Chapter 10
Molecular Structure
Chapter 10 Molecular Structure
10.1 Covalent Bonds
10.2 Hybridization of Atomic Orbitals
10.3 Valence Shell Electron-Pair
Repulsion Theory (VSEPR)
10.5 Intermolecular Forces
Review of Chemical Bonds
 What is a chemical bond?
the attractive force that holds atoms together in
multi-atom elements or in compounds
 Why does it come about?
Octet Rule: atoms combine to form bonds either by
transferring electrons to form ionic bonds, or by
sharing electrons in covalent bonds, until each
atom is surrounded by eight valence electrons
Problem
 How can an atom alter its electron configuration
to obtain an octet (or duet) of electrons of a
noble gas?
Solution
 A metal  e
a cation


e
 A nonmetal 

an anion
ionic compounds
Na + Cl
+
Na
 Sharing of electrons between nonmetals
molecular compounds
Cl
-
Types of Chemical Bonds:
 Ionic Bonds: transferring electrons
 Covalent Bonds: sharing electrons
 Metallic Bonds: Valence electrons are detached
from atoms, and spread in an 'electron sea'
that "glues" the ions together
Ionic Bond
 is a strong electrostatic attraction between a
positive ion and a negative ion
 electron is fully transferred from metal to
nonmetal
 non-directional, magnitude of bond equal is all
directions
 typically occur between a metal and a reactive
non-metal
n Na (3s1) -n e n Na+ (2s22p6)
n Cl (3s23p5) +n e n Cl- (3s23p6)
n NaCl(s)
Covalent Bond
 Cooperative sharing
of valence electrons
 Covalent bonds are
HIGHLY directional
 typically occur
between non-metal
 have a relatively low
melting and boiling
point
10.1 Covalent Bonds
Key points:
 Why and how do two atoms bond together ?
 What’s the difference between them?
 σbond andπbond
 single bond, double bond and triple bond
 polar and nonpolar covalent bond
1. Describing Covalent Bonds
 Definition: A shared pair of electrons between two
potential energy →
atoms is known as a covalent bond.
 Example: H2 H + H
HH
repulsion Maximum
attraction
no overlap;
no attraction
0
74pm Bond dissociation energy
Bond length
Distance between nuclei →
 Covalent Bond Theory :
 Bonds are formed by atom sharing two electrons
in overlapping atomic orbitals.
 the number of bonds formed = the number of
unpaired electrons usually
NH3
 Orbitals bond in the same axis to obtain maximum
overlap. e.g.:HCl
Cl 1s22s22p6 3s23p5
H 1s1
x
z
x
2. Bond Properties
 Bond Energy: the energy required to break one
mole of bonds in a gaseous species.
NH3(g)=NH2(g)+H(g)
D1=435.3kJ·mol-1
NH2(g)=NH(g)+H(g)
D2=397.5kJ·mol-1
NH(g)=N(g)+H(g)
D3=338.9kJ·mol-1
NH3(g)=N (g)+3H(g)
D=1171.5kJ·mol-1
EN-H=(D1+D2+D3)/3=390.5 kJ·mol-1→
 The larger the bond energy, the stronger the
chemical bond.
 Bond Length: distance between nuclears of
two bonded atoms
 Bond lengths from x-ray analysis:
143
122
 Bond Angle
104045‘
113
3. Types of Bonds
Ionic Bond: electrostatic attraction
Chemical Bond
Covalent Bond: sharing of electrons
Coordinate Covalent Bond:
both electrons from one atom (CO)
C 1s22s22p2
Bond order
one electron from each atom
orbital overlap
Coordinate
Bond Polarity
Covalent Bond
σbond
πbond
Polar Covalent Bond
Nonpolar Covalent Bond
O 1s22s22p4
Single Bond
Double Bond
Triple Bond
Covalent Bond:
 sigma bond and Pi bond
Sigma (σ) bond = end-to-end overlap
☆ head-on overlap = end-to-end overlap
☆ s-s , s-px , px-px *
 s + s sigma overlap
Sigma () bonds
 s + p sigma overlap
 p + p sigma overlap
Pi (π) bond = side-by-side overlap
☆ sideways overlap
Pi () bond
☆ pz-pz , py-py *
+
p orbital p orbital
 bond
σ键
π键




