Chemistry 281(01) Winter 2016
CTH 277 10:00-11:15 am
Instructor: Dr. Upali Siriwardane
E-mail: upali@latech.edu
Office: 311 Carson Taylor Hall ; Phone: 318-2574941;
Office Hours: MTW 8:00 am - 10:00 am;
TR 8:30 - 9:30 am & 1:00-2:00 pm.
January 12, 2016 Test 1 (Chapters 1&,2),
February 2, 2016 Test 2 (Chapters 3 &4)
February 26, 2016, Test 3 (Chapters 5 & 6),
Comprehensive Final Make Up Exam: March 1, 2016
9:30-10:45 AM, CTH 311.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-1
Molecular structure and bonding
Lewis structures
2.1 The octet rule
2.2 Resonance
2.3 The VSEPR model
Valence-bond theory
2.4 The hydrogen molecule
2.5 Homonuclear diatomic molecules
2.6 Polyatomic molecules
Molecular orbital theory
2.7 An introduction to the theory
2.8 Homonuclear diatomic molecules
2.9 Heteronuclear diatomic
2.10 Bond properties
Chemistry 281, Winter 2016, LA Tech
Chapter-2-2
Lewis Theory of Bonding
Octet Rule
All elements except hydrogen ( hydrogen have a
duet of electrons) have octet of electrons once
they from ions and covalent compounds.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-3
Noble gas configuration
The noble gases are noted for
valence etheir chemical stability and
He
2
existence as monatomic
Ne
8
Ar
8
molecules.
Kr
8
Except for helium,
Xe
8
They share a common electron
Rn
8
configuration
that is very stable.
This configuration has 8 valence-shell electrons.
All other elements reacts to achieve Noble Gas Electron
Configurations.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-4
The octet rule
• Atoms are most stable if they have a filled or
empty outer layer of electrons.
• Except for H and He, a filled layer contains 8
electrons - an octet.
• Two atoms will
gain or lose
(ionic compounds)
share
(covalent compounds)
Many atoms with fewer electrons will
share
(metallic compounds)
Chemistry 281, Winter 2016, LA Tech
Chapter-2-5
What changes take place during this
process of achieving closed shells?
a) sharing leads to covalent bonds and
molecules
b) gain/loss of electrons lead to ionic bond
c) Sharing with many atoms lead to
metallic bonds
Chemistry 281, Winter 2016, LA Tech
Chapter-2-6
Lewis Electron Dot symbols
Basic rules
Draw the atomic symbol.
X
Treat each side as a box that can
hold up to two electrons.
Count the electrons in the valence
shell.
Start filling box - don’t make pairs
unless you need to.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-7
Lewis symbols
Lewis symbols of second period elements
Li
N
Chemistry 281, Winter 2016, LA Tech
Be
O
B
F
C
Ne
Chapter-2-8
What is a Lewis Structure (electron-dot
formula) of a Molecule?
• A molecular formulas with dots around atomic
symbols representing the valence electrons
• All atoms will have eight (octet) of electrons
(duet for H) if the molecule is to be stable.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-9
Single covalent bonds
H
H
H
H
C H
F
F
H
Do atoms (except H) have octets?
Chemistry 281, Winter 2016, LA Tech
Chapter-2-10
Lewis structures
• This is a simple system to help keep track of
electrons around atoms, ions and molecules invented by G.N. Lewis.
• If you know the number of electrons in the
valence-shell of an atom, writing Lewis
structures is easy.
• Lewis structures are used primarily for s- and
p-block elements.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-11
How do you get the Lewis Structure from
Molecular formula?
• Add all valence electrons and get valence
electron pairs
• Pick the central atom: Largest atom normally or
atom forming most bonds
• Connect central atom to terminal atoms
• Fill octet to all atoms (duet to hydrogen)
Chemistry 281, Winter 2016, LA Tech
Chapter-2-12
Lewis Structure of H2O
Chemistry 281, Winter 2016, LA Tech
Chapter-2-13
Types of electrons
Bonding pairs
Two electrons that are shared between two
atoms. A covalent bond.
Unshared (nonbonding ) pairs
A pair of electrons that are not shared
between two atoms. Lone pairs or
nonbonding electrons.
oo
H Cl
oo
oo
Unshared
pair
oo
Bonding pair
Chemistry 281, Winter 2016, LA Tech
Chapter-2-14
Lewis Structure of H2O
2 bond pairs= 2 x 2 =
4
2 lone pairs = 2 x 2 =
4
Total
8 = 4 pairs
Bond pairs: an electron pair shared by two atom in
a bond. E.g. two pairs between O--H in water.
