HW#1-chem481-chapter

Chemistry 481(01) Spring 2015

Instructor: Dr. Upali Siriwardane e-mail: upali@latech.edu

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, Spring 2014, LA Tech Chapter-2-1

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 diatomic

2.10 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


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)


1. Draw Lewis structure for SbF

5

, ClF

3

, and IF

6

+ :


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


•

•

•

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.

SbF

5

, ClF

3

, and IF

6

+ :


Formal Charges

Formal 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

An arithmetic formula for calculating formal charge.

Formal charge = group number in periodic table

– number of bonds

– number of unshared electrons


Electron counts" and formal

" charges in NH

4

+ and BF

4

-


Resonance structures of SO

2

They both are!

O - S = O O = S - O

O S O

This 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 CO

3

2-

ion


Resonance structures of C

6

H

6

• Benzene, C

6

H

6

, 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 Orbitals

Symmetry 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, sp 2 , sp 3 , sp 3 d, sp 3 d 2


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 sp sp 2 sp 3 sp 3 d sp 3 d 2 geometry linear trigonal planar

# of orbital

2

3 tetrahedral 4 trigonal bipyramid 5 octahedral 6


What is hybridization?

Mixing of atomic orbitals on the central atoms valence shell (highest n orbitals)

Bonding: s p d

P x

P y

P z d z

2 sp, sp 2 , sp 3 , sp 3 d, sp 3 d 2

Chemistry 481, Spring 2014, LA Tech d x

2 y

2

Chapter-2-22

Possible hybridizations of s and p sp-hybridization:



1



2

= 1/



2

 s

= 1/



2

 s

- 1/



2

 p

+ 1/



2

 p sp 2 -hybridization:



1



2



3

= 1/



3

 s

= 1/



3

 s

= 1/



3

 s

+ 1/



6

 px

+ 1/



6

 px

- 2/



6

 px

+ 1/



2

 py

- 1/



2

 py sp 3 -hybridization:



1



2



3



4

= 1/



4

 s

= 1/



4

 s

= 1/



4

 s

= 1/



4

 s

+ 1/



4

 px

- 1/



4

 px

+ 1/



4

 px

- 1/



4

 px

+ 1/



4

 py

- 1/



4

 py

+ 1/



4

 pz

+ 1/



4

 pz

- 1/



4

 py

+ 1/



4

 py

- 1/



4

 pz

-1/



4

 pz


Possible hybridizations of s and p sp-hybridization:


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 p atomic orbitals d bond double 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 PF

5

5 electron pairs around S in SF

4

6 electron pairs around S in SF

6


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

O?


An example: SO

4

2-

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 electrons around.

Chemistry 481, Spring 2014, LA Tech

O

O S O

O

O

||

O S O

||

O

Chapter-2-28

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) b) c) d)

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

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 bonds s bonds p bond


What are d bonds d bond double and triple bond resulting from lateral or side way overlap of d atomic orbitals


Kinds of hybrid orbitals

Hybrid sp sp 2 sp 3 sp 3 d sp 3 d 2 geometry linear trigonal planar

# of orbital

2

3 tetrahedral 4 trigonal bipyramid 5 octahedral 6


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

Co(NH

3

)

6

3+ .


Hybridization involving d orbitals

• Co(NH

3

)

6

3+ ion Co 3+ : [Ar] 3 d 6

• Co 3+ : [Ar] 3 d 6 4 s 0 4 p 0

• Concentrating the 3 d electrons in the d xy

, d xz

, and d yz orbitals in this subshell gives the following electron configuration hybridization is sp 3 d 2


5. What is the oxidation state of metal in (a)

Co(NH

3

)

6

3+ ion (b) PtCl

4

2ion.

a) [Co(NH

3

)

6

] 3+

Co 3+ and NH

3 is neutral

Oxidation Sate of Co 3+ is +3 and NH

3 is 0

Therefore sum of the oxidation should be equal to

+3

+3= Co(NH

3

)

6

= (Co)3+6((NH

3

)0)= +3

Co is +3 in [Co(NH

3

)

6

] 3+ b) Pt is +2 in [PtCl

4

] 2because 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 (

Exclusion Principle )

Pauli

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: H

2 and He

2



Period 2 Diatomic Molecules and Li

2 and Be

2




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, O diamagnetic.

