
Shapes of molecules (molecules actually have
3D structures that affect how they bond in
reactions)

Based on electrons and their repulsion from
each other
 Valence Shell Electron Pair Repulsion (VSEPR)

Focused on general model ABx
 A = central atom, B = terminal (attached) atoms

Recall: Electron Pairs can be bonding pairs
or lone pairs
 Bonding pairs may be single bonds or multiple
bonds

An electron domain is a lone pair or a
bonded pair
 (double and triple bonds only count for ONE
electron domain)

How many electron domains do these
molecules have?

Geometric shapes based on electron domains
AB2
Linear
180° (provided no lone pairs on central)
Ex: CO2
AB3
Trigonal Planar
120° (provided no lone pairs on central)
Ex: BF3

Geometric shapes based on electron domains
AB4
Tetrahedral
109.5° (provided no lone pairs on central)
Ex: CH4
AB5
Trigonal Bipyramidal
120° (equatorial atoms) & 90° (axial atoms)
Ex: PCl5

Geometric shapes based on electron domains
AB6
Octahedral
90° between all atoms (provided no lone
pairs on central)
Ex: SF6

It is helpful to draw molecules on paper with
3-Dimensionality to better visualize the shape

Electron Domain: based on electron
domains (bonded and unbonded)
 Basically, it looks at ALL the electrons

Molecular Geometry: based on atoms ONLY
(bonded pairs only)
 Basically only looks at electrons in bonds (or
atoms)
1)
Draw Lewis Structure
2)
Count electron domains around central
atom
3)
Determine Electron Domain Geometry
4)
Determine Molecular Geometry from atoms
only

Consider the following three compounds:
 CH4, NH3, H2O

What are their electron domain geometries?

What are their molecular geometries?

Some electron domains repel MORE than
others

Lone pairs repel more
 Because not bonded, they experience less nuclear
attraction

Double and Triple bonds repel more
 Due to higher electron density

Other examples

Determine the electron domain geometry
and the molecular geometry for the
following. Include bond angles as well.

SeO2

SF6

State the electron domain geometry, the
molecular geometry and the bond angle (if the
bond angle is not ideal, state “less than _____°”)

PCl3

SeCl4

A central atom is one where it is surrounded
by 2 or more atoms

Determine the multiple shapes in the
molecule around each central atom

Ex:

Recall: A bond between 2 atoms is polar
covalent when difference in electronegativity
is greater than 0.5

A molecule OVERALL can also be polar or
nonpolar
 Based on geometry of molecule and
 Individual bonds in molecule

CO2 and H2O

First, check if there are lone pairs of electrons
 If yes, it is POLAR (unless they are “balanced”)

If no lone pairs, check if there are polar bonds
 If there are no polar bonds, is is NON POLAR

If there are polar bonds, check for symmetry
 If all terminal atoms are the same, it is NON POLAR
 If terminal atoms are not the same, but there is
symmetry, it is NON POLAR
 If terminal atoms are not the same, and there is no
symmetry, it is POLAR

Determine if the molecule is POLAR or NON
POLAR

CH2Cl2

For each hypothetical molecule below, draw
one structure that is polar and one that is non
polar

H2SF4
&
&
XeF4
H2PCl3

For each molecule below, determine
 the electron domain geometry,
 the molecular geometry,
 draw the 3D-structure on paper,
 state if it is polar or nonpolar


SbCl5
I 3-

Definition: atoms share electrons when an
atomic orbital on one atom overlaps with an
atomic orbital on the other

Explains why atoms share electrons
 Potential energy of resulting molecule is lower
than combined energies of isolated atoms
 A.K.A. when bonds form, energy given off
 Therefore, energy required to break bonds


Consider BeCl2
How can it bond with
2 atoms???
Hybridization: mixing of atomic orbitals

Hybrid orbitals are always shown in GREEN

Consider BF3

Consider CH4

Consider PCl5

Consider SF6
1)
2)
3)
4)
5)
Draw Lewis Structure
Count # of electron domains (this is # of
hybrid orbitals needed)
Determine the hybrid orbital needed
Draw ground state configuration of central
atom
Place electrons in the hybrid orbitals one at
a time

Use hybrid orbital theory to explain the
bonding in BrF5

Show the hybridization of the following
molecule. Be sure to include a) what type of
hybrid orbital is needed and 2) the orbital
diagram of the hybrid orbital

XeF4


How many electron domains around the
carbons?
What hybrid orbitals are needed?

Bonds that are made using hybrid orbitals (or
along the internuclear axis) are called sigma
(σ) bonds

Bonds that are made using remaining p
orbitals and overlap on the sides are called pi
(π) bonds