Lone pair : an electron pair found solely on a single
atom. E.g. two pairs found on the O atom at the
top and the bottom.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-15
Lewis Structure of H2S
Chemistry 281, Winter 2016, LA Tech
Chapter-2-16
Lewis Structure of CCl4
Chemistry 281, Winter 2016, LA Tech
Chapter-2-17
What is the Lewis Structure?
• CO2
• NH3 (PH3)
•
PCl3 (PF3, NCl3)
Chemistry 281, Winter 2016, LA Tech
Chapter-2-18
Lewis structure and multiple bonds
O C O
This arrangement needs
too many electrons.
How about making some double bonds?
O=C=O
That works!
=
is a double bond,
the same as 4 electrons
Chemistry 281, Winter 2016, LA Tech
Chapter-2-19
Multiple bonds
So how do we know that multiple bonds really
exist?
The bond energies and lengths differ!
Bond
Bond Length Bond energy
type
order pm kJ/mol
C C
C C
C C
Chemistry 281, Winter 2016, LA Tech
1
2
3
154
134
120
347
615
812
Chapter-2-20
Formal Charges
Formal charge = valence electrons - assigned
electrons
•There are two possible Lewis structures for a molecule.
Each has the same number of bonds. We can determine
which is better by determining which has the least formal
charge. It takes energy to get a separation of charge in the
molecule
•(as indicated by the formal charge) so the structure with
the least formal charge should be lower in energy and
thereby be the better Lewis structure
Chemistry 281, Winter 2016, LA Tech
Chapter-2-21
Formal Charge Calculation
An arithmetic formula for calculating formal charge.
Formal charge =
group number
in periodic table
Chemistry 281, Winter 2016, LA Tech
number of
–
bonds
–
number of
unshared electrons
Chapter-2-22
Electron counts" and" formal
charges in NH4+ and BF4-
Chemistry 281, Winter 2016, LA Tech
Chapter-2-23
What is Resonance Structures?
•Several Lewis structures that need to be drawn
for molecules with double bonds
•One Lewis structure alone would not describe
the bond lengths of the real molecule.
•E.g. CO32-, NO3-, NO2-, SO3
Chemistry 281, Winter 2016, LA Tech
Chapter-2-24
Resonance structures
Sometimes we can have two or more
equivalent Lewis structures for a molecule.
O-S=O
O=S-O
They both - satisfy the octet rule
- have the same number of bonds
- have the same types of bonds
Which is right?
Chemistry 281, Winter 2016, LA Tech
Chapter-2-25
Resonance structures of SO2
They both are!
O -S=O
O
O =S- O
S
O
This results in an average of 1.5 bonds
between each S and O.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-26
Resonance structures of CO32- ion
Chemistry 281, Winter 2016, LA Tech
Chapter-2-27
Resonance structures of NO3- ion
Chemistry 281, Winter 2016, LA Tech
Chapter-2-28
Resonance structures of SO3
Chemistry 281, Winter 2016, LA Tech
Chapter-2-29
Resonance structures of NO2- ion
Chemistry 281, Winter 2016, LA Tech
Chapter-2-30
Resonance structures of C6H6
• Benzene, C6H6, is another example of a
compound for which resonance structure
must be written.
•
All of the bonds are the same length.
or
Chemistry 281, Winter 2016, LA Tech
Chapter-2-31
Exceptions to the octet rule
Not all compounds obey the octet rule.
• Three types of exceptions
• Species with more than eight electrons around
an atom.
• Species with fewer than eight electrons around
an atom.
• Species with an odd total number of electrons.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-32
Atoms with more than eight electrons
• Except for species that contain hydrogen, this is
the most common type of exception.
• For elements in the third period and beyond, the
d orbitals can become involved in bonding.
Examples
• 5 electron pairs around P in PF5
• 5 electron pairs around S in SF4
• 6 electron pairs around S in SF6
Chemistry 281, Winter 2016, LA Tech
Chapter-2-33
An example: SO42O
1. Write a possible
arrangement.
O
S
O
O
2. Total the electrons.
6 from S, 4 x 6 from O
add 2 for charge
total = 32
O
||
3. Spread the electrons
around.
Chemistry 281, Winter 2016, LA Tech
O - S- O
||
O
Chapter-2-34
Atoms with fewer than eight electrons
Beryllium and boron will both form
compounds where they have less
than 8 electrons around them.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-35
Atoms with fewer than eight electrons
Electron deficient. Species other than
hydrogen and helium that have fewer than 8
valence electrons.
They are typically very reactive species.
F-
F
|
B
|
F
Chemistry 281, Winter 2016, LA Tech
+
H
|
:N - H
|
H
F H
| |
F-B-N–H
| |
F H
Chapter-2-36
What is VSEPR Theory
Valence Shell Electron Pair Repulsion
This theory assumes that the molecular structure is
determined by the lone pair and bond pair
electron repulsion around the central atom
Chemistry 281, Winter 2016, LA Tech
Chapter-2-37
What Geometry is Possible around
Central Atom?
