Chapter 9 Powerpoint - UIC Department of Chemistry

Chemistry and Chemical Reactivity

6th Edition

1

John C. Kotz

Paul M. Treichel

Gabriela C. Weaver

CHAPTER 9

Bonding and Molecular Structure:

Fundamental Concepts

Lectures written by John Kotz

© 2006 Brooks/Cole - Thomson

Cocaine

CHEMICAL

BONDING

2

Chemical Bonding

Problems and questions —

How is a molecule or polyatomic ion held together?

Why are atoms distributed at strange angles?

Why are molecules not flat?

Can we predict the structure?

How is structure related to chemical and physical properties?


3

Forms of Chemical Bonds

• There are 2 extreme forms of connecting or bonding atoms:

• Ionic —complete transfer of

1 or more electrons from one atom to another

• Covalent —some valence electrons shared between atoms

• Most bonds are somewhere in between.

4


Covalent Bonding

The bond arises from the mutual attraction of 2 nuclei for the same electrons.

Electron sharing results. (Screen 9.6)

H

A

+

H

B

H

A

H

B

Bond is a balance of attractive and repulsive forces.

5


Chemical Bonding:

Objectives

Objectives are to understand:

1. valence e- distribution in molecules and ions.

2. molecular structures

3. bond properties and their effect on molecular properties.


6

Bond and Lone Pairs

• Valence electrons are distributed as shared or

BOND PAIRS

and unshared or

LONE PAIRS

.


••

H Cl

••

•

• lone pair (LP) shared or bond pair

This is called a

LEWIS

ELECTRON DOT structure.

7

Valence Electrons

Electrons are divided between electrons

core

and valence

B atomic #5 has a total of 5 electrons. Two are in the

Core, [He] ; the remaining 3 are valence = 2s 2 2p 1

These valence electrons are available for bonding; we show them as “dots.”

8

Br [Ar] 3d 10 4s 2 4p 5

Core = [Ar] 3d 10 , valence = 4s 2 4p 5


Rules of the Game

No. of valence electrons of a main group atom = Group number

I II III IV V VI VII VIII

1 2 3 4 5 6 7 8

9


How many valence electrons are in …

• H

2

2(1)= 2 valence electrons

PCl

3

5+3(7)= 26 valence elctrons

CH

2

Cl

2

2(4)+2(1)+2(7)= 24 valence electrons


10

Lewis Structures

• Learn to find total number of valence electrons when given a chemical formula.

• Find which group each element belongs to.

• Group number = number of valence electrons.

• Include any charge in the total electron count

Each element in a Lewis structure has at least 8 electrons around it.

Except H, which has only 2.

This observation is called the

OCTET RULE


11

Building a Dot Structure

Ammonia, NH

3

1. Decide on the central atom; never H.

Central atom is atom of lowest affinity for electrons.

Therefore, N is central

2. Count valence electrons

H = 1 and N = 5

Total = (3 x 1) + 5

= 8 electrons / 4 pairs

12


Building a Dot Structure

3.

Form a single bond between the central atom and each surrounding atom

H N

H

H

4.

Remaining electrons form

LONE PAIRS to complete octet as needed.

H

••

N

3 BOND PAIRS and 1 LONE PAIR.

Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

H

H


13

Sulfite ion, SO

3

2-

How many valence electrons?

6+3(6) +2 = 26

Draw the skeleton with S in the middle.

Draw in all single bonds as lines

O

10 pairs of electrons are now left.

O S


O

14

Sulfite ion, SO

3

2-

Step 1. Central atom = S

Step 2. Count valence electrons

S = 6

3 x O = 3 x 6 = 18

Negative charge = 2

TOTAL = 26 e- or 13 pairs

Step 3. Form bonds

O

10 pairs of electrons are now left.

O S


O

15

Sulfite ion, SO

3

2-

Remaining pairs become lone pairs, first on outside atoms and then on central atom.

•

•

••

O •

•

•

•

••

••

O S

••

••

O

••

•

•

Each atom is surrounded by an octet of electrons.


16


17

Which of the following is NOT a correct Lewis dot structure?

18

1.

1

2.

2

3.

3

4.

4


0%

1

0%

2

0%

3

0%

4

Which of the following is NOT a correct

Lewis dot structure?

1.

2.

19

3.

4.

0%

1

0%

2

0%

3

0%

4


Double and even triple bonds are commonly observed for C,

N, P, O, and S

H

2

CO

SO

3

20


C

2

F

4


21

Sulfur Dioxide, SO

2

1. Central atom = S

2. Valence electrons = 18 or 9 pairs

3. Form double bond so that S has an octet

— but note that there are two ways of doing this.

bring in left pair

••

•

• O

••

••

S

OR bring in

•• right pair

O

••

•

•


22

Sulfur Dioxide, SO

2

This leads to the following structures.

These equivalent structures are called

RESONANCE STRUCTURES . The true electronic structure is a HYBRID of the two.


23

1.

3.

2.

4.


0%

1

0%

2

0%

3

0%

4

24


25

Formal Atom Charges

• Atoms in molecules often bear a charge (+ or -).

• The predominant resonance structure of a molecule is the one with charges as close to 0 as possible.

• Formal charge

= Group number

– (no. of bonds)

- (no. of LP electrons)

26


Carbon Dioxide, CO

2

27


• •

•

• O

+6 - ( 1 / 2 ) ( 4 ) - 4 = 0

C

• •

O

•

•

+4 - ( 1 / 2 ) ( 8 ) - 0 = 0

Formal charges on the formate ion:

29

1.

0, 0, 0

2.

+1, -1, -1

3.

+1, 0, -1

4.

0, 0, -1


0%

1

0%

2

0%

3

0%

4

Violations of the Octet Rule

Usually occurs with B and elements of higher periods.

30


BF

3

SF

4

Boron Trifluoride, BF

3

•

• F •

•

+

1

•

•

••

••

F B -

1

•

•

••

F •

•

What if we form a B —F double bond to satisfy the B atom octet?

31


Is There a B=F Double Bond in BF

3

Calc’d partial charges in BF

3

F is negative and B is positive

32


Sulfur Tetrafluoride, SF

4

• Central atom =

• Valence electrons = ___ or ___ pairs.

• Form sigma bonds and distribute electron pairs.

5 pairs around the S atom. A common occurrence outside the

2nd period.


33

MOLECULAR

GEOMETRY

MOLECULAR GEOMETRY

VSEPR

• V alence

S hell

E lectron

P air

R epulsion theory.

• Most important factor in determining geometry is relative repulsion between electron pairs.

Molecule adopts the shape that minimizes the electron pair repulsions.

35

Chemistry NOW


Rules to determine electronic structure

1. Look at Lewis Structure

2. Count electron domains a. Bonded atoms b. Lone pairs of electrons

3. Add # bonded atoms + # lp

4. Identify shape from the total # e- domains.


36

Electron Pair Geometries

Active Figure 9.8

37

Notice the bond angles for these.



Electron Pair Geometries

Active Figure 9.8

38

Structure Determination by

VSEPR

Ammonia, NH

3

1. Draw electron dot structure

2. Count BP’s and LP’s = 4

H

••

N

H

H

3. The 4 electron pairs are at the corners of a tetrahedron .

39 lone pair of electrons in tetrahedral position

N

H

H

H



VSEPR

Ammonia, NH

3

There are 4 electron pairs at the corners of a tetrahedron.

lone pair of electrons in tetrahedral position

N

H

H

H

The ELECTRON PAIR GEOMETRY is tetrahedral .


40


VSEPR

H

Ammonia, NH

3

The electron pair geometry is tetrahedral.

lone pair of electrons in tetrahedral position

N

H

H

The MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDAL .


41


VSEPR

Water, H

2

O


2. Count BP’s and LP’s = 4

3. The 4 electron pairs are at the corners of a tetrahedron.

The electron pair geometry is

TETRAHEDRAL.


42


VSEPR

Water, H

2

O

The electron pair geometry is

TETRAHEDRAL

43


The molecular geometry is

BENT .


Geometries for

Four Electron Pairs

Figure 9.9

44


VSEPR

•

•

Formaldehyde, CH

2

O

O •

•


H C H

2. Count BP’s and LP’s at C

3. There are 3 electron “lumps” around C at the corners of a planar triangle.

45

•

• O •

•

C

The electron pair geometry is PLANAR TRIGONAL with

120 o bond angles.

H H



VSEPR

H

Formaldehyde, CH

2

O

•

• O

•

•

The electron pair geometry is PLANAR

C TRIGONAL

H

The molecular geometry is also planar trigonal.

46



VSEPR

Methanol, CH

3

OH

Define H-C-H and C-O-H bond angles

Both the C atom and the

O atom are surrounded by 4 electron pairs.

H-C-H = 109 o

C-O-H = 109 o

H

109˚

••

H—C—O—H

••

H 109˚

47



VSEPR

Acetonitrile, CH

3

CN

Define unique bond angles

H-C-H = 109 o

C-C-N = 180 o

109˚

H

H—C—C N

H 180˚

One C is surrounded by 4 electron “lumps” and the other by 2 “lumps”

48


Fast Method to determine molecular structure!!

1. Count total valence electrons. SO

2

=18

2. Multiply the number of bonded atoms times 8. (H is multiplied by 2.) 2 x 8 = 16

3. Subtract these two numbers to get the number of lone pairs.

18-16 = 2 = 1 lp

4. Add #bonded atoms + # lp to get electronic geometry.

2 bonded + 1 lp

=3= trigonal planal; 120 º angles


49

Find electronic geometry for:

(IF

4

) -

Total valence electrons =

5(7)+1 =36

# lp =

36 - 4(8)= 4 = 2 lp

Electonic geometry and bond angles =

4+2=6; octahedral, 90 º


50

Find electronic geometry for:

(BrO

4

) -

Total valence electrons =

7+4(6)+1 =32

# lp =

32 - 4(8)= 0 = 0 lp

Electonic geometry and bond angles =

4+0=4; tetrahedral, 109 º


51

1.

3.

2.

4.


0%

1

0%

2

0%

3

0%

4

52

Structures with Central Atoms with More Than or Less Than 4

Electron Pairs

53

Often occurs with Group

3A elements and with those of 3rd period and higher.


Molecular

Geometries for

Five Electron

Pairs

Figure 9.11

All based on trigonal bipyramid

54


A certain molecule has five structural electron pairs and the molecule structure is linear. How many lone pairs are present on the central atom in this molecule?

55

25% 25% 25% 25%

1.

none

2.

one

3.

two

4.

three

1 2 3 4


A certain molecule has six structural electron pairs and the molecule structure is a square pyramid. How many lone pairs are present on the central atom in this molecule?

25% 25% 25% 25%

1.

none

2.

two

3.

one

4.

three

56

1 2 3 4


1.

3.

2.

4.

25% 25% 25% 25%

59

1 2 3 4


Bond Properties

• What is the effect of bonding and structure on molecular properties?

60

Free rotation around C –C single bond


No rotation around

C=C double bond

Bond Order

# of bonds between a pair of atoms

Double bond

Single bond

61


Acrylonitrile

Triple bond

Bond distances measured in

Angstrom units where 1 A =

10 -2 pm.


Bond Length

Bond length depends on bond order .

62

Molecular Polarity

Water

Boiling point =

100 ˚C

Methane

Boiling point

= 161 ˚C

63

Why do water and methane differ so much in their boiling points?

Why do ionic compounds dissolve in water?


+ d

d

••

H Cl

••

••

Bond Polarity

HCl is

POLAR because it has a positive end and a negative end.

Cl has a greater share in bonding electrons than does H.

64

Cl has slight negative charge (d

) and H has slight positive charge (+ d

)



Bond Polarity

• Three molecules with polar, covalent bonds.

• Each bond has one atom with a slight negative charge

(d

) and and another with a slight positive charge (+ d

)

65


Bond Polarity

This model, calc’d using CAChe software for molecular calculations, shows that H is +

(red) and Cl is (yellow). Calc’d charge is + or - 0.20.

66

Electronegativity,



 is a measure of the ability of an atom in a molecule to attract electrons to itself.

68

Concept proposed by Linus Pauling 1901-1994


Linus Pauling, 1901-1994

69

The only person to receive two unshared Nobel prizes

(for Peace and Chemistry).

Chemistry areas: bonding, electronegativity, protein structure


Electronegativity


70


Electronegativity

Figure 9.14

71

Electronegativity and Bond Polarity

Difference in electronegativity is a gauge of bond polarity. If differences are:

72

• around 0…then the covalent bond is non-polar

• around 2…then the covalent bond is polar

• around 3…then the bond is ionic

There is no sharp distinction between bonding types.

The positive end (or pole) in a polar bond is represented d

+ and the negative pole d

-.


Bond Polarity

Which bond is more polar (or DIPOLAR)?

O —H



3.5 - 2.1



1.4

O —F

3.5 - 4.0

0.5

-

OH is more polar than OF d

O

+ d

H

+ d

O

d

F and polarity is “reversed.”


73

Which of the following groups of elements is arranged correctly in order of increasing electronegativity?

74

25% 25% 25% 25%

1.

B < O < Al < F

2.

B < O < F < Al

3.

Al < B < O < F

4.

F < O < B < Al

1 2 3 4


Which of the following pairs of bonded atoms would be expected to have the greatest bond polarity?

