Chapter-2-1 Chemistry 481, Spring 2014, LA Tech Instructor: Dr. Upali Siriwardane e-mail: [email protected] Office: CTH 311 Phone 257-4941 Office Hours:

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


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


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


•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?


1. Draw Lewis structure for SbF5, ClF3, and IF6+:


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


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


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)


2. Predict geometry of central atom using VSEPR and the hybridization in problem 1.

SbF5, ClF3, and IF6+:


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


Formal Charge Calculation

Formal charge =

group number

in periodic table

number of

bonds

number of

unshared electrons

––

An arithmetic formula for calculating formal charge.


Electron counts" and formal charges in NH4

+ and BF4-

"


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


Resonance structures of CO32- ion


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


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.


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


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


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


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


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


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


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


Possible hybridizations of s and p

sp-hybridization:


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


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


3. Why hypervalent compounds are formed by elements such as Si, P and S, but not by C,N and O?


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


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:

::

::

::


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


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


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.





Hybridization



General rule




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.


What are p and s bondss bonds

p bond


What are d bonds

d bonddouble and triple bond resulting from lateral or side way overlap of d atomic orbitals


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


5. Using valence-bond (VB) theory to explain the bonding in the coordination complex ion, Co(NH3)6

3+.


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


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





Hybridization



General rule




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)


Molecular Orbital Theory

• Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule.


Bonding and Anti-bobding Molecular Orbital


Bond Order • Calculating Bond Order


Homo Nuclear Diatomic Molecules Period 1 Diatomic Molecules: H2 and He2


Homo Nuclear Diatomic Molecules Period 2 Diatomic Molecules and Li2 and Be2


Homo Nuclear Diatomic Molecules


Molecualr Orbital diagram for

B2, C2 and N2


Molecualr Orbital diagram for

O2, F2 and Ne2


7. Using molecular orbital theory and diagrams, explain why, O2 is a paramagnetic whereas N2 is diamagnetic.


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


Electronic Configuration and bond order


Electronic Configuration and bond order


Hetero Nuclear Diatomic Molecules Carbon monoxide CO


8. Draw molecular orbital diagrams for HF, CO, NO, NO+. Calculate their bond order and predict magnetic properties.


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.


Walsh Diagram for H3:


9. Draw a molecular orbital diagram for triangular H3

+ and describe the bonding.


10. Draw a Walsh diagram (orbital correlation diagram) and show that triangular H3

+ is more stable than linear H3

+.


Conjugated and aromatic molecules• trans-1,3-Butadiene• Allyl radical • Cyclopropenium ion: C3H3

+

• Cyclobutadiene• Cyclopentadiene• Benzene• C7H7

+ (tropyllium) and C8H82+


trans-1,3-Butadiene


Allyl radical


Cyclopropenium ion: C3H3+


Cyclopentadiene


Benzene


Aromatic Rings


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+


The Isolobal Analogy• Different groups of atoms can give rise to

similar shaped fragments.



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


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.


Linear Combination of Atomic Orbitals


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


Linear Combination of Atomic Orbitals



Band Theory of Metals


13. Describe metallic bonding and properties in terms of:

a) Electron-sea model of bonding:

b) Band Theory:


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.


15. Describe the metallic properties of sodium in terms of band theory.


16. Using a band diagram, explain how magnesium can exhibit metallic behavior even though its 3s band is completely full.


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



17. Draw a Band diagram for carbon/silicon/germanium/tin, and label valence band, conduction band and band gap?


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):


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):


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):


22. What the difference between a transistor (semiconductor device) and vacuum tube?


What is a transistor?


21. What is a transistor with emitter (E), collector(C) and base (B), and how it works?


23. Using the diagram explain how a diode work.


Superconductors• When Onnes cooled mercury to 4.15K, the

resistivity suddenly dropped to zero


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.


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.


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.

Documents

Chapter-2-1 Chemistry 481, Spring 2014, LA Tech Instructor: Dr. Upali Siriwardane e-mail: [email protected] Office: CTH 311 Phone 257-4941 Office Hours: