chapter 9 - georgia southern university-armstrong campus€¦ · chapter 9 • structure and...

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CHAPTER 9

• Structure and Molecular Bonding

2

Introduction

• Bonds - Attractive forces that hold atoms together in compounds

• Valence Electrons - The electrons involved in bonding are in the outermost (valence) shell.

3

Valence ElectronsValence ElectronsValence Electrons

Electrons are divided between Electrons are divided between corecore and and

valence electronsvalence electronsB 1sB 1s22 2s2s22 2p2p11

Core = [He]Core = [He] , valence = 2s, valence = 2s22 2p2p11

Br [Br [ArAr] 3d] 3d1010 4s4s22 4p4p55

Core = [Core = [ArAr] 3d] 3d1010 , valence = 4s, valence = 4s22 4p4p55

4

Rules of the GameRules of the GameRules of the GameNo. of valence electrons of a No. of valence electrons of a main group atom = Group main group atom = Group numbernumber

••For Groups 1AFor Groups 1A--4A, no. of bond pairs = 4A, no. of bond pairs =

group number.group number.

•• For Groups 5A For Groups 5A --7A, BP7A, BP’’s = 8 s = 8 -- GrpGrp. No.. No.

5

Rules of the GameRules of the Game••No. of valence electrons of an atom = Group No. of valence electrons of an atom = Group

numbernumber

••For Groups 1AFor Groups 1A--4A, no. of bond pairs = group 4A, no. of bond pairs = group

numbernumber

•• For Groups 5A For Groups 5A --7A, BP7A, BP’’s = 8 s = 8 -- GrpGrp. No. . No.

••Except for H (and sometimes atoms of Except for H (and sometimes atoms of

3rd and higher periods), 3rd and higher periods),

BPBP’’s + LPs + LP’’s = 4s = 4

This observation is called theThis observation is called the

OCTET RULEOCTET RULE

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Lewis Dot Formulas of Atoms

Li Be B C N O F Ne.... .. ..

..HeH

.

.. . .

.. ..

..

...

..

.. ..

.

...

.

. ...

7

Ionic Bonding

Formation of Ionic Compounds• An ion is an atom or a group of atoms

possessing a net electrical charge.

1. positive (+) ions or cations• These atoms have lost 1 or more electrons.

2. negative (-) ions or anions• 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.– NH4

+ - cation

– NO2-,CO3

2-, SO42- - anions

9


• 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) →+

10


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[ ]........

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• General trend

• Cations become isoelectronic with the preceding noble gas.

• Anions become isoelectronic with the following noble gas.

12


• In general for the reaction of 1 metals and 17 nonmetals, the reaction equation is:

2 M(s) + X2→ 2 M+ X-(s)

– where M is the metals Li to Cs

– and X is the nonmetals F to I.

Electronically this is occurring.

ns np ns np

M ↑↑↑↑ → M+

X ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↑↑↑ → X- ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓

13


• Next we examine the reaction of 2 metals with 17 nonmetals.

• One example is the reaction of Be and F2.

Be(s) + F2(g) →BeF2(g)

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• The valence electrons in these two elements are reacting in this fashion.

2s 2p 2s 2p

Be [He] ↑↓↑↓↑↓↑↓ → Be2+

F [He] ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↑↑↑ → F- ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓

15


..

..

... F..

. F....

Be.. Be

2+

2 F

........

• The remainder of the 2 metals and 17nonmetals react similarly.

• Symbolically this can be represented as:

M(s) + X2→ M2+ X

2

-

M can be any of the metals Be to Ba.

X can be any of the nonmetals F to Cl.

16


• For the reaction of 1 metals with 16 nonmetals, a good example is the reaction of lithium with oxygen.

• The reaction equation is:

( )-2

s22(g)(s) O Li2O Li4 +→+

17


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

You do it!You do it!

2s 2p 2s 2p

Li [He] ↑↑↑↑ → Li1+

O [He] ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↑↑↑ ↑↑↑↑ → O2- ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓↑↓ Draw the Lewis dot formula representation of

this reaction.

