molecular geometry and VSEPR theory

1. LEWIS STRUCTURES

2. Lewis Structures
• The development of periodic table and the
concept of electronic configuration gave
chemists a rationale for molecule and
compound formation.
• Lewis structures (also known as Lewis dot
diagrams, electron dot diagrams and
electron dot structures) are diagrams that
show the bonding between atoms of a
molecule and the lone pairs of electrons that
may exist in the molecule.

3. Covalent Bonding
• Covalent bonds share electrons in order to
form a stable octet around each atom in the
molecules.
• According to Lewis, the elements undergo
bonding in order to attain the stable noble
gas configuration (i.e., 8 valence electrons in
their outer most shell) of the nearest
elements in a periodic table. This observation
led to the guideline known as octet rule.

4. The Octet Rule
• The octet rule dictates that atoms are most
stable when their valence shells are filled
with eight electrons.
• It is based on the observation that the
atoms of the main group elements have a
tendency to participate in chemical
bonding in such a way that each atom of
the resulting molecule has eight electrons
in the valence shell.
• The octet rule is only applicable to the
main group elements.

5. Why Elements Follow the
Octet Rule
Atoms follow the octet rule because they
always seek the most stable electron
configuration. Following the octet rule results
in completely filled s- and p- orbitals in an
atom's outermost energy level. Low atomic
weight elements (the first 20 elements) are
most likely to adhere to the octet rule.

6. Limitations of Octet Rule:
(1)Hydrogen with 1 electron attains stability by
sharing, gaining or losing 1 valence electron. It
does not need to complete octet to attain
stability. Also, He has only 2 electrons and is
stable.
(2)Incomplete octet: In certain molecules such as
BeH2, BeCl2, BH3, BF3, the central atom has
less than 8 electrons in its valence shell, yet the
molecule is stable.
(3)Expanded octet: In certain molecules such as
PF5, SF6, IF7, H2SO4, the central atom has
more than 8 valence electrons, yet the molecule
is stable

7. 7
Drawing Lewis Structures
1) Count the total valence electrons for
the molecule: To do this, find the
number of valence electrons for each
atom in the molecule, and add them up.
Example: CO2
Valence electrons: C = 4, (O = 6)x2 ➔ 16
Total number of valence electrons is important,
not where they came from
To start: Draw Lewis symbols of all individual atoms in formula

8. 8
Drawing Lewis Structures
2a) Write the symbols for each atom
b) show which atom is attached to which
• 2 elements: central atom is first
• more than 2: order of connected atoms as written in
formula
c) Use a single dash to show bonds between
atoms:
CO2: 2 atoms, C is central, both O are
attached to C by single bond
O—C—O

9. 9
Drawing Lewis Structures
3) Complete octets around all atoms
bonded to central atom by adding
electrons.
(H does not get an octet!)
CO2: .. ..
: O – C – O :
.. ..
Count electrons now: 16 valence electrons
shown, both O have octet, C?

10. 10
Drawing Lewis Structures
4) Any leftover valence electrons after
step 3?
Place them on the central atom.
Even if doing so gives that atom more
than 8 electrons (there are exceptions
to octet rule)
CO2: no leftover valence electrons

11. 11
Drawing Lewis Structures
5) If central atom does not have an octet of electrons
after step 4:
Try multiple bonds!
Use unshared electrons already shown in step 3 and 4
and move them between central atom and other
atoms.
CO2

12. 12
Drawing Lewis Structures
6)
Some handy bond rules to remember for molecules:
– Hydrogen and the halogens (F, Cl, Br, I) form a single bond.
– The family oxygen (O, S, Se) forms 2 single bonds (or 1 double
bond)
– The family nitrogen (N, P, As) forms 3 single bonds, or 1 single
bond and a double bond or a triple bond. So does boron.
– The family carbon (C, Si) forms 4 single bonds, or 2 double
bonds, or…..
• A good thing to do is to bond all the atoms together by
single bonds, and then add the multiple bonds until the
rules above are followed.

13. 13
Formal Charges
The formal charge of an atom in a molecule is the
hypothetical charge the atom would have if we could
redistribute the electrons in the bonds evenly between the
atoms.

