VSEPR
What shape are your molecules in?
Background you need…
 Lewis structures
 How many bonds do each element make?
 What can expand?
 Bonding (covalent)
 Polarity
 Electronegativity and determining bond type
 Resonance v. Isomers
 Formal charge
Let’s review now…..
Lewis Structures
 Remember that Lewis structures want a
full outer shell
 Remember that for a given Lewis
structure, the number of electrons around
the atoms must equal the total number of
electrons individually assigned.
 Ex: C has 4, H has 1, so CH4 must have 8
total
Isomers
 Same formula, different arrangement of
atoms
 Physically break bonds and MOVE
atoms
Resonance Structures
 Have the same
alignment of
atoms, but
different bonding
(electrons ONLY
are moved, both
in bonds and
lone pairs)
Determining formal charge
Formal charge can be determined by:
Normal number of electrons in outer shell
[(1/2 the number of bonded electrons)
+
lone electrons]
_____________________________________
= formal charge
Example: N in NH4
FC =5- [(1/2 of 8)+ 0]= +1
Formal charge and
stability
 The most “happy” molecules tend to have
no formal charges
 However, molecules may be “happy” if
they have not NET charge on them (if
there is 1+ and 1-, so a net of +1 + (1)=0)
 Resonance structures that are the best
have a minimal formal charge and a full
octet around each atom
What is VSEPR?






Valence
Shell
Electron
Pair
Repulsion
Theory
Why?
 The shape of molecules influences their
characteristic:
 Physical properties
 polarity
 boiling point
 melting point
 state of matter at room temperature
 Chemical properties
 What it will bond with and energy associated with the
bond
 Biological properties
 Enzyme specificity (induced fit model require close
shapes)
Valence Bond (VB) Theory
 Deals with the overlap of the atomic orbitals (AO) of the
participating atoms to form a chemical bond. Due to the
overlapping, electrons are found in the bond region.
 However, the atomic orbitals for bonding may not be
"pure" atomic orbitals directly from the orbitals of the
atoms involved. Often, the bonding atomic orbitals
have a character of several possible types of orbitals
(say s, p, and d).
 The methods to get an AO with the proper character for
the bonding is called hybridization. The resulting
atomic orbitals are called hybridized atomic orbitals
or simply hybrid orbitals.
Valence Bond Theory
And VSEPR Notation
How does Lewis theory explain the bonds in H2 and F2?
Sharing of two electrons between the two atoms.
Bond Dissociation Energy
Bond Length
Overlap Of
H2
436.4 kJ/mole
74 pm
1s orbitals
F2
150.6 kJ/mole
142 pm
2p orbitals
Valence bond theory: bonds are formed by
sharing of e- from overlapping atomic orbitals.
Valence bond method
According to this model, the H-H bond forms
as a result of the overlap of the 1s orbitals
from each atom.
74 pm
9 - 12
Valence bond method
Hybrid orbitals are need to account for the
geometry that we observe for many molecules.
Example - Carbon
Outer electron configuration of 2s2 2px1 2py1
We know that carbon will form four equivalent
bonds such as in CH4, CH2Cl2 , CCl4.
The electron configuration appears to indicate
that only two bonds would form and they would
be at right angles -- not tetrahedral angles.
9 - 13
Hybridization
• To explain why carbon forms four identical
single bonds, we assume the original
orbitals will blend together.
• This lowers the energy
• Lower energy is more favorable
energy
2p
2sp3
2s
Unhybridized
Hybridized
9 - 14
The parent geometries: all others come from these
Some new conventions for
shapes….dashes and wedges
Steric Number
 The number of e- groups, or “things”
sprouting off of an atom
 These can be either
 Bonds
 Of any order (1, 2, or 3)
Or
 Lone pairs of electrons
Steric Number Examples
 Ex #1: CH4
 There are 4 H’s branching
off , so the steric number
is 4
 SN=4
 Ex #2: H2O
 SN= 4
 Explain why
 Ex #3: CO2
 SN= 2
 Explain why
General Formulas
 All molecules with a shared general
formula have a shared geometry
 we use them to help note shape
 Formulas are typically written with A’s,
X’s, and E’s
The letters stand for:
 A= the central atom
 X *= the number of atoms attached to the central
atom
 E= the number of lone pairs of electrons attached
to the central atom
 *Some sources use B’s in place of X’s
General Formula Examples
 Ex #1: CH4
 AX4
 Ex #2: H2O
 AX2E2
 Ex #3: CO2
 AX2
Linear
 AX2
Trigonal planar
 AX3
Tetrahedral
 AX4
Pyramidal (Trigonal or
tetrahedral)
 Tetrahedral parent
shape
 1 lone pair of
electrons
 AX3E
Bent
 Tetrahedral
parent shape
 2 lone pair of
electrons
 AX2E2
When determining polarity
it is important to look at
the dipole moments- do
they cancel out?
Expanded Octets
 Atoms with expanded CENTRAL octets
are not limited to having only 4 atoms
attached.
 Shapes that require expanded octets
are
 Trigonal bipyramidal (TBP) parent
 See-saw/ teeter-totter derivative
 Octahedral parent
 Square pyramidal, square planar, and Tshaped derivatives
Trigonal bipyramidal
 AX5
Seesaw
a.k.a.
Teeter-totter
 Trigonal bipyramidal
parent shape
 1 lone pair of
electrons
 AX4E
T-shaped
 Trigonal bipyramidal
parent shape
 2 lone pair of
electrons
 AX3E2
Linear
 Trigonal bipyramidal
parent shape
 3 lone pair of
electrons
 AX2E3
Octahedral
 AX6
Square pyramidal
 Octahedral parent
shape
 1 lone pair of
electrons
 AX5E
Square planar
 Octahedral parent
shape
 2 lone pair of
electrons
 AX4E2
T-shaped
 Octahedral parent
shape
 1 lone pair of
electrons
 AX3E3
 See the grid handout for more specifics
on each!
Summary of shapes
ID these VSEPR shapes…
Sweet drill and practice
web site
 Given generic shapes to ID:
 http://www.chemistry-drills.com/VSEPR1.php?q=1
 Given molecules to draw out:
 Basic: http://www.chemistrydrills.com/VSEPR-1.php?q=2
 Advanced: http://www.chemistrydrills.com/VSEPR-1.php?q=3