Molecular Structures
John W. Moore
Conrad L. Stanitski
Peter C. Jurs
C2H6O structural isomers:
H H
|
| ..
H–C–C–O
.. – H
|
|
H H
http://academic.cengage.com/chemistry/moore
Chapter 9
Molecular Structures
Stephen C. Foster • Mississippi State University
m.p./ °C
b.p./ °C
H
H
| .. |
H–C–O
–C–H
..
|
|
H
H
ethanol
dimethyl ether
-114
+78
-142
-25
Molecular shape is important!
Small structural changes cause large property
changes.
Using Molecular Models
Using Molecular Models
Physical models of 3D-structures:
Hand-drawn molecules:
Going back into
the screen
ball and stick
space filling
In the plane of
the screen
H
C
H
Computer versions:
H
H
Predicting Molecular Shapes: VSEPR
Coming out of
the screen
Predicting Molecular Shapes: VSEPR
The Valence Shell Electron Pair Repulsion model is
used to predict shapes. Key ideas:
1. e- pairs stay as far apart as possible.
• Repulsions are minimized.
2. Molecule shape is governed by the number of
Linear
Triangular planar
Tetrahedral
bond pairs and lone pairs present.
3. Treat multiple bonds like single bonds.
• Each is a single e- group.
4. Lone pairs occupy more volume than bonds.
Triangular bipyramidal
Octahedral
1
Predicting Molecular Shapes: VSEPR
Shapes that minimize repulsions:
Predicting Molecular Shapes: VSEPR
If a molecule contains:
• bonding pairs only – these angles are correct:
triangular
planar
linear
tetrahedral
triangular
bipyramidal
octahedral
Bonds and lone pairs determine shape.
Use the notation AXnEm
m lone pairs on
central atom A
n atoms bonded to
central atom A
Predicting Molecular Shapes: VSEPR
Molecules may be described by their:
• electron-pair (e- pair) geometry
• molecular geometry (molecular shape)
• These angle change (a little) if any “X” is replaced by a
lone pair:
• lone pair/lone pair repulsions are largest.
• lone pair/bond pair are intermediate in strength.
• bond/bond interactions are the smallest.
Predicting Molecular Shapes: VSEPR
AXnEm: 2 e- group central atoms (m + n = 2)
Linear e- pair geometry
2 e- groups
bond
pairs
These geometries may be different.
lone
pairs
2
0
AX2E0
1
AX1E1
linear
1
..
• Atoms can be “seen”, lone pairs are invisible.
linear
molecular geometry
Predicting Molecular Shapes: VSEPR
AX2E0 examples:
Predicting Molecular Shapes: VSEPR
AXnEm: 3 e- group central atoms (m + n = 3)
180.0°
Cl Be Cl
Triangular planar e- pair geometry
Linear.
3 e- groups
bond
pairs
“2” bonds, 0 lone pairs on C.
(treat double bonds as 1 bond)
Linear.
O
3
0
AX3E0
triangular planar
2
1
AX2E1
..
180.0°
O C
lone
pairs
angular (bent)
H C C
H
Each H-C-C unit is linear.
1
2
AX1E2
..
180.0°
linear
180.0°
molecular geometry
2
Predicting Molecular Shapes: VSEPR
AX3E0 examples:
Predicting Molecular Shapes: VSEPR
AXnEm: 4 e- group central atoms (m + n = 4)
4 e- groups
bond
pairs
120°
Cl B
Cl
lone
pairs
4
Triangular planar.
Tetrahedral e- pair
geometry
0
AX4E0
tetrahedral
Cl
C
C H
H
H
Each C is AX3E0 = triangular planar.
Predicting Molecular Shapes: VSEPR
H
AX4E0
All angles = tetrahedral angle.
H C H
AX3E1
triangular pyramidal
..
H
1
..
3
molecular
geometry
AX1E3?
All molecules with only
1 bond are linear!
2
2
AX2E2
angular (bent)
Predicting Molecular Shapes: VSEPR
VSEPR applies to each atom in a molecule.
• Alkanes: each C is tetrahedral.
H
AX3E1
Lone-pair/bond > bond/bond
repulsion: H-N-H angle is
reduced.
H N H
H
AX2E2
Two lone pairs: H-O-H angle is
even smaller.
H O
H
Predicting Molecular Shapes: VSEPR
Lactic acid:
H
Tetrahedral O
Triangular planar C
C
C
H
H
C
..
O
..
Central atoms with five or six e- pairs:
Bond pairs Lone pairs
Shape
5
0
Triangular bipyramidal
4
1
Seesaw
3
2
T-shaped
2
3
Linear
..
..
..
..
H
O
H
Expanded Octets
Tetrahedral C
Tetrahedral C
Tetrahedral O
O
H
6
5
4
3
0
1
2
3
Octahedral
Square pyramidal
Square planar
T-shaped
• lone pairs repel the most.
• they get as far apart as possible.
