bonds, vertices and endpoints represent carbon atoms, and hydrogen atoms on carbon are implied (not drawn). Heteroatoms (O, N, etc.) and their hydrogens are always shown explicitly.
Count the total number of valence electrons from all atoms in the molecule. For ions, add electrons for negative charges and subtract for positive charges.
Arrange atoms with the least electronegative atom in the centre (hydrogen is always terminal).
Place single bonds between each pair of connected atoms. Each bond uses 2 electrons.
Distribute the remaining electrons as lone pairs on outer atoms first, satisfying octets (duets for hydrogen).
If the central atom lacks an octet, convert lone pairs on outer atoms into additional bonds (double or triple).
Atom | Typical bonds | Lone pairs |
|---|---|---|
C | 4 | 0 |
N | 3 | 1 |
O | 2 | 2 |
F / halogens | 1 | 3 |
H | 1 | 0 |
These are quick checks, not laws. Charged atoms or unusual bonding environments can change the count.
Total valence electrons: 2(1) + 6 = 8. Oxygen is central. Place O–H, O–H (4 electrons used). Remaining 4 electrons go as two lone pairs on oxygen. Every atom has its target: O has 8 electrons around it, each H has 2.
Total valence electrons: 4 + 2(6) = 16. Carbon is central. Starting with single bonds (O–C–O) uses 4 electrons and leaves 12. Placing 6 on each O gives them octets but leaves carbon with only 4. Convert one lone pair from each O into a bond: O=C=O. Each atom now has 8 electrons.
Formal charge = (valence electrons of the free atom) – (lone pair electrons) – (½ × bonding electrons).
The sum of all formal charges in a molecule must equal the molecule's overall charge.
Oxygen in H₃O⁺ has 1 lone pair (2 electrons) and 3 bonds (6 bonding electrons). Formal charge on O = 6 – 2 – ½(6) = +1. Each hydrogen: 1 – 0 – ½(2) = 0. Total formal charge: +1, matching the ion's charge.
Oxygen has 3 lone pairs (6 electrons) and 1 bond (2 bonding electrons). Formal charge on O = 6 – 6 – ½(2) = –1. Hydrogen: 1 – 0 – ½(2) = 0. Total: –1, matching the ion's charge.
Some molecules cannot be accurately represented by a single Lewis structure. The real molecule is a hybrid (a blend) of two or more resonance contributors. Resonance structures differ only in electron placement; the atoms stay in the same positions.
The total number of valence electrons must remain the same across all structures.
No atom may exceed 8 electrons in its valence shell (the octet limit).
Atomic positions (the molecular skeleton) do not change. Only electrons move.
Curved arrows track where electron pairs move when converting one resonance structure to another. The tail of the arrow starts at the electrons; the head points to where they go.
Not all resonance structures contribute equally. Preference goes to structures with:
Complete octets on every atom.
The greatest number of covalent bonds.
The least separation of unlike charges.
Negative formal charge on the more electronegative atom and positive formal charge on the less electronegative atom.
Formaldehyde (H₂CO) has a single Lewis structure with a C=O double bond. Formate (HCO₂⁻) has two equivalent resonance structures: one with a C=O double bond to the top oxygen and a C–O single bond to the bottom, and one with the bonding pattern reversed. In the real ion, both C–O bonds are identical at 1.26 Å, intermediate between a typical C–O single bond (1.4 Å) and a C=O double bond (1.2 Å).
Formic acid has two resonance structures, but they are not equivalent. The major contributor has the C=O double bond and no formal charges. The minor contributor places a positive formal charge on carbon and a negative one on the singly bonded oxygen. The major contributor is preferred because it has more covalent bonds and no charge separation.
N₂O has 16 valence electrons and several possible resonance structures. The major contributor places a triple bond between the two nitrogen atoms (N≡N–O) with formal charges of –1 on the terminal N and +1 on the central N, and –1 on O. This is preferred because it has the most bonds, complete octets, and places the negative charge on the more electronegative atom (oxygen). The real molecule is linear with bond lengths of 1.13 Å (N–N) and 1.19 Å (N–O).
Electron groups (bonds and lone pairs) around a central atom repel each other. The molecule adopts the geometry that places these groups as far apart as possible, maximising the angle between them.
Electron regions | Geometry | Bond angle | Examples |
|---|---|---|---|
4 | Tetrahedral | 109.5° | CH₄, NH₃ (with one lone pair, shape is trigonal pyramidal) |
3 | Trigonal planar | 120° | BH₃, H₂C=CH₂ (each sp² carbon) |
2 | Linear | 180° | CO₂, HC≡CH |
The electron geometry counts all electron groups (bonds + lone pairs).
The molecular shape describes only the arrangement of atoms (what you see).
NH₃ has 4 electron groups (tetrahedral electron geometry) but its molecular shape is trigonal pyramidal because one of the four groups is a lone pair.
For 3D drawings: a solid wedge (▸) means a bond coming out of the plane towards you; a dashed wedge (╌) means a bond going behind the plane away from you; a plain line is in the plane.
Boron often does not follow the octet rule. In BH₃, boron has only 6 electrons and adopts a trigonal planar geometry with 120° bond angles.
