Putting it all together
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Molecular Shapes
1
SeO2
One more example:
1st question:
What’s the total # of valence electrons?
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What’s the Lewis Dot Structure for SeO2?
2
3
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Dot structure for SeO2
SeO2
Next question:
What’s the central atom?
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18 total valence electrons.
4
Dot structure for SeO2
Now?
Fill the octets
Se — O
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O—
5
..
..
..
:O — Se — O:
¨
¨
¨
Once we’ve filled the octets, what do we do?
Check the total # of valence electrons
20 total electrons – too many!
So, what do we do?
Make a bond!
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Dot structure for SeO2
6
Dot structure for SeO2
Are we done?
Check the formal charges.
FC(left O) = 6 – 2 – 4 = 0
FC(Se) = 6 – 3 – 2 = 1
FC(right O) = 6 – 1 – 6 = -1
Acceptable. Are we done yet?
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..
:O = Se — O:
¨
¨
¨
7
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RESONANCE
8
..
:O = Se — O:
¨
¨
¨
..
:O — Se = O:
¨
¨
¨
Resonance is always good!
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Dot structure for SeO2
9
SO2
One final example:
1st question:
What’s the total # of valence electrons?
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What’s the Lewis Dot Structure for SO2?
10
11
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Dot structure for SO2
SO2
Next question:
What’s the central atom?
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18 total valence electrons.
12
Dot structure for SO2
Now?
Fill the octets
S— O
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O—
13
..
..
..
:O — S — O:
¨
¨
¨
Once we’ve filled the octets, what do we do?
Check the total # of valence electrons
20 total electrons – too many!
So, what do we do?
Make a bond!
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Dot structure for SO2
14
Dot structure for SO2
Are we done?
Check the formal charges.
FC(left O) = 6 – 2 – 4 = 0
FC(S) = 6 – 3 – 2 = 1
FC(right O) = 6 – 1 – 6 = -1
Acceptable. Is it the best?
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..
:O = S — O:
¨
¨
¨
15
Dot structure for SO2
S can have an expanded octet, we can make a bond out of 2 electrons
from O without eliminating any S electrons?
Check the formal charges.
FC(O) = 6 – 2 – 4 = 0
FC(S) = 6 – 4 – 2 = 0
EVEN BETTER THAN THE PREVIOUS STRUCTURE!
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:O = S = O:
¨
¨ ¨
16
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Once you have a Lewis Structure, then you can determine the 3D Molecular structure – AND ITS EASY!!!!
17
3-D Molecular Structures
What holds molecules together?
What is on the outside of all molecules?
ELECTRONS
What do we know about electrons?
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ELECTRONS
THEY HATE EACH OTHER
18
Joe hates Jane
Joe and Jane both go to Bob’s party. If I’m in the kitchen,
where’s Jane?
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Joe and Jane hate each other (the reasons are unclear, but Jane
never got a phone call much less the jewelry she thought she
deserved!  )
19
ANYWHERE BUT THE KITCHEN!
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(unless she wants to make out with my brother right in front of
me just to try and p*&& me off!)
20
3-D Molecular Models
One rule – VSEPR
Electrons hate each other, they stay as far away
from each other as possible!!!
One corollary to the rule: non-bonding pairs hate
each other more than bonds.
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Valence Shell Electron Pair Repulsion
21
Consider CO2 – what’s the LDS?
..
..
:O = C = O :
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The rest is all geometry! And only the geometry of central
atoms matters.
22
Look at carbon, how many electron groups does it
have around it?
2 – 2 sets of double bonds (all bonds, whether
single, double, or triple count as a single group)
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..
..
:O = C = O :
23
2 electron groups that hate each other (like Joe and Jane) –
what’s the farthest apart they can get?
Completely opposite sides of the molecule!
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..
..
:O = C = O :
24
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Move either C-O bond and it gets
closer to the other one!
25
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Space-filling model
26
What about CH2O
Central atom?
Carbon
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Total Valence electrons?
4 + 2*1 + 6 = 12