3
Expanded Octets
Expanded Octets
AXnEm: m + n = 5
Triangular bipyramidal e- pair geometry.
Triangular
bipyramidal
Seesaw
T-shaped
F
F
F
F P F
F S F
F Cl F
F
F
F
F Xe F
Linear
The atoms are non-equivalent.
Green atoms are axial
axial; blue atoms are equatorial
equatorial.
Expanded Octets
Expanded Octets
AXnEm: m + n = 6 Octahedral e- pair geometry:
F
All atoms are
equivalent in
AX6E0
F
F S F
F
Octahedral
Square pyramid
Lewis dot + VSEPR predict molecular shapes, but…
but
How do atomic orbitals (s, p…) lead to these shapes?
Valence bond theory:
theory: bonds occur when partiallyoccupied atomic orbitals overlap.
F Br F
F
Cl I
Cl
Cl
HF – H(1s) overlaps F(2p)
Valence Bond Theory
This works for H2 and HF, but why does…
• Be form compounds?
• Be (1s2 2s2).
• No unpaired e- to share.
• Experiments show: linear BeH2, BeCl2, …
• C form 4 bonds at tetrahedral angles?
•
•
•
•
74 pm
Cl
F
Square planar
Orbitals Consistent with Molecular Shapes
H2 – H(1s) overlaps H(1s)
F
F
C (1s2 2s2 2p2).
2px1 2py1 Two bonds?
p orbitals are at 90° to each other
Experiments show: tetrahedral CH4, CCl4, …
109 pm
4
Orbitals Consistent with Molecular Shapes
Atomic orbitals (AOs) can be hybridized (mixed).
sp Hybrid Orbitals
Be compounds (BeH2, BeF2 …):
• Sets of identical hybrid orbitals form identical bonds.
• Number of hybrids formed = number of AOs mixed.
One s orbital + one p orbital → two sp hybrids.
Each sp hybrid (180° apart) holds one e-.
Two equivalent covalent bonds form.
sp2 Hybrid Orbitals
B forms three sp2 hybrid orbitals:
• One s orbital mixes with two p orbitals.
• One p orbital remains unmixed.
sp2 Hybrid Orbitals
B compounds (BH3, BF3 …):
Each sp2 hybrid (120° apart) holds one e-.
Three equivalent covalent bonds form.
sp3 Hybrid Orbitals
C forms four sp3 hybrid orbitals:
• One s orbital mixes with three p orbitals.
• All p orbitals are mixed.
sp3 Hybrid Orbitals
N and O compounds (NH3, H2O…) have more e-:
In C, each sp3 hybrid (109.5° apart) holds one e-.
Four equivalent covalent bonds form.
5
sp3 Hybrid Orbitals
Hybridization
e- pair
“Octet rule” molecules have tetrahedral
shape.
• sp3 hybridized (CH4, NH3, H2O, H2S, PH3, …)
Head-to-head bond = a sigma bond (σ
σ bond).
bond
There are:
H
• 4 σ bonds in CH4
σ bond
• 3 σ bonds in NH3
• 2 σ bonds in H2O
C
H
H
Summary:
Mixed
s+p
s+p+p
s+p+p+p
Hybrids (#)
sp (2)
sp2 (3)
sp3 (4)
Remaining Geometry
p+p
Linear
p
Triangular planar
Tetrahedral
d orbitals can also form hybrids:
Mixed
Hybrids (#)
Remaining Geometry
s+p+p+p+d
sp3d (5)
d+d+d+d
Triangular bipyramid
s+p+p+p+d+d
sp3d2 (6)
d+d+d
Octahedral
H
Hybridization in Molecules with Multiple Bonds
Carbon atoms form:
• tetrahedral centers (CH4, CHF3 , C2H6…) = sp3
• triangular-planar centers (H2CO, C2H4 …) = sp2
H
C
C H
H
H
Hybridization in Molecules with Multiple Bonds
• a σ bond – head-to-head overlap of sp2 on each C atom.
• a π bond – sideways overlap of p AOs on the C atoms.
H
C
C
H
H
Unhybridized C p orbitals each contain one
e-.
H
C σ bond C
The double bond in ethene is composed of:
H
C (sp2) + C (sp2) overlap (σ bond):
H
H
C
overlap
H
C
H
H
Sideways overlap forms one π bond
• the lobes above and below the plane together equal 1 bond
Hybridization in Molecules with Multiple Bonds
Hybridization in Molecules with Multiple Bonds
Formaldehyde is similar:
C also forms linear centers:
• C2H2 (acetylene) = sp hybridized
H
C
C H
The triple bond is:
• one σ bond
• two π bonds
• sp hybridization leaves two unmixed p orbitals
on each C.
6
Hybridization in Molecules with Multiple Bonds
Hybridization in Molecules with Multiple Bonds
π bonds prevent bond rotation:
σ bond: C (sp) + C (sp) overlap:
H
C
C
H
Two p orbitals on each C contain a single e-.