A molecule's overall polarity depends on two things: whether its individual bonds are polar, and whether the molecular geometry allows those bond dipoles to cancel.
Each O–H bond is polar (Δ ≈ 1.4). Water has a bent shape (about 104.5°), so the two bond dipoles do not point in opposite directions. They add together to produce a net dipole moment. Water is a polar molecule, which is why it is a liquid at room temperature.
Each C=O bond is polar (Δ = 1.0). CO₂ is linear (180°), so the two bond dipoles point in exactly opposite directions and cancel out. The molecule has no net dipole moment and is non-polar, which is why CO₂ is a gas at room temperature.
Students often think resonance structures represent the molecule flipping back and forth between two forms. They do not. The molecule exists as a single hybrid at all times; the individual structures are just our imperfect way of drawing it.
Students frequently forget to check that the total electron count is the same across resonance structures. If your new structure has gained or lost electrons, it is not a valid resonance contributor.
Students sometimes draw resonance structures that move atoms. Only electrons move in resonance. If you have to break and reform a sigma bond to a different atom, you have drawn an isomer, not a resonance structure.
Students often confuse electron geometry with molecular shape. A molecule with 4 electron groups always has tetrahedral electron geometry, but its molecular shape depends on how many of those groups are lone pairs.
⚠️ Drawing Lewis structures and assigning formal charges is tested heavily and is the foundation for every mechanism problem in the course.
⚠️ Expect questions asking you to draw all reasonable resonance structures for a molecule and identify the major contributor using the four ranking rules.
⚠️ VSEPR questions will ask you to predict bond angles and molecular geometry from a Lewis structure. Know the angles: 109.5° (tetrahedral), 120° (trigonal planar), 180° (linear).
⚠️ You may be asked whether a molecule is polar or non-polar. The answer requires both bond polarity and geometry.
⚠️ Line drawings are the standard notation for organic chemistry. Be able to convert between a full Lewis structure and a skeletal (line) drawing, keeping track of implied hydrogen atoms.
True or false: in a Lewis structure, a line represents a lone pair.
Fill in the blank: the formal charge formula is ___ minus ___ minus ½ × ___.
True or false: resonance structures can differ in the positions of atoms.
Fill in the blank: a molecule with 4 electron groups around a central atom has a bond angle of approximately ___°.
True or false: CO₂ has polar bonds but is a non-polar molecule.
Answers: 1. False (a line represents a bonding pair / shared pair). 2. Valence electrons of the free atom, lone-pair electrons, bonding electrons. 3. False (only electron placement changes). 4. 109.5. 5. True (the bond dipoles cancel due to linear geometry).
Q: Draw the Lewis structure for HNO₃. How many total valence electrons does the molecule have?
A: Total valence electrons = 1 + 5 + 3(6) = 24. Nitrogen is the central atom. One valid structure places a double bond between N and one oxygen, single bonds to the other two oxygens (one of which carries the H). Nitrogen bears a +1 formal charge and the singly-bonded oxygen without H bears a –1 formal charge.
Q: For the formate ion (HCO₂⁻), draw both resonance structures and explain why the two C–O bond lengths are equal.
A: One structure has C=O to the top oxygen and C–O⁻ to the bottom; the other reverses this. Because the two structures are equivalent, they contribute equally to the hybrid. The real C–O bond length (1.26 Å) is between a single bond (1.4 Å) and a double bond (1.2 Å), and both C–O bonds are identical.
Q: Predict the molecular geometry of NH₃ and explain why it is not trigonal planar.
A: Nitrogen in NH₃ has 4 electron groups: three N–H bonds and one lone pair. VSEPR says 4 groups adopt a tetrahedral electron arrangement. Because one group is a lone pair (not visible as a bond), the molecular shape is trigonal pyramidal, not trigonal planar.
Q: Is CHCl₃ (chloroform) a polar molecule? Explain.
A: Yes. Carbon is tetrahedral with three C–Cl bonds and one C–H bond. The three C–Cl bond dipoles do not cancel due to the asymmetric arrangement (one position is occupied by H instead of Cl), so the molecule has a net dipole moment.
Q: Convert CH₃CH₂OH (ethanol) into a line (skeletal) drawing.
A: Draw a zig-zag with two line segments (representing two carbon atoms). At the end of the second segment, write OH. The hydrogens on carbon are implied and not shown.
Lewis structures and formal charge are prerequisites for understanding acid-base chemistry in Chapter 2. Resonance appears throughout the course whenever you need to explain charge delocalisation, particularly in aromatic systems, carboxylate anions, and amides. VSEPR and hybridisation (covered in the next set of notes) are tightly linked: the geometry predicted by VSEPR maps directly onto sp³, sp², and sp hybridisation.
Lewis structure, Lewis dot diagram, formal charge, resonance, resonance hybrid, resonance contributor, curved arrow, arrow pushing, electron pushing, VSEPR, valence shell electron pair repulsion, molecular geometry, bond angle, tetrahedral, trigonal planar, linear, trigonal pyramidal, bent, dipole moment, polar molecule, non-polar molecule, line drawing, skeletal structure, line-angle drawing, wedge and dash, lone pair, bonding pair, octet rule, CHEM 25500, organic chemistry Purdue, Chapter 1