1 Chapter 9 Covalent Bonding Molecular Compounds

1

Chapter 9

Covalent Bonding

Molecular Compounds

2

How does H2 form? The nuclei repel

++

3

How Does H2 Form?

++

The nuclei repel But they are attracted to electrons So they share the electrons

4

Covalent Bonds Nonmetals hold onto their valence

electrons They don’t give away electrons to bond Still want noble gas configuration Get it by sharing valence electrons with

each other By sharing both atoms get to count the

electrons toward their noble gas configuration

5

Covalent Bonding Fluorine has seven valence electrons

F

6

Covalent Bonding Fluorine has seven valence electrons A second atom also has seven

F F

7

Covalent Bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons

F F

8


F F

9


F F

10


F F

11


F F

12

Covalent Bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with a full outer orbital

F F

13


F F8 Valence electrons

14


F F8 Valence electrons

15

Single Covalent Bond A sharing of two valence electrons Only nonmetals and Hydrogen form

covalent bonds Different from an ionic bond because

they actually form molecules Two specific atoms are joined In an ionic solid you can’t tell which

atom the electrons moved from or to

16

Show How They Formed It’s like a jigsaw puzzle. I have to tell you what the final formula

is. You put the pieces together to end up

with the right formula. For example- let’s look at how water

H2O is formed with covalent bonds.

17

Water

H

O

Each hydrogen has 1 valence electron

Each hydrogen wants 1 more

The oxygen has 6 valence electrons

The oxygen wants 2 more

They share to make each other happy

18

Water Put the pieces together The first hydrogen is happy The oxygen still wants one more

H O

19

Water The second hydrogen attaches Every atom has full energy levels

H OH

20

Multiple Bonds Sometimes atoms share more than one

pair of valence electrons. A double bond is when atoms share two

pair (4 total) of electrons. A triple bond is when atoms share three

pair (6 total) of electrons.

21

Carbon Dioxide CO2 - Carbon is central

atom ( I have to tell you)

Carbon has 4 valence electrons

Wants 4 more Oxygen has 6 valence

electrons Wants 2 more

O

C

22

Carbon Dioxide Attaching 1 oxygen leaves the oxygen 1

short and the carbon 3 short

OC

23

Carbon Dioxide Attaching the second oxygen leaves

both oxygen 1 short and the carbon 2 short

OCO

24

Carbon Dioxide The only solution is to share more

OCO

25


OCO

26


OCO

27


OCO

28


OCO

29


OCO

30

Carbon Dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in

the bond

OCO

31


the bond

OCO8 valence electrons

32


the bond

OCO8 valence electrons

33


the bond movie

OCO

8 valence electrons

34

How to Draw Them Add up all the (total have) valence

electrons Count up the total number of electrons

(total want) to make all atoms happy Subtract total have from total want Divide by 2 Tells you how many bonds - draw them. Fill in the rest of the valence electrons

to fill atoms up

35

NH3 Example NH3 – 1 N and 3 H’s One N - has 5 valence electrons

and wants 8 total Each H’s - has 1 valence electron

and wants 2 total NH3 has 5+3(1) = 8 Total Have

NH3 wants 8+3(2) = 14 Total Want (14-8)/2= 3 bonds 4 atoms with 3 bonds

N

HH

H

36

NHH

NH3 Example Draw in the 3 bonds Account for all 8 valence electrons Now everything has a full octet

H

37

HCN Example With HCN - C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8 H - has 1 valence electrons wants 2 HCN has 5+4+1 = 10

HCN wants 8+8+2 = 18

(18-10)/2= 4 bonds 3 atoms with 4 bonds -will require multiple

bonds - not to H

38

HCN Put in single bonds Need 2 more bonds Must go between C and N

NH C

39

HCN Put in single bonds Need 2 more bonds Must go between C and N (triple bond) Uses 8 electrons - 2 more to add

NH C

40

HCN Put in single bonds Need 2 more bonds Must go between C and N (triple bond) Uses 8 electrons of 10 - 2 more to add 2 must go on N to fill octet

NH C

41

Another Way of Indicating Bonds – Dots and Bars

Can use a bar (line) to indicate a bond Called a Lewis Structure or

Structural Formula Each bar ( ) represents a bonded pair

of valence electrons

H HO = H HO

42

Structural Examples

H C N

C OH

H

C has 8 electrons because each line is 2 electrons

Same for N

The same for C in this case

Same for O with two lines and 4 electrons

43

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond Carbon monoxide CO

OC

44



OC

45



OC

46

How Do We Know if Coordinate Covalent?

Have to draw the diagram and see what happens.

Often happens with polyatomic ions and acids.

47

Resonance When more than one valid dot diagrams

with the same connections are possible. Consider the two ways to draw ozone

(O3). Which one is it? Does it go back and forth? Is it a hybrid of both, like a mule?

Consider the three ways to draw NO3-

48

It is a Mule!

The hybrid bonds are shorter than single bonds, but longer than double bonds.

49

VSEPR Valence Shell Electron Pair Repulsion Predicts three dimensional geometry of

molecules Name tells you the theory: Valence shell - outside electrons Electron Pair Repulsion - electron pairs

try to get as far away as possible Can determine the angles of bonds Movie

50

VSEPR Based on the number of pairs of

valence electrons both bonded and unbonded.

Unbonded pairs are called lone pairs. CH4 - draw the structural formula

Has 4 + 4(1) = 8 wants 8 + 4(2) = 16 (16-8)/2 = 4 bonds

51

VSEPR Single bonds fill

all atoms. There are 4 pairs

of electrons pushing away.

The furthest they can get away is 109.5º.

