1

Chapter 7

Chemical Bonding

2

Nature of Chemical Bond• Atoms are held together by electrostatic attraction between

positively charged nuclei and negatively charged electron clouds.

• Chemical Bond: a link between atoms that result from mutual attraction of their nuclei for electrons.

• Bond energy: the energy required to break a bond.

• Forces in substances:– Attractive: between electron clouds and respective bonding

nuclei (of the two atoms that bond)– Repulsive: between all the electron clouds in the bonding

atoms

3

Formation of Bonds• Bonding involves only the valence electrons (those in

the highest energy level).

• Use the periodic chart to guide determination of valence electrons

• WHEN BONDING OCCURS:

– Atoms attain an OCTET: a stable Noble Gas configuration.

– the resulting system is at the lowest possible potential energy level.

– The process of bonding is, therefore, exothermic: energy is being released. If the energy released is Large we get a strong bond; small ΔE bond is weak

4

Types of Bonds

• Ionic bond: formed by transfer of electrons from the valence energy level of one atom to another’s

• Covalent bond: formed when atoms share electrons.

• Metallic bond: ions of metals are surrounded by sea of electrons that bind all ions together.

5

Types of Bonds: Ionic

• Ionic Bond: results from electrostatic attraction between positive and negative ions.

• Produced by TRANSFER of electrons from valence energy level of one atom to another

• Occurs between metal and nonmetals , or polyatomic ions.

6

Types of Covalent Bonding

Types

• Nonpolar

• Polar

• Coordinate Covalent Bond

• Network covalent

7

Electronegativity

• A measure of how strongly the atoms attract electrons in a bond.

• The bigger the electronegativity difference the more polar the bond.

• 0.0 - 0.3 Covalent nonpolar• 0.3 - 1.0 Covalent moderately polar• 1.0 -1.7 Covalent polar• >1.7 Ionic• Use table 6.6, page 171 in your textbook

8

Pg 335

Table 8-1Representative Electronegativity

Differences

Covalent: = 0

Polar:0.3 < < 1.7

Ionic > 1.7

9

Nonpolar Covalent Bonding

• Nonpolar covalent bond: electrons are shared equally by atoms.– Electronegativity difference <0.3

• Examples:H-H O2

Cl-Cl Si-HN2 Ge-H

10

Polar Covalent Bonding

• Polar covalent bonds: the electrons are not shared equally between the bonding atoms. The more electronegative atom attracts the electrons.– Electronegativity difference between >0.3

and <1.7• One end is slightly positive, the other negative. • The charge distribution is indicated using small

delta and

11

A + indicates that the atom that doesn’t hog the electron density (less electronegative) has a partial positive charge. A - indicates the atom that hogs the electrons in the molecule (is more electronegative)

H F FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms

electron richregion

electron poorregion e- riche- poor

+ -

9.5

12

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

13

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

+-

14

H - F+ -

H - F+ -

H - F+ - H - F

+ -

H - F+ -

H - F+ -

H - F+ -

H - F+ -

- +

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Examples

• Determine the nature of the bond (nonpolar covalent, polar covalent, or ionic) for the following two elements:

1. Na and Cl 6. N and O

2. Si and F 7. S and F

3. S and Br 8. S and H

4. C and H 9. Te and I

5. Al and F 10. K and I

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Network Covalent Bond

• Occurs in compounds (elements) where all the atoms are bonded with covalent bonds .

• Examples: diamond, graphite

diamond

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AllotropesDifferent Structural forms of the same

element

Carbon BuckyballsAmorphous

Graphite

18

Coordinate Covalent Bond

• Additional Examples:– NH3 + H+ → NH4

+

– AlCl3 + Cl-1 → AlCl4-1

– Al(OH)3 + OH- → Al(OH)4-1

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Lewis Theory: An Overview

• Valence e- play a fundamental role in chemical bonding.

• e- transfer leads to ionic bonds.

• Sharing of e- leads to covalent bonds.

• e- are transferred or shared to give each atom a noble gas configuration – the octet.

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Lewis Symbols (Structures)

• A chemical symbol represents the nucleus and the core e-.