Note here, that the extent of overlap between the
orbitals is less in a pie bond as compared to a
sigma bond. Thus, the force of attraction
between the shared electron pair and the nuclei
is comparatively less than that in a sigma bond.
Therefore, a pie bond is always weaker than a
sigma bond,and it can only coexist with sigma
bond.
A single covalent bond is a sigma bond.
A double covalent bond is made up of one sigma
and one pie bond.
A triple covalent bond is made up of one sigma
and two pie bonds.
Bond Order and Multiple Bonds
☆ single bond => 1 pairs shared
☆ double bond => 2 pairs shared
☆ triple bond => 3 pairs shared N2
Relationship Between Order, Length, and Energy
C—C
C
C
C
C
Bond order =
1
2
3
length (pm) =
154
134
121
energy (kJ/mol) = 347
611
837
Conclusion: Bond lengths SHORTEN as bond order
INCREASES. A double bond is STRONGER than a
single bond.
10.1 Covalent Bonds
• Polar Covalent Bonds and Nonpolar Covalent
Bonds
Problem
• The electrons in the covalent bond are shared
equally?
Solution
• The electronegativity (EN) differences of the
elements are used to determine the extent of this
unsharing.
Nonpolar Bond
diff. EN = 0
Electrons shared
equally.(atoms
with same
electronegativity)
H2: H H
Polar Bond
diff. EN > 0
Electrons shared
unequally.(atoms
with different
electronegativity)
HCl:
+
Ionic Bond
diff. EN > 1.7
Electron
transferred.
-
+
¦Ä Na¦Ä + ClNaCl: Na
H Cl
Notice: bond polarity≠molecular polarity
Cl
-
10.2 Hybridization of Atomic Orbitals




