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Chapter 9 Chemical Bonding I: Lewis Theory. Read/Study: Chapter 9 MGC Homework: Due April 17, 2008 at 11:50 p.m. MGC Quiz: Due by April 20, 2008 at 11:50 p.m. Chapters 7, 8, and 9. Chapters 9 and 10. Chemical Bond Types Ionic Pure Covalent Polar Covalent Coordinate Covalent

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Chapter 9

Chemical Bonding I: Lewis Theory

Read/Study:Chapter 9

MGC Homework: Due April 17, 2008 at 11:50 p.m.

MGC Quiz: Due by April 20, 2008 at 11:50 p.m.

Chapters 7, 8, and 9


Chapters 9 and 10
Chapters 9 and 10

  • Chemical Bond Types

    • Ionic

    • Pure Covalent

    • Polar Covalent

    • Coordinate Covalent

    • Metallic

  • Properties of Chemical Bonds

    • Bond Length

    • Bond Strength

    • Bond Angles

    • Bond Order


Chapters 9 and 101
Chapters 9 and 10

  • Chemical Bond Types

    • Ionic

    • Pure Covalent

    • Polar Covalent

    • Coordinate Covalent

    • Metallic


Chapter 9

  • Your Advanced Chemical Bonding Tool Kit

    • Valence Shell Electron Pair Repulsion

  • Valence Bond Theory (VB)

  • Molecular Orbital Theory (MO)

Chapter 10


  • Key Questions

    • Why do atoms combine to form

      compounds and molecules?

    • Why do they combine in definite

      proportions by mass?

    • Why do molecules have characteristic

      shapes?

  • Important Terms

    • Valence - The combining capacity of an

      element.

    • Chemical Bond - A “strong” link among

      atoms in a molecule or crystal.


Overview of Bonding Types

Ionic Bonding - An ionic bond is a chemical

bond that results from an electrostatic attraction

among oppositely charged ions in a compound.

They form when electrons are transferred from

one atom to another to form ions.

Na [Ne] 3s1 + Cl [Ne] 3s2 3p5

Na [Ne]+ + Cl [Ne]-

Cl-

Chapter 9

Na+


Overview of Bonding Types

Covalent Bonding - A covalent bond is a

chemical bond that results from a sharing of

electrons among the atoms in a compound.

e-

e-

+

+

+

Energy

H H

Chapter 9

e-

H2

+

e-

+


Overview of Bonding Types

Polar Covalent Bonding - A covalent bond that

occurs when the atoms unequally share one or

more pairs of electrons. This happens when the

atoms have different electronegativities.

e- e-

e-

e-

+

F

H

e-

e-

Energy

e- e-

e- e-

Chapter 9

e-

e-

e-

e-

d+

d-

H

F

e- e-


Overview of Bonding Types

Coordinate Covalent Bonding (Dative) - A dative

bond is a covalent bond that occurs when the two

shared electrons are donated to the bond bythe

same atom. The donating atom is the donor or

Lewis Base and the accepting atom is the acceptor

or Lewis Acid.

..

..

:F:

F:

:F:

B:F:

:F:

..

..

..

..

..

..

:F: B:F:

..

:

..

..

..

..

..

:F:

Donor

..

..

Acceptor

Tetrafluoroborate

ion


Overview of Bonding Types

Metallic Bonding - The force of attraction that

holds metal atoms together in a metallic lattice.

It results from the fact that the valence electrons

(outer shell electrons) are NOT bound to a

particular atom but are free to be shared by all

of the atoms - they are delocalized.

