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Transcript of Chapter-2-1 Chemistry 481, Spring 2014, LA Tech Instructor: Dr. Upali Siriwardane e-mail:...
Chapter-2-1Chemistry 481, Spring 2014, LA Tech
Instructor: Dr. Upali Siriwardane
e-mail: [email protected]
Office: CTH 311 Phone 257-4941
Office Hours:
M,W 8:00-9:00 & 11:00-12:00 am;
Tu,Th, F 9:30 - 11:30 a.m.
April 7 , 2015: Test 1 (Chapters 1, 2, 3)
April 30, 2015: Test 2 (Chapters 5, 6 & 7)
May 19, 2015: Test 3 (Chapters. 19 & 20)
May 19, Make Up: Comprehensive covering all Chapters
Chemistry 481(01) Spring 2015
Chapter-2-2Chemistry 481, Spring 2014, LA Tech
Molecular structure and bonding Lewis structures
2.1 The octet rule 2.2 Structure and bond properties 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 481, Spring 2014, LA Tech
What changes take place during this process of achieving closed shells?
a) sharing leads to covalent bonds and molecules Covalent Bond: each atom gives one electron
Coordinative bond: two electron comes from one atom
b) gain/loss of electrons lead to ionic bond Cations and anions: Electrostatic attractions
c) Sharing with many atoms lead to metallic bonds: delocalization of electrons
Chapter-2-4Chemistry 481, Spring 2014, 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-2-6Chemistry 481, Spring 2014, 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-2-7Chemistry 481, Spring 2014, 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-2-8Chemistry 481, Spring 2014, LA Tech
Possible Molecular Geometry
• Linear (180)• Trigonal Planar (120)• T-shape (90, 180)• Tetrahedral (109)• Square palnar ( 90, 180)• Sea-saw (90, 120, 180)• Trigonal bipyramid (90, 120, 180)• Octahedral (90, 180)
Chapter-2-9Chemistry 481, Spring 2014, LA Tech
2. Predict geometry of central atom using VSEPR and the hybridization in problem 1.
SbF5, ClF3, and IF6+:
Chapter-2-10Chemistry 481, Spring 2014, LA Tech
Formal ChargesFormal charge = valence electrons - assigned electrons
If 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-2-11Chemistry 481, Spring 2014, 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-2-12Chemistry 481, Spring 2014, LA Tech
Electron counts" and formal charges in NH4
+ and BF4-
"
Chapter-2-13Chemistry 481, Spring 2014, 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.Ave. Bond order= total pairs shared/ # bonds= 3/2=1.5
Resonance structures of SO2
Chapter-2-15Chemistry 481, Spring 2014, 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-2-16Chemistry 481, Spring 2014, 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-2-17Chemistry 481, Spring 2014, LA Tech
Valence-bond (VB) theory
VB theory combines the concepts of atomic orbitals,
hybrid orbitals, VSEPR, resonance structures, Lewis
structures and octet rule to describe the shapes and
structures of some common molecules.
It uses the overlap of atomic orbitals or hybrid orbitals of the
to from sigma ( )s , pi ( )p bonds and ( )d bonds
Chapter-2-18Chemistry 481, Spring 2014, LA Tech
Linear Combination of Atomic OrbitalsSymmetry Adapted
Linear Combination of Atomic Orbitals –LCAO
Atomic orbitals on single atom:
Hybridization
Atomic orbitals in a molecule with more than one atom:
Molecular Orbital (MO) formation
General rule
Number of Hybrid Orbital produced = # hybridized
Number of MO produced = # orbitals combined
Chapter-2-19Chemistry 481, Spring 2014, 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-2-20Chemistry 481, Spring 2014, 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-2-21Chemistry 481, Spring 2014, 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-2-22Chemistry 481, Spring 2014, 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-2-23Chemistry 481, Spring 2014, LA Tech
Possible hybridizations of s and p
sp-hybridization:
y1 = 1/Ö2ys - 1/Ö2yp
y2 = 1/Ö2ys + 1/Ö2yp
sp2
-hybridization:
y1 = 1/Ö3ys + 1/Ö6ypx + 1/Ö2ypy
y2 = 1/Ö3ys + 1/Ö6ypx - 1/Ö2ypy
y3 = 1/Ö3ys - 2/Ö6ypx
sp3
-hybridization:
y1 = 1/Ö4ys + 1/Ö4ypx + 1/Ö4ypy + 1/Ö4ypz
y2 = 1/Ö4ys - 1/Ö4ypx - 1/Ö4ypy + 1/Ö4ypz
y3 = 1/Ö4ys + 1/Ö4ypx - 1/Ö4ypy - 1/Ö4ypz
y4 = 1/Ö4ys - 1/Ö4ypx + 1/Ö4ypy -1/Ö4ypz
Chapter-2-24Chemistry 481, Spring 2014, LA Tech
Possible hybridizations of s and p
sp-hybridization:
Chapter-2-25Chemistry 481, Spring 2014, LA Tech
What are p and s bondss bondssingle bond resulting from head to head overlap of atomic orbital
p bonddouble and triple bond resulting from lateral or side way overlap of p atomic orbitals
d bonddouble and triple bond resulting from lateral or side way overlap of d atomic orbitals
Chapter-2-26Chemistry 481, Spring 2014, 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-2-27Chemistry 481, Spring 2014, LA Tech
