Source: Chemistry by Overby, Chapters 9.1, 9.6, 9.8, 10.1, 10.2
Difficulty: Intermediate | Prerequisites: Chapter 2 naming, periodic table groups, electron configurations, valence electrons.
This material bridges naming (Chapter 2) and quantitative chemistry (Chapter 3). Once you know what a compound is called, the next question is: how are the atoms actually connected, and what shape does the molecule take? Lewis structures show you the electron arrangement. VSEPR theory uses that arrangement to predict three-dimensional shape. Shape then determines whether the molecule is polar or nonpolar, which governs physical properties like boiling point and solubility. If you skipped valence electrons or electron configurations, go back and review those first, because every step here depends on counting electrons correctly.
Lewis dot symbols show valence electrons around an atom. Lewis structures show how atoms in a molecule share or hold lone pairs of electrons, and some molecules have resonance structures (multiple valid arrangements). VSEPR theory predicts 3D shape by minimising electron-pair repulsion. Shape plus bond polarity tells you whether the whole molecule is polar or nonpolar.
Lewis dot symbol
A representation of an atom showing its chemical symbol surrounded by dots for each valence electron. Think of it as a quick visual count of the electrons available for bonding.
Lewis structure (Lewis formula)
A diagram showing all valence electrons in a molecule or ion as bonding pairs (lines) and lone pairs (dots). In simple terms, it is the blueprint of how atoms share electrons.
Bonding pair
A pair of electrons shared between two atoms, represented as a line in a Lewis structure.
Lone pair (nonbonding pair)
A pair of valence electrons on an atom that is not shared with another atom. Think of it as electrons that belong to one atom alone and take up space around it.
Resonance structures
Two or more valid Lewis structures for the same molecule or ion that differ only in the placement of electrons (not atoms). The real structure is a blend (resonance hybrid) of all of them. In simple terms, the electrons are delocalised, and no single drawing captures the full picture.
VSEPR theory (valence-shell electron-pair repulsion)
A model that predicts molecular shape by assuming that electron groups around a central atom arrange themselves as far apart as possible to minimise repulsion. Think of it as: electron clouds act like balloons tied together, pushing each other away.
Electron geometry (parent structure)
The arrangement of all electron groups (bonding and lone pairs) around the central atom.
Molecular geometry (molecular shape)
The arrangement of only the atoms in a molecule (ignoring lone pairs for the visual shape). This is what you report as the "shape" of the molecule.
Bond angle
The angle formed between two bonds originating from the same central atom.
Dipole moment
A measure of the separation of positive and negative charge in a molecule. A molecule with a net dipole moment is polar; one where dipoles cancel is nonpolar.
Polar molecule
A molecule with an uneven distribution of electron density, resulting in a net dipole moment. In simple terms, one end is slightly positive and the other slightly negative.
Nonpolar molecule
A molecule where bond dipoles cancel out due to symmetry, giving no net dipole moment.
Each dot represents one valence electron. The number of valence electrons equals the group number (for main-group elements using the 1A-8A system).
Place dots singly on four sides of the symbol before pairing any. This matters for understanding how many bonds an atom can form.
For monatomic ions: Add dots for gained electrons (anions) or remove dots for lost electrons (cations), then show the charge.
Step-by-step method:
Count total valence electrons. Add up valence electrons for every atom. For polyatomic anions, add electrons equal to the charge. For cations, subtract.
Place the least electronegative atom in the centre. Hydrogen and fluorine are always terminal (outer) atoms. Carbon is almost always central.
Connect outer atoms to the central atom with single bonds. Each bond uses 2 electrons.
Distribute remaining electrons as lone pairs. Start with the outer atoms, giving each an octet (hydrogen gets 2). Place leftover electrons on the central atom.
If the central atom lacks an octet, form multiple bonds. Convert lone pairs from an outer atom into shared pairs (double or triple bonds) until the central atom has 8 electrons.
Shared (bonding) electrons: Electrons in a bond (line). A single bond = 2 shared electrons, a double bond = 4, a triple bond = 6.
Unshared (lone pair) electrons: Dots sitting on one atom, not between two atoms.
Resonance occurs when you can draw more than one valid Lewis structure by moving electrons (not atoms).
Classic examples: ozone (O3), nitrate (NO3-), carbonate (CO3 2-).
The actual molecule is a resonance hybrid: all resonance structures contribute. Bond lengths and strengths are intermediate between the extremes.
To spot resonance, look for a situation where a double bond could be placed in more than one position equivalently.
The seven diatomic elements (H2, N2, O2, F2, Cl2, Br2, I2) each have characteristic Lewis structures. N2 has a triple bond, O2 has a double bond, and the halogens have single bonds with three lone pairs on each atom.
Count the number of electron groups around the central atom. An electron group is any of: a single bond, a double bond, a triple bond, or a lone pair. (A double bond counts as one group, not two.)
The electron groups arrange to be as far apart as possible.
Electron Groups | Lone Pairs | Electron Geometry | Molecular Shape | Bond Angle |
|---|---|---|---|---|
2 | 0 | Linear | Linear | 180° |
3 | 0 | Trigonal planar | Trigonal planar | 120° |
3 | 1 | Trigonal planar | Bent | < 120° |
4 | 0 | Tetrahedral | Tetrahedral | 109.5° |
4 | 1 | Tetrahedral | Trigonal pyramidal | < 109.5° |
4 | 2 | Tetrahedral | Bent | < 109.5° |
You must know the shapes and names in Tables 10.1 and 10.2 (up through tetrahedral). Lone pairs compress bond angles slightly below the ideal values because lone pairs repel more strongly than bonding pairs.
Three-step process:
Draw the Lewis structure and determine the molecular shape via VSEPR.