• What is Electronic or Basic Structure?
• Arrangement of electron pairs around the central
atom is called the electronic or basic structure
• What is Molecular Structure?
• Arrangement of atoms around the central atom is
called the molecular structure
Chemistry 281, Winter 2016, LA Tech
Chapter-2-38
Possible Molecular Geometry
1. Linear (180)
2. Trigonal Planar (120)
3. T-shape (90, 180)
4. Tetrahedral (109)
5. Square palnar ( 90, 180)
6. Sea-saw (90, 120, 180)
7. Trigonal bipyramid (90, 120, 180)
8. Octahedral (90, 180)
Chemistry 281, Winter 2016, LA Tech
Chapter-2-39
Molecular Structure from VSEPR
Theory
• H2O
• Bent or angular
• NH3
• Pyramidal
• CO2
• Linear
Chemistry 281, Winter 2016, LA Tech
Chapter-2-40
Molecular Structure from VSEPR Theory
•
•
•
•
•
•
SF6
Octahedral
PCl5
Trigonal bipyramidal
XeF4
Square planar
Chemistry 281, Winter 2016, LA Tech
Chapter-2-41
What is a Polar Molecule?
• Molecules with unbalanced electrical charges
• Molecules with a dipole moment
• Molecules without a dipole moment are called
non-polar molecules
Chemistry 281, Winter 2016, LA Tech
Chapter-2-42
How do you a Pick Polar Molecule?
• Get the molecular structure from VSEPR theory
• From c (electronegativity) difference of bonds see
whether they are polar-covalent.
• If the molecule have polar-covalent bond, check
whether they cancel from a symmetric
arrangement.
• If not molecule is polar
Chemistry 281, Winter 2016, LA Tech
Chapter-2-43
Which Molecules are Polar
• H 2O
• Bent or angular, polar-covalent bonds,
asymmetric molecule-polar
• NH3
• Pyramidal, polar-covalent bonds,
asymmetric molecule-polar
• CO2
• Linear, polar-covalent bonds,
symmetric molecule-polar
Chemistry 281, Winter 2016, LA Tech
Chapter-2-44
What is hybridization?
Mixing of atomic orbitals on the central atoms
valence shell (highest n orbitals)
Bonding: s
Px
p
Py
d
Pz
dz2
dx2- y2
sp,
sp2,
sp3,
sp3d,
sp3d2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-45
What is hybridization?
Mixing of atomic orbitals on the central atom
Bonding
a hybrid orbital could over lap with another ()atomic orbital
or () hybrid orbital of another atom to make a covalent
bond.
possible hybridizations: sp, sp2, sp3, sp3d, sp3d2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-46
What is Valence Bond Theory
• Describes bonding in molecule using atomic
orbital
• orbital of one atom occupy the same region
with a orbital from another atom
• total number of electrons in both orbital is
equal to two
Be Cl2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-47
sp2 and sp3 Hybridization
BF3
Chemistry 281, Winter 2016, LA Tech
Chapter-2-48
What are p and s bonds
s bonds
single bond resulting from head to head overlap of
atomic orbital
p bond
double and triple bond resulting from lateral or side
way overlap of atomic orbitals
Chemistry 281, Winter 2016, LA Tech
Chapter-2-49
How do you tell the hybridization of a central
atom?
•Get the Lewis structure of the molecule
•Look at the number of electron pairs on the central
atom. Note: double, triple bonds are counted as
single electron pairs.