+ + + + + +

+ + + + + +

+ + + + + +


Summary of Bonding Types

Ionic

Bonding

Polar Covalent

Bonding

Covalent

Bonding

Electropositive

+

Electronegative

Electronegative

+

Electronegative

Metallic

Bonding

Electropositive

+

Electropositive


Assignment:

State whether each of the following has ionic,

covalent, polar covalent, or metallic bonding:

H2O NaI S4 HCl

Polar

Ionic

Covalent

Polar

K NH4Cl Rb C (diamond)

Metallic

Dative

Polar

and

ionic

Metallic

Covalent

Network


Chapter 9

  • Your Advanced Chemical Bonding Tool Kit

    • Valence Shell Electron Pair Repulsion

  • Valence Bond Theory (VB)

  • Molecular Orbital Theory (MO)

Chapter 10


Lewis Electron Dot Symbols

A symbol for an atom or ion consisting of the

chemical symbol for the element surrounded

by a number of dots equal to the number of

valence electrons in the atom.

Alkaline Earth Metals!!

H 1s1 H •

Li [He] 2s1 Li •

Na [Ne] 3s1Na •

K [Ar] 4s1 K •

Rb [Kr] 5s1 Rb •

Cs [Xe] 6s1Cs •

Be [He] 2s2Be:

Mg [Ne] 3s2Mg:

Ca [Ar] 4s2 Ca:

Sr [Kr] 5s2 Sr:

Ba [Xe] 6s2Ba:


Lewis Electron Dot Symbols

..

Assignment:

F [He] 2s22p5:F:

Cl [Ne] 3s23p5:Cl:

Br [Ar] 4s24p5:Br:

I [Kr] 5s25p5:I:

.

Draw the Lewis Symbols

for the following -

..

.

..

C P

Tl Ge

As Po

.

..

.

Used Primarily with

Main Group Elements


Ionic Bonding

Ionic Bonding - An ionic bond is a chemical

bond that results from an electrostatic attraction

among oppositely charged ions in a compound.

They form when electrons are transferred from

one atom to another to form ions.

Na [Ne] 4s1 + Cl [Ne] 3s2 3p5

Na[Ne]+ + Cl [Ne]-

Cl-

Na+


Ionic Bonding - “What’s yours is mine.”

“What’s mine you can have”

1. Metal Ion Formation - Metals in Groups 1 and 2

tend to give up one and two electrons, respectively,

to form 1+ and 2+ ions that are isoelectronic with

the preceding noble gas.

Positive ions

are called

“cations”.

Li•

Li+ + e-

[He] 2s1

[He]+

Sr2+ + 2 e-

Sr:

2+

[Kr] 5s2

[Kr]


2. Negative Ion Formation - The atoms of non-

metals in Groups 15, 16, and 17 tend to gain

electrons as to fill their valence shell s and p orbitals.

Thus, they become isoelectronic with the noble gas

at the end of their respective period.

..

: Br : + e-

..

: Br :

-

.

..

2-

..

: Se :

..

: Se : + 2 e-

..

.

: P : + 3 e-

..

: P :

3-

..


3. The Octet Rule - A “rule” expressing the

tendency of some main group elements to obtain

a total of 8 electrons in their valence shell. There

are MANY exceptions!

: Sn2+ + 2e-

[Kr] 4d10 5s2

: Sn :

[Kr] 4d10 5s2 5p2

: Tl+ + e-

[Xe] 4f14 5d106s2

: Tl •

[Xe] 4f14 5d106s2 6p1


  • 4. Formation of Ionic Bonds

    • Transfer of electrons to form ions.

    • Electrostatic attraction among the ions to

      form an ionic crystal lattice.

..

Ca : + : O :

..

Ca2+ + : O :

2-

..

-

+

-

-

6:6 Coordination;

Octahedral Arrangement

-

+

-

-

-



Covalent Bonding - “Share and share alike!”

Covalent bonding involves the sharing of electrons.

1.Lewis Structures of Molecules:

H · + · H H:H

H · + · O · + · H H:O:H

Octet

..

..

..

..

Shared Lone

Pairs Pairs

The total number of electrons in a Lewis structure

of a molecule is the sum of the valence electrons in

the individual atoms.


H

..

H

:

:

H:O:H C::C H:C:::C:H

..

:

:

H

H

Single

Bonds

Double

Bond

Triple

Bond

H

H

H-O-H C=C H-CC-H

H

H

  • Find the number of electrons in the Lewis

    structure by adding up the valence electrons

    of all the atoms in the molecule or ion; add one

    extra electron for each negative charge; subtract

    one electron for each positive charge.


1 C @ 4 valence electrons = 4 V.E.

2 O @ 6 valence electrons = 12 V.E.

Total Valence Electrons = 16 V.E.

CO2

  • Draw a skeletal structure of the molecule or

    ion by arranging atoms and putting one single

    bond between atoms that are bonded to each

    other.

O-C-O or O:C:O

  • Distribute the remaining electrons so as to

    satisfy the octet rule as closely as possible.

: O=C=O:

..

..


..

But what about….

:O:::C:O: ???

..

Both structures obey the octet rule. Which one

is RIGHT???

2. Formal Charges - A positive, negative, or zero

value assigned “formally” to the atoms in a Lewis

structure; it is calculated for each atom using the

following formula:

Formal Charge =

V.E. - [(unshared electrons) + 1/2(shared electrons)]

Purpose - To provide a method for predicting the

“best” Lewis structure for molecules with more

than one bonding possibility.


..

..

..

..

..

..

or :O=S-O:

:O-S=O:

..

..

FCOL = 6 - [6 + 1/2(2)] = 6 - [6 + 1] = 6 - 7 = -1

FCS = 6 - [2 + 1/2(6)] = 6 - [2 + 3] = 6 - 5 = +1

FCOR = 6 - [4 + 1/2(4)] = 6 - [4 + 2] = 6 - 6 = 0

..

..

..

..

..

..

:O-S=O:

:O=S-O:

..

..

-1 +1

+1 -1

Resonance Structures


  • Rules for Using Formal Charges

  • The most stable structures have the least

    formal charge.

  • Structures in which adjacent atoms have

    formal charges of the same sign tend to be

    unstable.

  • Structures in which positive charges are on

    more electronegative atoms are not as stable.

Since both of the SO2 structures are equivalent

(not “identical”), they both contribute equally and

should be shown as resonance forms.


  • 3. Resonance Structures - Two or more equivalent

    • ways to depict the bonding in a molecule or ion;

    • two or more equivalent and legitimate Lewis

    • structures for the same molecule or ion.

Assignment:

Draw all appropriate resonance structures for

the nitrate ion, NO3-, the carbonate ion, CO3-,

and the sulfur trioxide molecule, SO3.

Class Exercise:Construct appropriate Lewis

structures for CO2, calculate formal charges for

the atoms, and determine if resonance structures

are important.


  • 4. Exceptions to the Octet Rule

    • Atoms with more than 8 electrons - this is

      possible for elements in row 3 and beyond in

      the periodic table due to the fact that

      d-orbitals become available to handle addi-

      tional electrons.

..

..

:Cl:

:Cl:

P Cl:

:Cl:

:Cl:

..

:F:

:F F:

S

:F F:

:F:

..

..

..

..

..

..

..

..

..

..

..

..

..


..

:F:

B

:F: :F:

+1

- 1

:F:

B

:F: :F:

..

..

..

..

..

..

..

..

:Cl Be Cl: :Cl = Be = Cl:

..

..

+1 - 2 +1


..

..

..

:O - Cl - O:

..

.

..


  • Polar Covalent Bonding - “All are equal but

  • some are more ‘equal’ than others”

  • Bond Polarity - The result of an uneven charge

    distribution between two atomic nuclei that

    bonded to each other. It is due to the fact that

    different elements have different levels of

    attraction - electronegativity - for the electron

    pairs being shared.

  • Non-Polar Bond - A “pure” covalent bond

    wherein the two atoms sharing the electrons

    have identical attraction for them - they have

    the same electronegativity.


.. ..

H:H :Cl:Cl: :N:::N:

.. ..