Questions:
Why can they hybrid?
What kinds of orbitals are used for
hybridization?
Why do they need to hybrid?
How to predict the shapes of molecules?
Problem
 Experiment shows that the
methane molecule (CH4) is
tetrahedral, with four equivalent
bonds.
Solution
2 2
2
1s22s12p3
C 1s 2s 2p
2p
2p
2s
2s
1s
ground state
1s
excited state
2p
2s
1s
excited state
mutually
perpendicular
spherically
symmetrical
“atomic orbitals” vs. “hybrid orbitals”
 atomic orbitals
☆ on the same atom (central atom only )
☆ in the same principal energy level or,
occasionally, in adjacent energy level
 hybrid orbitals
☆ the shapes and directional properties of
new hybrid orbitals ≠the orbitals used in
constructing them
☆ the number of hybrid orbitals formed =
number of atomic orbitals used
10.2 Hybridization of Atomic
Orbitals
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4]
solution
2+
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
[Cu(NH3) 4] 2+
solution
After the paint is mixed, there are still the same
numbers of cups of paint, but the color has changed.
The mixing of orbitals is analogous. sp3
10.2 Hybridization of Atomic
Orbitals
The directional characteristics of the new hybrid
orbitals will be decided by the type and number
of atomic orbitals using in the mixing.
Table 3 Three Types of s-p Hybrid
Atomic
orbitals
one s, one p
one s, two p
one s, three
p
Hybrid
Orbitals
two sp
three sp2
four sp3
Geometric
Arrangement
Linear*
Trigonal planar*
Tetrahedral
10.2 Hybridization of Atomic
Orbitals
sp
10.2 Hybridization of Atomic
Orbitals
2
sp
10.2 Hybridization of Atomic
Orbitals
3
sp
10.2 Hybridization of Atomic
Orbitals
Needed to form 4
sigma bonds
The mutual repulsion
among the electron
pairs will orient them
toward the apices of a
tetrahedron.
sp3
four sp3 hybrid orbitals
1s
hybrided
sp3
C-H sp3-s sigma bonds
1s
in CH4
2 2 2
C 1s 2s 2p
1s22s12p3
2p
2p
2s
2s
1s
ground state
1s
excited state
sp3
four sp3 hybrid orbitals
1s
hybrided
sp3
C-H sp3-s sigma bonds
1s
in CH4
1. An Outline of Hybrid Orbital Theory
I.
Why can they hybrid?
We can mathematically mix two or more of
these wave functions that describe the
electron, and produce an equal number
of wave functions that have different
shapes and orientation. However, the
number of new hybrid orbitals will be the
same as the number of orbitals used to
form them.
What kind of orbitals are used for
hybridization?
Hybrids are formed among orbitals lying in
the same principal energy level or,
occasionally, in adjacent energy levels.
Example ns np
( n-1) d ns np or ns np nd
II.
10.2 Hybridization of Atomic
Orbitals
III.
Why do they need to hybrid?
Each sp3 hybrid orbital has a large lobe pointing
in one direction and a small lobe pointing in the
opposite direction. The enlarged lobe of the
resulting orbital can give more favorable
overlap with the orbital of another atom and
thus form a stronger bond than can either the p
or s orbitals alone.
＋ ＋＋
＋ ＋
_
_
＋
_
＋＋
Mix
Mix
_
＋＋
10.2 Hybridization of Atomic
Orbitals
IV. Why do they form a geometry ?
The mutual repulsion among the electron pairs
will orient hybrid orbitals toward the apices
of a geometry.
Hybrid Number of Geometric
Orbitals Orbitals
Arrangement
sp
sp2
sp3
sp3d
2
3
4
5
Linear
Trigonal
Tetrahedral
Trigonal bipyramidal
sp3d2
6
Octahedral
10.2 Hybridization of Atomic
2. Application- Beryllium Chloride, BeCl2 : sp
Orbitals
2 2
Be 1s 2s 2p
2p
2s
2s
1s
ground state
1s
excited state
Use sp
hybrid
orbitals
two unchanged and
vacant p orbitals
2p
2p
sp
sp
1s
hybridized
Be-Cl bonds:
sp-p sigma bond
1s
in BeCl2
Two electron pairs produce a
linear arrangement. Cl—Be—Cl
10.2 Hybridization of Atomic
Orbitals
2 2 1
B 1s 2s 2p 2p
2p
2s
2s
1s
ground state
1s
excited state
2p
BF3 : sp2
2p
sp2
Use
sp2
hybrid
orbitals
1s
hybridized
sp2
B-F bonds F
1s
B
F
F
in BF3
Three electron pairs produce a
triangular planar arrangement .
10.2 Hybridization of Atomic
Orbitals
3
sp :CH4 NH3 H2O
2 2
2s
2p 2p
C
CH4
2p
2s
2s
sp3
hybrid orbitals
equivalent hybridization
NH3
2 3
N 2s 2p 2p
2s
H2O
2 4
O 2s 2p 2p
2s
nonequivalent hybridization
104045‘
C2H6
C2H2
C2H4
What would be the name of
the hybrid orbitals created
by joining 1 s type, 3 p
type and 2 d type orbitals?
a. spd b. sp3d2 c. s3p2d
10.3 Valence Shell Electron-Pair
(VSEPR)
Repulsion Theory
Key Points:
 What’s the difference between the
arrangement of electron pairs and the
molecular geometry ?
 How to predict the shapes of molecules?
PbCl2 Pb: 5d106s26p2 Pb2+→ sp？linear？
HgCl2 Hg: 5d106s2 Hg2+→ sp？linear？
10.3 Valence Shell Electron-Pair
Repulsion Theory
1. VSEPR model
 The arrangement of electron pairs describes the
arrangement of all electron pairs, shared or
unshared, around a central atom.
The molecular geometry of a molecule is the
geometry described by the bonded atoms and does
not include the unshared pairs of electrons.
 The minimum repulsion occurs when the electron
pairs are as far apart as possible.
10.3 Valence Shell Electron-Pair
Repulsion Theory
Table 1 Arrangement of electron pairs about an atom
Number of electron Pairs
Arrangement of Electron Pairs
2
Linear
3
Trigonal planar
4
Tetrahedral
5
Trigonal
bipyramidal
6
Octahedral
10.3 Valence Shell Electron-Pair
Repulsion Theory
 If more than one structure is possible, the most
stable structure can often be determined using the
following rules:
 Bond angles: 90º>120º>180º
 Lone pairs (LP): LP-LP>LP-BP>BP-BP
 Bond order：triple >double >single bond
 Electronagetivity of the donor atom：Xsmall > Xlarge
 Select the stable structure which has the smallest
number of lone pair-lone pair repulsions (usually
90o).
10.3 Valence Shell Electron-Pair
Repulsion Theory
2. How to predict molecular structures?
 Determine the number of electron pairs (both
bonded pairs and lone pairs) around the central
atom.
the number of electron-pairs=1/2 (the number of
valence electrons on the central atom + the
number of electrons furnished by donor atom)
 the number of valence electrons on the central
atom ＝group number
 hydrogen and the halogens (Group VIIA) donate
one electron each for sharing，
10.3 Valence Shell Electron-Pair
Repulsion Theory
 Elements in the oxygen group (Group VIA) are
considered to donate no electrons.
 Positive ion: subtract the positive charge on the
ion.
Negative ion: add the negative charge on the ion.
 In the case of an odd number of electrons, treat
the extra electron (one-half of an electron-pair) as
if it were an electron-pair (lone pair).
 Count a multiple bond as one pair.
10.3 Valence Shell Electron-Pair
Repulsion Theory
 Identify the molecular geometry
No. of Arrangement of
e- pairs pairs
2（sp）
Molecular geometry of ABn
Linear
3（sp2） Trigonal planar
AB2
AB3
4（sp3） Tetrahedral
AB4
5
Trigonal
3
（sp d ） bipyramidal
AB 5
AB4
AB6
AB5
6
(sp3d2)
Octahedral
AB2
AB3
AB2
AB3
AB4
AB2
10.3 Valence Shell Electron-Pair
Repulsion Theory
3. Applying VSEPR Theory
 Using the VSEPR theory we could predict the
geometry of an AXn molecule or ion .
 Using the VSEPR theory we could predict bond
angles and molecular geometry.
 Using the VSEPR theory we could predict whether
the molecule is polar or nonpolar.
 Using the VSEPR theory we could describe the
hybrid orbitals used by the central atom.
10.3 Valence Shell Electron-Pair
Repulsion Theory
 Using the VSEPR theory Central atom: Cl
3s23p5
we could predict the
geometry of an AXn
F
F
molecule or ion .
Cl F
F
F
Predict the geometry of the
LP–LP 0
following molecules or
LP–BP 4
ions.
BP–BP 2
+
BF3 PCl5 H2O H3O
NH3 NH4+ NO2 NO2SF4 ClF3 XeF2 IF5
XeF4 SO42- O3
F
F
Cl
Cl
F
1
3
2
0
6
0
F
10.3 Valence Shell Electron-Pair
Repulsion Theory
Procedure
☆Determine the number of electron pairs.
※a multiple ≌ one pair ≌ an unpaired
electron
☆ Arrange the electron pairs.
☆ Identify the molecular geometry.
with lone pairs: the arrangement of electron
pairs ≠ the molecular geometry
10.3 Valence Shell Electron-Pair
Repulsion Theory
 Using the VSEPR theory we could predict bond
angles and molecular geometry.
H
H N H
O
CH4 : H C H NH3 :
H2O:
H
H
H
H
No lone pair
One lone pair
Two lone pairs
Tetrahedral
Trigonal pyramidal Bent
Notice: When lone pairs are present, the
bond angles are smaller than the perfect angles.
10.3 Valence Shell Electron-Pair
Repulsion Theory
 Using the VSEPR theory we could predict whether
the molecule is polar or nonpolar.
Problem
How dose bond polarity relate to molecular polarity?
Solution
☆ For diatomic molecules, if the bond is polar then
the molecule is polar.
☆ For all other molecules, we need to know the
molecular shape to be able to predict the polarity.
10.3 Valence Shell Electron-Pair
Repulsion Theory
If two people of exactly equal strength pull on
the box in exactly opposite directions, their
efforts cancel and there is no movement.
Direction of movement
10.3 Valence Shell Electron-Pair
Repulsion Theory
nonpolar molecule
1. Polar or nonpolar
covalent bond
2. has a perfect geometry
(no lone pairs )
3. dipole moments (μ) = 0
4. Example:
O C O
CO2 :
Linear
polar bonds
nonpolar molecule
polar molecule
Polar covalent bond
does not have a perfect
geometry (with lone pairs )
μ≠0
SO2 :
O
S
Bent
polar bonds
polar molecule
O
10.3 Valence Shell Electron-Pair
Repulsion Theory
Conclusion
☆ For diatomic molecules, if the bond is polar
then the molecule is polar.
☆ For all other molecules, we need to know the
molecular shape to be able to predict the
polarity.
※ Molecules with lone pairs of electrons on a
central atom are generally polar molecules.
※ Nonpolar molecules that contain polar bonds
are observed when there are no lone pairs on the
central atom, and all of the atoms bonded to the
central atom are identical.
10.3 Valence Shell Electron-Pair
Repulsion Theory
 Use the VSEPR model to obtain the
arrangement of electron pairs.
 From the arrangement of the electron pairs,
deduce the type of hybrid orbitals.
Arrangement of electron pairs Hybrid Orbitals
Linear
Sp(2)
Trigonal planar
Tetrahedral
Trigonal bipyramidal
Sp2(3)
Sp3(4)
Sp3d(5)
Octahedral
Sp3d2(6)
10.5 Types of Intermolecular Forces
Key Points:
 What kinds of attractions exist between
molecules?
 What’s the difference between the
intermolecular forces and the
intramolecular forces ?
10.5 Types of Intermolecular Forces
 Intramolecular forces hold atoms together in
molecules.
Ionic
Covalent
Metallic
Cation-anion Nuclei-shared e- pair Cations-delocalized e-
 Intermolecular forces are those
between
Polar bond to H-dipole
molecules.
H bond
Dipole-dipole
Dipole-induced dipole
Dispersion (London)
charge (high EN of N, O, F)
Dipole charges
Dipole chargepolarizable e- cloud
Polarizable e- clouds
10.5 Types of Intermolecular Forces
2. Explaining Liquid Properties