14. 14
Lewis Structures
The Lewis structure of carbon tetrachloride
• Provides information about connectivities
• Provides information about valence orbitals
• Provides information about bond character
However, the Lewis structure provides no
information about the shape of the molecule

15. 15
Lewis Structures
– All atoms are drawn in the same plane (the
paper).
– Do not show the shape of the molecule.
3-D Drawing
Lewis Structure

16. MOLECULAR GEOMETRY
AND
THE VSEPR THEORY

17. 17
Molecular Structure
The structure of a molecule is defined by:
• Bond length: the distance between two
atoms held together by a chemical bond
– Bond length decreases as the number of
bonds between two atoms increases.
• Single bond is the longest.
• Triple bond is the shortest.
• Bond angle: the angle made by the
“lines” joining the nuclei of the atoms in a
molecule
H
O
H
104.5o

18. 18
Molecular Geometry
• Bonding pairs of electrons: electrons
that are shared between two atoms
Cl
Cl C Cl
Cl
CCl4 has 4 bonding pairs,
C has 4 electron domains
Bonding
pairs
Bonding
pairs

19. 19
Molecular Geometry
• Nonbonding (lone) pairs of electrons:
electrons that are found principally on one
atom, not in between atoms
= unshared electrons
H N H
H
Nonbonding
pair

20. How do we determine the shapes
of molecules and ions?
VSEPR Theory
Valence Shell Electron Pair
Repulsion Theory

21. 21
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best arrangement of a given
number of electron domains is the
one that minimizes the repulsions
among them.”
The Valence Shell Electron Pair
Repulsion model predicts the shapes of
molecules and ions by assuming that the
valence shell electron pairs are arranged
as far from one another as possible to
minimize the repulsion between them.

22. 22
VSEPR Model
Electron Pair Geometry
(electron domain
geometry) – is determined
by the number and
arrangement of all electron
pairs (bonding and lone)
around the central atom.
Molecular geometry – is
determined by the
arrangement of atoms (or
bonding electron pairs only)
around the central atom.
In molecules with no lone pairs,
Electron Domain Geometry = Molecular Geometry
N
H
H
:
H
Bonding
electrons
Non-bonding electrons
or LONE PAIR

23. 23
VSEPR Theory
(1) The shape of the molecule is determined by repulsions
between all of the electron present in the valance shell
(2) Electron pairs repel each other and try to remain as
far apart from one another as possible to minimize this
repulsion.• Lone pair electrons takes up more space
round the central atom than a bondpair.
Repulsion strengths
Lone pair -Lone pair > Lone pair -Bond pair > Bond pair-
Bond pair
(3) The positions of atoms in space (molecular geometry or
configuration about a central atom) is determined by
the relative positions that electron pairs achieve after
taking repulsions into account.
(4) The arrangement of pairs about a central atom
depends on the number of pairs that exist about the
atom.

24. 24
Predicting Molecular Geometry
◼ The AXE shorthand notation is used
AXE shorthand notation:
• A - central atom
• X - terminal atoms
• E - lone pair electrons
◼ Only five basic shapes.
◼ When a lone pair replaces an atom, the molecular
geometry changes as well as the angles.
AX3E0
Linear
Trigonal
Planar
Tetrahedral
Octahedral
Trigonal
Bipyramidal

25. 25
Predicting Molecular Geometry
1. Draw the Lewis structure.
2. Determine how many electron pairs (bonded and
non-bonded) are around the central atom.
**Treat a multiple bond like a single bond when
determining a shape - Count a multiple bond as
one pair.
3. Write the AXE shorthand notation.
4. Determine the electron pair geometry (**one
of the five basic shapes).
5. If the molecule has lone pairs around the
central atom, then determine the molecular
geometry. (This is a subset of the electron
geometry.)

26. 26
Predicting Molecular Geometry
The 5 Basic Shapes
• Two electron pairs in the valence orbital are
arranged linearly
• Three electron pairs are organized in a
trigonal planar arrangement
• Four electron pairs are organized in a
tetrahedral arrangement
• Five electron pairs are arranged in a trigonal
bipyramidal arrangement
• Six electron pairs are organized in an
octahedral arrangement

27. 27
Predicting Molecular Geometry
# of e-
pairs
2 3 4 5 6
# of
bonding
pairs
2 3 4 5 6
# of lone
pairs
0 0 0 0 0
Type AX2 AX3 AX4 AX5 AX6
e- pair
geometry
/domain
LINEAR
TRIGONAL
PLANAR
TETRAHED
RAL
TRIGONAL
BIPYRAMIDAL
OCTAHEDR
AL
Geometry LINEAR
TRIGONAL
PLANAR
TETRAHED
RAL
TRIGONAL
BIPYRAMIDAL
OCTAHEDR
AL
Example
BeCl2, BF3, CO3
2−,
−
CH4, PO4
3−,
2− PCl5 SF6

28. 28
Predicting Molecular Geometry
• In molecules with no lone pairs,
Electron Domain Geometry = Molecular Geometry
• The electron-domain geometry is often not the shape of the
molecule, however.
• The molecular geometry is that defined by the positions of
only the atoms in the molecules, not the nonbonding pairs.
• Molecular geometry is a consequence of electron-domain
geometry.

