Chapter-2-1 Chemistry 281, Winter 2015, LA Tech CTH 277 10:00-11:15 am Instructor: Dr. Upali...

Chapter-2-1Chemistry 281, Winter 2015, LA Tech

CTH 277 10:00-11:15 am

Instructor: Dr. Upali Siriwardane

E-mail: [email protected]

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

mailto:[email protected]


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


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.


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


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)


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


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.


Lewis symbols

Li Be B C

N O F Ne

Lewis symbols of second period elements


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.


Single covalent bonds

HH

H

CH H

H

Do atoms (except H) have octets?

F F


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.


•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?


Lewis Structure of H2O


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


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


Lewis Structure of H2S


Lewis Structure of CCl4


• CO2

• NH3 (PH3)

• PCl3 (PF3, NCl3)

What is the Lewis Structure?


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


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


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


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-

"


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


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


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


Resonance structures of CO32- ion


Resonance structures of NO3- ion


Resonance structures of SO3


Resonance structures of NO2- 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.


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


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.


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 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

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)


Molecular Structure from VSEPRTheory

• H2O

• Bent or angular• NH3

• Pyramidal• CO2

• Linear


Molecular Structure from VSEPR Theory

• SF6

• Octahedral• PCl5

• Trigonal bipyramidal• XeF4

• Square planar


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?• 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


•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


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


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


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


sp2 and sp3 Hybridization

BF3


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


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


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


Molecular Orbital Theory

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


Bonding and Anti-bobding Molecular Orbital


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)


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

O2, F2 and Ne2


Molecualr Orbital diagram for

B2, C2 and N2


Homonuclear Diatomic Molecules2nd Period


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)


Electronic Configuration and bond order


Hetero Nuclear Diatomic Molecules HF molecule


Hetero Nuclear Diatomic Molecules Carbon monoxide CO


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


Linear Combination of Atomic Orbitals


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


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.

Chapter-2-1 Chemistry 281, Winter 2015, LA Tech CTH 277 10:00-11:15 am Instructor: Dr. Upali...

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