Source: Comprehensive Guide to Chemical Principles, CHM 111
Tags: Lewis structures, VSEPR, molecular geometry, bond angles, electronegativity, polarity, polar covalent, ionic bonds, resonance structures, electron pairs, valence electrons, octet rule, dipole moment
Difficulty: Introductory Prerequisites: Familiarity with the periodic table, basic electron configuration, and the concept of valence electrons.
This topic is the foundation for nearly everything else in general chemistry. Before you can understand how reactions happen, why certain substances dissolve or conduct electricity, or how to predict the properties of a compound, you need to know how atoms bond and what shapes the resulting molecules take. If you are coming into CHM 111 cold, start here. The ideas in this section, particularly Lewis structures and VSEPR theory, will appear on almost every exam in some form.
Atoms bond by sharing or transferring electrons. Lewis structures show you where the electrons are. VSEPR theory uses those electron positions to predict the 3D shape of a molecule, which in turn determines whether it is polar or nonpolar.
Lewis structure
A diagram that shows the bonding between atoms in a molecule and any lone pairs of electrons that remain. Think of it as a map of where every valence electron lives.
Valence electrons
The electrons in an atom's outermost shell, and the only ones involved in bonding. For main-group elements, the group number on the periodic table tells you the count.
Octet rule
Most atoms (except hydrogen, which wants 2) are stable when surrounded by 8 electrons. In simple terms, atoms "want" a full outer shell, and bonding is how they get there.
VSEPR theory (Valence Shell Electron Pair Repulsion)
A model for predicting molecular shape. Electron groups around a central atom repel each other and spread out as far as possible, like balloons tied together at a single point.
Electronegativity
A measure of how strongly an atom pulls shared electrons towards itself. Fluorine is the most electronegative element; metals at the bottom left of the periodic table are the least.
Nonpolar covalent bond
A bond where electrons are shared equally between two atoms. This happens when the atoms have the same or very similar electronegativities (e.g., H₂, Cl₂).
Polar covalent bond
A bond where electrons are shared unequally because one atom is more electronegative than the other (e.g., the O–H bonds in water).
Ionic bond
A bond formed by the transfer of electrons from one atom to another, creating oppositely charged ions that attract (e.g., NaCl).
Dipole moment
A measure of the separation of positive and negative charge in a molecule. A molecule with a net dipole moment is polar.
Resonance structures
Two or more valid Lewis structures for the same molecule, where the actual electron distribution is a blend (hybrid) of all of them. Think of it as the molecule averaging its electron positions.
Resonance hybrid
The real structure of a molecule that has resonance. It is not any single Lewis structure but a weighted average of all contributing structures.
Count the total number of valence electrons for every atom in the molecule. Add electrons for negative charges, subtract for positive charges.
Identify the central atom. The least electronegative atom (excluding hydrogen, which is always terminal) goes in the centre.
Draw single bonds (one pair of electrons each) between the central atom and each surrounding atom.
Distribute remaining electrons as lone pairs on outer atoms first, completing their octets.
If the central atom still lacks an octet, convert lone pairs on outer atoms into double or triple bonds.
Worked example: SBr₂
Total valence electrons: S contributes 6, each Br contributes 7. That gives 6 + 2(7) = 20 electrons.
Sulfur is less electronegative than bromine, so S is the central atom.
Draw single bonds: S–Br on each side uses 4 electrons. 16 remain.
Place 3 lone pairs on each Br (6 electrons each = 12 electrons used). 4 electrons remain.
Place the remaining 4 electrons (2 lone pairs) on S. Every atom now has an octet.
Worked example: CO₂
Total valence electrons: C contributes 4, each O contributes 6. Total = 16.
Carbon is the central atom. Place single bonds to each O: 4 electrons used, 12 remain.
Distribute to outer O atoms: 3 lone pairs each = 12, but that leaves C with only 4 electrons (no octet).
Convert one lone pair from each O into a bonding pair: the result is two C=O double bonds. Each O retains 2 lone pairs. Total electrons: 16. Octets satisfied. The shape is linear.
VSEPR predicts geometry by counting electron groups (bonds plus lone pairs) around the central atom and arranging them to minimise repulsion.
2 electron groups: Linear geometry, 180° bond angle.
3 electron groups: Trigonal planar geometry, 120° bond angles. If one group is a lone pair, the molecular shape is bent.
4 electron groups: Tetrahedral geometry, 109.5° bond angles. With one lone pair the shape is trigonal pyramidal; with two lone pairs it is bent.
5 electron groups: Trigonal bipyramidal geometry, 90° and 120° bond angles.
6 electron groups: Octahedral geometry, 90° bond angles.
A critical distinction: electron geometry counts all electron groups, including lone pairs. Molecular geometry describes only the positions of the atoms (lone pairs are invisible to the shape name). This is one of the most common exam slip-ups.
Electronegativity differences between bonded atoms determine bond type:
0 to ~0.4: Nonpolar covalent (electrons shared equally).
~0.4 to ~1.7: Polar covalent (electrons shared unequally).
Greater than ~1.7: Ionic (electron effectively transferred).
A molecule's overall polarity depends on two things: whether its bonds are polar, and whether those bond dipoles cancel out due to the molecule's shape.