3. Why hypervalent compounds are formed by elements such as Si, P and S, but not by C,N and O?
Chapter-2-28Chemistry 481, Spring 2014, 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-2-29Chemistry 481, Spring 2014, LA Tech
Atoms with fewer than eight electrons
Beryllium and boron will both form compounds where they have less than 8 electrons around them.
:Cl:Be:Cl: ::
::
:F:B:F:
:F:
::
::
::
Chapter-2-30Chemistry 481, Spring 2014, 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-2-31Chemistry 481, Spring 2014, 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-2-32Chemistry 481, Spring 2014, LA Tech
How do you a Pick Polar Molecule?a) Get the molecular structure from VSEPR theory
b) From c (electronegativity) difference of bonds see whether they are polar-covalent.
c) If the molecule have polar-covalent bond, check whether they cancel from a symmetric arrangement.
d) If not molecule is polar
Predicting symmetry of molecule and the polarity will be discussed in detail in Chapter 7.
Chapter-2-33Chemistry 481, Spring 2014, LA Tech
Linear Combination of Atomic OrbitalsSymmetry Adapted
Linear Combination of Atomic Orbitals –LCAO
Atomic orbitals on single atom:
Hybridization
Atomic orbitals in a molecule with more than one atom:
Molecular Orbital (MO) formation
General rule
Number of Hybrid Orbital produced = # hybridized
Number of MO produced = # orbitals combined
Chapter-2-34Chemistry 481, Spring 2014, LA Tech
6. Draw a diagram to illustrate each described overlap:
a) s bonding overlap of two p orbitals
b) d bonding overlap of two d orbitals
c) p bonding overlap of a p orbital and a d orbital
d) s antibonding overlap of a p and a d orbital
e) d antibonding overlap of two d orbitals.
Chapter-2-36Chemistry 481, Spring 2014, LA Tech
What are d bonds
d bonddouble and triple bond resulting from lateral or side way overlap of d atomic orbitals
Chapter-2-37Chemistry 481, Spring 2014, 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-2-38Chemistry 481, Spring 2014, LA Tech
5. Using valence-bond (VB) theory to explain the bonding in the coordination complex ion, Co(NH3)6
3+.
Chapter-2-39Chemistry 481, Spring 2014, 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-2-40Chemistry 481, Spring 2014, LA Tech
5. What is the oxidation state of metal in (a) Co(NH3)6
3+ ion (b) PtCl42- ion.
a) [Co(NH3)6] 3+
Co3+ and NH3 is neutral
Oxidation Sate of Co3+ is +3 and NH3 is 0
Therefore sum of the oxidation should be equal to +3 +3= Co(NH3)6 = (Co)3+6((NH3)0)= +3
Co is +3 in [Co(NH3)6]3+
b) Pt is +2 in [PtCl4]2- because Cl- is -1
Chapter-2-41Chemistry 481, Spring 2014, LA Tech
Linear Combination of Atomic OrbitalsSymmetry Adapted
Linear Combination of Atomic Orbitals –LCAO
Atomic orbitals on single atom:
Hybridization
Atomic orbitals in a molecule with more than one atom:
Molecular Orbital (MO) formation
General rule
Number of Hybrid Orbital produced = # hybridized
Number of MO produced = # orbitals combined
Chapter-2-42Chemistry 481, Spring 2014, 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-2-43Chemistry 481, Spring 2014, LA Tech
Molecular Orbital Theory
• Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule.
Chapter-2-46Chemistry 481, Spring 2014, LA Tech
Homo Nuclear Diatomic Molecules Period 1 Diatomic Molecules: H2 and He2
Chapter-2-47Chemistry 481, Spring 2014, LA Tech
Homo Nuclear Diatomic Molecules Period 2 Diatomic Molecules and Li2 and Be2
Chapter-2-51Chemistry 481, Spring 2014, LA Tech
7. Using molecular orbital theory and diagrams, explain why, O2 is a paramagnetic whereas N2 is diamagnetic.
Chapter-2-52Chemistry 481, Spring 2014, 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: (2 ) s 2(2s*) 2 (2 ) p 4 (2p*) 2
