Transcript Slide 1

Lecture 7
Molecular Bonding Theories
1) Valence Shell Electron Pair Repulsion (VSEPR) theory
Simple theory for qualitative prediction of geometry of polyatomic species
1)
Draw a reasonable Lewis structure. Count lone pairs and atoms directly attached to the central
atom (N).
2)
Point lone pairs and bonds at the central atom (A) to the vertices of a polygon or polyhedron.
Start from the most symmetrical shapes shown below for 2 < N < 8 (B = an atom or a lone pair):
Non-planar arrangements (N = 4, 5, 6, 7, 8):
Linear or planar arrangements (N =2 or 3):
B
B
A B
linear
A
B
B
trigonal planar
B
B
B
A B
B
tetrahedron
B
A B
B
B
trigonal bipyramid
B
B A B
B
B
B
octahedron
B
B
B
B
B A
A
B
B
B
B
B
B
B
B
pentagonal bipyramid tetragonal antiprysm
B
B
3)
Consider distortions. Distortions of an initially more symmetrical shape are due to the fact that:
•
•
•
Lone pairs need more space than bonds.
Multiple bonds need more space than single bonds.
Bonds formed with more electronegative elements occupy less space.
4)
Smaller “objects” will tend to occupy axial positions in trigonal bipyramids and equatorial
positions in pentagonal bipyramids.
5)
For 4th and higher row elements one lone pair tends to be stereochemically inactive.
2) Most common molecular shapes
Formula
Lone
pairs
Example
Geometry
Formula
Lone pairs
Example
Geometry
AB2
0
BeH2
Linear
AB5
0
PF5
Trigonal
bipyramid
1
CH2
Bent
1
2
H2O
Bent
IF5
Tetragonal
pyramid
3
XeF2
Linear
0
WF6
Octahedral
0
BF3
Trigonal planar
1
XeF6
Distorted
octahedral
1
NH3
Trigonal
pyramid
AB7
0
IF7
Pentagonal
bipyramid
AB8
0
ZrF84-
Tetragonal antiprism
1
Mo(CN)84-
Trigonal
dodecahedron
0
ReH92-
Tricapped
trigonal prism
AB3
2
AB4
AB6
T-shaped
IF3
-
Tetrahedral
0
BH4
1
SF4
Butterfly
2
XeF4
Square planar
H
C
H
O H
H
H
F
F
Xe
I
F
F
F
N H
H
F
S
F
F
F
F
F
AB9
Xe F
F
F
F
F
I
F
F
H
H
H
H
H
Re
H
H
H
H
trigonal dodecahedron
3) Concept of hybridization
•
Describes geometry of polyatomic species ABx, but predicts degeneracy
that does not exist.
•
Assumes that before an atom A forms x s-bonds, x non-equivalent atomic
orbitals yat combine to build a set of the same number x of equivalent
hybrid orbitals, yhyb. Similarly hybrid orbitals for p-bonding can be formed.
Orbitals suitable for the combination can be found by applying the group
theory.
•
Each hybrid orbital yhyb, j is a linear combination of atomic orbitals, yat, i.
y hyb, j   cijy at ,i
i
The probability to find an electron on an j-th hybrid orbital yhyb, j
•
y
•
2
hyb, j
dv    c y at ,i dv  c  1
2
ij
2
 y at , jy at ,k dv  0;  y at , j dv  1
2
2
ij
since
The probability to find an electron on an i-th atomic orbital yat, i
 cij2  1
i
i
j
4) sp-Hybrid orbitals
•
Linear molecules AB2 (BeH2). sp-hybridization.
yhybcs1ys + cp1ypx
ys
ypx
x
x
x
yhybcs2ys - cp2ypx
2 atomic orbitals
2 hybrid orbitals
•
Calculating the coefficients cij:
cs12 + cs22 = 1
cP12 + cP22 = 1
cS1 = cS2 = (1/2)1/2
cP1 = cP2 = (1/2)1/2
5) Concept of hybridization. sp2-Hybrid orbitals
•
Trigonal planar molecules AB3 (BF3). sp2- or d2s-hybridization.
d2s
sp2
yhyb
ypy
y
yhyb
ypx
y
y dx
xy
y
ys
y dx y
dx2-y2
yd
ys
y
xy
yhyb
yd
yhyb
x
x
x
3 atomic orbitals
•
3 atomic orbitals
3
1
y
yhyb
2
2
+ 0ypx +
2y
py
3
yhyb, 2 =
1
y
3 s
+
1
y
2 px
-
1
y
6 py
yhyb, 3 =
1
y
3 s
-
1
y
2 px
-
1
6
ypy
3
3 hybrid orbitals
Calculating the coefficients cij (sp2; for d2s use dxy instead of px and dx2-y2
instead of py):
1
y
3 s
yhyb
1
x
3 hybrid orbitals
yhyb, 1 =
2
2
y
6) dsp2-Hybrid orbitals
•
Square planar molecules AB4 ([PtCl4]2-). dx2-y2sp2- or dx2-y2dz2p2-hybridization.
yhyb
yhyb
ys
ypy
y
ypx
x
yhyb
yhyb
ydx2-y2
y
x
y
4 hybrid orbitals
4 atomic orbitals
•
Calculating the coefficients cij:
yhyb, 1 =
1
y
2 s
+ 0 ypx +
yhyb, 2 =
1
y
2 s
+
yhyb, 3 =
1
ys
2
+ 0 ypx -
yhyb, 4 =
1
y
2 s
-
1
2
1
2
1
2
ypy +
ypx + 0 ypy 1
2
ypy +
ypx + 0 ypy -
1
2
ydx2-y2
1
y
2 dx2-y2
1
ydx2-y2
2
1
y
2 dx2-y2