• Dots around the symbol represent valence e-.

Si•

••

•

N••

••

• P••

••

• As••

••

• Sb••

••

• Bi••

••

•

••Al••

• Se••

•••

Ar••

••

••I •••

••

••

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Ionic and Molecular Compounds

• Formation of sodium chloride (ionic):

• Formation of hydrogen chloride (covalent):

A metal and a nonmetal transfer electrons to form an ionic compound. Two nonmetals share electrons to form a molecular compound.

Na + Na+ [ ]Cl

Cl

H + Cl

Cl

H

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Ionic Compounds

Ionic compounds consist of

a lattice

of positive

and negative ions.

Lattice: three dimensional array of ions

NaCl:

23

Properties of Ions

• When will a stable bond be formed?• Octet Rule: both ions attain noble gas

configuration

• Example: NaCl versus Na+Cl-

Na: [Ne]3s1 Cl: [Ne]3s23p5

Na+: [Ne] Cl-: [Ne]3s23p6 = [Ar]

STRONG ELECTROSTATIC ATTRACTION

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Ionic Bonding Formation(1)• An atom with a low ionization energy reacts with

an atom with high electron affinity. • Between active metal (positive ion) and active

nonmetal (negative ion), or between polyatomic ions.

• Opposite charges hold the atoms together.

• Electronegativity difference – Electronegativity difference between atoms

>1.7 – 50% or more ionic character (check PT)

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Ionic Bond Formation(2)

• Both the positive and negative ions acquire noble gas configuration.

• The charge of the ion (oxidation state) is determined by the number of electrons lost or gained.

• The grater the electronegativity difference between the elements, the greater the ionic character of the bond. (>1.7 , 50% ionic).

• Only valence electrons participate in bond formation ( exception: transition elements).

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Lewis Structures for Ionic Compounds

• CHECK THE BLACKBOARD (Lewis structures and Orbital Notations)

• NaCl• MgCl2• Al2O3

• MgO• AlF3

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Properties of Ionic Compounds1. Crystalline structure.

A regular repeating arrangement of ions in the solid (lattice).

2. Ions are strongly bonded.3. Structure is rigid.4. High melting points

due to strong forces between ions.5. Oxidation states (charges) are determined by the #

of valence electrons (Group #)6. Conduct electricity when molten and in aqueous

solution7. Do not reflect light – therefore are white

(exceptions: transition elements ions)8. Solubility in water depends on lattice energy and

nature of solvent.

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Crystalline structure

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Ionic solids are brittle

+ - + -+- +-

+ - + -+- +-

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Ionic solids are brittle

+ - + -

+- +-+ - + -

+- +-

• Strong Repulsion breaks crystal apart.

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Metallic Bond

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Metallic Bond

• Metals consist of crystalline lattice in which positive ions (kernels) are arranged in fixed patterns.

• The valence electrons are free to move and they belong to the entire crystal.

• “Electron Sea” model

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Metallic Bond

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COVALENT BONDS

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What about covalent compounds?

The electrons in each atom are attracted to the nucleus of the other.

The electrons repel each other,The nuclei repel each other.

The atoms reach a distance with the lowest possible energy.

The distance between the atoms is the bond length.

36

When Atoms Combine to make Molecules

Fig 8-1

Atoms contain both positive and negative charges. When they come Together they arrange themselves so that the attractive forges of oppositeCharges is greater than the repulsive forces of like charges

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Fig 8-3 Pg 330

The interaction energy of a pair of hydrogen atoms varies with internuclear separation.

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How does H2 form?

• The nuclei repel

++

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How does H2 form?

