valence bond theory pptx

11-1

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Theories of Covalent Bonding

11.1 Valence Shell Electron Pair Repulsion Theory

11.2 Valence Bond (VB) Theory and Orbital Hybridization

11.3 Molecular Orbital (MO)Theory and Electron Delocalization

Chapter 9Prentice Hall © 2003

• Lewis structures and VSEPR do not explain why a bond forms. How do we account for shape in terms of quantum mechanics?

• What are the orbitals that are involved in bonding?

• We use Valence Bond Theory:• Bonds form when orbitals on atoms overlap.

• A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.

• There are two electrons of opposite spin in the orbital overlap.

Covalent Bonding and Orbital Overlap



Hydrogen, H2

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Hydrogen fluoride, HF

Fluorine, F2


• As two nuclei approach each other their atomic orbitals overlap.

• As the amount of overlap increases, the energy of the interaction decreases.

• At some distance the minimum energy is reached.• The minimum energy corresponds to the bonding

distance (or bond length).• As the two atoms get closer, their nuclei begin to repel

and the energy increases.



• At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).



• Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding.

• Hybridization is determined by the electron domain geometry.

sp Hybrid Orbitals

• Consider the BeF2 molecule (experimentally known to exist):

Hybrid Orbitals


sp Hybrid Orbitals• Be has a 1s22s2 electron configuration.

• There is no unpaired electron available for bonding.

• We conclude that the atomic orbitals are not adequate to describe orbitals in molecules.

• We know that the F-Be-F bond angle is 180 (VSEPR theory).

• We also know that one electron from Be is shared with each one of the unpaired electrons from F.

Hybrid Orbitals


sp Hybrid Orbitals• We assume that the Be orbitals in the Be-F bond are 180

apart.• We could promote and electron from the 2s orbital on Be to the

2p orbital to get two unpaired electrons for bonding.• BUT the geometry is still not explained.

• We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital.

• The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital.

Hybrid Orbitals


sp Hybrid Orbitals• The lobes of sp hybrid orbitals are 180º apart.• Since only one of the Be 2p orbitals has been used in

hybridization, there are two unhybridized p orbitals remaining on Be.

Hybrid Orbitals

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Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.

atomic orbitals

hybrid orbitals

orbital box diagrams

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Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued).

orbital box diagrams with orbital contours

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Figure 11.3 The sp2 hybrid orbitals in BF3.

sp2 and sp3

Hybrid Orbitals

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Figure 11.4 The sp3 hybrid orbitals in CH4.

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Figure 11.5 The sp3 hybrid orbitals in NH3.

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Figure 11.5 continued The sp3 hybrid orbitals in H2O.

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Figure 11.6 The sp3d hybrid orbitals in PCl5.

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Figure 11.7 The sp3d2 hybrid orbitals in SF6.

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Hybrid Orbitals

The number of hybrid orbitals obtained equals the number of atomic orbitals mixed.

The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.

Key Points

sp sp2 sp3 sp3d sp3d2

Types of Hybrid Orbitals

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Figure 11.8

The conceptual steps from molecular formula to the hybrid orbitals used in bonding.

Molecular formula

Lewis structure

Molecular shape and e- group arrangement

Hybrid orbitals

Step 1 Step 2 Step 3

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SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule

SOLUTION:

PROBLEM: Use partial orbital diagrams to describe mixing of the atomic orbitals of the central atom leads to hybrid orbitals in each of the following:

PLAN: Use the Lewis structures to ascertain the arrangement of groups and shape of each molecule. Postulate the hybrid orbitals. Use partial orbital box diagrams to indicate the hybrid for the central atoms.

(a) Methanol, CH3OH (b) Sulfur tetrafluoride, SF4

(a) CH3OH H

CH H

OH

The groups around C are arranged as a tetrahedron.

O also has a tetrahedral arrangement with 2 nonbonding e- pairs.

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SFF

F

F

SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule

continued

2p

2s single C atom

sp3

hybridized C atom

2p

2s single O atom

sp3

hybridized O atom

(b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs.

3p

3s

3d

S atomsp3d

3d

hybridized S atom


Hybridization Involving d Orbitals• Since there are only three p-orbitals, trigonal bipyramidal

and octahedral electron domain geometries must involve d-orbitals.

• Trigonal bipyramidal electron domain geometries require sp3d hybridization.

• Octahedral electron domain geometries require sp3d2 hybridization.

• Note the electron domain geometry from VSEPR theory determines the hybridization.

Hybrid Orbitals


Summary

1. Draw the Lewis structure.

2. Determine the electron domain geometry with VSEPR.

3. Specify the hybrid orbitals required for the electron pairs based on the electron domain geometry.

Hybrid Orbitals


Multiple Bonds

-Bonds: electron density lies on the axis between the nuclei.

• All single bonds are -bonds. -Bonds: electron density lies above and below the plane

of the nuclei.• A double bond consists of one -bond and one -bond.• A triple bond has one -bond and two -bonds.• Often, the p-orbitals involved in -bonding come from

unhybridized orbitals.

