CHAPTER 8 Bonding and Molecular Structure. 2 Introduction Bonds: Attractive forces that hold atoms...

CHAPTER 8

Bonding and Molecular Structure

2

Introduction

• Bonds: Attractive forces that hold atoms together in compounds

• Valence Electrons: the outermost electrons– -These e- are involved in bonding

3

Electrons are divided between core and valence electrons

B 1sB 1s22 2s 2s22 2p 2p11

Core = [He]Core = [He] , valence = 2s , valence = 2s22 2p 2p11

Br [Ar] 3dBr [Ar] 3d1010 4s 4s22 4p 4p55

Core = [Ar] 3dCore = [Ar] 3d1010 , valence = 4s , valence = 4s22 4p 4p55

Valence Electrons

4

-The number of valence electrons of a main group atom is the Group number

-For Groups IA-IVA, number of bonding -For Groups IA-IVA, number of bonding (unpaired) (unpaired) electrons is equal to the group numberelectrons is equal to the group number-For Groups VA -VIIA, number of bonding (unpaired) -For Groups VA -VIIA, number of bonding (unpaired)

electrons is equal to 8 - group numberelectrons is equal to 8 - group number

Valence Electrons

5

-Except for H (and sometimes atoms of -Except for H (and sometimes atoms of the 3the 3rdrd group and higher) group and higher)

-The total number of valence electrons -The total number of valence electrons around a given atom in a molecule will around a given atom in a molecule will be eight:be eight:

OCTET RULE- (with the exception of hydrogen) atoms in molecules prefer to be surrounded by 8 electrons (or have 4 bonds = 8 electrons)

Valence Electrons

6

Lewis Dot Formulas of Atoms

H He

Li Be B C N O F Ne

IA IIA IIIA IVA VA VIA VIIA VIIIA

7

Ionic Bonding

An ion is an atom or a group of atoms possessing a net electrical charge

• -cations: positive (+) ions• These atoms have lost 1 or more

electrons

1. -anions: negative (-) ions• These atoms have gained 1 or more

electrons

8

Formation of Ionic Compounds• Monatomic ions consist of one atom

– Examples:• Na+, Ca2+, Al3+ - cations• Cl-, O2-, N3- -anions

• Polyatomic ions contain more than one atom

Examples:• NH4

+ - cation• NO2

-,CO32-, SO4

2- - anions

9

Formation of Ionic Compounds• General trend:

– metals become isoelectronic with the preceding noble gas electron configuration

– nonmetals become isoelectronic with the following noble gas electron configuration

10

Formation of Ionic Compounds

• Reaction of Group IA Metals with Group VIIA Nonmetals

point melting

C842an with gas solid

solid whiteyellow silver

LiF 2 F Li 2

nometal 17 -G metal 1-G

o

(s)2(g)(s)

11


• 1s 2s 2p Li F

These atoms form ions with these configurations.

Li+ same configuration as [He] F- same configuration as

[Ne]Li + F.

..

.... .

Li+

F[ ]...... ..

12

Formation of Ionic Compounds• In general:

the reaction of IA metals and VIIA nonmetals:

2 M(s) + X2 2 MX(s)

– where M is the metals Li to Cs– and X is the nonmetals F to I

Electronically it looks like: ns np ns npM M+ __ __ __ __X X-

13


reaction of IIA metals with VIIA nonmetals:

Be(s) + F2(g) BeF2(g)

14


The valence electrons in these two elements react like:

2s 2p 2s 2p Be [He] Be2+ __ __ __

__F [He] F-

Lewis dot structure representation:

15


The remainder of the IIA metals and VIIA nonmetals react similarly:

M(s) + X2 MX

2

M can be any of the metals Be to BaX can be any of the nonmetals F to I

16

Formation of Ionic CompoundsFor the reaction of IA metals with

VIA nonmetals:

-2s22(g)(s) O Li2O Li4

Draw the valence electronic configurations for Li, O, and their appropriate ions

17

Formation of Ionic Compounds• Draw the electronic configurations

for Li, O, and their appropriate ionsYou do it!You do it!

2s 2p 2s 2p Li [He] Li1+

O [He] O2-

Draw the Lewis dot formula representation of this reaction

18


Simple Binary Ionic Compounds Table• Reacting Groups General Formula Example

IA + VIIA MX NaF IIA + VIIA MX2 BaCl2IIIA + VIIA MX3 AlF3

IA + VIA M2X Na2O

IIA + VIA MX BaOIIIA + VIA M2X3 Al2S3

19


• Reacting Groups General Formula Example

IA + VA M3X Na3N

IIA + VA M3X2 Mg3P2

IIIA + VA MX AlN

-H forms ionic compounds when bound to metals (IA and IIA metals

For example: LiH, KH, CaH2, and BaH2

-When H is bound to nonmetals, the compounds are covalent in nature

20

Formation of Covalent Bonds• -potential energy of an H2 molecule as

a function of the distance between the two H atoms

21

Covalent Bonding• Atoms share electrons

• If the atoms share: • 2 electrons a single covalent bond is

formed• 4 electrons - a double covalent bond• 6 electrons - a triple covalent bond

Atoms have a lower potential energy when bound…this is a more favorable situation (why?)