• The nuclei repel

• But they are attracted to electrons

• They share the electrons

++

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Covalent Bond Formation

Covalent bond forms by overlap of orbitals.• Two types of bonds

Sigma bond: all single bonds are sigma bonds ( along the internuclear

plane)

Pi bond: in multiple bonds: the first one is sigma, all other bonds are pi.( above/below and front/back)

There areSingle bondsMultiple bonds (double and triple only)

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Two types of Bonds

• Sigma bonds(σ) from overlap of orbitals along the axis connecting the nuclei between the atoms

• Pi bond (): perpendicular overlap of p-orbitals above and below the axis connecting the atoms

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Sigma Bond Formation

• s-s overlap: Overlap of two s orbitals (s-s) : H-H bond

• p-p overlap: Overlap of two p orbitals (p-p) facing each other along the same axis (x-axis): Cl-Cl bond

• s-p overlap: overlap of s-orbital and p-orbital along the same axis. H-Cl bond

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Sigma Bond Formation: s-s Orbital Overlap

2 single atoms start overlap

Overlap complete

45

Sigma Bond Formation p-p overlap

• P-orbital overlap

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Sigma bond: p – p overlap

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Pi Bond

• Forms by a vertical overlap of two p orbitals (p-p vertical overlap)

• Exists only when there are multiple covalent bonds

• Example H2 -C=C-H2: the double

bonds contains one sigma and one pi bond

48

Sideways overlap of p-orbitals to form a pi - bond

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Pi Bond

50

51

Pi-bond formation in ethene.

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Movies of Bond Formation

• http://www.google.com/imgres?imgurl=http://www.cem.msu.edu/~harrison/johnston/movies/nitrogen/sigma7icon.jpg&imgrefurl=http://www.cem.msu.edu/~harrison/johnston/nitrogen.html&h=92&w=99&sz=5&tbnid=aJAk-oo0W7Oe7M:&tbnh=71&tbnw=77&hl=en&start=11&prev=/images%

• http://www.google.com/imgres?imgurl=http://www.cem.msu.edu/~harrison/johnston/movies/nitrogen/sigma7icon.jpg&imgrefurl=http://www.cem.msu.edu/~harrison/johnston/nitrogen.html&h=92&w=99&sz=5&tbnid=aJAk-oo0W7Oe7M:&tbnh=71&tbnw=77&hl=en&start=11&prev=/images%

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Depiction of Covalent Compounds

• Covalent compounds can be described by

– Molecular formula (molecular compounds)

– Structural formula (depicts the arrangement of the atoms in space)

– Lewis structure: depicts the arrangement of the electrons around the atoms in a molecule

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Writing Lewis Structures

• The Lewis Structures will be written for molecules that obey the Octet Rule

• AND

• Exceptions to the Octet Rule– Electron deficiency– Expanded Octet Rule– Odd electron molecules

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Writing Lewis Structures

Blackboard

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Resonance

• When more than one dot diagram with the same connections are possible.

• Use double arrows to indicate it is the “average” of the structures.

• NO2-

• Which one is it?• Does it go back and forth.• It is a mixture of both, like a mule.• NO3

- CO3-2 SO3

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Molecular Geometry

58

Molecular Geometry

• Lewis structures tell us how the atoms are connected to each other.

• They don’t tell us anything about shape.

• The shape of a molecule can greatly affect its properties.

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Molecular Geometry

Three theories to explain:

• VSEPR: Valence shell electron pair repulsion theory.

• Valence Bond Theory: Hybridization theory

• Molecular Orbital Theory: advanced course

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Importance of Molecular Shape

Three dimensional structure of a molecule can have a profound effect on its reactivity and biological activity.

These two molecules have identical formulas and shape, but they are mirror images of each other and they have different pharmacological activity (PA) .

Enantiomers and optical isomers

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Hemoglobin

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caffeine

penicillin

AmmoniaWater

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VSEPR• Lewis structures tell us how the atoms are

connected to each other.

• They don’t tell us anything about shape.

• The shape of a molecule can greatly affect its properties.

• Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

64

VSEPR Theory

• Electron Pair Repulsion Theory: electron pairs (both shared and unshared) try to orient themselves as far away as possible in the space around the central atom.

65

VSEPR Theory• Uses Lewis structures and shared and

unshared pair of electrons to predict geometry.

• All electrons are in their original atomic orbitals.

• Predicts three dimensional geometry of molecules. Can predict the angles of bonds.

