This cheat sheet covers VSEPR molecular geometry, the model chemists use to predict the three-dimensional shapes of molecules and ions. Students need it because Lewis structures show connections, but VSEPR explains the actual shape around a central atom. It is especially useful for predicting bond angles, polarity, and common molecular shapes from an electron-domain count.
The main idea is that electron groups repel each other and spread out as far apart as possible around a central atom. AXE notation uses for the central atom, for bonded atoms, and for lone pairs, so has three bonded atoms and one lone pair. The steric number is , and it determines the electron geometry before lone pairs are hidden to name the molecular geometry.
Key Facts
- VSEPR stands for valence shell electron pair repulsion, and it predicts shapes by placing electron domains as far apart as possible.
- The steric number is , where is the number of bonded atoms and is the number of lone pairs on the central atom.
- For , the electron geometry is linear, the common AXE form is , and the ideal bond angle is .
- For , the electron geometry is trigonal planar, has bond angles of , and is bent with angles less than .
- For , the electron geometry is tetrahedral, has bond angles of , is trigonal pyramidal, and is bent.
- For , the electron geometry is trigonal bipyramidal, with ideal angles of , , and .
- For , the electron geometry is octahedral, with ideal angles of and .
- Lone pairs repel more strongly than bonding pairs, so the repulsion order is and lone pairs usually compress bond angles.
Vocabulary
- VSEPR theory
- A model that predicts molecular shape by arranging electron domains around a central atom to minimize repulsion.
- Electron domain
- A region of electron density around a central atom, including a single bond, double bond, triple bond, or lone pair.
- Steric number
- The total number of bonded atoms and lone pairs around the central atom, calculated as .
- AXE notation
- A shorthand system where is the central atom, is the number of bonded atoms, and is the number of lone pairs.
- Electron geometry
- The arrangement of all electron domains around the central atom, including both bonds and lone pairs.
- Molecular geometry
- The shape formed by the atoms only, after lone pairs are considered but not shown as vertices of the shape.
Common Mistakes to Avoid
- Counting double or triple bonds as multiple domains is wrong because any bond between the same two atoms counts as one electron domain in VSEPR.
- Naming electron geometry instead of molecular geometry is wrong when lone pairs are present because lone pairs affect the shape but are not included as atoms in the molecular shape name.
- Ignoring lone pairs on the central atom is wrong because lone pairs increase , change , and often reduce bond angles.
- Assuming all tetrahedral-domain molecules have bond angles is wrong because and have compressed angles due to lone pair repulsion.
- Using molecule polarity without considering shape is wrong because polar bonds can cancel in symmetric shapes such as linear or tetrahedral when the outer atoms are identical.
Practice Questions
- 1 A molecule has a central atom with bonded atoms and lone pair. Find , write the AXE notation, and name the molecular geometry.
- 2 Predict the electron geometry, molecular geometry, and ideal bond angle for a molecule with AXE notation .
- 3 A central atom has electron domains and no lone pairs. What is the AXE notation, electron geometry, and set of ideal bond angles?
- 4 Explain why is nonpolar but is polar, even though both molecules contain polar bonds.
Understanding VSEPR Molecular Geometry Reference
A reliable VSEPR prediction starts with a correct Lewis structure. Count all valence electrons first. Include extra electrons for a negative charge and remove electrons for a positive charge.
Choose the central atom, which is usually the least electronegative atom except for hydrogen. Connect atoms with single bonds, finish the outer atoms, then place any remaining electrons on the central atom. After that, count regions around the central atom rather than counting individual electron pairs.
A single bond, double bond, or triple bond each occupies one region. Multiple bonds contain more electron density, so they can push nearby bonds a little more strongly, but they do not create extra regions.
Lone pairs change more than the name of the shape. Their electron density stays closer to the central atom because it is not shared with another nucleus. This gives lone pairs more room to push bonded atoms together.
The result is smaller bond angles than the ideal values. In arrangements with five regions, lone pairs usually choose an equatorial position. An equatorial lone pair has fewer close interactions than an axial lone pair.
This detail explains why some five-region molecules have shapes that seem less obvious from a flat Lewis structure. For six regions, opposite positions are especially important because they lie straight across from each other.
Molecular shape helps predict polarity, but shape alone is not enough. Each polar bond has an uneven pull on electrons, called a bond dipole. In a symmetrical molecule, these pulls can cancel in all directions.
Carbon dioxide has polar carbon oxygen bonds, yet its straight shape makes the whole molecule nonpolar. Water has polar oxygen hydrogen bonds and a bent shape, so its bond dipoles do not cancel. It is polar.
This difference affects melting and boiling points, solubility, and how substances interact. Water dissolves many ionic or polar substances because of its strong uneven charge distribution.
Students often lose marks by mixing up electron geometry with molecular geometry. Electron geometry includes lone pairs because they take up space. Molecular geometry describes only the positions of atoms.
It helps to sketch the three-dimensional arrangement using wedges and dashed bonds when needed. A solid wedge points toward the viewer, while a dashed bond points away. VSEPR is a useful model, not a perfect measurement tool.
Actual angles vary because atoms differ in size, bonds differ in electron density, and surrounding groups can pull on electrons. Transition metal compounds often need more advanced bonding models. For common main-group molecules in school chemistry, careful electron counting and clear separation of lone pairs from bonded atoms usually lead to the right prediction.