22

Writing Lewis Formulas:

• 1. Add the number of valence electrons for all the atoms that are present in the molecule

• 2. Add or subtract electrons based on the molecule’s (or ion’s) charge

• 3. Identify the central atom and draw a skeletal structure: – -the one that requires the most e- to complete octet – -the less electronegative

• 4. Place a bond between each atom (1 bond = 2 e-)

• 5. Fill in octet of outer atoms first• 6. Finish by completing the octet of central atom

– – if you run out of e- then multiple bonds must be created between the central atom and atoms bound to it

23

Writing Lewis Formulas

octet rule: representative elements usually attain stable noble gas electron configurations (8 valence e-) in most compounds

You must distinguish the difference between: – -bonding electrons and nonbonding

electrons -shared (paired) and unshared

(unpaired) electrons

24

Formation of Covalent Bonds

• Lewis dot structures:

• 1. H2 molecule formation:

2. HCl molecule formation:

25

Lewis Structures• Homonuclear diatomic molecules

– 1. Two atoms of the same element, H2:

H HorH H..

2. Fluorine, F2:

3. Nitrogen, N2:

26

Lewis Structures

heteronuclear diatomic molecules

1. hydrogen fluoride, HF

2. hydrogen chloride, HCl

3. hydrogen bromide, HBr

or ··H Cl··

··H Cl..

······

27

Lewis Structures

• Water, H2O

•Ammonia molecule , NH3

28

Lewis Structures

• Polyatomic ions:

• ammonium ion NH4+

Notice that the N-atom in this molecule has eight electrons around them (H does not)

29

Writing Lewis Formulas

• Sulfite ion, SO32-.

30

Double and even triple bonds are commonly observed for C, N, P, O, and S

••O OC

•• ••

••

••O OC

•• ••

••

HH22COCO

SOSO33

CC22FF44

31

Lewis Structures

• Example: Write Lewis dot and dash formulas for sulfur trioxide, SO3

32

Resonance

• There are three possible structures for SO3:

O S

O

O·· ······ ··

······

OS

O

O·· ···· ·· ··

··

······

O S

O

O·· ····

·· ··

····

-Two or more Lewis formulas are necessary to show the bonding in a molecule -use equivalent resonance structures to show the molecule’s structure

-Double-headed arrows are used to indicate resonance formulas

33

Resonance

Resonance is a flawed method of representing molecules– -There are no single or double bonds in

SO3

SO O

O

34

Sulfur Dioxide, SOSulfur Dioxide, SO22Sulfur Dioxide, SOSulfur Dioxide, SO22

1. Central atom =1. Central atom =

2. Valence electrons = ___2. Valence electrons = ___

or ___ pairsor ___ pairs

4. Form double bond so 4. Form double bond so that S has an octet — but that S has an octet — but note that there are two note that there are two ways of doing this.ways of doing this.

3. Write the Lewis 3. Write the Lewis structurestructure

35

Limitations of the Octet Rule

• There are some molecules that violate the octet rule:

1. - Be2. - Group IIIA3. -Odd number of total electrons.4. -Central element must have a share of more

than 8 valence electrons to accommodate all of the substituents. (i.e. S and P)

36


• Example: Write Lewis formula for BBr3.

37

Sulfur Tetrafluoride, SFSulfur Tetrafluoride, SF44Sulfur Tetrafluoride, SFSulfur Tetrafluoride, SF44

Central atom = Central atom =

Valence electrons = ___ or ___ Valence electrons = ___ or ___ pairs.pairs.

Form sigma bonds and Form sigma bonds and distribute electron pairs.distribute electron pairs.

F

••

••

••

F

F

S••

••

••

••

•• F

••

••

••

••

••

F

••

••

••

F

F

S••

••

••

••

•• F

••

••

••

••

•• 5 pairs around the S 5 pairs around the S atom. A common atom. A common occurrence outside the occurrence outside the 2nd period. 2nd period.

5 pairs around the S 5 pairs around the S atom. A common atom. A common occurrence outside the occurrence outside the 2nd period. 2nd period.

38


• Example: Write dot structures for AsF

5.

39

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

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

• Formal charge = Group number – 1/2 (# of bonding electrons) - (# of Lone electrons)

• • = Group number – (# of bonds) • – (# of Lone electrons)

Formal Atomic Charges

40

Formal Charge

CO CO22

Formal Charge

CO CO22

. .