H
C
H
C
overlap
H
C
C
H
Molecule
ethane (CH3–CH3)
ethene (CH2=CH2)
ethyne (HC≡CH)
C-C bonding
σ
σ+π
σ+π+π
C-C rotation
yes
no
no
Non-rotating double bonds allow cis-trans isomerism
to occur.
Two π bonds
• above and below overlaps are 1 bond.
• front and back overlaps are a second bond.
Molecular Polarity
Molecular Polarity
• Most bonds are polar (e.g. C-O)
• Water is polar (bond dipoles do not cancel)
• O is δ-, C is δ+ (ENO = 3.5, ENC = 2.5)
• But many molecules are nonpolar (e.g. CO2).
O=C=O
δ-
2δ+
O
H
arrow points to δ-,
the + shows δ+
H
+
Net
dipole
Dipole, μ = 1.85 D
δ-
• The dipoles cancel because of CO2’s shape.
• have equal size but point in opposite directions.
Molecular Polarity
Dipole moment (μ) is a measure of
molecule polarity:
Units: coulomb meter (Cm)
Debye (D)
weakly polar
highly polar
nonpolar (μ=0)
Molecular Polarity
Molecule
H2
HF
HCl
HBr
HI
CH4
CH3Cl
CH2Cl2
CHCl3
CCl4
μ (D)
0
1.78
1.07
0.79
0.38
0
1.92
1.60
1.04
0
A molecule is nonpolar if it is:
• AXnE0 and all X are identical.
CO2
AX2E0
linear
CH4
AX4E0
tetrahedral
CCl4
AX4E0
tetrahedral
PF5
AX5E0
triangular bipyramidal
• “divisible” into nonpolar AXnE0 shapes
PCl3F2 triangular planar (PCl3) + linear (PF2 )
XeF4
linear (XeF2) + linear (XeF2 )
7
Molecular Polarity
AXnEm molecules are polar if they don’t divide into
nonpolar shapes, and::
• m ≠ 0:
H2O
AX2E2
bent
polar
NH3
AX3E1
pyramidal
polar
H
F
+
Molecular Polarity
No net dipole
C
F
C
F
F
Net
dipole
F
F
F
CHF3 is polar
CF4 is non polar
• The X in AXnE0 differ:
CH2Cl2
PF4Cl
AX4E0
AX5E0
tetrahedral
polar
triangular bipyramidal polar
How polar? It depends on the number, type, and
geometry of the polar bonds.
Molecular Polarity
Noncovalent Interactions
Non polar
AX5E0 and “X” differ.
BUT divisible into
nonpolar shapes:
linear + triangular
planar
+
PCl3F2
PCl5
Non polar
AX5E0; identical X
PCl4F
Polar
AX5E0
“X” differ
Polar
PF3Cl2
AX5E0 and “X” differ.
Doesn’t divide into
nonpolar shapes
London Forces
Also called dispersion forces.
• Random e- motion produces a temporary dipole in
one molecule which induces a dipole in another.
↔ 40 kJ/mol):
Small molecule = few e- = weak attraction.
Large molecule = many e- = stronger attraction.
• Occur between all atoms and molecules.
The only force between nonpolar molecules.
Molecules attract each other.
Intermolecular forces:
• also called noncovalent interactions.
• are small (compared to bonding forces).
• do not include ionic or metallic-bonding forces.
Three types:
• London forces.
• dipole-dipole attraction.
• hydrogen bonding.
London Forces
Noble Gas
# of ebp (°C)
He
2
−269
F2
Halogen
Hydrocarbon
# of e- bp (°C)
# of e- bp (°C)
18
−188
CH4
10
−161
Ne
10
−246
Cl2
34
−34
C2H6
18
−88
Ar
18
−186
Br2
70
+59
C3H8
26
−42
Kr
36
−152
I2
106
+184
C4H10
34
0
• Strength (0.05
More e- = larger attraction = higher b.p.
8
Dipole-Dipole Attractions
Polar molecules attract each other.
Strength: 5 ↔ 25 kJ/mol.
Dipole-Dipole Attractions
Nonpolar Molecules
# of e- bp (°C)
SiH4
18
−112
GeH4
36
−90
Br2
70
+59
Relative importance of dipole/dipole and London is
hard to predict:
HI
HCl
Hydrogen Bonding
An especially large dipole-dipole attraction.
• 10 ↔ 40 kJ/mol.
• Occurs when H bonds directly to F, O or N.
Polar Molecules
# of e- bp (°C)
PH3
18
−88
AsH3
36
−62
ICl
70
+97
Dipole
small (0.38 D)
large (1.07 D)
London
large (54 e-)
small (18 e-)
bp (°C)
−36
−85
stickier
Hydrogen Bonding
H on one molecule
interacts with O on
another molecule.
F, O & N are small with large electronegativities.
• results in large δ+ and δ- values.
H-bonds are usually drawn as dotted lines.
Hydrogen Bonding
Water is a liquid at
room T (not a gas).
9