Non-Polar Bonds (Pure Covalent Bonds)

  • Electronegativity - The general tendency of an

    atom or group of atoms to attract SHARED

    electrons.

    1. Electronegativity Scales - Relative values

    assigned to the elements in the Periodic Table

    to represent their attraction for electrons in

    a CHEMICAL BOND. (This is NOT the same

    as ELECTRON AFFINITY.)


2. Electronegativity Differences - The greater the

difference in electronegativity between two

atoms that are bonded together, the more polar

the bond will be.

E.N. for Cl = 3.2*

E.N. for H = 2.2

D E.N. for HCl = 1.0

 +

 -

H - Cl

*Pauling Scale

  • 3. Polar Molecules - A molecule is polar IF it has

    • one or more polar bonds AND is unsymmetrical

    • in charge distribution.


Assignment:

State whether or not each of the following

molecules is polar:

N2

H2O

SF6

CCl4

SO2

Non-Polar

Polar

Non-Polar

Non-Polar

Polar


Coordinate Covalent Bonding (Dative Bonding) -

“I have plenty so you may share”

This type of covalent bonding involves a sharing

of two electrons that have BOTH been provided

by only one of the atoms.

H

O: H+ H3O+

H

..

H3N:BF3

H3N: BF3


Chapter 10

Chemical Bonding II:

Molecular Shapes, Valence Bond

Theory, and Molecular Orbital Theory

Read/Study:Chapter 10

MGC Homework: Due April 23, 2008 at 11:50 p.m.

MGC Quiz: Due April 25, 2008 at 11:50 p.m.


Chapter 9

Chapter 10

  • Your Advanced Chemical Bonding Tool Kit

    • Valence Shell Electron Pair Repulsion

  • Valence Bond Theory (VB)

  • Molecular Orbital Theory (MO)


Valence Shell Electron Pair Repulsion : A model

of chemical bonding that allows the shapes of

molecules to be predicted by making the logical

assumption that electron pairs in molecules tend to stay as far apart as possible.

VSEPR is a sophisticated use of Lewis structures

to determine the geometry of polyatomic ions and

molecules.


  • Central Atom - An atom in a molecule or

    ion that is bonded to two or more other

    atoms.

  • Molecules with one Central Atom

    1. Write the Lewis structure.

..

..

:O = C = O:

2. Count the VSEPR electron pairs on the

  • Central Atom.

  • A. Count each bond - single, double, or

    • triple - as ONEVSEPR pair.


..

..

Total of 2 VSEPR

pairs on Carbon!

:O = C = O:

  • C. Place VSEPR pairs around the Central

    • Atom so that they are as far apart as

    • possible.

: C :

180o


  • D. Using ATOMS, not electron pairs,

    • determine the geometry.

O - C - O

Geometry is

Linear!

  • E. In determining geometry, take into

    • consideration the electron pair re-

    • pulsions:

    • (1) Bond Pair - Bond Pair Repulsion

    • (2) Bond Pair - Lone Pair Repulsion

    • (3) Lone Pair - Lone Pair Repulsion

BP-BP < BP-LP < LP-LP

Greater Repulsion


Geometry Types

# VSEPR

Pairs Geometry

2 Linear

3 Trigonal Planar

4 Tetrahedral

5 TrigonalBipyramidal

6 Octahedral


Assignment:

Determine the geometry of the following species:

Ammonia - NH3 Sulfur Hexafluoride - SF6

Nitrate Ion - NO3- Water - H2O

Ammonium Ion - NH4+ Sulfur Dioxide - SO2

Iodine Heptafluoride - IF7 Chlorite Ion - ClO2-


Chapter 9

Chapter 10

  • Your Advanced Chemical Bonding Tool Kit

    • Valence Shell Electron Pair Repulsion

  • Valence Bond Theory (VB)

  • Molecular Orbital Theory (MO)


Valence Bond Theory

  • 1. Limitations of Lewis Structures and VSEPR

    • Gives only information about geometry.