. .

. .

. .

41

Thiocyanate Ion, SCN-

Thiocyanate Ion, SCN-

••

•

••S NC

•••

••

•

••S NC

•••

•••

••S NC

•••

Which is the most stable resonance form?Which is the most stable resonance form?

Formal Charge

42

Theories of Covalent Bonding

• Valence Shell Electron Pair Repulsion Theory– Commonly designated as VSEPR– Principal originator

• R. J. Gillespie in the 1950’s

• Valence Bond Theory (Chapter 9)– Involves the use of hybridized atomic

orbitals– Principal originator

• L. Pauling in the 1930’s & 40’s

43

VSEPR Theory

electron densities around the central atom are arranged as far apart as possible to minimize repulsions (why?)

• Five basic molecular shapes:• Linear, trigonal planar, tetrahedral,

trigonal bipyramidal, octahedral

44

VSEPR Theory

1. Two regions of high electron density around the central atom.

45

VSEPR Theory2. Three regions of high electron density around the

central atom.

46

VSEPR Theory3. Four regions of high electron density around

the central atom.

47

VSEPR Theory

4. Five regions of high electron density around the central atom.

48

VSEPR Theory

5. Six regions of high electron density around the central atom.

49

VSEPR Theory

1. Electronic geometry(family): locations of regions of electron density around the central atom(s)

2.2. Molecular geometry:Molecular geometry: arrangement of atoms around the central atom(s)

Electron pairs are not used in the molecular geometry determination

50

VSEPR Theory

Lone pairs (unshared pairs) of electrons require more volume than shared pairs– -there is an ordering of repulsions of lone

electrons around central atom

Criteria for the ordering of the repulsions:1. Lone pair to lone pair is the strongest repulsion.2. Lone pair to bonding pair is intermediate repulsion.3. Bonding pair to bonding pair is weakest repulsion.

51

Molecular Shapes and Molecular Shapes and BondingBonding

• Symbolism:A = central atomB = bonding pairs around central atomU = lone pairs around central atom

• For example:AB3U designates that there are 3 bonding

pairs and 1 lone pair around the central atom

52

Linear Electronic Geometry: AB2

Some examples of molecules with this geometry:BeCl

2, BeBr

2, BeI

2, HgCl

2, CdCl

2

53

Trigonal Planar Electronic Geometry: AB3

Some examples of molecules with this geometry are:

BF3, BCl3

54

Tetrahedral Electronic Geometry: AB

4

Some examples of molecules with this geometry are: CH

4, CF

4, CCl

4,

SiH

4,

SiF

4

55

VSEPR Theory

• An example of a molecule that has the same electronic and molecular geometries is methane (CH4)– -Electronic and molecular geometries are

tetrahedral

H

C

HHH

56

Tetrahedral Electronic Geometry: AB4

57

Tetrahedral Electronic Geometry: AB3U

Some examples of molecules with this geometry are: NH3, NF3, PH3, PCl3, AsH3

– -trigonal pyramidal-electronic and molecular geometries are different.

. .

107.5°

. .

58

• Some examples of molecules with this geometry are: H2O, OF2, H2S– -bent

-electronic and molecular geometries are different

Tetrahedral Electronic Geometry: AB2U2

104.5°

59

VSEPR Theory• An example of a molecule that has

different electronic and molecular geometries is water (H2O)– -Electronic geometry is tetrahedral– -Molecular geometry is bent or angular

H

C

HHH

60

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and

AB2U3

Some examples of molecules with this geometry are: PF5, AsF5, PCl5

axial

axial

equatorial

61

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

If lone pairs are incorporated into the trigonal bipyramidal structure, there are three possible new shapes:1. One lone pair - Seesaw shape2. Two lone pairs - T-shape3. Three lone pairs – linear

The lone pairs occupy equatorial positions first: -they are 120o from each other

-90o from the axial positions– Results in decreased repulsions compared to

lone pair in axial positionaxial

axial

equatorial

62


AB2U3

• AB4U molecules have:

1. trigonal bipyramid electronic geometry

2. seesaw shaped molecular geometry 3. polar

• One example of an AB4U molecule is

SF4

63


AB2U3

H

C

HHH

64

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

• AB3U2 molecules have: 1. 1. trigonal bipyramid electronic

geometry 2. T-shaped molecular geometry 3. polar

• One example of an AB3U2 molecule is

IF3

65


AB2U3

H

C

HHH

66


AB2U3

• AB2U3 molecules have:

1.trigonal bipyramid electronic geometry

2.linear molecular geometry 3.nonpolar


BrF2-

67


AB2U3

H

C

HHH

68

Octahedral Electronic Geometry: AB6, AB5U, and AB4U2

• Some examples of molecules with this geometry are: SF6, SeF6, SCl6, etc.