Chapter-2-55Chemistry 481, Spring 2014, LA Tech
Hetero Nuclear Diatomic Molecules Carbon monoxide CO
Chapter-2-56Chemistry 481, Spring 2014, LA Tech
8. Draw molecular orbital diagrams for HF, CO, NO, NO+. Calculate their bond order and predict magnetic properties.
Chapter-2-57Chemistry 481, Spring 2014, LA Tech
MO Correlation Diagrams ( Walsh Diagrams)
• The correlation diagram clearly indicates that the molecular orbital energy levels changes as the H3 changes from linear to cyclic (equilateral triangle) structure. In the case of
• linear H3 the overlap between two terminal H is minimal, where as in the case of cyclic H3 the overlap is substantial. This will bring the lowest MO (bonding) and the highest MO (antibonding) down in energy. At the same time, the non-bonding MO (middle one) will
• go up in energy, leading to a degenerate set of levels. Thus H3
+ (two electrons) will be triangular.
Chapter-2-59Chemistry 481, Spring 2014, LA Tech
9. Draw a molecular orbital diagram for triangular H3
+ and describe the bonding.
Chapter-2-60Chemistry 481, Spring 2014, LA Tech
10. Draw a Walsh diagram (orbital correlation diagram) and show that triangular H3
+ is more stable than linear H3
+.
Chapter-2-61Chemistry 481, Spring 2014, LA Tech
Conjugated and aromatic molecules• trans-1,3-Butadiene• Allyl radical • Cyclopropenium ion: C3H3
+
• Cyclobutadiene• Cyclopentadiene• Benzene• C7H7
+ (tropyllium) and C8H82+
Chapter-2-68Chemistry 481, Spring 2014, LA Tech
11. Using molecular orbital diagrams for pi (p) orbitals explain the relative stabilities of the following:
(a) C3H3 and C3H3+
(b) C4H4 and C4H4+
(c) C5H5 and C5H5-
(d) C6H6 and C6H6+
(e) C7H7 and C7H7+
Chapter-2-69Chemistry 481, Spring 2014, LA Tech
The Isolobal Analogy• Different groups of atoms can give rise to
similar shaped fragments.
Chapter-2-71Chemistry 481, Spring 2014, LA Tech
12. Pick the isolobal fragments among the following:a) Co3(CO)9Co(CO) 3, Co3(CO)9PR, Co3(CO)9CH
b) H3CCl, Mn(CO)5H, Re(CO) 5Cl
c) R2SiH2, Fe(CO)4H2, H2CH2
Chapter-2-72Chemistry 481, Spring 2014, 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-2-74Chemistry 481, Spring 2014, 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-2-78Chemistry 481, Spring 2014, LA Tech
13. Describe metallic bonding and properties in terms of:
a) Electron-sea model of bonding:
b) Band Theory:
Chapter-2-79Chemistry 481, Spring 2014, LA Tech
14. Draw the s band (molecular orbitals) for ten Na on a line (one dimensional) and show bonding and anti-bonding molecular orbitals and fill electrons.
Chapter-2-80Chemistry 481, Spring 2014, LA Tech
15. Describe the metallic properties of sodium in terms of band theory.
Chapter-2-81Chemistry 481, Spring 2014, LA Tech
16. Using a band diagram, explain how magnesium can exhibit metallic behavior even though its 3s band is completely full.
Chapter-2-82Chemistry 481, Spring 2014, 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-2-84Chemistry 481, Spring 2014, LA Tech
17. Draw a Band diagram for carbon/silicon/germanium/tin, and label valence band, conduction band and band gap?
Chapter-2-85Chemistry 481, Spring 2014, LA Tech
18. Draw a band diagrams to show the difference between(Band gaps: C = 5.47, Si = 1.12, Ge = 0.66, Sn = 0)
Conductor (Sn):
Insulator (C):
Semiconductor (Ge):
Chapter-2-86Chemistry 481, Spring 2014, LA Tech
19. Draw a band diagram for thermal/photo (Intrinsic) and doped (Extrinsic) semiconductors and explain the origin of semicondictivity?
Thermal/photo (Intrinsic) (Ge):
Doped (Extrinsic) (Si/As):
Chapter-2-87Chemistry 481, Spring 2014, LA Tech
20. Draw a band diagram for a p-type (Si/Ga) and n-type (Si/As) semiconductors and show holes and electrons that is responsible for semiconductivity.
p-type(Si/Ga):
n-type(Si/As):
Chapter-2-88Chemistry 481, Spring 2014, LA Tech
22. What the difference between a transistor (semiconductor device) and vacuum tube?
Chapter-2-90Chemistry 481, Spring 2014, LA Tech
21. What is a transistor with emitter (E), collector(C) and base (B), and how it works?
Chapter-2-92Chemistry 481, Spring 2014, LA Tech
Superconductors• When Onnes cooled mercury to 4.15K, the
resistivity suddenly dropped to zero
Chapter-2-93Chemistry 481, Spring 2014, 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-2-94Chemistry 481, Spring 2014, 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-2-95Chemistry 481, Spring 2014, 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.