Assess individual bond polarities. A bond between atoms with different electronegativities is polar. The dipole arrow points from the less electronegative atom towards the more electronegative one.
Check whether bond dipoles cancel. In symmetrical shapes (linear with identical outer atoms, trigonal planar with identical outer atoms, tetrahedral with identical outer atoms), the dipoles cancel and the molecule is nonpolar. In asymmetric shapes or molecules with lone pairs on the central atom, dipoles usually do not cancel, making the molecule polar.
CO2: linear, two equal C=O dipoles point in opposite directions, cancel. Nonpolar.
H2O: bent shape, two O-H dipoles do not cancel. Polar.
CCl4: tetrahedral, four identical C-Cl dipoles cancel by symmetry. Nonpolar.
CHCl3: tetrahedral but not all bonds identical, dipoles do not cancel. Polar.
Total valence electron count formula:
Total e- = (sum of valence electrons for each atom) + (charge added for anions) - (charge removed for cations)
Octet rule: Most main-group atoms in a Lewis structure will be surrounded by 8 electrons (4 pairs). Hydrogen is the exception: it needs only 2.
Electronegativity trend (for polarity): Increases going up and to the right on the periodic table. Fluorine is the most electronegative element.
Molecular shape determines how drugs bind to receptors in your body: a molecule with the wrong geometry simply will not fit. Polarity determines whether a substance dissolves in water (polar solvent) or in oil (nonpolar solvent), which is the basis of everything from soap chemistry to drug delivery. Engineers designing new materials need to predict polarity to control intermolecular forces, melting points, and solubility.
Counting a double bond as two electron groups. A double bond (or triple bond) counts as one electron group in VSEPR. Only the number of groups matters, not the number of electron pairs within each group.
Confusing electron geometry with molecular geometry. Electron geometry includes lone pairs in the arrangement. Molecular geometry describes only where the atoms sit. A molecule with four electron groups but one lone pair has tetrahedral electron geometry but trigonal pyramidal molecular geometry.
Assuming all molecules with polar bonds are polar. CO2 has two polar C=O bonds, but the linear shape means the dipoles cancel. Symmetry can make a molecule with polar bonds nonpolar overall.
Forgetting to add or subtract electrons for ions. When drawing Lewis structures for polyatomic ions, students frequently forget to adjust the electron count for the charge. This throws off the entire structure.
⚠️ Drawing Lewis structures is almost guaranteed on the exam. Practise the five-step method until it is automatic.
⚠️ Expect a question asking you to identify whether a molecule has resonance and to draw all resonance structures.
⚠️ VSEPR: you need to go from Lewis structure to electron geometry to molecular shape to bond angle. Practise the full chain, not just one step.
⚠️ Polarity questions require the whole pipeline: Lewis structure, shape, bond dipoles, then the final polar/nonpolar call. A question that asks "is this molecule polar?" is really testing all of Chapters 9 and 10 at once.
⚠️ Know the geometry table up through tetrahedral cold. If you can recall it instantly, you save time on every VSEPR question.
True or false: A triple bond counts as three electron groups in VSEPR. (False. It counts as one.)
Fill in the blank: A molecule with four electron groups and no lone pairs has a ______ shape. (Tetrahedral.)
True or false: H2O is nonpolar because it contains only two bonds. (False. The bent shape gives it a net dipole.)
Fill in the blank: Resonance structures differ only in the placement of ______, not atoms. (Electrons.)
True or false: The molecular geometry of NH3 is trigonal pyramidal. (True.)
Q: Draw the Lewis structure for CO2 and state its molecular geometry.
A: Carbon is central with double bonds to each oxygen. Each oxygen has two lone pairs. Molecular geometry: linear (two electron groups, no lone pairs, 180° bond angle).
Q: How many total valence electrons are in the nitrate ion, NO3-?
A: N has 5, each O has 6, plus 1 for the negative charge = 5 + 18 + 1 = 24 valence electrons.
Q: Does NO3- exhibit resonance? If so, how many resonance structures?
A: Yes. Three equivalent resonance structures, each with the double bond on a different N-O position.
Q: Predict the molecular geometry of H2O and state whether it is polar or nonpolar.
A: Oxygen has four electron groups (two bonding pairs, two lone pairs). Electron geometry is tetrahedral; molecular geometry is bent. The two O-H bond dipoles do not cancel in a bent shape, so H2O is polar.
Q: Is CCl4 polar or nonpolar? Explain.
A: Nonpolar. It has a tetrahedral shape with four identical C-Cl bonds. The symmetry causes all bond dipoles to cancel.
Q: What is the electron geometry and molecular geometry of a molecule with three electron groups and one lone pair on the central atom?
A: Electron geometry is trigonal planar. Molecular geometry is bent (the lone pair occupies one position but is not "seen" as an atom).
Lewis structures build on the naming material in Chapter 2: you need to know whether a compound is ionic or molecular before deciding whether to draw a Lewis structure (ionic compounds use electron transfer, not shared pairs). Molecular geometry and polarity connect forward to intermolecular forces, which govern boiling points, solubility, and phase behaviour in later chapters. The electron-counting skills here are also essential for Chapter 3 mole calculations, because you need correct formulas to compute molar mass.
Tags: Lewis dot structure, Lewis formula, electron dot diagram, valence electrons, bonding pair, lone pair, nonbonding pair, resonance structures, resonance hybrid, VSEPR theory, electron geometry, molecular geometry, molecular shape, bond angle, linear, trigonal planar, bent, tetrahedral, trigonal pyramidal, dipole moment, polar molecule, nonpolar molecule, polarity, electronegativity, octet rule, diatomic elements, CHM 11100, Purdue general chemistry, Overby chemistry, chemical bonding