29. 29
Linear Electron Domain
• In the linear electron domain, two electron pairs are around
the central atom
• In the linear domain, there is only one molecular geometry:
linear.
NOTE: If there are only two atoms in the molecule, the
molecule will be linear no matter what the
electron domain is.
# of e-
pairs
e- domain
geometry
number of lone pairs
0 (AX2) 1(AXE) 2 3
2
linear
BeCl2, HgCl2,
CO2
Linear/diatomic
N2
none none

30. 30
Trigonal Planar Electron Domain
There are three molecular geometries:
• Trigonal planar, if all the electron domains
are bonding,
• Bent, if one of the domains is a lone pair.
• Linear, if two of the domains are lone pairs
# of e-
pairs
e- domain
geometry
number of lone pairs
0 (AX2) 1(AXE) 2 3
3
Trigonal
planar
Trigonal planar
BCl3
Bent
SO2
Linear
CO2
none

31. 31
Tetrahedral Electron Domain
• There are three molecular geometries:
– Tetrahedral, if all are bonding pairs,
– Trigonal pyramidal if one is a nonbonding pair,
– Bent if there are two nonbonding pairs.
# of e-
pairs
e- domain
geometry
number of lone pairs
0 (AX4) 1(AX3E) 2(AX2E2) 3
4 Tetrahedral
Tetrahedral
CH4
Trigonal
pyramidal
NH3 H2O
none
Bent

32. 32
Trigonal Bipyramidal ē Domain
• There are four distinct molecular geometries in this
domain:
– Trigonal bipyramidal
– Seesaw
– T-shaped
– Linear
# of e-
pairs
e- domain
geometry
number of lone pairs
0 (AX5) 1(AX4E) 2(AX3E2) 3(AX2E3)
5
Trigonal
bipyramidal
Trigonal bipyramidal
PCl5
Seesaw
SF4
T-shaped
ClF3
Linear
XeF2

33. 33
Octahedral Electron Domain
• All positions are equivalent in the octahedral domain.
• There are three molecular geometries:
– Octahedral
– Square pyramidal
– Square planar
# of e-
pairs
e- domain
geometry
number of lone pairs
0 (AX6) 1(AX5E) 2(AX4E2) 3
6
Octahedral Octahedral
SF6
Square
Pyramidal
IF5
Square planar
XeF4

34. 34
Molecular Geometries (summary)

35. 35
VSEPR (example)
2. Two electron pairs around the central atom.
Two bonding pairs and Zero lone pairs. AX2E0
Example 1: BeCl2
electron pair geometry = molecular geometry
Geometry is Linear. Bond angle is 180o.
AX2E0
Cl Be Cl
1. Draw the Lewis structure

36. 36
VSEPR (example)
Three electron pairs around the
central atom.
Three bonded and Zero lone pairs.
AX3E0
Example 2: BF3
F:
:F
B
:F:
.. ..
..
..
triangular planar
(or trigonal planar)
AX3E0

37. 37
VSEPR (example)
Example 3: SO2
AX3E0 AX2E1
electron geometry = triangular planar.
molecular geometry = bent or angular.
O S O
S
O
O
Three electron pairs around
the central atom.
Two bonded and One lone
pairs.
AX2E1

38. 38
VSEPR (example)
Example 4: CH4
H
H
H C
H
Four electron pairs around
the central atom. Zero lone
pairs. AX4E0
AX4E0
tetrahedral
electron pair geometry = molecular geometry

39. 39
VSEPR (example)
Example 5: NH3
AX4E0 AX3E1
electron geometry = tetrahedral.
molecular geometry = triangular pyramidal
H
H N
H
H N H
H
Four electron pairs around the
central atom.
Three bonding and One lone
pair. AX3E1

40. 40
VSEPR (example)
Example 6: H2O
electron geometry = tetrahedral
molecular geometry = angular or bent
H O H
Four electron pairs around the
central atom.
Two bonding and Two lone
pairs.
AX4E0 AX2E2
O
H H