27
..
:O:
|
H— C — H
¨
Total electrons – 14 electrons
What’s the solution?
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Connect the atoms and fill
octets
28
:O:
||
H— C — H
Total electrons – 12 electrons
Formal charges are all ZERO
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LDS
29
:O:
||
H— C — H
Joe, Jane, and Amanda all hate each other (boy
was that a bad night of Tequila!  ) If they all
end up at the same bar, what’s the farthest apart
they can get?
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3-D Geometry
30
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3-D Geometry – Trigonal
planar
31
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Not all bonds are the same
32
H
|
H — C —H
|
H
4 electron groups around the central atom.
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Methane – CH4
33
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Tetrahedral Geometry, a little harder
to see…
34
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It’s basically a pyramid
35
You don’t need to predict it…
2 electron groups – linear
3 electron groups – trigonal planar
4 electron groups – tetrahedral
Geometry demands it!
And, since 4 electron groups is an octet, that’s the
whole story, right?
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…it’s automatic!
36
Expanded Octets
BUT, they are also automatic based on the number of electron
groups that hate each other.
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For S, P that have expanded octets, you can have more
complicated geometries.
37
PCl5
Another one of my favorite molecules!
Cl
Cl — P — Cl
|
Cl
Joe and Amanda and Jane and Brad and
Laura…never mind, you know the drill!
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Cl
38
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Trigonal Bipyramidal
39
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Two pyramids, base-to-base
40
SF6 – love this molecule also!
F
F — S—F
F
F
6 electron groups that hate each other!
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F
41
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Octahedral Geometry
42
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Two 4-sided pyramids, base-to-base
43
2 electron groups – linear
3 electron groups – trigonal planar
4 electron groups – tetrahedral
5 electron groups – trigonal bipyramidal
6 electron groups – octahedral
And the angles are “sort-of” predictable
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Geometry is 3000 years old!
44
Geometry is 3000 years old!
2 electron groups – linear - 180º
3 electron groups – trigonal planar - 120º
5 electron groups – trigonal bipyramidal - 90º, 120º
6 electron groups – octahedral - 90º
This ASSUMES every position is identical – the real world
has nuances
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4 electron groups – tetrahedral – 109.5º
45
Lone pairs vs. bonding pairs
..
H—N—H
|
H
H
|
H — C —H
|
H
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NH3 vs. CH4
46
47
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Water
48
MOLECULAR geometry vs.
ELECTRON geometry
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The molecular geometry is just the geometry of the bonds,
ignoring non-bonding electrons.
49
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Ammonia
50
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Water
51
52
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5
6
6
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Expanded octets
5
5
5
6
53
What is the ELECTRON geometry of NO2-?
A. Tetrahedral
B. Trigonal planar
C. Bent
D. Linear
E. Octahedral
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Clicker question
54
A.
B.
C.
D.
E.
90 degrees
120 degrees
117 degrees
107 degrees
109.5 degrees
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What is the bond angle in NO2-?
55
Clicker question
A.
B.
C.
D.
E.
Tetrahedral
Trigonal planar
Bent
Linear
Octahedral
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What is the molecular geometry of NO2-?
56
Polarity of Molecules
•
in order for a molecule to be polar it must
1) have polar bonds
electronegativity difference - theory
bond dipole moments - measured
2) have an unsymmetrical shape
•
•
polarity affects the intermolecular forces of
attraction
•
therefore boiling points and solubilities
•
•
vector addition
like dissolves like
nonbonding pairs affect molecular polarity, strong
pull in its direction
57
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•
•
What holds molecules together?
Bonds
Bonds are made up of?
Electrons
How do the electrons hold atoms together?
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58
Two ways:

Ionic Bonds – attraction between ions of
opposite charges
Na+ Cl-

Covalent Bonds – sharing of electrons
between adjacent atoms
PF3
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59
Are they really different?
Let’s share a pie!
Mine
Yours
Mine
Yours
Which pie are we actually sharing?
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60
Sharing doesn’t have to be equal!
Mine
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61
Ionic and covalent are part of a
continuum
Ionic
Covalent
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62
Two extremes
Mine
Yours
Ours
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63
Something in the middle
Mine
Yours
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Mine
Yours
64
Ionic and covalent are part of a
continuum
Ionic
Uneven sharing
Polar
Equal sharing
Non-polar
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65
So, consider a bond, any bond:
Cl – Cl
Which case is this?
Equal sharing!
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66
So, consider a bond, any bond:
H-Cl
Which case is this?
Unequal sharing! How do you know?
They are on opposite sides of the Periodic
table!
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67
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68
A metal + a non-metal =
An ionic compound!
Non-metals love electrons, metals don’t!
There is a periodic trend for “electron love”:
electronegativity or electron affinity.
Electronegativity increases to the right and
going up (F is most electronegative, Fr is
least)
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70
Electronegativity determines polarity
The polarity of a bond is determined by the
difference in electronegativity between the
atoms at either end of the bond.
 E.N. = Larger E.N. – smaller E.N.
If  E.N. < 0.4, its considered non-polar
If 0.4 <  E.N. < 2.0, its considered polar
If  E.N. > 2.0, its considered ionic
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71
Cl – Cl
 E.N. = 3.0 – 3.0 = 0
Non-polar
H-Cl
 E.N. = 3.0 – 2.1 = 0.9
Polar
NaCl
 E.N. = 3.0 – 0.9 = 2.1
Ionic
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72
Polarity is represented as an arrow…
…pointing toward the more negative atom.
Cl – Cl
H-Cl
NaCl
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73
The H-Cl bond is polar. The bonding electrons are
pulled toward the Cl end of the molecule. The net result
is a polar molecule.
74
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Molecule Polarity
H2O vs. CO2
O=C=O
What can we say about the bonds?
E.N. = 3.5 – 2.5 = 1.0
Polar
..
H–O–H
¨
What can we say about the bonds?
 E.N. = 3.5 – 2.1 = 1.4
Polar
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75
Does that make sense?
What do you know about water?
Yup, it’s a polar solvent.
What about CO2? Would you expect it to be
polar?
Would it surprise you to learn that CO2 is
NONPOLAR?!?!?!?
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76
It’s all in the geometry.
CO2 has two polar
bonds, but…
the polarity is
equal and
oppositely
directed, so it
cancels!
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77
H2O
Water also has two
equally polar
bonds, but they
aren’t pointing in
opposite
directions.
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78
H2O
Polarity is a “vector”,
it has size and
direction. You can’t
separate the two.
Think of it as travel
directions.
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79
If I leave my house and go 1 mile North and
then 1 mile South, where am I?
1 mile South
1 mile North
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80
If I leave my house and go 1 mile North,
and then 1 mile West, where am I?
1 mile West
1.414 mi NW
1 mile North
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81
H2O
A polarity vector is
just the direction
that a proton would
go (toward the
negative), and the
length of the vector
is its magnitude.
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82
H2O
The polarity of the
molecule is distinct
from the polarity of
the bonds in the
molecule.
Net
Dipole
Non-polar bonds =
Non-polar molecule
Polar bonds…depends
on the geometry!
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83
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Vector Addition
84
85
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The O-C bond is polar. The bonding electrons are
pulled equally toward both O ends of the molecule. The
net result is a nonpolar molecule.
86
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Molecule Polarity
The H-O bond is polar. The both sets of bonding
electrons are pulled toward the O end of the
molecule. The net result is a polar molecule.
87
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Molecule Polarity
The H-N bond is polar. All the sets of bonding
electrons are pulled toward the N end of the
molecule. The net result is a polar molecule.
88
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Molecule Polarity
N=3.0, H = 2.1
They don’t cancel, all 3 are
pointing up and nothing is
pointing down!
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EN=0.9 (polar)
89
Clicker
Is CCl4…
A.
B.
C.
Polar
Non-polar
Ionic
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90
Clicker
Is CH2Cl2…
A.
B.
C.
D.
E.
Slightly Polar
Really polar
Non-polar
I hate Christmas
I hate you!
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91
Lewis structure
Cl
Cl
C
Cl
Cl
The bonds are…
Polar – C=2.5, Cl=3.0 EN=0.5 (weakly polar)
The molecular geometry is…
Tetrahedral
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92
Do the polarities cancel?
A 3-D model helps…
Cl
Cl
C
Cl
Cl
Cl
C
C
Cl
Doesn’t completely cancel because they aren’t pointing directly opposite.
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93
Clicker
Is HCN…
A.
B.
C.
D.
E.
F.
Polar
Non-polar
Ionic