66

Properties of Shared and Unshared Electron Pairs(1)

• Shared pair of electron: the electron pair is attracted by both nuclei – shape slender, like a cigar.

• Unshared pairs: take a lot of space as electrons repulse each other. Shaped like a pear, or Mickey Mouse ears.

67

Properties of Electron Pairs (2)

• The strength of repulsions between pair of electrons:

unshared-unshared > shared - unshared>

shared - shared • Molecular shape: repulsions between

charge cloud determines the arrangement

68

VSEPR• Molecules take a shape that puts electron

pairs as far away from each other as possible.

• Draw the Lewis structure to determine electron pairs.

• Determine:–bonding–nonbonding lone pair

• Lone pair take more space.• Multiple bonds count as one pair.

69

VSEPR

• The number of pairs determines

–bond angles

–underlying structure

• The number of atoms determines

–actual shape

70

Electron-Group Geometries

• Differentiate between electron-group geometries and molecular geometry

• 2 electron groups: linear

• 3 electron groups: trigonal planar

• 4 electron groups: tetrahedral

• 5 electron groups: trigonal bipyramidal

• 6 electron groups: octahedral

71http://www.mpcfaculty.net/mark_bishop/tri-plan.htm

Electron Group Geometries

72

Electron Group Geometries

Electronpairs

BondAngles

UnderlyingShape

2 180° Linear

3 120° Trigonal Planar

4 109.5° Tetrahedral

590° &120°

Trigonal Bipyramidal

6 90° Octagonal

73

Valence shell electron pair repulsion (VSEPR) model:N0 LONE Electron Pairs around the Central Atom

Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.

AB22 0

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

10.1

linear linear

B B

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Two electron-group geometries

75

Molecular Geometry

No Lone Pairs of Electrons on Central Atom

76

AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: all pairs are shared (0 lonely pairs)

AB3 3 0trigonal planar

trigonal planar

10.1

7710.1

3 electron-group geometry

Other examples: SO3

78

AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: all pairs are shared

AB3 3 0trigonal planar

trigonal planar

10.1

AB4 4 0 tetrahedral tetrahedral

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Tetrahedral Structure

8010.1

4 electron-group geometry

Other examples: NH4+1, SO4

-2

81

AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: All electron pairs are shared

AB3 3 0trigonal planar

trigonal planar

10.1

AB4 4 0 tetrahedral tetrahedral

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

8210.1

5 electron-group geometry

83

AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: all electron pairs are shared

AB3 3 0trigonal planar

trigonal planar

10.1

AB4 4 0 tetrahedral tetrahedral

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB6 6 0 octahedraloctahedral

8410.1

6 electron-pair geometry

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Tetrahedral Structure

8610.1

5 electron-group geometry

8710.1

6 electron-pair geometry

88

Molecular Geometry

Lone Electron Pairs on Central Atom

89

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: THREE ELECTRON-GROUPS

AB3 3 0trigonal planar

trigonal planar

AB2E 2 1 same bent

10.1

90

91

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: FOUR ELECTRON-GROUPS

AB3E 3 1

AB4 4 0 tetrahedral tetrahedral

tetrahedraltrigonal

pyramidal

10.1

92

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: FOUR ELECTRON-GROUPS

AB4 4 0 tetrahedral tetrahedral

10.1

AB3E 3 1 tetrahedraltrigonal

pyramidal

AB2E2 2 2 tetrahedral bent

H

O

H

93

94bonding-pair vs. bonding

pair repulsionlone-pair vs. lone pair

repulsionlone-pair vs. bonding

pair repulsion> >

Effect of Unshared Electron Pair on Bond Angles

95

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: FIVE ELECTRON-GROUPS

10.1

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal

bipyramidaldistorted

tetrahedron

See-saw

96

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: FIVE ELECTRON-GROUPS

10.1

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal

bipyramidaldistorted

tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

ClF

F

F

97

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: FIVE ELECTRON-GROUPS ( LONE PAIRS)

10.1

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal

bipyramidaldistorted

tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

AB2E3 2 3trigonal

bipyramidallinear

I

I

I

98

Five Electron Pairs on Central Atom

99

100

101

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: SIX ELECTRON-GROUPS ( LONE PAIRS)

