chapter-2-1 chemistry 281, winter 2015, la tech cth 277 10:00-11:15 am instructor: dr. upali...
Post on 26-Dec-2015
224 Views
Preview:
TRANSCRIPT
Chapter-2-1Chemistry 281, Winter 2015, LA Tech
CTH 277 10:00-11:15 am
Instructor: Dr. Upali Siriwardane
E-mail: upali@latech.edu
Office: 311 Carson Taylor Hall ; Phone: 318-257-4941;
Office Hours: MTW 8:00 am - 10:00 am;
Th,F 8:30 - 9:30 am & 1:00-2:00 pm.
January 13, 2015 Test 1 (Chapters 1&,2),
February 3, 2015 Test 2 (Chapters 2 & 3)
February 26, 2015, Test 3 (Chapters 4 & 5),
Comprehensive Final Make Up Exam: March 3
Chemistry 281(01) Winter 2015
Chapter-2-2Chemistry 281, Winter 2015, LA Tech
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 diatomic2.10 Bond properties
Chapter-2-3Chemistry 281, Winter 2015, LA Tech
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.
Chapter-2-4Chemistry 281, Winter 2015, LA Tech
Noble gas configuration
The noble gases are noted for
their chemical stability and
existence as monatomic
molecules.
Except for helium,
They share a common electron
configuration
that is very stable.
This configuration has 8 valence-shell electrons.
All other elements reacts to achieve Noble Gas Electron
Configurations.
valence e-
He 2
Ne 8
Ar 8
Kr 8
Xe 8
Rn 8
valence e-
He 2
Ne 8
Ar 8
Kr 8
Xe 8
Rn 8
Chapter-2-5Chemistry 281, Winter 2015, LA Tech
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)
Chapter-2-6Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-7Chemistry 281, Winter 2015, LA Tech
Lewis Electron Dot symbols
X
Basic rules
Draw the atomic symbol.
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.
Chapter-2-8Chemistry 281, Winter 2015, LA Tech
Lewis symbols
Li Be B C
N O F Ne
Lewis symbols of second period elements
Chapter-2-9Chemistry 281, Winter 2015, LA Tech
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.
Chapter-2-10Chemistry 281, Winter 2015, LA Tech
Single covalent bonds
HH
H
CH H
H
Do atoms (except H) have octets?
F F
Chapter-2-11Chemistry 281, Winter 2015, LA Tech
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.
Chapter-2-12Chemistry 281, Winter 2015, LA Tech
•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)
How do you get the Lewis Structure from Molecular formula?
Chapter-2-13Chemistry 281, Winter 2015, LA Tech
Lewis Structure of H2O
Chapter-2-14Chemistry 281, Winter 2015, LA Tech
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.
H Cl
oo
oo
oo
oo
Bonding pair
Unshared
pair
Chapter-2-15Chemistry 281, Winter 2015, LA Tech
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.
Lewis Structure of H2O
Chapter-2-16Chemistry 281, Winter 2015, LA Tech
Lewis Structure of H2S
Chapter-2-17Chemistry 281, Winter 2015, LA Tech
Lewis Structure of CCl4
Chapter-2-18Chemistry 281, Winter 2015, LA Tech
• CO2
• NH3 (PH3)
• PCl3 (PF3, NCl3)
What is the Lewis Structure?
Chapter-2-19Chemistry 281, Winter 2015, LA Tech
Lewis structure and multiple bonds
O=C=O
This arrangement needs
too many electrons.
How about making some double bonds?
That works!
O C O
= is a double bond,
the same as 4 electrons
Chapter-2-20Chemistry 281, Winter 2015, LA Tech
Multiple bonds
So how do we know that multiple bonds really exist?
The bond energies and lengths differ!
BondBond Length Bond energy
type order pm kJ/mol
C C 1 154 347
C C 2 134 615
C C 3 120 812
Chapter-2-21Chemistry 281, Winter 2015, LA Tech
Formal ChargesFormal 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
Chapter-2-22Chemistry 281, Winter 2015, LA Tech
Formal Charge Calculation
Formal charge =
group number
in periodic table
number of
bonds
number of
unshared electrons
––
An arithmetic formula for calculating formal charge.
Chapter-2-23Chemistry 281, Winter 2015, LA Tech
Electron counts" and formal charges in NH4
+ and BF4-
"
Chapter-2-24Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-25Chemistry 281, Winter 2015, LA Tech
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?
Resonance structures
Chapter-2-26Chemistry 281, Winter 2015, LA Tech
They both are!
O - S = O O = S - O
O S OThis results in an average of 1.5 bonds
between each S and O.
Resonance structures of SO2
Chapter-2-27Chemistry 281, Winter 2015, LA Tech
Resonance structures of CO32- ion
Chapter-2-28Chemistry 281, Winter 2015, LA Tech
Resonance structures of NO3- ion
Chapter-2-29Chemistry 281, Winter 2015, LA Tech
Resonance structures of SO3
Chapter-2-30Chemistry 281, Winter 2015, LA Tech
Resonance structures of NO2- ion
Chapter-2-31Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-32Chemistry 281, Winter 2015, LA Tech
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.
