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Chapter 10 Notes
Chemical Bonding II
Molecular Shape,
Valence Bond Theory, and
Molecular Orbital Theory
Sections 10.1 – 10.8
I. Artificial Sweeteners: Fooled by Molecular Shape
Artificial sweeteners such as aspartame (Nutrasweet) taste sweet like
sugar but do not have the calories of sugar. This is because taste and
nutritional value are independent properties. The caloric value of
food depends on the amount of energy released when the food is
metabolized in the body. Many artificial sweeteners are not even
metabolized by the body- they just pass straight though.
The taste of a food begins with specialized cells on the tongue that
detect different molecules of food. These cells are so sensitive that
the tongue can detect one molecule of sugar out of thousands of
different molecules in a bite of food. We experience certain tastes
when molecules of food fit into a special part of a taste receptor on
our tongue, much in the same way a key fits into a lock. Artificial
sweeteners “trick” our tongues into tasting sweetness because these
molecules mimic the molecular shape of the sugar molecule and fit
snuggly into our taste receptors.
- used to determine the 3- D shape of a molecule
- based on the idea that electrons groups (bonds & lone pair e─)
repel one another to varying degrees
- the combination of repulsions on the central atom of a molecule
determines its 3-D shape
• VSEPR THEORY (VALENCE SHELL ELECTRON PAIR REPULSION):
II. VSEPR Theory: The Five Basic Shapes
*The basic idea behind VSEPR is that repulsions between electron
groups determine molecular geometry. The preferred geometry is
the one in which the electron groups have the maximum separation
(and therefore the minimum energy) possible.
Preview
We will first look at molecular geometries where there are two to
six electron groups around a central atom and all of the electron
groups are bonding groups (single or double bonds). We will then
look at what happens to the shapes when some of the electron
groups become lone pair electrons.
A. TWO ELECTRONS GROUPS: LINEAR GEOMETRY
Bond Angles: 180°
Other examples:
SiO2
CS2
Si O O
C S S
16
16
B. THREE ELECTRONS GROUPS: TRIGONAL PLANAR
GEOMETRY
Bond Angles: 120° Other examples:
BH3 SO3
B
H
H H
S
O
O O
In CH2O, the bond angles deviate slightly from
120° because there is more electron density in a
double bond than in a single bond.
Different electron groups repel each other in
slightly different ways.
6 24
C. FOUR ELECTRONS GROUPS: TETRAHEDRAL GEOMETRY
Bond Angles: 109.5°
Other examples:
SiCl4 CF4
Si
Cl
Cl Cl
Cl
C
F
F F
F
32 32
D. FIVE ELECTRONS GROUPS: TRIGONAL BIPYRAMIDAL
GEOMETRY
Bond Angles: 90°, 120°
Other examples:
SOF4 ClO2F3
S
O
F F
F
F
40 40
Cl
O
F O
F
F
E. SIX ELECTRONS GROUPS: OCTAHEDRAL GEOMETRY
Bond Angles: 90°
Other examples:
SCl6
S
Cl
Cl Cl
Cl
Cl
Cl
48
III. VSEPR Theory: The Effect of Lone Pairs
Preview
We will now look at molecular geometries where there are four to
six electron groups around a central atom and some of the electron
groups are lone pairs. Keep in mind that these shapes are all
variations of the geometries from the last section, but now one or
more bonding groups have been replaced with lone pairs. The
difference in shapes results from the fact that lone pair electrons
repel other lone pair and bonding electrons to a greater extent than
bonding electrons repel one another.
THE ORDER OF ELECTRON PAIR REPULSION:
lone pair - lone pair > lone pair- bonding pair > bonding pair- bonding pair
Most repulsive Least repulsive
* Want to be as far away
from each other as
possible
• In this section, don’t get confused between electron and molecular
geometries.
• In the last section, the electron and molecular geometries were the same
(same name for each).
• In this section, they are different.
• There are only 5 electron geometries, but there are 11 molecular
geometries.
THE DIFFERENCE BETWEEN:
Electron Geometry:
Molecular Geometry:
The geometrical arrangement of electron groups
The geometrical arrangement of the atoms
A. FOUR ELECTRON GROUPS WITH LONE PAIRS
(TETRAHEDRAL ELECTRON GEOMETRY)
Bond Angles: 107°
TRIGONAL PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR
Other examples:
PF3 NCl3
P
F F F
N
Cl Cl Cl
26 26
Bond Angles: 104.5°
BENT (MOLEC. GEO.): TWO LONE PAIRS
Other examples:
SCl2 H2S
A. FOUR ELECTRON GROUPS WITH LONE PAIRS
(TETRAHEDRAL ELECTRON GEOMETRY)
S Cl Cl
S H H
18 8
Summarizing Tetrahedral Electron Geometries
B. FIVE ELECTRON GROUPS WITH LONE PAIRS
(BIPYRAMIDAL ELECTRON GEOMETRY)
Bond Angles: 90°, 120°
SEESAW (MOLEC. GEO.): ONE LONE PAIR
Other examples:
SeCl4 IOF3
Se
Cl
Cl
Cl
Cl
I
O
F
F
F
34
Lone pair must go equatorial
34
B. FIVE ELECTRON GROUPS WITH LONE PAIRS
(BIPYRAMIDAL ELECTRON GEOMETRY)
Bond Angles: 90°, 120°
T-SHAPED OR LINEAR (MOLEC. GEO.): TWO OR THREE
LONE PAIRS
Other examples:
ClF3 I3
-
Cl
F
F
F I
I
I
28 28
Lone pairs must go equatorial
C. SIX ELECTRON GROUPS WITH LONE PAIRS
(OCTAHEDRAL ELECTRON GEOMETRY)
Bond Angles: 90°
SQUARE PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR
Other examples:
XeOF4 BrF4-
Xe
O
F
F
F
F
Br F
F
F
F
SQUARE PLANAR (MOLEC. GEO.): TWO LONE PAIRS
Lone pairs must go axial
42 36
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