I hate Hanukkah
I hate Kwanzaa
I hate you.
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94
H
C
N:
H = 2.1
C = 2.5
N = 3.0
C-H is nonpolar (EN=0.4)
C-N is polar (EN=0.5)
Molecule is weakly polar.
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95
A very brief review of geometry
We need:
1. Electron configuration – to get the
number of valence electrons
2. Lewis structure – to determine “where”
the valence electrons are.
3. 3-D Molecular structure – VSEPR
determines shape of the molecules by
determining position of electron groups
4. Polarity – unequal sharing of electron
density creates polar bonds and (maybe)
polar molecules.
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96
Let’s try NOCl – polar or non-polar?
First we need the Lewis structure
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97
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98
Let’s try NOCl – polar or non-polar
First we need the Lewis structure
N has 5 valence electrons (2s22p3)
O has 6 valence electrons (2s22p4)
Cl has 7 valence electrons (3s23p5)
That’s a total of 18 valence electrons!
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99
What’s the most likely central atom?
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100
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101
Leftmost or down most.
N is the farthest to the left of the periodic
table. N is the most likely central atom.
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102
Stick everything to N and fill octets.
N
O
Cl
Check the number of valence electrons used
20! That’s not 18!
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103
How do I get rid of electrons?
Make a double bond!
But where…?
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104
When in doubt, try them all!
O
N
Cl
O
N
Cl
So, which one is it? Or is it resonance?!?!?
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105
Let formal charges sort them out!
O
6-1-6=-1
O
6-2-4=0
N
5-3-2=0
N
5-3-2=0
Cl
7-2-4=+1
Cl
7-1-6=0
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106
How many electron groups around
N?
A.
B.
C.
D.
E.
3
4
None of the above
Both A and B
Neither A Nor B nor C
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107
We have a winner!
O
N
Cl
Is it polar or non-polar?
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108
We need to consider the bonds and
the shape.
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109
First…the bonds.
EN
O = 3.5
N = 3.0
Cl = 3.0
••
•O
•
••
N
••
Cl ••
••
Calculate the electronegative difference for EACH
bond
N=O ΔEN = 3.5-3.0 = 0.5 (barely polar)
N-Cl ΔEN = 3.0-3.0 = 0.0 (not polar)
So I have a polar bond and nothing to cancel it, so
the molecule must be weakly polar. I really don’t
even need to consider
shape, but let’s look anyway!
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3 Electron groups around the central atom
••
N
••
•O
•
••
Cl ••
••
Electron geometry – Trigonal planar
Molecular geometry - bent
3.0
3.0
Cl
N
O3.5
1) polar bonds, N-O
2) asymmetrical shape
polar
111
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Clicker question
What is the O-N-Cl bond angle?
A.180 degrees
B.123 degrees
C.120 degrees
D.117 degrees
E.90 degrees
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112
Let’s try one more
SO3 – polar or non-polar?
First we need the Lewis structure
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113
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114
Let’s try SO3 – polar or non-polar
First we need the Lewis structure
S has 6 valence electrons (3s23p4)
O has 6 valence electrons (2s22p4)
That’s a total of 24 valence electrons!
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115
What’s the most likely central atom?
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116
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117
Leftmost or down most.
S and O are equally to the left of the
periodic table. S is below O. S is the most
likely central atom.
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118
Stick everything to S and fill octets.
S
O
O
O
Check the number of valence electrons used
26! That’s not 24!
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119
How do I get rid of electrons?
Make a double bond!
But where…?
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120
When in doubt, try them all!
O
S
O
O
O
S
O
O
O
S
O
O
So, which one is it? Or is it resonance?!?!?
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121
Check the formal charges
O
S
O
O
O
S
O
O
O
S
O
O
All equivalent S: 6-4-0=+2
O: 6-2-4=0
O:6-2-6=-1
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122
It’s resonance!
O
S
O
O
O
S
O
O
O
S
O
O
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123
We need to consider the bonds and
the shape.
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124
First the bonds:
EN
O = 3.5
S = 2.5
••
•O
•
••
••
•O•
• •
S
••
O ••
ΔEN=3.5-2.5=1.0 (all 3 bonds are polar!)
125
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But let’s check the geometry!
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126
Ah…symmetry!
••
•O
•
••
3.5
3.5
••
•O•
• •
S
Trigonal
Planar
O
S
••
O ••
2.5
3.5
O
O
1) polar bonds, all S-O
2) symmetrical shape
nonpolar
127
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This is the starting point…
Polarity influences how molecules interact with each
other.
Molecular interactions influence physical properties of
materials.
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128
• polar molecules are attracted to other polar
molecules
• since water is a polar molecule, other polar
molecules dissolve well in water
• and ionic compounds as well
• some molecules have both polar and nonpolar
parts
129
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Molecular Polarity Affects
Solubility in Water