Many of the physical properties of liquids (and
certain solids) can be explained in terms of
intermolecular forces.
 play a key role in determining the conditions
under which a substance is a solid, a liquid, or a
gas
 affect physical properties, such as the melting
points of solids and the boiling points of liquids.
10.5 Types of Intermolecular Forces
3. Dipole Moment
I.
A polar molecule and dipole moment

A polar molecule
has permanent dipole

Polar molecules are aligned in an electric
field
10.5 Types of Intermolecular Forces
I.


Dipole moment 
= qx
, dipole moment
q, charge (or fractional charges)
x, distance between charges
a quantitative measure of the degree of charge
separation in a molecule.
expresses a molecule’s polarity. Molecules with
large dipole moments are highly polar.
Nonpolar molecules have a zero dipole moment.
10.5 Types of Intermolecular Forces
II.



Instantaneous And Induced Dipoles
In describing electronic structures we
speak of electron charge density or the
probability that an electron is in a
certain region at a given time.
One probability is that at some
particular instant―purely by
chance―electrons are concentrated in
one region of an atom or molecule. This
displacement of electrons causes a
normally nonpolar species to become
momentarily polar.
After this, electrons in a neighboring
atom or molecule may be displaced to
also produce a dipole, called an
induced dipole.
10.5 Types of Intermolecular Forces
4. Types of Intermolecular Forces
Dipole-Dipole Forces
Dipole- Induced Dipole Forces
London Forces
Hydrogen Bonding
Van Der
Waals Forces
10.5 Types of Intermolecular Forces
I.


