1. The document provides an overview of key concepts for Lewis structures and VSEPR theory, including how to draw Lewis structures for covalent compounds and ions. 2. It explains valence shell electron pair repulsion theory (VSEPR) which is used to predict the 3D shape of molecules based on electron pairs around a central atom. 3. A table is included that lists common molecular geometries determined by VSEPR theory based on the number of electron regions around the central atom.

1. AP Chemistry - Core Concept Cheat Sheet
15: Lewis Structures & VSEPR Theory
Key Chemistry Terms
Arranging Atoms in Lewis Structures
• Valence Shell: Electrons in the outermost shell that are
involved in bonding.
• Valence Bond Theory: Overlap of atomic orbitals form a
bond.
• Covalent Bond: A bond formed between two nonmetals
that involves shared electrons.
• Octet Rule: Atoms are most stable with a full valence shell
(with in most cases is “8”).
• Lewis Structure: A 2D representation of a molecule and
its bonds.
• Lone Pair: Pair of electrons not being shared in a bond.
• Bonding Pair: Pair of electrons that are a bond. Both
atoms sharing the electrons can “count” them in their
valence shell.
• Single bond: One pair of shared electrons.
• Double bond: Two pairs of shared electrons. Shorter and
stronger than a single bond.
• Triple bond: Three pairs of shared electrons. Shorter and
stronger than a double bond.
• Polyatomic Ion: Group of atoms covalently bonded
together that have a net charge.
• Ionic Compound: Electrons are transferred from one
atom to another. Transfer of electrons results in net
charges, which bond with electrostatic attraction.
• Valence Shell Electron Pair Repulsion Theory
(VSEPR): Is used to predict the 3D shape of a molecule
that is based on electron-pair electrostatic repulsion.
Electrons considered include: bonding pairs, nonbonding
pairs and single electron.
• Electron Geometry: 3D structure of a molecule
determined by counting the electron regions around a
central atom.
• Electron Region: Each bond (single, double or triple) and
lone pair count as “1” electron region.
• Molecular Geometry: 3D structure determined by the
atoms bonded to the central atom.
• Ligand: Atoms bonded to the central atom.
1.
2.
3.
4.
5.
Drawing Covalent Lewis Structures
1.
2.
3.
4.
5.
6.
Arrange the atoms as above.
Determine the # of valence electrons for each atom.
Draw the valence electrons—do not double up where a
bond is going to form between two atoms.
Count to see if all atoms have full valences (noting
exceptions above).
If two atoms adjacent to each other do not have full
valences, move in an electron from each to form a double
bond. Repeat for triple bond if necessary.
If two atoms that are not adjacent to each other need to
double bond, try moving a hydrogen to one of them to
cause two atoms adjacent to each other to need the
double bond.
Single bonds:
Double bonds:
Polyatomic anion:
Drawing Ionic Lewis Structures
Transfer the electrons from the metal to the nonmetals until
all have full valences and the net charge = 0.
Ionic compound:
Another Approach to Lewis Structures
Determining # of Valence Electrons
Determining valence electrons from the periodic table:
1 2
For molecules with only 2 elements, arrange the atoms
symmetrically.
“COOH” is a carboxylic acid (both O’s bond to the C and
the H goes on one of the O’s).
Hydrogen and halogens cannot go in the middle.
Write the remaining atoms in the order they appear in the
formula.
Write the hydrogen and halogen atoms around the
element they are written next to in the formula.
1.
2.
3.
4.
3 4 5 6 7 8
5.
Arrange the atoms as above.
Determine the total # of valence electrons for the whole
molecule.
Put 1 bonding pair between a set of atoms to be bonded.
Place remaining electrons in lone pairs, starting with the
most electronegative element.
If atoms do not have full valence shells, move a lone pair
from an adjacent atom in to double, or triple, bond.
VSEPR Theory and Geometry
Drawing Lewis Structures for an Element
A = central atom; X = ligands; E = lone pairs
Molecular
Formula
2
AX2
2.
3
AX3
3.
3
AX2E
4
AX4
Exceptions to the Octet Rule
4
AX3E
4
AX2E2
1. Hydrogen and Helium can only hold 2 electrons (they
5
AX5
only have a 1s orbital).
5
AX4E
2. Boron and Beryllium can be full with 6 electrons.
5
AX3 E2
3. Any element in period 3 or below can have more than 8
electrons (they have empty d orbitals that can hold
5
AX2E3
them).
6
AX6
6
AX5E
6
AX4 E2
How to Use This Cheat Sheet: These are the keys related this topic. Try to read through it carefully
blank sheet of paper. Review it again before the exams.
1.
Use the element symbol to represent the nucleus and
core electrons.
Determine # of valence electrons.
Draw valence electrons—placing one on each side before
doubling up.
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Electron
regions
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Name
Linear
Trigonal Planar
Bent
Tetrahedron
Trigonal pyramidal
Bent
Trigonal bipyramidal
See-saw
T-shaped
Linear
Octahedron
Square pyramidal
Square planar
twice then rewrite it out on a