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A Soap Molecule
Sodium Stearate
130
2 minutes on Chapter 11
It’s all just dipoles – but not always from polar
molecules!
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Intermolecular Forces!
131
Instantaneous dipole-induced dipole
forces – not the strongest but ALWAYS
present
Br
δ-
Instantaneous dipole
δ+
Br
Br
δ+
δ-
Br
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Induced dipole
The electron cloud is
mobile.
Charge density is
constantly moving
132
around
133
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A.
B.
C.
D.
E.
Oil molecules hate water molecules.
Oil molecules love water molecules
Oil molecules are indifferent to water molecules
Save this question for Chapter 11
Oil molecules
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Which of the following is most
true?
134
Dipole – Dipole Interactions
δ-
Permanent dipole
δ+
H
H
δ+
δ-
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Br
Permanent dipole
Br
135
Permanent dipole-induced dipole
forces
Br
δ-
Permanent dipole
δ+
H
Br
δ+
δ-
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Induced dipole
Br
136
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A Soap Molecule
Sodium Stearate
137
• polar molecules are attracted to other polar
molecules
• since water is a polar molecule, other polar
molecules dissolve well in water
• and ionic compounds as well
• some molecules have both polar and nonpolar
parts
138
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Molecular Polarity Affects
Solubility in Water
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Back to our regularly
scheduled program…
139
• Lewis theory gives good first approximations of the bond angles in
molecules, but usually cannot be used to get the actual angle
• Lewis theory cannot write one correct structure for many molecules
where resonance is important
• Lewis theory often does not predict the correct magnetic behavior of
molecules
• e.g., O2 is paramagnetic, though the Lewis structure predicts it is
diamagnetic
140
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Problems with Lewis Theory
Valence Bond Theory
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• Linus Pauling and others applied the principles of quantum mechanics
to molecules
• they reasoned that bonds between atoms would arise when the
orbitals on those atoms interacted to make a bond
• the kind of interaction depends on whether the orbitals align along
the axis between the nuclei, or outside the axis
141
Orbital Interaction
• the interaction energy between atomic orbitals is negative when the
interacting atomic orbitals contain a total of 2 electrons
142
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• as two atoms approached, the partially filled or empty valence atomic
orbitals on the atoms would interact to form molecular orbitals
• the molecular orbitals would be more stable than the separate atomic
orbitals because they would contain paired electrons shared by both
atoms
Orbital Diagram for the
Formation of H2S
1s
↑
+ ↑↓
1s
↑
3s
↑
↑ ↑↓ S
3p
↑↓
H-S bond
↑↓
H-S bond
H
Predicts Bond Angle = 90°
Actual Bond Angle = 92°
143
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H
Valence Bond Theory - Hybridization
• C = 2s22px12py12pz0 would predict 2 or 3 bonds that are 90° apart, rather
than 4 bonds that are 109.5° apart
• to adjust for these inconsistencies, it was postulated that the valence
atomic orbitals could hybridize before bonding took place
• one hybridization of C is to mix all the 2s and 2p orbitals to get 4 orbitals
that point at the corners of a tetrahedron
144
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• one of the issues that arose was that the number of
partially filled or empty atomic orbital did not predict
the number of bonds or orientation of bonds
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Unhybridized C Orbitals Predict the
Wrong Bonding & Geometry
145
1.
2.
the valence electrons in an atom reside in the quantum
mechanical atomic orbitals or hybrid orbitals
a chemical bond results when these atomic orbitals overlap and
there is a total of 2 electrons in the new molecular orbital
a)
3.
the electrons must be spin paired
the shape of the molecule is determined by the geometry of the
overlapping orbitals
146
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Valence Bond Theory
Main Concepts
Hybridization
• some atoms hybridize their orbitals to maximize bonding
• better explain observed shapes of molecules
• same type of atom can have different hybridization depending on the
compound
• C = sp, sp2, sp3
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• hybridizing is mixing different types of orbitals to make a new set of
degenerate orbitals
• sp, sp2, sp3, sp3d, sp3d2
• more bonds = more full orbitals = more stability
A.
B.
C.
D.
E.
1
2
3
4
6
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How many hybrid orbitals does
water need?
148
• H cannot hybridize!!
• the number of standard atomic orbitals combined =
the number of hybrid orbitals formed
• the number and type of standard atomic orbitals
combined determines the shape of the hybrid orbitals
• the particular kind of hybridization that occurs is the
one that yields the lowest overall energy for the
molecule
• in other words, you have to know the structure of the
molecule beforehand in order to predict the hybridization
149
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Hybrid Orbitals
Orbital Diagrams with
Hybridization
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• place electrons into hybrid and unhybridized valence orbitals
as if all the orbitals have equal energy
• when bonding, s bonds form between hybrid orbitals and p
bonds form between unhybridized orbitals that are parallel
150
Carbon Hybridizations