10.1

AB6 6 0 octahedraloctahedral

AB5E 5 1 octahedral square pyramidal

Br

F F

FF

F

102

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

VSEPR: SIX ELECTRON-GROUPS ( LONE PAIRS)

10.1

AB6 6 0 octahedraloctahedral

AB5E 5 1 octahedral square pyramidal

AB4E2 4 2 octahedral square planar

Xe

F F

FF

103

104

105

106

107

Angles in the Different Geometries

108

Angles in the Different Geometries

109

VSEPRElectron

pairsBond

AnglesUnderlyingShape

2 180° Linear

3 120° Trigonal Planar

4 109.5° Tetrahedral

590° &120°

Trigonal Bipyramidal

6 90° Octagonal

110

Actual shape

ElectronPairs

BondingPairs

Non-Bonding

Pairs Shape

2 2 0 linear

3 3 0 trigonal planar

3 2 1 bent4 4 0 tetrahedral4 3 1 trigonal pyramidal4 2 2 bent

111

Actual Shape

ElectronPairs

BondingPairs

Non-Bonding

Pairs Shape

5 5 0 trigonal bipyrimidal

5 4 1 See-saw

5 3 2 T-shaped5 2 3 linear

112

Actual Shape

ElectronPairs

BondingPairs

Non-Bonding

Pairs Shape

6 6 0 Octahedral

6 5 1 Square Pyramidal

6 4 2 Square Planar6 3 3 T-shaped6 2 1 linear

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VSEPR: Examples

• CH4 - draw the structural formula

• Determine the number of shared and unshared pairs of electrons.

• Write the Lewis structure

114

VSEPR

• Single bonds fill all atoms.

• There are 4 pairs of electrons pushing away.

• The furthest they can get away is 109.5º (REMEMBER: three dimensional space.

C HH

H

H

115

4 atoms bonded

• Basic shape is tetrahedral.

• A pyramid with a triangular base.

• Same shape for everything with 4 pairs.

CH HH

H109.5º

116

3 bonded - 1 lone pair

N HH

H

NH HH

<109.5º

• Still basic tetrahedral but you can’t see the electron pair.

• Shape is calledtrigonal pyramidal.

117

2 bonded - 2 lone pair

OH

H

O HH

<109.5º

• Still basic tetrahedral but you can’t see the 2 lone pair.

• Shape is calledbent.

118

Geometry Using VSEPR Theory

• Vocabulary: – A: central atom– B: bonded atom– E: unshared pair of electrons

• CHECK BLACKBOARD FOR TYPE OF GEOMETRY AND EXAMPLES.

119

Linear Geometry ( 3 atoms)

AB B

Designation: AB2

120

Three Electron Densities

A

B

B

B

B

B

A

E

AB3, trigonal planar AB3E, Bent

121

Four Electron Densities

AB4 AB3E AB2E2

122

Five Electron Densities

AB5

AB3E2

AB4E

AB2E3

123

Six Electron Densities

AB6 AB5E

AB4 E2

124

3 Atoms No Lone Pair (double Bond)

CH

HO

• The farthest you can the electron pair apart is 120º

125

3 atoms no lone pair

CH

HO

• The farthest you can place the electron pairs apart is 120º.

• Shape is flat and called trigonal planar.

C

H

H O

120º

126

2 atoms no lone pair

• With three atoms the farthest they can get apart is 180º.

• Shape called linear.

C OO180º

127

Valence Bond Theory: Hybrid Orbitals

Combines bonding with geometry

128

Hybridization

• The mixing of several atomic orbitals to form the same number of hybrid orbitals.

• All the hybrid orbitals that form are the same.

129

Types of Hybrid Orbitals

• sp3 : 1 s and 3 p orbitals mix to form 4 sp3 orbitals.

• sp2 :1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital intact.

• sp : s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals intact.

• sp3d: five orbitals• sp3d2 : six orbitals

130

Hybridization

• We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry.

• We combine one s orbital and 3 p orbitals.

