Difficulty: Introductory to Intermediate | Prerequisites: Periodic table familiarity, electron configurations, electronegativity basics.
Lewis structures are the foundation of how chemists visualise bonding. Once you can draw them, VSEPR (Valence Shell Electron Pair Repulsion) theory lets you predict the three-dimensional shape of any molecule or ion, which in turn explains polarity, reactivity, and physical properties. This material builds directly on electron configurations and valence electrons from earlier in the course. If you are comfortable counting valence electrons and understand electronegativity, you have what you need to start here.
Lewis structures show how valence electrons are distributed as bonds and lone pairs around atoms. VSEPR theory uses the number of electron groups (bonding + lone pairs) on a central atom to predict electron geometry and molecular geometry. Exceptions to the octet rule (expanded octets, incomplete octets, odd-electron species) come up for elements in Period 3+ or electron-deficient atoms like boron.
Lewis structure
A diagram showing all valence electrons in a molecule or ion as dots or lines (bonds). Think of it as the blueprint for where electrons live.
Valence electrons
The outermost electrons of an atom, which participate in bonding. The group number on the periodic table tells you how many an atom has (for main-group elements).
Bonding pair
A pair of electrons shared between two atoms, drawn as a line. In simple terms, this is the glue holding two atoms together.
Lone pair (non-bonding pair)
A pair of valence electrons sitting on one atom, not shared. These still take up space and affect molecular shape.
Electron geometry
The arrangement of all electron groups (bonding pairs + lone pairs) around the central atom. This describes the full spatial layout, including invisible lone pairs.
Molecular geometry
The arrangement of only the atoms (ignoring lone pairs) around the central atom. This is the actual shape you would see if you could photograph the molecule.
VSEPR theory (Valence Shell Electron Pair Repulsion)
The model that predicts geometry by assuming electron groups repel each other and spread out as far as possible. In simple terms, electron clouds push away from each other like balloons tied together.
Resonance structures
Two or more valid Lewis structures for the same molecule or ion, differing only in the placement of electrons (not atoms). The real structure is a blend (hybrid) of all of them.
Formal charge
A bookkeeping charge assigned to each atom in a Lewis structure: formal charge = (valence electrons) minus (lone pair electrons) minus (half of bonding electrons). It helps you decide which Lewis structure is most representative.
Octet rule
The tendency for main-group atoms to surround themselves with eight electrons (four pairs) when forming bonds. Hydrogen is the exception: it aims for two.
Expanded octet
When a central atom has more than eight electrons around it. Only possible for atoms in Period 3 or below, which have accessible d orbitals. Examples: SF6, XeF4, PCl5.
Incomplete octet
When an atom has fewer than eight electrons around it and is still stable. Boron (as in BF3, BBr3) and beryllium compounds are the classic examples.
Odd-electron species (free radical)
A molecule or ion with an odd total number of valence electrons, meaning at least one electron is unpaired. CH3 (methyl radical) is an example.
Count total valence electrons. For ions, add electrons for negative charges, subtract for positive charges.
Identify the central atom (usually the least electronegative element, never hydrogen).
Connect each outer atom to the central atom with a single bond (each bond uses 2 electrons).
Distribute remaining electrons as lone pairs on outer atoms first, then on the central atom, to satisfy octets.
If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.
The number of electron groups on the central atom determines the electron geometry. Lone pairs then modify the molecular geometry.
2 electron groups: electron geometry = linear (180°)
0 lone pairs: molecular geometry = linear
3 electron groups: electron geometry = trigonal planar (120°)
0 lone pairs: molecular geometry = trigonal planar
1 lone pair: molecular geometry = bent
4 electron groups: electron geometry = tetrahedral (109.5°)
0 lone pairs: molecular geometry = tetrahedral
1 lone pair: molecular geometry = trigonal pyramidal
2 lone pairs: molecular geometry = bent
5 electron groups: electron geometry = trigonal bipyramidal
0 lone pairs: molecular geometry = trigonal bipyramidal
1 lone pair: molecular geometry = seesaw
2 lone pairs: molecular geometry = T-shaped
3 lone pairs: molecular geometry = linear
6 electron groups: electron geometry = octahedral
0 lone pairs: molecular geometry = octahedral
1 lone pair: molecular geometry = square pyramidal
2 lone pairs: molecular geometry = square planar
PF4+ (Question 1)
Phosphorus has 5 valence electrons; four F atoms each contribute 7; subtract 1 for the positive charge: total = 5 + 4(7) - 1 = 32 electrons. P is central with 4 single bonds to F, using 8 electrons. The remaining 24 go as 3 lone pairs on each F. P has 4 bonding pairs, 0 lone pairs.
Electron geometry: tetrahedral
Molecular geometry: tetrahedral
AsBr3 2- (Question 1)
Arsenic has 5 valence electrons; three Br atoms contribute 7 each; add 2 for the 2- charge: total = 5 + 3(7) + 2 = 28 electrons. As is central with 3 single bonds to Br (6 electrons used). Remaining 22 electrons: 3 lone pairs on each Br (18 electrons) plus 2 lone pairs on As (4 electrons). As has 3 bonding pairs + 2 lone pairs = 5 electron groups.
Electron geometry: trigonal bipyramidal
Molecular geometry: T-shaped
NO2- (Question 1)
Nitrogen has 5 valence electrons; two O atoms contribute 6 each; add 1 for the negative charge: total = 5 + 2(6) + 1 = 18 electrons. N is central. One N=O double bond and one N-O single bond uses 6 electrons. Distribute remaining 12 as lone pairs. N ends up with 1 lone pair. This gives 3 electron groups on N (1 double bond + 1 single bond + 1 lone pair; remember, a double bond counts as one group).