 
2p
2s
sp hybridized
 
2sp
 
2p
sp2 hybridized

 
2sp2
sp3 hybridized



2sp1513

2p

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Unhybridized
sp3 Hybridization
• atom with 4 areas of electrons
• tetrahedral geometry
• 109.5° angles between hybrid orbitals
H
s
H
sp3 •• sp3
C
N
H
s
H
H
152
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• atom uses hybrid orbitals for all bonds and lone pairs
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sp3 Hybridization of C
153
154
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A.
B.
C.
D.
E.
6
5
4
3
2
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How many electron groups in
CO2
155
sp3 Hybridized Atoms
Orbital Diagrams
2s

2s
 
2p
C

  
2p
N

156


  
2sp3
  
2sp3
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sp3 hybridized atom
Unhybridized atom
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Methane Formation with sp3 C
157
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3
Ammonia Formation with sp
N
158
CH3NH2 Orbital Diagram
s

s


1s H
1s H

s



s
1s H
159

s
1s H
1s H
sp3 N
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



C
s

sp3
Practice - Draw the Orbital Diagram for
the sp3 Hybridization of Each Atom

3s

2s
  
3p
Cl
  
2p
O
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Unhybridized atom
160
Practice - Draw the Orbital Diagram for
the sp3 Hybridization of Each Atom
Unhybridized atom

2s
Cl
   
3sp3
  
2p
O
   
2sp3
161
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3s
  
3p


sp3 hybridized atom
Types of Bonds
• a sigma (s) bond results when the bonding atomic
orbitals point along the axis connecting the two
bonding nuclei
• s-to-s, p-to-p, hybrid-to-hybrid, s-to-hybrid, etc.
• a pi (p) bond results when the bonding atomic
orbitals are parallel to each other and perpendicular
to the axis connecting the two bonding nuclei
• between unhybridized parallel p orbitals
• the interaction between parallel orbitals is not as
strong as between orbitals that point at each other;
therefore s bonds are stronger than p bonds
162
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• either standard atomic orbitals or hybrids
163
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164
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165
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166
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• because orbitals that form the s bond point along the internuclear
axis, rotation around that bond does not require breaking the
interaction between the orbitals
• but the orbitals that form the p bond interact above and below the
internuclear axis, so rotation around the axis requires the breaking of
the interaction between the orbitals
167
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Bond Rotation
168
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169
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sp2
• atom with 3 areas of electrons
• trigonal planar system
• C = trigonal planar
•
•
H
s
N = trigonal bent
O = “linear”
••
O ••
C
sp2
• 120° bond angles
• flat
• atom uses hybrid orbitals for s bonds and lone
pairs, uses nonhybridized p orbital for p bond
170
••
O
•• 3
sp
H
s
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sp2
171
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p
sp2
sp2
Hybrid orbitals overlap to form s bond
Unhybridized p orbitals overlap to form p bond
172
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s
+
173
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sp2 Hybridized Atoms
Orbital Diagrams
Unhybridized atom

2s
  
2p

N
2s
1p

174
 
2sp2

2p
 
2sp2

2p
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2s
 
2p
C
3s
1p


sp2 hybridized atom
CH2NH Orbital Diagram


s




s
s


C
s
pN

sp2

1s H
1s H
1s H
175
sp2 N
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
pC
p
Practice - Draw the Orbital Diagram for
the sp2 Hybridization of Each Atom. How many s
and p bonds would you expect each to form?
Unhybridized atom

B
2s
2p

  
2p
2s
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
O
176
Practice - Draw the Orbital Diagram for
the sp2 Hybridization of Each Atom. How many s
and p bonds would you expect each to form?
Unhybridized atom
2s
2p