2 is a paramagnetic whereas N

2 is


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] 6 s 2 4 f 14 5 d 10 6 p 2

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 O follows.

2 as

O

2

: (

2 s)

2 (2 s

*) 2 (2 p)

4 (2 p

*

)

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 H

3 changes from linear to cyclic (equilateral triangle) structure. In the case of linear H

3 the overlap between two terminal H is minimal, where as in the case of cyclic H

3 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 H

3

+ (two electrons) will be triangular.


Walsh Diagram for H

3

:


9. Draw a molecular orbital diagram for triangular H

3

+ and describe the bonding.


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

3

+ stable than linear H

3

+ .

is more


Conjugated and aromatic molecules

•

•

•

•

•

•

• trans-1,3-Butadiene

Allyl radical

Cyclopropenium ion: C

3

H

3

+

Cyclobutadiene

Cyclopentadiene

Benzene

C

7

H

7

+ (tropyllium) and C

8

H

8

2+


trans -

1,3-Butadiene


Allyl radical


Cyclopropenium ion: C

3

H

3

+


Cyclopentadiene


Benzene


Aromatic Rings


11. Using molecular orbital diagrams for pi (p) orbitals explain the relative stabilities of the following:

(a) C

3

H

3 and C

3

H

3

+

(b) C

4

H

4 and C

4

H

4

+

(c) C

5

H

5 and C

5

H

5

-

(d) C

6

H

6 and C

6

H

6

+

(e) C

7

H

7 and C

7

H

7

+


•

The Isolobal Analogy

Different groups of atoms can give rise to similar shaped fragments.



12. Pick the isolobal fragments among the following: a) Co

3

(CO)

9

Co(CO)

3

, Co

3

(CO)

9

PR, Co

3

(CO)

9

CH b) H

3

CCl, Mn(CO)

5

H, Re(CO)

5

Cl c) R

2

SiH

2

, Fe(CO)

4

H

2

, H

2

CH

2


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.




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





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

C

Si

Ge

Sn

Band Gap

5.47 eV

1.12 eV

0.66 eV

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 ntype (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 - electronphonon 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 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.


HW#1-chem481-chapter

Molecular structure and bonding

What is VSEPR Theory

Possible Molecular Geometry

Resonance structures of CO

ion

Resonance structures of C

H

Exceptions to the octet rule

Kinds of hybrid orbitals

An example: SO

Kinds of hybrid orbitals

Hybridization involving d orbitals

Bond Order

Homo Nuclear Diatomic Molecules

Hetero Nuclear Diatomic Molecules

Conjugated and aromatic molecules

trans -

1,3-Butadiene

Allyl radical

Cyclopropenium ion: C

H

Cyclopentadiene

Benzene

Aromatic Rings

The Isolobal Analogy

Metallic Bonding

Bonding Models for Metals

Types of Materials

Superconductors

The Meissner Effect

Theory of Superconduction

High Temperature Superconduction

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HW#1-chem481-chapter

Molecular structure and bonding

What is VSEPR Theory

Possible Molecular Geometry

Resonance structures of CO

ion

Resonance structures of C

H

Exceptions to the octet rule

Kinds of hybrid orbitals

An example: SO

Kinds of hybrid orbitals

Hybridization involving d orbitals

Bond Order

Homo Nuclear Diatomic Molecules

Hetero Nuclear Diatomic Molecules

Conjugated and aromatic molecules

trans -

1,3-Butadiene

Allyl radical

Cyclopropenium ion: C

H

Cyclopentadiene

Benzene

Aromatic Rings

The Isolobal Analogy

Metallic Bonding

Bonding Models for Metals

Types of Materials

Superconductors

The Meissner Effect

Theory of Superconduction

High Temperature Superconduction

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