C HH

H

H

52

4 Atoms Bonded to Central Basic shape is called

Tetrahedral. Like a pyramid with a

triangular base. Same shape for

everything with 4 pairs. CH H

H

H109.5º

53

Tetrahedral Shape

54

3 Bonded - 1 Lone Pair

N HH

H

NH HH

<107.3º

Still basic tetrahedral but has an unshared electron pair.

Shape is calledTrigonal Pyramidal.

55

Trigonal Pyramidal

56

2 Bonded - 2 Lone Pair

OH

H

O HH

<104.5º

Still basic tetrahedral but has 2 lone pair of electrons.

Shape is calledBent.

57

Bent Shape

Early Microwave Oven

58

Consequences of Consequences of HH22O PolarityO Polarity

Microwave Microwave ovenoven

59

3 Atoms No Lone Pairs

CH

HO

If the atoms all are the same, the farthest you can move the shared electrons apart is 120º

BF

FF

60

3 Atoms No Lone Pairs

BF

FF

The farthest you can move the electron pair apart is 120º.

Shape is flat and called Trigonal Planar.

B

F

F F

120º

61

Trigonal Planar

62

2 Atoms No Lone Pairs With three atoms the farthest they can

get apart is 180º. Shape called Linear.

C OO180º

63

Molecular Shapes

64

Two Types of Bonds Sigma bonds (σ bond) forms from the overlap of orbitals Sigma is centered between the atoms Pi bond ( bond) above and below atoms Between adjacent p orbitals. The two bonds of a

double bond are

one sigma and

one pi bond.

65

CCH

H

H

H

66

Bond Polarity Electrons are pulled, as in a tug-of-war,

between the atoms nuclei The greater the difference in

electronegativity between two bonded atoms, the more polar the bond.

(-)S N (+) If the difference is great enough,

electrons are transferred from the less electronegative atom to the more electronegative one to form a bond that has ionic character.- Ionic bondIonic bond.

67

Polar Bonds In equal sharing (such as diatomic

molecules), the bond that results is called a nonpolar covalent bond

When two different atoms are connected, the electrons may not be shared equally.

This unequal sharing is a polar covalent bond or simply a polar bond.

How do we measure how strong the atoms pull on electrons?

68

Electronegativity A measure of how strongly the atoms

attract electrons in a bond. The bigger the electronegativity

difference the more polar the bond. Fig 9-15 gives electronegativity values 0.0 - 0.4 Covalent nonpolar 0.41 - 1.7 Covalent polar >1.7 Ionic Fig 9-16 shows bond character

69

70

Polar and Non-Polar

Like Dissolves Likes!

Polar solutions dissolve polar compounds.

Same for Non-Polar

71

Electronegativity Chart

72

73

74

How to show a bond is polar Isn’t a whole charge just a partial charge means a partially positive means a partially negative

The Cl pulls harder on the electrons The electrons spend more time near the Cl

H Cl

75

Polar Molecules

Molecules with Charged Ends

76

Polar Molecules Polar Molecules have a positive and a

negative end This requires two things to be true The molecule must contain polar bonds

This can be determined from differences in electronegativity.

Symmetry can not cancel out the effects of the polar bonds.

Must determine geometry to check polarity.

77

Is it polar? HF H2O

NH3

CCl4

CO2

78

Bond Dissociation Energy The energy required to break a CH

bond is C - H + 393 kJ C + H We get the Bond dissociation energy

back when the atoms are put back together

If we add up the BDE of the reactants and subtract the BDE of the products we can determine the energy of the reaction (H)

79

Find the Energy Change for the Reaction

CH4 + 2O2 CO2 + 2H2O

For the reactants we need to break 4 C-H bonds at 393 kJ/mol and 2 O=O bonds at 495 kJ/mol= 2562 kJ/mol

For the products we form 2 C=O at 736 kJ/mol and 4 O-H bonds at 464 kJ/mol = 3328 kJ/mol

reactants - products = 2562-3328 = -766kJ

80

Endothermic and Exothermic Endothermic reactions occur greater

energy needed to break old bonds then is released forming new bonds.

Exothermic reactions occur when more energy is released making new bonds than is needed to break the old bonds.

81

Intermolecular Forces

What Holds Molecules to Each Other?

82

Intermolecular Forces “Inter” means between or among So these are forces between molecules These forces of attraction make solid

and liquid molecular compounds possible.

Intermolecular forces are also called van der Waal’s forces. We will talk about three types.

With nonpolar molecules we have dispersion forces or London forces. These are temporary and weak forces.

83

Dispersion Forces Compressing a gas (colliding atoms) causes

a temporary shift in the electron cloud and slight polarity.

Also depends on the # of electrons More electrons mean stronger forces Bigger molecules more electrons greater

attraction• Fluorine is a gas• Bromine is a liquid• Iodine is a solid

84

Dipole Interactions A stronger type are the dipole-dipole

forces between polar molecules Occur when polar molecules are

attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely

hooked like in ionic solids.

85

Dipole Interactions Occur when polar molecules are

attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely

hooked like in ionic solids.

H Cl

H Cl

86

Dipole Interactions

87

Hydrogen Bonding A special dipole-dipole attraction force

is caused by hydrogen being bonded to F, O, or N.

F, O, and N are very electronegative so it is a very strong dipole.

The hydrogen partially share with the lone pair in the molecule next to it.

The strongest of the intermolecular forces. Movie

88

Hydrogen Bonding

HH

O+ -

+

H HO+-

+

89

Hydrogen bonding

HH

O H HO

HH

O

H

H

OH

HO

H

HO HH

O

Documents

1 Chapter 9 Covalent Bonding Molecular Compounds