    • Is based on the “octet rule” which has many

      exceptions.

    • Cannot adequately explain bonding in

      species such as Li2 and H2+.

    • Does not reflect the QUANTUM nature of

      electrons.


  • 2. Quantum Mechanical Theories of Bonding

    • Valence Bond Theory - Involves the over-

      lapping of atomic orbitals from THE SAME

      atom.

    • Molecular Orbital Theory - Involves the

      formation of MOLECULAR orbitals around

      two or MORE nuclei in a molecule. The

      molecular orbitals are formed by the over-

      lapping of atomic orbitals from DIFFERENT

      atoms.


3. Valence Bond Theory

  • A. The Simplest View - Bonds are formed by

    • the simple overlap of atomic orbitals from

    • two different atoms.

Chapter 10

H

H

H 1s1 H-H s1s2 H 1s1

I 5 p1x

I 5 p1x

I-I s5p2


A single bond consists of 2 electrons of opposite

spin. The electrons are in a Sigma Bond (s).

Sigma Bond - A bond resulting from the overlap

of two atomic orbitals from DIFFERENT atoms,

resulting in the build-up of electron density along

the interatomic axis.

+500

Repulsion

H + H

Energy

0

ATTRACTION

H + H

74 pm; - 436 kJ/mol


B. Orbital Hybridization

Cl Be Cl

[He] 2s2

[Ne] 3s2 3p5

[Ne] 3s2 3p5

How do we explain the bonding in BeCl2 using

Valence Bond Theory?

We must invoke….


Orbital Hybridization!!!

An Imaginary mixing process in which the orbitals

of an atom rearrange to form new atomic orbitals

called Hybrid Orbitals.

2p

2p

Energy

sp hybrid

s + p

sp Hybridized Be Atom

2s


To explain what is known about the bonding in

BeCl2 using VB theory, we must find a way to

“un-pair” the paired electrons in Be to make it

“suitable” for bonding with Cl.

  • 1. Hybridize the same number of atomic orbitals

    • as there are bonds to explain.

  • 2. Promote the appropriate number of electrons

    • into the hybrid orbitals.

  • 3. Form bonds with other atoms by overlapping

    • the hybrid orbitals with either simple atomic

    • orbitals or hybrid orbitals on the other atoms.


BF3

sp2 hybridization:

Unhybridized p

2p

Energy

s + 2 p

sp2 hybrids

sp2 Hybridized B Atom

2s


sp3 hybridization:

:NH3

Lone Pair

2p

Energy

s + 3 p

sp3 hybrids

sp3 Hybridized N Atom

2s


sp3d hybridization:

PF5

Unhybridized d orbitals

3d

3p

Energy

s + 3 p + d

sp3d hybrids

sp3d Hybridized P Atom

3s


sp3d2 hybridization:

SF6

Unhybridized d orbitals

3d

3p

Energy

s + 3 p + 2 d

sp3d2 hybrids

sp3d2 Hybridized S Atom

3s



  • C. Multiple Bonding

    • (1) Review - Single bonds are ALWAYS

      • s - Bonds. The electron density is

      • predominantly between the nuclei on the

      • interatomic axis.

s-s s-bond

s1

s1

s-p s-bond

s1

p1


p1

p1

p-p s-bond

sp sp

s-sp s-bonds

s1

s1

  • (2) Multiple Bonds - Consist of one s-bond

    • and one or more p-bonds: bonds that

    • form by “side-wise” or “lateral” overlap

    • of parallel p-orbitals.


Double Bond Formation:

H2C=CH2

Unhybridized p

orbital

2p

Energy

s + 2 p

sp2 hybrids

2s

The sp2 hybrid orbitals form a trigonal plane

perpendicular to the unhybridized p-orbital. They

form s-bonds while two parallel p-orbitals form

p-bonds.


Assignment:

Describe the structure and bonding in O2 using

Lewis structures and valence bond theory. Show all

steps. Do these structures show any unpaired

electrons in dioxygen?