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Simple Binary Ionic Compounds Table• Reacting Groups Compound General Formula

Example

1 + 17 MX NaF

2 + 17 MX2 BaCl23 + 17 MX3 AlF31 + 16 M2X Na2O

2 + 16 MX BaO

3 + 16 M2X3 Al2S3

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• Reacting Groups Compound General Formula

Example

1 + 15 M3X Na3N

2 + 15 M3X2 Mg3P23 + 15 MX AlN

H, a nonmetal, forms ionic compounds with 1 and 2 metals for example, LiH, KH, CaH2, and BaH2.

Other hydrogen compounds are covalent.

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• Coulomb’s Law – inverse square law

( )( )

ions ofcenter between distance d

ionson charge of magnitude q

ionsbetween attraction of force F

where

d

qqF

2

=

=

=

∝−+

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• Force - Small ions with high ionic charges>> large ions with small ionic charges

-1-22-2

3

3

2 ClKOCaOAl +++ >>

22

Covalent Bonding• Atoms share electrons.

• If the atoms share 2 electrons a singlecovalent bond is formed.

• 4 electrons - a double bond.

• 6 electrons - a triple bond.

• The atoms have a lower potential energy when bound.

23

Formation of Covalent Bonds• This figure shows the potential energy of an H2 molecule as a function of the distance between the two H atoms.

24

Writing Lewis Formulas:

• 1. Sum the number of valence electrons for atoms present.

• 2. Add or subtract electrons for the charge.

• 3. Identify the central atom (one that requires more e- to complete octet – less e-neg if in same group) and draw a skeletal structure.

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

• 5. Fill in octet of outer atoms.

• 6. Complete octet of central atom – if deficient make multiple bonds.

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Formation of Covalent Bonds

• Use Lewis dot formulas

• 1. H molecule formation representation.

+H. H . H H.. or H2

H Cl H Cl+.....

.

...

..

..

... or HCl

2. HCl molecule formation

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Writing Lewis Formulas:The Octet Rule

• Lewis octet rule - representative elements usually attain stable noble gas electron configurations in most of their compounds.

• Distinguish between bonding (or shared) electrons and nonbonding (or unshared or lone pairs) of electrons.

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Lewis Formulas for Molecules and Polyatomic Ions

• Homonuclear diatomic molecules.

– 1. Two atoms of the same element.

1. Hydrogen molecule, H2.

H HorH H..

F F....

.

...

..

.. ..F F..

.

.

.. ..

.. ..or

N N······

··

·· N N·

···or

2. Fluorine, F2.

3. Nitrogen, N2.

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• Next, look at heteronuclear diatomic molecules.– 1. hydrogen fluoride, HF

or ··H F

··

··H F..····

··

or ··H Cl

··

··H Cl..····

··

or ··H Br

··

··H Br..····

··

2. hydrogen chloride, HCl

3. hydrogen bromide, HBr

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• Water, H2O

•Ammonia molecule , NH3

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• Ammonia molecule , NH3

H

H

N········ H

H

H

O··

····

··•Water, H2O

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• Polyatomic ions.

• One example is the ammonium ion , NH4+.

H

H

N··

······ H

H +

•Notice that the atoms other than H in these

molecules have eight electrons around them.

32


• Example: Write Lewis dot and dash formulas for the sulfite ion, SO3

2-.

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• Sulfite ion, SO32-.

O S O

O··

··

····

····

··

··

····

··

····

2-O S

O

O·· ·· ··

··

···· ··

····

··

2-or

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Double and Double and

even triple even triple

bonds are bonds are

commonly commonly

observed for C, observed for C,

N, P, O, and SN, P, O, and S

•

•O OC•• ••

•

•

HH22COCO

SOSO33

CC22FF44

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Lewis Structures

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

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Lewis Structures

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

orO S O

O··

··

····

····

······

··

··

·· O S

O

O··

··

···· ··

····

··

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Resonance

• There are three possible structures for SO3.– .

O S

O

O··

··

···· ··

····

··

OS

O

O·· ·

····· ··

··

··

··

··

O S

O

O····

···· ··

····

oTwo or more Lewis formulas are necessary to show the

bonding in a molecule, we must use equivalent resonance

structures to show the molecule’s structure.

oDouble-headed arrows are used to indicate resonance formulas.

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Resonance

• Resonance is a flawed method of representing molecules.

– There are no single or double bonds in SO

3.

SO O

O

39

Sulfur Dioxide, SO2Sulfur Dioxide, SOSulfur Dioxide, SO221. Central atom = S1. Central atom = S

2. Valence electrons = 18 or 9 2. Valence electrons = 18 or 9

pairspairs •

•O OS••

••

••

••••

•

•

bring inleft pair

OR bring inright pair

3. Form double bond so that S has an octet 3. Form double bond so that S has an octet ——

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

•

• OS••

••

••

••••

•

•O

40

Sulfur Dioxide, SO2Sulfur Dioxide, SOSulfur Dioxide, SO22

This leads to the following resonance This leads to the following resonance

structures.structures.

•

•O OS••

••

••

••••

•

•

bring inleft pair

OR bring inright pair

•

•O OS••

••

••

••

•

••

•O OS••

••

••

••

•

•

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Writing Lewis Formulas:Limitations of the Octet Rule

• There are some molecules that violate the octet rule.– For these molecules the N - A = S rule does not apply:

1. - Be.

2. - Group - 13.

3. -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. some S, P)

42


• Example: Write Lewis formula for BBr3.

B··. Br·

···

··

.

BBr Br

Br

····

····

······

··

······

··

Br B

Br

Br····

··

·· ··

··

····

··

or

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Sulfur Tetrafluoride, SF4Sulfur Sulfur TetrafluorideTetrafluoride, SF, SF44•• Central atom = Central atom =

•• Valence electrons = ___ or ___ Valence electrons = ___ or ___

pairs.pairs.

•• Form sigma bonds and distribute Form sigma bonds and distribute

electron pairs.electron pairs.

F

••

•

•

••

F

F

S••

••

•

•

•

•

•

• F••

••

••

••

•• 5 pairs around the S atom. A common

occurrence outside the 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.

44


• Example: Write dot and dash formulas for AsF

5.

45

Formal Atom ChargesFormal Atom ChargesFormal 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 Formal charge = Group number = Group number –– 1/2 (no. of bonding electrons) 1/2 (no. of bonding electrons) -- (no. of LP electrons)(no. of LP electrons)

46

Calculated Partial Charges in CO2

CalculatedCalculated Partial Charges in Partial Charges in COCO22

Yellow = negative & red = positive

Relative size = relative charge

Yellow = negativeYellow = negative && red = positivered = positive

Relative size = relative chargeRelative size = relative charge

47

Thiocyanate Ion, SCN-ThiocyanateThiocyanate Ion, SCNIon, SCN--

6 - (1/2)(2) - 6 = -1 5 - (1/2)(6) - 2 = 0

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

••

•

•

• S NC•

•

•

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Thiocyanate Ion, SCN-ThiocyanateThiocyanate Ion, SCNIon, SCN--

••

•

•

• S NC•

•

•

••

•

•

•S NC•

•

•

•••

•

• S NC•

•

•

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

49

Calculated Partial Charges Calculated Partial Charges in SCNin SCN--

All atoms negative, but most on the S

All atoms negative, but All atoms negative, but

most on the Smost on the S••

•

•

• S NC•

•

•

50

Dipole Moments

• Asymmetric charge distribution

• The dipole moment has the symbol µ.

• µ is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q.

51

Dipole Moments

• For example, HF and HI:

units Debye 0.38 units Debye 1.91

I-H F-H

-- δδδδ ++

aa

52

Dipole Moments

• There are some nonpolar molecules that have polar bonds.

• There are two conditions that must be true for a molecule to be polar.

1. There must be at least one polar bond present or one lone pair of electrons.

2. The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

3. Examples (water, CF4, CO2, NH3, NH4+)

53

Polar Molecules: The Influence of Molecular Geometry

• Molecular geometry affects molecular polarity.

– Due to the effect of the bond dipoles and how they either cancel or reinforce each other.

A B A

linear molecule

nonpolar

A B

A

angular molecule

polar

54

Polar Molecules: The Influence of Molecular Geometry

• Polar Molecules must meet two requirements:

1. One polar bond or one lone pair of electrons on central atom.

2. Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel.

55

Polar and Nonpolar Covalent Bonds

• Covalent bonds in which the electrons are shared equally are designated as nonpolarcovalent bonds.

– Nonpolar covalent bonds have a symmetrical charge distribution.

N N······

··

·· N N·

···or H HorH H

.

.

56


• Polar covalent bonds - electrons are not shared equally -different electronegativities.

bondpolar very 1.9 Difference

4.0 2.1 ativitiesElectroneg

F H

1.9

=

43421

57


• Compare HF to HI.

bondpolar slightly 0.4 Difference

2.5 2.1 ativitiesElectroneg

I H

0.4

=

43421

58

Two Simple 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 10)

– Involves the use of hybridized atomic orbitals

– Principal originator

• L. Pauling in the 1930’s & 40’s

59

VSEPR Theory

• VSEPR - electron density around the central atom are arranged as far apart as possible to minimize repulsions.

• Five basic molecular shapes

60

VSEPR Theory

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

61

VSEPR Theory

2 Three regions of high electron density around the central atom.

62

VSEPR Theory

3 Four regions of high electron density around the central atom.

63

VSEPR Theory

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

64

VSEPR Theory

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

65

VSEPR Theory

1.1. Electronic geometryElectronic geometry - locations of regions

of electron density around the central atom(s).

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

Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used.

66

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

HH

H

67

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

HH

H

68

VSEPR Theory

• Lone pairs of electrons (unshared pairs) require more volume than shared pairs.

– Consequently, there is an ordering of repulsions of electrons around central atom.

• Criteria for the ordering of the repulsions:

69

VSEPR Theory

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.

• Mnemonic for repulsion strengths

lp/lp > lp/bp > bp/bp

• Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o.

70

Molecular Shapes and BondingMolecular Shapes and Bonding

• In the next sections we will use the following terminology:

A = central atom

B = bonding pairs around central atom

U = lone pairs around central atom

• For example:

AB3U designates that there are 3 bonding pairs and 1 lone pair around the central atom.

71

Linear Electronic Geometry:AB2

Species (No Lone Pairs of Electrons on A)

• Some examples of molecules with this geometry are: BeCl

2, BeBr

2, BeI

2, HgCl

2, CdCl

2

• All of these examples are linear, nonpolar molecules.

• Important exceptions occur when the two substituents are not the same!BeClBr or BeIBr will be linear and polar!

72

Trigonal Planar Electronic Geometry: AB3 Species (No Lone Pairs of

Electrons on A)

• Some examples of molecules with this geometry are:

BF3, BCl3

• All of these examples are trigonal planar, nonpolar molecules.

• Important exceptions occur when the three substituents are not the same!

BF2Cl or BCI2Br will be trigonal planar and polar!

73

Tetrahedral Electronic Geometry: AB4

Species (No Lone Pairs of Electrons on A)


CH4, CF

4, CCl

4, SiH

4, SiF

4

• All of these examples are tetrahedral, nonpolar molecules.

• Important exceptions occur when the four substituents are not the same!

CF3Cl or CH2CI2 will be tetrahedral and polar!

74

Tetrahedral Electronic Geometry: AB4Species (No Lone Pairs of Electrons on

A)

75

Tetrahedral Electronic Geometry: AB3U Species (One Lone Pair of

Electrons on A)

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

• These molecules are our first examples of central atoms with lone pairs of electrons.Thus, the electronic and molecular geometries are different.

All three substituents are the same but molecule is polarpolar.

• NH3 and NF3 are trigonal pyramidal, polar molecules.

76

Tetrahedral Electronic Geometry: AB2U2 Species (Two Lone Pairs of

Electrons on A)

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

• These molecules are our first examples of central atoms with two lone pairs of electrons.Thus, the electronic and molecular geometries are different.

Both substituents are the same but molecule is polarpolar.

• Molecules are angular, bent, or V-shaped and polar.

77

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

AB2U3� Some examples of molecules with this geometry are: PF5, AsF5, PCl5, etc.

• These molecules are examples of central atoms with five bonding pairs of electrons.The electronic and molecular geometries are the same.

• Molecules are trigonal bipyramidal and nonpolar when all five substituents are the same.If the five substituents are not the same polarpolarmolecules can result, AsF4Cl is an example.

78


AB2U3

• If lone pairs are incorporated into the trigonalbipyramidal structure, there are three possible new shapes.

1. One lone pair - Seesaw shape

2. Two lone pairs - T-shape

3. Three lone pairs – linear

• The lone pairs occupy equatorial positions because they are 120o from two bonding pairs and 90o from the other two bonding pairs.

– Results in decreased repulsions compared to lone pair in axial position.

79


AB2U3• AB4U molecules have:

1. trigonal bipyramid electronic geometry

2. seesaw shaped molecular geometry

3. and are polar

• One example of an AB4U molecule is

SF4

80


AB2U3

Molecular Geometry

H

C

HH

H

81


AB2U3

• AB3U2 molecules have:

1. trigonal bipyramid electronic geometry

2. T-shaped molecular geometry

3. and are polar

• One example of an AB3U2 molecule is

IF3

82


AB2U3

Molecular Geometry

H

C

HH

H

83


AB2U3• AB2U3 molecules have:

1.trigonal bipyramid electronic geometry

2.linear molecular geometry

3.and are nonpolar


84


AB2U3

Molecular Geometry

H

C

HH

H

85

Octahedral Electronic Geometry: AB6, AB5U, and AB4U2


SF6, SeF6, SCl6, etc.

• These molecules are examples of central atoms with six bonding pairs of electrons.

• Molecules are octahedraloctahedral and nonpolarnonpolarwhen all six substituents are the same.

If the six substituents are not the same polarpolarmolecules can result, SF5Cl is an example.

86


• If lone pairs are incorporated into the octahedral

structure, there are two possible new shapes.

1. One lone pair - square pyramidal

2. Two lone pairs - square planar

• The lone pairs occupy axial positions because

they are 90o from four bonding pairs.

– Results in decreased repulsions compared to lone pairs

in equatorial positions.

87


• AB5U molecules have:

1.octahedral electronic geometry

2.Square pyramidal molecular geometry

3.and are polar.

• One example of an AB4U molecule is

IF5

88


Molecular Geometry

H

C

HH

H

89


• AB4U2 molecules have:

1.octahedral electronic geometry

2.square planar molecular geometry

3.and are nonpolar.


XeF4

90


Molecular GeometryPolarity

H

C

HH

H

91

Compounds Containing Double Bonds

• Ethene or ethylene, C2H4, is the simplest organic compound containing a double bond.

• Compound must have a double bond to obey octet rule.

92

Compounds Containing Double Bonds

Lewis Dot Formula

CC

H

HH

H

C CH

H

H

H··

··

···· ·

·

··

o r

93

Bond PropertiesBond Properties• What is the effect of bonding and

structure on molecular properties?

Free rotation Free rotation

around Caround C––C single C single

bondbond

No rotation No rotation

around C=C around C=C

double bonddouble bond

94

Bond Order# of bonds between a pair of atoms

Bond OrderBond Order# of bonds between a pair of atoms# of bonds between a pair of atoms

Double bondDouble bondDouble bondSingle bondSingle bondSingle bond

Triple bond

Triple Triple

bondbond

AcrylonitrileAcrylonitrileAcrylonitrile

95

Bond OrderBond Order

Fractional bond orders occur in molecules

with resonance structures.

Consider NO2-

Bond order = 3 e - pairs in N — O bonds

2 N — O bonds

Bond order = Total # of e - pairs used for a type of bond

Total # of bonds of that type

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

O O O O

N••

••••

••

••••••••••

••

••••

••N

96

Bond OrderBond Order

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

97

Bond LengthBond Length

• Bond length is the distance between the nuclei of two bonded atoms.

98


Bond length depends on size of bonded atoms.

HH——FF

HH——ClCl

HH——II

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

Bond distances measured Bond distances measured

in Angstrom units where 1 in Angstrom units where 1

A = 10A = 10--2 2 pm.pm.

99

Bond length depends on bond order.

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

Bond distances measured Bond distances measured

in Angstrom units where 1 in Angstrom units where 1

A = 10A = 10--2 2 pm.pm.


100

• —measured by the energy req’d to break a bond. See

Table 9.10.

• BOND STRENGTH (kJ/mol)

H—H 436

C—C 346

C=C 602

C≡C 835

N≡N 945

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

101

102


• Measured as the energy req’d to break a

bond. See Table 9.10

103

• Measured as the energy req’d to break a

bond. See Table 9.10.

• BOND STRENGTH (kJ/mol)

H—H 436

C—C 346

C=C 602

C≡C 835

N≡N 945

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


104

Molecular PolarityMolecular Polarity

Why do ionic compounds dissolve in Why do ionic compounds dissolve in

water?water?

WaterWater

Boiling point Boiling point

= 100 = 100 ˚̊CC

MethaneMethane

Boiling point Boiling point

= = --161 161 ˚̊CC

Why do water and Why do water and

methane differ so methane differ so

much in their much in their

boiling points?boiling points?

105

Bond PolarityBond PolarityBond Polarity

•• Three molecules with Three molecules with

polar, covalent bonds.polar, covalent bonds.

•• Each bond has one Each bond has one

atom with a slight atom with a slight

negative charge (negative charge (--δδ) ) and and another with and and another with

a slight a slight positivepositivecharge (+ charge (+ δδ))

106

ElectronegativityElectronegativity, , χχχχ is a measure of the ability of an is a measure of the ability of an atom in a molecule to attract atom in a molecule to attract electrons to itself.electrons to itself.

Concept proposed by Linus Pauling 1901-1994Concept proposed by Concept proposed by LinusLinus Pauling 1901Pauling 1901--19941994

107

ElectronegativityElectronegativityFigure 9.14Figure 9.14

108

Molecular PolarityMolecular PolarityMolecular PolarityMolecules will be polar ifMolecules will be polar if

a)a) bonds are polarbonds are polar

ANDAND

b)b) the molecule is NOT the molecule is NOT ““symmetricsymmetric””

All above are NOT polarAll above are NOT polar

109

Polar or Polar or NonpolarNonpolar??

Compare CO2 and H2O. Which one is polar?

110

Polar or Polar or NonpolarNonpolar??

• Consider AB3 molecules: BF3, Cl2CO, and NH3.

111

Molecular Polarity, BF3Molecular Polarity, BFMolecular Polarity, BF33

F

F F

B

F

F F

B

B atom is B atom is

positive and positive and

F atoms are F atoms are

negative.negative.

B—F bonds in BF3 are polar. BB——F bonds in BFF bonds in BF33 are polar. are polar.

But molecule is symmetrical and NOT polar

But molecule is symmetrical But molecule is symmetrical

and and NOTNOT polarpolar

112

Molecular Polarity, HBF2Molecular Polarity, HBFMolecular Polarity, HBF22

B atom is B atom is

positive but positive but

H & F atoms H & F atoms

are negative.are negative.

H

F F

B

H

F F

B

B—F and B—H bonds in HBF2

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

BB——F and BF and B——H bonds in HBFH bonds in HBF22

are polar. But molecule is NOT are polar. But molecule is NOT

symmetrical and is polar.symmetrical and is polar.

113

Is Methane, CH4, Polar?Is Methane, CHIs Methane, CH44, Polar?, Polar?

Methane is symmetrical and is NOT Methane is symmetrical and is NOT

polar.polar.

114

Is CH3F Polar?Is CHIs CH33F Polar?F Polar?

C—F bond is very polar. Molecule is not symmetrical and so is polar.

CC——F bond is very polar. F bond is very polar.

Molecule is not symmetrical Molecule is not symmetrical

and so is polar.and so is polar.

115

CHCH44 …… CClCCl44Polar or Not?Polar or Not?

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

chapter 9 - georgia southern university-armstrong campus€¦ · chapter 9 • structure and...

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