  
2p
2s

 
2sp2
2p
O
1s
1p
  
2sp2

2p
177
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
B
3s
0p


sp2 hybridized atom
sp
• atom with 2 areas of electrons
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• linear shape
• 180° bond angle
• atom uses hybrid orbitals for s bonds or lone pairs, uses
nonhybridized p orbitals for p bonds
H
s
178
C
sp
p
s
p
N
sp
179
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180
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sp Hybridized Atoms
Orbital Diagrams
2s

2s
 
2p
  
2p
N
1s
2p
181
 
2sp
 
2p

2sp
 
2p

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
C
2s
2p
sp hybridized atom

Unhybridized atom
HCN Orbital Diagram


s
s

sp C

1s H
182


pN


sp N
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
pC
2p
• atom with 5 areas of electrons around it
• trigonal bipyramid shape
• See-Saw, T-Shape, Linear
• 120° & 90° bond angles
• use empty d orbitals from valence shell
• d orbitals can be used to make p bonds
183
••
•• O
••
-1
••
•• F •• •
•
O•
•
I
••
••O
•• F •
•
••
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3
sp d
184
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sp3d Hybridized Atoms
Orbital Diagrams
sp3d hybridized atom

3s

3s
  
3p
  
3p
P

3d
  
3sp3d
S  
3d
(non-hybridizing d orbitals not shown)
185
 
3sp3d

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Unhybridized atom

sp3d2
• atom with 6 areas of electrons around it
• use empty d orbitals from valence shell
• d orbitals can be used to make p bonds
••F••
••
••F
•• •
•
186
••
•• F ••
Br •
•
•• F •
•
••
••F••
••
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• octahedral shape
• Square Pyramid, Square Planar
• 90° bond angles
187
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sp3d2 Hybridized Atoms
Orbital Diagrams
sp3d2 hybridized atom
Unhybridized atom
↑↓ ↑ ↑
3s
3p
↑↓
↑↓ ↑↓ ↑
5s
5p
S
↑ ↑ ↑ ↑ ↑ ↑
3sp3d2
3d
I
↑↓ ↑ ↑ ↑ ↑ ↑
5sp3d2
5d
(non-hybridizing d orbitals not shown)
188
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↑↓
189
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H
••
O
••
••
•O
•
B
H
••
O
••
H
H = can’t hybridize
B = 3 electron groups = sp2
O = 4 electron groups = sp3
190
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Example - Predict the Hybridization of
All the Atoms in H3BO3
Practice - Predict the Hybridization and
Bonding Scheme of All the Atoms in
NClO
••
N
••
Cl ••
••
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••
•O
•
191
••
•O
•
••
N
••
Cl ••
••
N = 3 electron groups = sp2
O = 3 electron groups = sp2
Cl = 4 electron groups = sp3
192
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Practice - Predict the Hybridization and
Bonding Scheme of All the Atoms in
NClO
1) Start by drawing the Lewis Structure
2) Use VSEPR Theory to predict the electron
group geometry around each central atom
3) Use Table 10.3 to select the hybridization
scheme that matches the electron group
geometry
4) Sketch the atomic and hybrid orbitals on the
atoms in the molecule, showing overlap of the
appropriate orbitals
5) Label the bonds as s or p
193
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Predicting Hybridization and
Bonding Scheme
Ex 10.8 – Predict the hybridization and
bonding scheme for CH3CHO
Draw the Lewis Structure
H O
H
C1 C2 H
H
Predict the electron group
geometry around inside
atoms
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C1 = 4 electron areas
 C1= tetrahedral
C2 = 3 electron areas
 C2 = trigonal planar
194
Ex 10.8 – Predict the hybridization and
bonding scheme for CH3CHO
Determine the hybridization C1 = tetrahedral
3
of the interior atoms

C1
=
sp
H O
C2 = trigonal planar
H C C H
H
 C2 = sp2
Sketch the molecule and
orbitals
1
2
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195
Ex 10.8 – Predict the hybridization and
bonding scheme for CH3CHO
Label the bonds
H O
H
C1 C2 H
H
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196
Problems with Valence Bond
Theory
• VB theory predicts many properties better than Lewis Theory
• bonding schemes, bond strengths, bond lengths, bond rigidity
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• however, there are still many properties of molecules it
doesn’t predict perfectly
• magnetic behavior of O2
197