• sp3 hybridization has tetrahedral geometry.

131

How we get to hybridization

• We know the geometry from experiment.

• We know the orbitals of the atom

• hybridizing atomic orbitals can explain the geometry.

• So if the geometry requires a tetrahedral shape, it is sp3 hybridized.

• This includes bent and trigonal pyramidal molecules because one of the sp3 lobes holds the lone pair.

132

Hybridization

• We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry.

• We combine one s orbital and 3 p orbitals.

• sp3 hybridization has tetrahedral geometry.

133

sp3 geometry

109.5º

• This leads to tetrahedral shape.

• Every molecule with a total of 4 atoms and lone pair is sp3 hybridized.

• Gives us trigonal pyramidal and bent shapes also.

134

135

136

sp2 hybridization

• C2H4, BF3

• double bond ( in C2H4 ) acts as one pair

• trigonal planar

• Have to end up with three blended orbitals

• use one s and two p orbitals to make sp2 orbitals.

• leaves one p orbital perpendicular

137

138

139

140

sp3 geometry

109.5º

• This leads to tetrahedral shape.

• Every molecule with a total of 4 atoms and lone pair is sp3 hybridized.

• Gives us trigonal pyramidal and bent shapes also.

141

sp2 hybridization in C2H4

• trigonal planar

• 120º angle

• one bond

• One sigma and one pi bond between the C-C atoms

142

Where is the P orbital?

• Perpendicular

• The overlap of orbitals makes a sigma bond ( bond)

143

CCH

H

H

H

144

What about sp

• one s and one p hybridize

• Linear

• C2H2

145

sp hybridization

• end up with two lobes 180º apart.

• p orbitals are at right angles

• makes room for two bonds and two sigma bonds.

• a triple bond or two double bonds

146

Hybridization and Geometry of Electrons

• sp3: tetrahedral• sp2: trigonal planar• sp: linear• sp3d: trigonal bipyramid• sp3d2: octahedral• Geometry of molecule determined by

number of shared and unshared electrons.

147

Types of Hybrid Orbitals

• sp3 : 1 s and 3 p orbitals mix to form 4 sp3 orbitals.

• sp2 :1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital intact.

• sp : s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals intact.

• sp3d: five orbitals• sp3d2 : six orbitals

148

Polar Bonds

• When the atoms in a bond are the same, the electrons are shared equally.

• This is a nonpolar covalent bond.

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

• This is a polar covalent bond.

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

149

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

150

Hybridization and Geometry of Electrons

• sp3: tetrahedral• sp2: trigonal planar• sp: linear• sp3d: trigonal bipyramid• sp3d2: octahedral• Geometry of molecule determined by

number of shared and unshared electrons.

151

Types of Hybrid Orbitals

• sp3 : 1 s and 3 p orbitals mix to form 4 sp3 orbitals.

• sp2 :1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital intact.

• sp : s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals intact.

• sp3d: five orbitals• sp3d2 : six orbitals

152

Polar Bonds

• When the atoms in a bond are the same, the electrons are shared equally.

• This is a nonpolar covalent bond.

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

• This is a polar covalent bond.

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

153

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

154

Polar Molecules

Molecules with endsDetermined by polarity of bonds

AndSymmetry of Molecules

155

Partial Ionic Compounds (cont.)

Covalent

Polar Covalent

Ionic

Increased Ionic Character

156

Criteria for Polarity of molecules

• 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 first.

157

Criteria for Polarity of molecules

• 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 first.

158

Geometry and polarity• Three shapes will cancel them out.

• Linear

159

Geometry and polarity• Three shapes will cancel them out.

• Planar triangles

120º

160

Geometry and polarity• Three shapes will cancel them out.

• Tetrahedral

161

Geometry and polarity• Others don’t cancel

• Bent

162

Geometry and polarity• Others don’t cancel

• Trigonal Pyramidal

163

Is it polar?

• HF

• H2O

• NH3

• CCl4

• CO2

• PCl5• CO

164

Bond Dissociation Energy

• The energy required to break a bond

• 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)

165

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