Dipole-Dipole Forces
A polar molecule has two
"poles". (Permanent Dipole )
Molecules with a permanent
dipole can align themselves
so that the negative end of
one molecule is attracted to
the positive end of another.
These attractions are known
as dipole-dipole attractions.
10.5 Types of Intermolecular Forces
I.




Dipole-Dipole Forces
Notice
only in permanence
somewhat stronger than
London forces because
permanent
occurs in all polar
molecules
10.5 Types of Intermolecular Forces
II.




Dipole-Induced Dipole
Interactions
Dipole H2O and nonpolar O2:
The dipole of H2O induces a
dipole in O2 by be distorting
the O2 electron cloud –
Polarizability
Interactions of polarizable
molecules are called induced
dipole interactions
Notice occurs in polar and
nonpolar molecules
10.5 Types of Intermolecular Forces
III.



Induced dipole - Induced dipole
Interactions (London forces )
Two nonpolar atoms or molecules
has no dipole moment .
Momentary attractions and
repulsions between nuclei and
electrons in neighboring
molecules lead to induced dipoles.
The attraction between these
induced dipoles then spreads
from molecule to molecule, thus
providing an attractive force
between molecules.
±
±
10.5 Types of Intermolecular Forces






Even nonpolar molecules and
uncombined atoms have attractive
forces between them.
Notice
all particles have London forces
London forces increase with
molecular weight.
only force between noble gases and
nonpolar compounds.
about 1/1000 as strong as a covalent
bond



Notice
All particles have London forces. London
forces increase with molecular weight.
about 1/1000 as strong as a covalent bond
polar-polar
polarnonpolar
nonpolarnonpolar
Dipole-Dipole Dipole- Induced
Forces
Dipole Forces
√
√
√
London
Forces
√
√
√
10.5 Types of Intermolecular Forces
V.

Hydrogen Bonding
Problem
Examination of the
boiling points of the
groups VA, VIA,
and VIIA reveals
that the first
compound in each
of these series has
an unexpectedly
high boiling point.

A hydrogen atom bonded to an electronegative
atom appears to be special. The electrons in the
O-H bond are drawn to the O atom, leaving the
dense positive charge of the hydrogen nucleus
exposed. It’s the strong attraction of this
exposed nucleus for the lone pair on an
adjacent molecule that accounts for the strong
attraction. A similar mechanism explains the
attractions in HF and NH3.
 Condition
X—H…Y(X,Y=F、O、N)
 X→ Xlarge、rsmall
 Y→ Xlarge、rsmall ，lone pair electron
10.5 Types of Intermolecular Forces

Notice

Rare, strong form of dipole-dipole interaction.

Require H bonded to highly electronegative atom
(N, O or Halogen): one of the following three
structures must be present. H-N O-H F-H
One molecule has a hydrogen atom attached by a
covalent bond to an atom of nitrogen, oxygen, or
fluorine. The other molecule has a nitrogen, oxygen,
or fluorine atom present that possesses one or more
lone pairs.

strongest of four intermolecular forces

are HIGHLY directional
10.5 Types of Intermolecular Forces


Types of Hydrogen Bonds
intermolecular hydrogen bonds
Increase boiling points

intramolecular hydrogen bonds
Decrease boiling points
10.5 Types of Intermolecular Forces
5. The Double Helix of DNA is held together by
hydrogen bonding

In the three-dimensional Watson-Crick
structure, two polynucleotide DNA strands
wind around each other to form a double helix.
Hydrogen bonds form between specific base
pairs. Adenine is hydrogen bonded to thymine,
and guanine is hydrogen bonded to cytosine to
form complementary base pairs (A on one
strand with T on the other strand, or C with G).

adenine腺嘌呤 thymine胸腺嘧啶

guanine鸟嘌呤cytosine胞嘧啶
10.5 Types of Intermolecular Forces
white = hydrogen
blue = nitrogen
black = carbon
red = oxygen
10.5 Types of Intermolecular Forces
 Two hydrogen bonds
occur between every
adenine and thymine
pair; three between
each guanine and
cytosine.