69


If lone pairs are incorporated into the octahedral structure, there are two possible new shapes:1. One lone pair - square pyramidal2. Two lone pairs - square planar

The lone pairs occupy any position because they are all 90o from all bonds positions:– -Additional lone pairs occupy the position

180º from the first set of lone pairs– -This results in decreased repulsions compared

to lone pairs in the other positions

70


• AB5U molecules have:

1.octahedral electronic geometry 2.Square pyramidal molecular

geometry 3.polar.

• One example of an AB4U molecule is

IF5

71


• AB4U2 molecules have:

1.octahedral electronic geometry 2.square planar molecular geometry 3.and are nonpolar.


XeF4

72

Polarity and Polarity and ElectronegativityElectronegativity

Figure 8.11Figure 8.11

73

Dipole Moments

• For example, HF and HI:

units Debye 0.38 units Debye 1.91

I-H F-H

--

74

Dipole Moments

some “nonpolar molecules” that have polar bonds

Two conditions to be polar:1. 1. There must be at least one polar bond

present or one lone pair of electrons2. 2. the molecule must be nonsymmetric

Examples: water, CF4, CO2, NH3, NH4+

75

Polar Molecules

• Molecular geometry affects molecular polarity– -they either cancel or reinforce each

other

A B A

linear molecule nonpolar

A B A

angular molecule

polar

76

Polar and Nonpolar Bonds

• Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds– -Nonpolar covalent bonds have a

symmetrical charge distribution (electron distribution)

N N········ ·· N N·· ··or H HorH H..

77


• Polar covalent bonds: electrons are not shared equally

• -they have different electronegativities

H FElectronegativities: 2.1 4.0

Difference = 1.9 very polar bond

78


• Compare HF to HI:

H IElectronegativities: 2.1 2.5

Difference = 0.4 slightly polar bond

more complicated geometries exist…

79

• Three molecules with polar covalent bonds:

• -Each bond has one atom with a slight negative charge (-)

• -another with a slight positive charge (+)

Bond Polarity

80

Polar or Nonpolar?Polar or Nonpolar?

AB3 molecules: BF3, Cl2CO, and NH3

81

Polar or Nonpolar?Polar or Nonpolar?CO2 and H2O

Which one is polar?

82

CHCH44 … CCl … CCl44Polar or Not?Polar or Not?

• Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”

83

Compounds Containing Double Bonds

• Ethene or ethylene, C2H4, is the simplest organic compound containing a double bond.– -has a double bond to obey octet ruleLewis Dot Formula

CC

H

HH

H

C CH

H

H

H····

·· ·· ··

··or

84

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

Free rotation Free rotation around C–C single around C–C single bondbond

No rotation No rotation around C=C around C=C double bonddouble bond

and

Double Bonds

85

Bond Order # of bonds between similar pairs of

atoms

Bond Order # of bonds between similar pairs of

atoms

Double bondDouble bondDouble bondDouble bondSingle bondSingle bondSingle bondSingle bond

Triple Triple bondbondTriple Triple bondbond

AcrylonitrileAcrylonitrileAcrylonitrileAcrylonitrile

86

Consider NO2-:

that typeof bound atoms of # Total

type-one of bonds of # Total = orderBond

that typeof bound atoms of # Total

type-one of bonds of # Total = orderBond

The N—O bond order = 1.5The N—O bond order = 1.5The N—O bond order = 1.5The N—O bond order = 1.5

O O O O

N••

••••

••••

••••••••••

••••

••N

Bond Order

87

Bond order is proportional to two important bond properties:

(a) bond strength(b) bond length

745 kJ745 kJ

414 kJ414 kJ 110 pm110 pm

123 pm123 pm

Bond Order

88

the distance between the nuclei of two bonded atoms

Bond Length

89

Bond LengthBond LengthBond length depends on size of bonded atoms

H—FH—F

H—ClH—Cl

H—IH—I

Bond distances measured Bond distances measured in Angstrom units where 1 in Angstrom units where 1 ÅÅ = 10 = 10-2 -2 pm.pm.


90

Bond length depends on bond order



91

• Measure of the energy required to break a bond

• See Table 9.10• BOND STRENGTH (kJ/mol) H—H 436 KJ C—C 346 KJ C=C 602 KJ CC 835 KJ NN 945 KJ

The GREATER the number of bonds (bond order) the The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the HIGHER the bond strength and the SHORTER the bond.bond.

The GREATER the number of bonds (bond order) the The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the HIGHER the bond strength and the SHORTER the bond.bond.

Bond Strength

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CHAPTER 8 Bonding and Molecular Structure. 2 Introduction Bonds: Attractive forces that hold atoms...

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