Symmetric shapes (linear CO₂, tetrahedral CCl₄) can have polar bonds but still be nonpolar molecules because the dipoles cancel.
Asymmetric shapes (bent H₂O, trigonal pyramidal NH₃) have dipoles that do not cancel, making the molecule polar overall.
Some molecules cannot be represented accurately by a single Lewis structure. Ozone (O₃) and the carbonate ion (CO₃²⁻) are classic examples.
Draw all valid Lewis structures by moving lone pairs into bonding positions (or vice versa). The atom positions stay fixed; only electron placement changes.
The true structure is the resonance hybrid: a blend of all contributing structures.
Experimental evidence supports this. In CO₃²⁻, all three C–O bonds are measured at the same length, which is intermediate between a single and a double bond. No single Lewis structure predicts that.
VSEPR geometry summary:
Electron Groups | Lone Pairs | Molecular Shape | Bond Angle |
|---|---|---|---|
2 | 0 | Linear | 180° |
3 | 0 | Trigonal planar | 120° |
3 | 1 | Bent | ~117° |
4 | 0 | Tetrahedral | 109.5° |
4 | 1 | Trigonal pyramidal | ~107° |
4 | 2 | Bent | ~104.5° |
5 | 0 | Trigonal bipyramidal | 90°, 120° |
6 | 0 | Octahedral | 90° |
Molecular polarity determines solubility ("like dissolves like"), which is why oil and water do not mix: oil molecules are nonpolar, water is polar. The same principle drives drug design, since a drug molecule's polarity affects how it moves through the body and crosses cell membranes.
VSEPR geometry is central to understanding enzyme–substrate interactions in biochemistry. The 3D shape of a molecule determines whether it fits into an enzyme's active site.
Students often confuse electron geometry with molecular geometry. Electron geometry counts all electron groups (including lone pairs). Molecular geometry describes only where the atoms sit. Methane (CH₄) and water (H₂O) both have tetrahedral electron geometry, but methane's molecular geometry is tetrahedral while water's is bent.
A molecule with polar bonds is not automatically a polar molecule. If the shape is symmetric enough for the dipoles to cancel (e.g., CO₂), the molecule is nonpolar overall.
Resonance structures are not different molecules flickering back and forth. The molecule exists as one hybrid all the time. The individual structures are just our limited way of drawing something that a single Lewis diagram cannot capture.
Hydrogen always goes on the outside of a Lewis structure, never as the central atom. It can only form one bond.
⚠️ Drawing Lewis structures is the gateway to VSEPR questions. If your Lewis structure is wrong, every answer that follows will be wrong too. Practise counting valence electrons until it is automatic.
⚠️ "What is the molecular geometry?" and "What is the electron geometry?" are different questions. Read carefully.
⚠️ Resonance structures appear frequently in exam questions involving formal charge or bond order. Know how to identify when resonance is possible (typically when you can draw multiple valid Lewis structures by shifting a double bond).
⚠️ Polarity questions often require you to combine VSEPR shape prediction with electronegativity reasoning. A two-part skill.
True or false: A molecule with four electron groups around its central atom always has a tetrahedral molecular shape.
Fill in the blank: The bond angle in a trigonal planar molecule is ____°.
True or false: CO₂ is a polar molecule because it contains polar C=O bonds.
Fill in the blank: To find the total valence electrons for an ion with a 2– charge, you ______ 2 electrons to the total.
True or false: Resonance structures differ in the positions of their atoms.
Answers: 1. False (lone pairs change the molecular shape). 2. 120. 3. False (the linear shape causes the dipoles to cancel). 4. Add. 5. False (only electron positions change; atoms stay put).
Q: Draw the Lewis structure for water (H₂O) and predict its molecular geometry.
A: Oxygen has 6 valence electrons; each hydrogen has 1. Total = 8. Place O in the centre with single bonds to each H (4 electrons used). The remaining 4 electrons go on O as 2 lone pairs. Oxygen has 4 electron groups (2 bonds + 2 lone pairs), so the electron geometry is tetrahedral and the molecular geometry is bent, with a bond angle of approximately 104.5°.
Q: Why is CCl₄ nonpolar even though each C–Cl bond is polar?
A: CCl₄ has a tetrahedral shape. The four polar C–Cl bond dipoles point symmetrically in all directions and cancel each other out, resulting in no net dipole moment.
Q: Explain why ozone (O₃) is described as a resonance hybrid rather than drawn with a single Lewis structure.
A: Two valid Lewis structures can be drawn for O₃, differing in which O–O bond is the double bond. The real molecule has two equivalent O–O bonds of intermediate length, which no single structure shows. The resonance hybrid averages both structures.
Q: What is the molecular geometry of a molecule with 3 bonding groups and 1 lone pair on the central atom?
A: The electron geometry is tetrahedral (4 groups total), but the molecular geometry is trigonal pyramidal because the lone pair is not "seen" in the shape name.
This material connects directly to the solutions and polarity topic: whether a solute dissolves in a given solvent depends on polarity, which you predict using the tools from this section. It also underpins stoichiometry, because writing correct molecular formulas requires understanding how atoms bond. Later in general chemistry, molecular shape and polarity become essential for intermolecular forces and phase behaviour.
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