Use valence bond theory to explain why the bond

angles in ammonia are NOT 90o.

H

|

:N - H

|

H

Why Not??


Chapter 9

Chapter 10

  • Your Advanced Chemical Bonding Tool Kit

    • Valence Shell Electron Pair Repulsion

  • Valence Bond Theory (VB)

  • Molecular Orbital Theory (MO)


  • 4. Molecular Orbital Theory

    • A. Assumptions

    • (1) During bonding, atomic orbitals from

    • DIFFERENT atoms are transformed

    • into new orbitals with different shapes,

    • energies, and electron density distri-

    • butions.

    • (2) This is brought about by the overlap-

    • ping of atomic orbitals among different

    • atoms.

Chapter 10


  • (3) Molecular Orbitals are the allowed

  • states for an electron moving in the

  • electric field generated by two or more

  • nuclei. The Aufbau principle, the

  • Pauli Exclusion principle, and Hund’s

  • Rule of Maximum Multiplicity are all

  • used to fill Molecular Orbitals.

  • B. Rules

    • (1) The total number of molecular orbitals

    • is the same as the number of atomic

    • orbitals combined.


s1s*

1s

1s

Atomic

Orbitals

Atomic

Orbitals

s1s

Molecular Orbitals


s1s*

H

H

1s

1s

s1s

  • (3) A molecule is stable with respect to its

  • atoms whenever the number of bonding

  • electrons is greater than the number of

  • antibonding electrons.


s1s*

H2 (s1s)2

1s

1s

s1s

H2+ (s1s)1

H2- (s1s)2 (s1s*)1


He2+ (s1s)2 (s1s*)1

1s

1s

He2 (s1s)2 (s1s*)2

Non-Bonding


  • C. Bond Order -

    • Bond Order = ½(#bonding electrons -

    • # antibonding electrons)

H2 (s1s)2 B.O. = 1/2(2 - 0) = 1

H2+ (s1s)1 B.O. = 1/2(1 - 0) = 1/2

H2- (s1s)2(s1s*)1 B.O. = 1/2(2 - 1) = 1/2

He2+ (s1s)2(s1s*)1 B.O. = 1/2(2 - 1) = 1/2

He2 (s1s)2(s1s*)2 B.O. = 1/2(2 - 2) = 0


D. Homonuclear Diatomic Molelcules

of n = 2

The n = 2 shell has n2 or 4 atomic orbitals.

Therefore, two identical atoms can form

8 molecular orbitals.

2p

2p

Li2 - N2

pattern

2s

2s

1s

1s


s2p*

ppy*, ppz*

ppy, ppz

O2 - Ne2

pattern

s2px

s2s*

s2s

O2 is paramagnetic!!

s1s*

s1s

(s1s)2 (s1s*) 2 (s2s) 2 (s2s*) 2 (s2p) 2 (p2p) 2 (p2p*) 2


The molecular orbital diagrams for hetero-

nuclear diatomic moleculessimilar to those of

homonuclear diatomic molecules. However, the

atomic orbital energy levels are different, thus

causing the molecular orbital diagrams to be

unsymmetrical.

2s

2s

O

atom

C

atom

CO

Molecule


Some Final Thoughts….

1. So which is the most accurate “picture” of the

chemical bond? Is it Lewis Structures with its

resonance forms, formal charges, and the often

violated “octet rule”?

2. Maybe VSEPR with its predictive powers is best!

3. But, of course, the Valence Bond theory uses quantum

mechanical concepts and gives us more information

about what happens to atomic orbitals during bond

formation.


4. Yes, but the Molecular Orbital Theory is more

sophisticated and “intellectually satisfying”.

5. Perhaps we should use them all in different contexts

for different purposes. After all, they each are a

“window on the world” of molecular reality. We

need to understand them ALL; their limitations and

their strengths.

Which one should you use? The simplest one that

will answer the question you are asking!!



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