75

25% 25% 25% 25%

1.

N-O

2.

K-F

3.

B-N

4.

S-Cl

1 2 3 4



Molecules —such as HI and H

2

O — can be POLAR (or dipolar).

They have a

DIPOLE MOMENT

. The polar HCl molecule will turn to align with an electric field.

76




The magnitude of the dipole is given in Debye units.

Named for Peter Debye

(1884 1966). Rec’d 1936

Nobel prize for work on xray diffraction and dipole moments.

77

Dipole Moments 78

Why are some molecules polar but others are not?



Molecules will be polar if a)bonds are polar

AND b) the molecule is NOT “symmetric”

79


All above are NOT polar


Polar or Nonpolar?

Compare CO

2 and H

2

O. Which one is polar?

80

To determine whether or not a molecule is polar,

1. Mark all bonds that are polar. These are vectors; they have direction and charge.

2. Equal and opposite vectors cancel out.

3. Determine the direction of any uncancelled vectors.

4. If molecule is symmetric, and all bonded atoms are the same, the molecule is nonpolar.


81

Dipole Moments of Polyatomic Molecules

Example: in CO

2, each C-O dipole is canceled because the molecule is linear. In H

2

O, the H-O dipoles do not cancel because the molecule is bent.

82


Carbon Dioxide

• CO

2 is NOT polar even though the CO bonds are polar.

• CO

2 is symmetrical.

83

Positive C atom is reason CO

2 and H

2 give H

2

O react to

CO

3

-0.75


+1.5

-0.75

Polar or Nonpolar?

• Consider AB

3 molecules: BF

3

, Cl

2

CO, and NH

3

.

84


F

Molecular Polarity, BF

3

F

B

F

B atom is positive and

F atoms are negative.

85

B —F bonds in BF

3 are polar.

But molecule is symmetrical and

NOT polar


F

H

Molecular Polarity, HBF

2

B

F

B atom is positive but H

& F atoms are negative.

86

B —F and B—H bonds in HBF

2 are polar. But molecule is NOT symmetrical and is polar.


Is Methane, CH

4

, Polar?

87

Methane is symmetrical and is NOT polar.


Is CH

3

F Polar?

88

C —F bond is very polar.

Molecule is not symmetrical and so is polar.


CH

4

… CCl

4

Polar or Not?

89

• Only CH

4 and CCl

4 are NOT polar. These are the only two molecules that are “symmetrical.”


Which are the polar molecules below? (Molecular shapes are indicated.

Lone pairs are not indicated)

90

25% 25% 25% 25%

1.

A, C, E

2.

B, C, E

3.

A, B, D

4.

C, D, E

1 2 3 4


Substituted Ethylene

91

• C —F bonds are MUCH more polar than

C —H bonds.

• Because both C —F bonds are on same side of molecule, molecule is POLAR .


Substituted Ethylene

92

• C —F bonds are MUCH more polar than C —H bonds.

• Because both C —F bonds are on opposing ends of molecule, molecule is NOT POLAR .


Chapter 9 Powerpoint - UIC Department of Chemistry

CHEMICAL

BONDING

Chemical Bonding

Forms of Chemical Bonds

• Most bonds are somewhere in between.

Covalent Bonding

Chemical Bonding:

Objectives

Bond and Lone Pairs

BOND PAIRS

LONE PAIRS

.

Valence Electrons

core

Rules of the Game

Lewis Structures

OCTET RULE

Building a Dot Structure

Building a Dot Structure

Sulfite ion, SO

Sulfite ion, SO

Sulfite ion, SO

Double and even triple bonds are commonly observed for C,

N, P, O, and S

Sulfur Dioxide, SO

Sulfur Dioxide, SO

Formal Atom Charges

Carbon Dioxide, CO

Violations of the Octet Rule

Boron Trifluoride, BF

F is negative and B is positive

Sulfur Tetrafluoride, SF

MOLECULAR

GEOMETRY

MOLECULAR GEOMETRY

VSEPR

Electron Pair Geometries

Electron Pair Geometries

Geometries for

Four Electron Pairs

Molecular

Geometries for

Five Electron

Pairs

Bond Properties

Electronegativity,

Linus Pauling, 1901-1994

Electronegativity

Carbon Dioxide

Molecular Polarity, BF

Molecular Polarity, HBF

Is Methane, CH

, Polar?

Is CH

F Polar?

CH

… CCl

Polar or Not?

Substituted Ethylene

Substituted Ethylene

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