Chapter-2-33Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-34Chemistry 281, Winter 2015, LA Tech
An example: SO42-
1. Write a possible arrangement.
2. Total the electrons.6 from S, 4 x 6 from O
add 2 for charge
total = 32
3. Spread the electronsaround.
S O
O
O
O
- - ||
||
S O
O
O
O
Chapter-2-35Chemistry 281, Winter 2015, LA Tech
Atoms with fewer than eight electrons
Beryllium and boron will both form compounds where they have less than 8 electrons around them.
Chapter-2-36Chemistry 281, Winter 2015, LA Tech
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
|
B
|
F
F - +
H
|
:N - H
|
H
F H
| |
F - B - N – H
| |
F H
Chapter-2-37Chemistry 281, Winter 2015, LA Tech
What is VSEPR TheoryValence 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
Chapter-2-38Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-39Chemistry 281, Winter 2015, LA Tech
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)
Chapter-2-40Chemistry 281, Winter 2015, LA Tech
Molecular Structure from VSEPRTheory
• H2O
• Bent or angular• NH3
• Pyramidal• CO2
• Linear
Chapter-2-41Chemistry 281, Winter 2015, LA Tech
Molecular Structure from VSEPR Theory
• SF6
• Octahedral• PCl5
• Trigonal bipyramidal• XeF4
• Square planar
Chapter-2-42Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-43Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-44Chemistry 281, Winter 2015, LA Tech
•H2O
•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
Which Molecules are Polar
Chapter-2-45Chemistry 281, Winter 2015, LA Tech
What is hybridization?
Mixing of atomic orbitals on the central atoms
valence shell (highest n orbitals)
Bonding: s p d
sp,
sp2,
sp3,
sp3d,
sp3d2
Px Py Pz dz2
dx2
- y2
Chapter-2-46Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-47Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-48Chemistry 281, Winter 2015, LA Tech
sp2 and sp3 Hybridization
BF3
Chapter-2-49Chemistry 281, Winter 2015, LA Tech
What are p and s bondss 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
Chapter-2-50Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-51Chemistry 281, Winter 2015, LA Tech
Kinds of hybrid orbitals
Hybrid geometry # of orbital
sp linear 2
sp2 trigonal planar 3
sp3 tetrahedral 4
sp3d trigonal bipyramid 5
sp3d2 octahedral 6
Chapter-2-52Chemistry 281, Winter 2015, LA Tech
Hybridization involving d orbitals•Co(NH3)6
3+ 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
Chapter-2-53Chemistry 281, Winter 2015, LA Tech
Molecular Orbital Theory
• Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule.
Chapter-2-54Chemistry 281, Winter 2015, LA Tech
Bonding and Anti-bobding Molecular Orbital
Chapter-2-55Chemistry 281, Winter 2015, LA Tech
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)
Chapter-2-56Chemistry 281, Winter 2015, LA Tech
Bond Order • Calculating Bond Order
Chapter-2-57Chemistry 281, Winter 2015, LA Tech
Homo Nuclear Diatomic Molecules Period 1 Diatomic Molecules: H2 and He2
Chapter-2-58Chemistry 281, Winter 2015, LA Tech
Homo Nuclear Diatomic Molecules Period 2 Diatomic Molecules and Li2 and Be2
Chapter-2-59Chemistry 281, Winter 2015, LA Tech
Homo Nuclear Diatomic Molecules
Chapter-2-60Chemistry 281, Winter 2015, LA Tech
Molecualr Orbital diagram for
O2, F2 and Ne2
Chapter-2-61Chemistry 281, Winter 2015, LA Tech
Molecualr Orbital diagram for
B2, C2 and N2
Chapter-2-62Chemistry 281, Winter 2015, LA Tech
Homonuclear Diatomic Molecules2nd Period
Chapter-2-63Chemistry 281, Winter 2015, LA Tech
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)
Chapter-2-64Chemistry 281, Winter 2015, LA Tech
Electronic Configuration and bond order
Chapter-2-65Chemistry 281, Winter 2015, LA Tech
Hetero Nuclear Diatomic Molecules HF molecule
Chapter-2-66Chemistry 281, Winter 2015, LA Tech
Hetero Nuclear Diatomic Molecules Carbon monoxide CO
Chapter-2-67Chemistry 281, Winter 2015, LA Tech
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.
Chapter-2-68Chemistry 281, Winter 2015, LA Tech
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
Chapter-2-69Chemistry 281, Winter 2015, LA Tech
Linear Combination of Atomic Orbitals
Chapter-2-70Chemistry 281, Winter 2015, LA Tech
Linear Combination of Atomic Orbitals
Chapter-2-71Chemistry 281, Winter 2015, LA Tech
Chapter-2-72Chemistry 281, Winter 2015, LA Tech
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 Band Gap C 5.47 eVSi 1.12 eVGe 0.66 eVSn 0 eV
Chapter-2-73Chemistry 281, Winter 2015, LA Tech
Chapter-2-74Chemistry 281, Winter 2015, LA Tech
Superconductors• When Onnes cooled mercury to 4.15K, the
resistivity suddenly dropped to zero
Chapter-2-75Chemistry 281, Winter 2015, LA Tech
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.
Chapter-2-76Chemistry 281, Winter 2015, LA Tech
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 - electron-phonon interaction: the
electron is attracted to the positive charge density (red
glow) created by the first electron distorting the lattice
around itself.
Chapter-2-77Chemistry 281, Winter 2015, LA Tech
High Temperature Superconduction•BCS theory predicted a theoretical maximum to Tc of around 30-40K. Above this, thermal energy would cause electron-phonon 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.
top related