Electron geometry: trigonal planar
Molecular geometry: bent
This ion has resonance structures (the double bond can be drawn to either oxygen).
SCl6: expanded octet. Sulphur has 6 bonds to chlorine, giving it 12 electrons around it. Possible because S is in Period 3.
CH3: odd-electron species (free radical). Total valence electrons = 4 + 3(1) = 7, an odd number. Carbon has only 7 electrons around it, with one unpaired.
XeF4: expanded octet. Xenon has 4 bonds to F and 2 lone pairs, giving it 12 electrons. Xe is in Period 5.
BBr3: incomplete octet. Boron has 3 bonds to Br with no lone pairs, giving it only 6 electrons around it. Boron commonly has an incomplete octet.
The azide ion N3- has a total of 16 valence electrons (3 x 5 + 1). Three valid resonance structures exist:
Structure 1: N=N=N with formal charges -1, +1, -1 (left to right). Each terminal N has 2 lone pairs.
Structure 2: N-N≡N with formal charges -2, +1, 0. The left N has 3 lone pairs; the right N has 1 lone pair.
Structure 3: N≡N-N with formal charges 0, +1, -2. The mirror image of Structure 2.
All three contribute to the resonance hybrid. Structure 1 is generally considered the best single representation because it minimises formal charges.
\text{Formal Charge} = (\text{valence electrons}) - (\text{lone pair electrons}) - \frac{1}{2}(\text{bonding electrons})Alternatively: Formal charge = (valence electrons) - (dots) - (lines), where dots = lone pair electrons on that atom and lines = number of bonds to that atom.
Molecular geometry determines whether a molecule is polar, which affects solubility, boiling point, and biological activity. Drug design relies heavily on understanding molecular shapes: a drug molecule needs the right 3D shape to fit into an enzyme's active site, much like a key fitting a lock.
Students often confuse electron geometry with molecular geometry. Electron geometry includes lone pairs in the count; molecular geometry describes only atom positions. A molecule with 4 electron groups but 1 lone pair is tetrahedral in electron geometry but trigonal pyramidal in molecular geometry.
Students often think a double bond counts as two electron groups. It does not. A double bond (or triple bond) counts as one electron group for VSEPR purposes.
Students often assume that all atoms follow the octet rule. Elements in Period 3 and beyond can expand their octet (e.g. S, P, Xe), and some elements like boron are stable with fewer than 8 electrons.
Students often think resonance structures are different molecules that the ion flips between. They are not. The actual molecule is a single hybrid, a blend of all the structures at once.
⚠️ Expect to draw Lewis structures for ions and name both electron and molecular geometry. The exam gives you ions with charges, so remember to adjust the electron count accordingly.
⚠️ Octet rule exceptions are directly tested. Know the three categories (expanded octet, incomplete octet, odd-electron) and be able to identify which applies to a given molecule.
⚠️ Resonance structures with formal charges are tested. You must be able to draw all valid structures and assign formal charges to every atom, including when formal charge is zero.
⚠️ You do not need to show formal charges on the basic Lewis structure questions (Q1), but you do for the resonance question (Q4). Read each question carefully.
True or False: A molecule with 4 electron groups and 2 lone pairs on the central atom has a tetrahedral molecular geometry.
Answer: False. The electron geometry is tetrahedral, but the molecular geometry is bent.
Fill in the blank: The total number of valence electrons in NO2- is ____.
Answer: 18.
True or False: A double bond counts as two electron groups in VSEPR theory.
Answer: False. A double (or triple) bond counts as one electron group.
Fill in the blank: BBr3 is an example of an ____ to the octet rule.
Answer: Incomplete octet (boron has only 6 electrons around it).
True or False: Resonance structures differ in the positions of atoms.
Answer: False. Resonance structures differ only in the placement of electrons; atom positions stay the same.
Q: Draw the Lewis structure for PF4+ and name its electron geometry and molecular geometry.
A: P is central with 4 single bonds to F. Each F has 3 lone pairs. Total valence electrons = 32. Electron geometry: tetrahedral. Molecular geometry: tetrahedral.
Q: What exception to the octet rule does XeF4 demonstrate? How many lone pairs does Xe have?
A: Expanded octet. Xenon has 4 bonds to fluorine and 2 lone pairs, giving it 12 electrons around it. Electron geometry is octahedral; molecular geometry is square planar.
Q: Draw all resonance structures for N3- and assign formal charges to every atom.
A: Three structures exist. (1) N=N=N with charges -1, +1, -1. (2) N-N≡N with charges -2, +1, 0. (3) N≡N-N with charges 0, +1, -2. The sum of formal charges in each structure equals -1, matching the ion's overall charge.
Q: For AsBr3 2-, determine the electron geometry and molecular geometry.
A: As has 3 bonding pairs and 2 lone pairs (5 electron groups total). Electron geometry: trigonal bipyramidal. Molecular geometry: T-shaped.
Q: CH3 has 7 valence electrons. What type of octet rule exception is this?
A: Odd-electron species (free radical). With an odd number of electrons, it is impossible for every electron to be paired, so carbon has an incomplete octet with one unpaired electron.
Molecular geometry connects directly to polarity (covered in intermolecular forces): a symmetric geometry like tetrahedral with identical outer atoms is nonpolar, while an asymmetric geometry like trigonal pyramidal is polar. This in turn determines boiling points, solubility, and other physical properties.
Hybridisation (sp, sp2, sp3, etc.) is determined by the same electron group count used in VSEPR. A central atom with 4 electron groups is sp3 hybridised; 3 groups is sp2; 2 groups is sp. This links to sigma/pi bond labelling.
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