Transcript XI: Advanced Bonding Concepts

Topic 12 : Advanced Bonding Concepts
LECTURE SLIDES
• VSEPR shape theory
• Bond Polarity
• Molecular Polarity
Kotz & Treichel, 9.6- 9.8
Molecular and Polyatomic Ion Shapes
Once a Lewis structure is drawn, the three dimensional geometry of the species can easily
be determined by utilizing the “valence shell electron
pair repulsion theory” called “VSEPR”:
“VSEPR” theory is based on the tendency of negatively
charged regions to repel each other and align as far
apart as possible, resulting in predictable shapes for
any covalently bonded species.
To utilize “VSEPR”, the number of regions of electron
density around the central atom in the species is
counted.
Count as “one region”:
• Single Bonds
• Unshared Pairs
• Multiple bonds between same two atoms
Examples of “four regions”:
A
A
“two regions”:
“three regions”:
A
A
A
A
A
Basic Shapes predicted by VSEPR:
Two regions:
o
180
Three Regions:
Bond Angles
A
A
linear
120o
Geometry
trigonal planar
Five Regions:
Four Regions:
109.5o
90o
120o
A
A
Six Regions:
trigonal-bipyramidal
tetrahedral
90o
A
octahedral
trigonal
trigonal-bipyramidal
tetrahedral
octahedral
Before we begin, some guidelines about forming
double and triple bonds in Lewis structures:
C, N, O, S form double and triple bonds and
never show incomplete octets (less than 8
e’s)
Metals, metalloids, and halogens do not as a
rule form multiple bonds. Compounds
containing these elements will often show an
incomplete octet around the central
atom.
Type One: Two Regions
Examples: BeCl2, CO2, NO2+
BeCl2
Cl
Be
Be 2
2Cl 14
Cl
16 e's/2= 8 prs
Cl
Be
Cl
(octet violator)
Number of regions around CENTRAL ATOM: 2
Cl
Be
Cl
shape : LINEAR
bond angles: 180o
CO2
O
C
C 4
2O 12
O
16 e's/2= 8 prs
O
C
O
O
C
O
Number of regions around CENTRAL ATOM: 2
O
C
O
shape : LINEAR
bond angles: 180o
NO 2
+
O
N
N 5
2O 12
+1 -1e
O
16 e's/2= 8 prs
O
N
O
O
N
O
Number of regions around CENTRAL ATOM: 2
+
O
N
O
shape : LINEAR
bond angles: 180o
Type Two: Three Regions
NO3-, NO2-
N
5
3O 18
11
O
NO 3-
O
N
O
24 e's/2=12 prs
O
O
N
-
O
O
O
N
O
(three resonance structures)
Number of regions around CENTRAL ATOM: 3
-
O
shape : TRIGONAL PLANAR
bond angles: 120o
N
O
O
NO 2-
O
N
N
5
2O 12
11
O
18 e's/2=9 prs
O
N
O
O
N
O
(two resonance structures)
Number of regions around CENTRAL ATOM: 3
shape : TRIGONAL PLANAR
bond angles: 120o
N
O
O
N
O
O
Black orbital
indicates pair of
unshared e’s
NOTE: “molecular geometry” (bonds only): BENT
Group Work 12.1: Geometry, 2,3 Regions
1. Do a Lewis Structure for HCN and CH2O.
2. Draw each molecule “to shape”
3. Describe geometry and bond angles for each
Type Three: Four Regions
CH2Cl2, NH3, H2O, NH4+
C
2Cl
2H
H
CH2Cl2
H
C
Cl
4
14
2
20 e's/2=10 prs
Cl
H
H
C
Cl
Cl
Number of regions around CENTRAL ATOM: 4
H
shape : TETRAHEDRAL
bond angles: 109.5o
C
H
Cl
Cl
N
3H
NH 3
H
N
H
5
3
8e's/2=4 prs
H
H
N
H
H
Number of regions around CENTRAL ATOM: 4
shape : TETRAHEDRAL
bond angles: < 109.5o
N
H
H
H
Note: molecular geometry, trigonal pyramid
As is turns out, unshared pairs of electrons around
the central atom are not held in place between two
atoms as bonded pairs are.
They tend to occupy more space and to be somewhat
more “repulsive” than bonded pairs.
When grouped with bonded pairs to tiny atoms like H,
they tend to distort the bond angles, pushing the
bonded pairs closer together.
The bond angles in ammonia are closer to 107o.
O
2H
H2O
H
O
H
6
2
8e's/2=4 prs
H
O
H
Number of regions around CENTRAL ATOM: 4
O
H
shape : TETRAHEDRAL
bond angles: < 109.5o
(~105o)
H
Note: molecular geometry: BENT
N
4H
1+
H
NH 4+
H
N
H
5
4
-1
8e's/2=4 prs
H
Number of regions around CENTRAL ATOM: 4
H
+
N
H
H
H
shape : TETRAHEDRAL
bond angles: 109.5o
GROUP WORK 12.2: Geometry, 3, 4 Regions
Do Lewis structure and assign shape and bond angles:
CO32-, SiCl4
Type Four: Five Regions
PF5 , ClF3 , IF2-, SF4
P
5F
F
F
PF5
F
P
F
F
5
35
40e's/2=20 prs
Number of regions around CENTRAL ATOM: 5
F
F
P
shape : Trigonal Bipyramidal
bond angles: 120o, 90o
F
F
F
F
F
Bond angles
in “equatorial”
position 120o
F
P
F
F
Bond angles,
each “axial” F,
90o from trigonal
plane
Cl
3F
F
ClF3
Cl
F
7
21
28e's/2=14 prs
F
Number of regions around CENTRAL ATOM: 5
F
Cl
F
shape : Trigonal Bipyramidal
bond angles: 90o
F
Note: “T-shaped”; unshared pairs always “equatorial”
I
7
2F 14
11
F
IF2-
I
F
22e's/2=11 prs
Number of regions around CENTRAL ATOM: 5
F
I
shape : Trigonal Bipyramidal
bond angles: 180o
F
Note: “linear” molecular geometry
S
6
4F 28
F
SF4
S
F
F
34e's/2=17 prs
F
Number of regions around CENTRAL ATOM: 5
F
F
S
shape : Trigonal Bipyramidal
bond angles: 90o, 120o
F
F
Note: “Seesaw” molecular geometry
Type Five: Six Regions
SF6, IF5, XeF4
F
F
SF6
F
S
6
6F 42
S
F
F
F
48e's/2=24prs
Number of regions around CENTRAL ATOM: 6
F
F
F
S
F
F
F
shape : OCTAHEDRAL
bond angles: 90o
F
IF5
F
I
F
F
F
I
7
5F 35
42e's/2=21 prs
Number of regions around CENTRAL ATOM: 6
F
F
I
F
shape : OCTAHEDRAL
bond angles: 90o
F
F
Note: molecular geometry “square pyramidal”
F
XeF4
F
Xe
F
F
Xe
8
4F 28
36e's/2=18 prs
Number of regions around CENTRAL ATOM: 6
F
F
Xe
F
F
Note: “Square planar”
shape : OCTAHEDRAL
bond angles: 90o
GROUP WORK 12.3: Geometry, 5, 6 Regions
Do Lewis structure and assign shape and bond angles:
ICl4+, XeOF2, ICl4Note: O in XeOF2 is equatorial, experimental
evidence
To see relevance of “shape work”, let’s turn next
to bond and molecular polarity. To help examine
this topic we turn back to the property of
“electronegativity”:
ELECTRONEGATIVITY
The trends in ionization energies and electron affinities
can be thought of as summarized in a single property
called “electronegativity” (en or X).
Electronegativity is a unit-less set of assigned values
on a scale of 0 --> 4 describing the ability of an atom to
attract electrons to itself.
The values reaches a maximum at fluorine, with an X =4.
Nonmetals have the largest values, metals the lowest.
Noble gases have no assigned X value.
Most active non-metals
INCREASES
>2
1.5-1.9
.8-1.4
Most active metals
ELECTRONEGATIVITY
ELECTRONEGATIVITY VALUES, MAIN GROUP
ELEMENTS
H
2.1
Li
1.0
Na
1.0
K
0.9
Rb
0.9
Cs
0.8
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Ba
1.0
B
2.0
Al
1.5
Ga
1.7
In
1.6
Tl
1.6
C
2.5
Si
1.8
Ge
1.9
Sn
1.8
Pb
1.7
N
3.0
P
2.1
As
2.1
Sb
1.9
Bi
1.8
O
3.5
S
2.5
Se
2.4
Te
2.1
F
4.0
Cl
3.0
Br
2.8
I
2.5
We have classified bonds “ionic” and “covalent”,
depending on whether electron pairs are shared or
electrons are completely transferred from one atom
to another.
In actuality, there is no sharp dividing line between the
two types but rather a continuum:
Evenly shared
electrons
Unevenly shared
electrons
Transferred
electrons
To determine where a bond lies in this “continuum”,
it is useful to consider the difference in electronegativity
(
X) between the two atoms making up the bond:
When the difference ( X) is close to zero, sharing is fairly
even and electrons are not much closer to one atom than
the other. The bonds are considered “non-polar.”