•Follow the following chart
Chemistry 281, Winter 2016, LA Tech
Chapter-2-50
Kinds of hybrid orbitals
Hybrid
sp
sp2
geometry
# of orbital
linear
2
trigonal planar
3
sp3
sp3d
sp3d2
tetrahedral
trigonal bipyramid
octahedral
Chemistry 281, Winter 2016, LA Tech
4
5
6
Chapter-2-51
Hybridization involving d orbitals
•Co(NH3)63+ ion
Co3+: [Ar] 3d6
•Co3+: [Ar] 3d6 4s0 4p0
•Concentrating the 3d electrons in the dxy, dxz, and
dyz orbitals in this subshell gives the following
electron configuration hybridization is sp3d2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-52
Molecular Orbital Theory
• Molecular orbitals are obtained by combining the
atomic orbitals on the atoms in the molecule.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-53
Bonding and Anti-bobding Molecular Orbital
Chemistry 281, Winter 2016, LA Tech
Chapter-2-54
Basic Rules of Molecular Orbital Theory
The MO Theory has five basic rules:
• The number of molecular orbitals = the number of atomic
orbitals combined
• Of the two MO's, one is a bonding orbital (lower energy)
and one is an anti-bonding orbital (higher energy)
• Electrons enter the lowest orbital available
• The maximum # of electrons in an orbital is 2 (Pauli
Exclusion Principle)
• Electrons spread out before pairing up (Hund's Rule)
Chemistry 281, Winter 2016, LA Tech
Chapter-2-55
Bond Order
• Calculating Bond Order
Chemistry 281, Winter 2016, LA Tech
Chapter-2-56
Homo Nuclear Diatomic Molecules
Period 1 Diatomic Molecules: H2 and He2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-57
Homo Nuclear Diatomic Molecules
Period 2 Diatomic Molecules and Li2 and Be2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-58
Homo Nuclear Diatomic Molecules
Chemistry 281, Winter 2016, LA Tech
Chapter-2-59
Molecualr Orbital diagram for
O2, F2 and Ne2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-60
Molecualr Orbital diagram for
B2, C2 and N2
Chemistry 281, Winter 2016, LA Tech
Chapter-2-61
Homonuclear Diatomic Molecules
2nd Period
Chemistry 281, Winter 2016, LA Tech
Chapter-2-62
Electronic Configuration of molecules
When writing the electron configuration of an
atom, we usually list the orbitals in the order in
which they fill.
Pb: [Xe] 6s2 4f14 5d10 6p2
We can write the electron configuration of a
molecule by doing the same thing.
Concentrating only on the valence orbitals, we
write the electron configuration of O2 as
follows.
O2: (2s2s) 2(2s*2s) 2 (2s2p) 2 (2p2p) 4 (2p*2p) 2 ( 2s*2p)
Chemistry 281, Winter 2016, LA Tech
Chapter-2-63
Electronic Configuration and bond order
Chemistry 281, Winter 2016, LA Tech
Chapter-2-64
Hetero Nuclear Diatomic Molecules
HF molecule
Chemistry 281, Winter 2016, LA Tech
Chapter-2-65
Hetero Nuclear Diatomic Molecules
Carbon monoxide CO
Chemistry 281, Winter 2016, LA Tech
Chapter-2-66
Metallic Bonding
• Metals are held together by delocalized
bonds formed from the atomic orbitals of all
the atoms in the lattice.
• The idea that the molecular orbitals of the
band of energy levels are spread or
delocalized over the atoms of the piece of
metal accounts for bonding in metallic
solids.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-67
Bonding Models for Metals
•Band Theory of Bonding in Solids
•Bonding in solids such as metals,
insulators and semiconductors may be
understood most effectively by an
expansion of simple MO theory to
assemblages of scores of atoms
Chemistry 281, Winter 2016, LA Tech
Chapter-2-68
Linear Combination of Atomic Orbitals
Chemistry 281, Winter 2016, LA Tech
Chapter-2-69
Linear Combination of Atomic Orbitals
Chemistry 281, Winter 2016, LA Tech
Chapter-2-70
Chemistry 281, Winter 2016, LA Tech
Chapter-2-71
Types of Materials
• A conductor (which is usually a metal) is a
solid with a partially full band
• An insulator is a solid with a full band and
a large band gap
• A semiconductor is a solid with a full band
and a small band gap
• Element
C
Si
Ge
Sn
Chemistry 281, Winter 2016, LA Tech
Band Gap
5.47 eV
1.12 eV
0.66 eV
0
eV
Chapter-2-72
Chemistry 281, Winter 2016, LA Tech
Chapter-2-73
Superconductors
• When Onnes cooled mercury to 4.15K, the
resistivity suddenly dropped to zero
Chemistry 281, Winter 2016, LA Tech
Chapter-2-74
The Meissner Effect
•Superconductors show perfect diamagnetism.
•Meissner and Oschenfeld discovered that a
superconducting material cooled below its critical
temperature in a magnetic field excluded the
magnetic flux.Results in levitation of the magnet in
a magnetic field.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-75
Theory of Superconduction
•BCS theory was proposed by J. Bardeen, L. Cooper
and J. R. Schrieffer. BCS suggests the formation of
so-called 'Cooper pairs'
Cooper pair formation - electronphonon interaction: the electron
is attracted to the positive charge
density (red glow) created by the
first electron distorting the lattice
around itself.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-76
High Temperature Superconduction
•BCS theory predicted a theoretical maximum to Tc of around
30-40K. Above this, thermal energy would cause electronphonon interactions of an energy too high to allow
formation of or sustain Cooper pairs.
• 1986 saw the discovery of high temperature
superconductors which broke this limit (the highest known
today is in excess of 150K) - it is in debate as to what
mechanism prevails at higher temperatures, as BCS cannot
account for this.
Chemistry 281, Winter 2016, LA Tech
Chapter-2-77