When the difference is above zero to about 1.7, the
electrons are closer to the more electronegative atom
and partial charge buildup, polarization, develops.
When a metal or polyatomic cation is present and
the ( X) is 1.7 or higher, ionic bonding becomes
the more likely bond type and valence electrons
are transferred to the more electronegative atom.
X
X
BOND
TYPE:
Cl
Cl
H
H
H
Cl
Na+
3.0
3.0
2.1
2.1
2.1
3.0
1.0
0
(NONPOLAR)
COVALENT
0
.9
POLAR
COVALENT
Cl3.0
2.0
IONIC
So, we need to consider a third more specialized
type of bond, “the polar covalent bond:”
This type of bond will be the important factor to
be considered when we look at molecular polarity,
which arises from molecular shape and bond
polarity.
The polar molecular in turn will exhibit different
solubilities and boiling points than non polar
molecules.
Let us consider the bond between H and Cl in a
molecule of hydrogen chloride (only hydrochloric
acid when in water!):
Orbital between H and Cl
H
+
Cl
H
Cl
H
E pair closer to Cl,
more electronegative
Cl
The electron cloud from the pair of shared
electrons is more dense closer to the chlorine,
and much less dense closer to the hydrogen.
The bond has become “polarized”: it has
developed a region (or “pole”) of partial
positive charge buildup and a region (or
“pole”) of partial negative buildup.
H
Cl
Major portion of
electron density
“partially
positive”
~+
H
“partially
negative”
~Cl
or
~+
H
~Cl
Arrow to indicate polar
bond, pointing to more
(-) atom
The molecule has only one bond, and it is polar.
This makes the entire molecule a “dipole”, one which
has a positive and negative pole and will align in an
electrical or magnetic field:
H
Cl
All diatomic
molecules with
polarized bonding
between the two
atoms are DIPOLES.
Other examples of diatomic dipoles:
CARBON MONOXIDE
C
O
2.5
X
X
3.5
X
I
F
2.5
4.0
X
1.0
C
IODINE FLUORIDE
O
1.5
I
F
MOLECULAR POLARITY, LARGER MOLECULES
All diatomic molecules with polar bond(s) are dipoles,
but the situation is not so simple for larger molecules.
There are two factors to consider:
• Are the bonds polar?
• Are they arranged so that the center of positive
charge and the center of negative charge do not
“coincide”?
N
NH 3
H
N
H
H
H
H
H
X
X
3.0
2.1
0.9
BOND POLARITY
N
H
H
H
MOLECULAR
POLARITY
Cl
Cl
H
C
C
H
Cl
H
Cl
X
H
3.0
X
Center of +,charges
coincide,
center of
molecule
2.5
0.5 Polar
X
X
2.1
0.4 polar
H
Cl
C
Cl
2.5
H
Cl
C
H
H
Cl
NONPOLAR
MOLECULES
Center of
Charge
Center of
Charge
POLAR BONDS, NON-POLAR MOLECULE
H
CH2Cl2
C
H
Cl
Cl
C
X
X
H
C
2.5
2.1
X
X
0.4
Cl
2.5
3.0
0.5
H
H
Cl
Cl
DIPOLE
CO32-
2-
O
C
O
X
O
3.5
X
2.5
1.0
2-
O
O
~-
2-
C
C
O
Centers Coincide,
no dipole
O
O
~-
O
~-
O
CH2O
C
H
X
X
C
O
2.5
3.5
H
C
X
X
1.0
H
2.1
2.5
0.4
O
DIPOLE
C
H
H
H
H2O
X
X
O
H
3.5
2.1
1.4
O
H
DIPOLE
H
In conclusion, to be a dipole, a polar molecule (or
polyatomic ion), the presence of polar bonds is
required.
However, in addition, the polar bonds must be
arranged so that they are not canceling.
Molecular shape must be such that the center of the
negative charge buildup does not coincide with the
center of positive charge buildup.
Determine the polarity of each of the below:
-
Group Work
12.4
•Draw to shape
•Check  en
•Draw arrow, if
dipole
NO 2-
NO 2
+
O
N
O
+
O
N
O
-
O
NO 3O
N
O
Importance of Polarity
As it turns out, it is the difference in polarity which
determines, for the same sized species, whether
it is soluble in water and whether it is a gas at
room temperature or a volatile liquid which
evaporates quickly or a high boiling liquid which
does not evaporate at all.
The attractions between molecules which causes
these differences all arise from increasing polarity or
its complete lack...
End
Topic 12