Multiple Bonds

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Figure 11.9 The bonds in ethane(C2H6).

both C are sp3 hybridizeds-sp3 overlaps to bonds

sp3-sp3 overlap to form a bondrelatively even

distribution of electron density over all bonds


Multiple Bonds

Ethylene, C2H4, has:• one - and one -bond;

• both C atoms sp2 hybridized;

• both C atoms with trigonal planar electron pair and molecular geometries.

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Figure 11.10 The and bonds in ethylene (C2H4).

overlap in one position -

p overlap -

electron density


Multiple Bonds


Multiple Bonds

• When triple bonds form (e.g. N2) one -bond is always above and below and the other is in front and behind the plane of the nuclei.


Multiple Bonds

Consider acetylene, C2H2

• the electron pair geometry of each C is linear;

• therefore, the C atoms are sp hybridized;

• the sp hybrid orbitals form the C-C and C-H -bonds;

• there are two unhybridized p-orbitals;

• both unhybridized p-orbitals form the two -bonds;

• one -bond is above and below the plane of the nuclei;

• one -bond is in front and behind the plane of the nuclei.

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Figure 11.11 The and bonds in acetylene (C2H2).

overlap in one position -

p overlap -


Multiple Bonds


Multiple Bonds

Delocalized Bonding• So far all the bonds we have encountered are localized

between two nuclei.• In the case of benzene

• there are 6 C-C bonds, 6 C-H bonds,• each C atom is sp2 hybridized,• and there are 6 unhybridized p orbitals on each C

atom.


Multiple Bonds

Delocalized Bonding


Multiple Bonds

Delocalized Bonding• In benzene there are two options for the 3 bonds

• localized between C atoms or• delocalized over the entire ring (i.e. the electrons

are shared by all 6 C atoms).• Experimentally, all C-C bonds are the same length in

benzene.• Therefore, all C-C bonds are of the same type (recall

single bonds are longer than double bonds).


Multiple Bonds

General Conclusions• Every two atoms share at least 2 electrons.• Two electrons between atoms on the same axis as the

nuclei are bonds. -Bonds are always localized.• If two atoms share more than one pair of electrons, the

second and third pair form -bonds.• When resonance structures are possible, delocalization is

also possible.


Molecular Orbitals

• Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization. (E.g. why does O2 interact with a magnetic field?; Why are some molecules colored?)

• For these molecules, we use Molecular Orbital (MO) Theory.

• Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.

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SAMPLE PROBLEM 11.2 Describing the Bond in Molecules

SOLUTION:

PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO.

PLAN: Use the Lewis structures to ascertain the arrangement of groups and shape at each central atom. Postulate the hybrid orbitals taking note of the multiple bonds and their orbital overlaps.

H3C

C

CH3

O

sp3 hybridized

sp3 hybridized

CC

C

O

H

H

HHH

H

sp2 hybridized

bondsbond

CC

C

O

sp3

sp3

sp3

sp3

sp3

sp3

sp3

sp3

sp2 sp2

sp2

sp2

sp2sp2

H

HH

HH

H

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The Central Themes of MO Theory

A molecule is viewed on a quantum mechanical level as a collection of nuclei surrounded by delocalized molecular orbitals.

Atomic wave functions are summed to obtain molecular wave functions.

If wave functions reinforce each other, a bonding MO is formed (region of high electron density exists between the nuclei).

If wave functions cancel each other, an antibonding MO is formed (a node of zero electron density occurs between the nuclei).

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Amplitudes of wave functions added

Figure 11.14

An analogy between light waves and atomic wave functions.

Amplitudes of wave functions

subtracted.


Molecular Orbitals

• Molecular orbitals:• each contain a maximum of two electrons;

• have definite energies;

• can be visualized with contour diagrams;

• are associated with an entire molecule.

The Hydrogen Molecule• When two AOs overlap, two MOs form.


Molecular Orbitals

The Hydrogen Molecule• Therefore, 1s (H) + 1s (H) must result in two MOs for

H2:

• one has electron density between nuclei (bonding MO);

• one has little electron density between nuclei (antibonding MO).

• MOs resulting from s orbitals are MOs. (bonding) MO is lower energy than * (antibonding)

MO.


Molecular Orbitals

The Hydrogen Molecule


Molecular Orbitals

The Hydrogen Molecule• Energy level diagram or MO diagram shows the energies

and electrons in an orbital.• The total number of electrons in all atoms are placed in

the MOs starting from lowest energy (1s) and ending when you run out of electrons.• Note that electrons in MOs have opposite spins.

• H2 has two bonding electrons.

• He2 has two bonding electrons and two antibonding electrons.


Figure 11.15 The MO diagram for H2.

En

erg

y

MO of H2

*1s

1s

AO of H

1s

AO of H

1s

H2 bond order = 1/2(2-0) = 1

Filling molecular orbitals with electrons follows the same concept as filling atomic orbitals.


Second-Row Diatomic Molecules

Electron Configurations and Molecular Properties

• Two types of magnetic behavior:• paramagnetism (unpaired electrons in molecule): strong

attraction between magnetic field and molecule;

• diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule.

• Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field:

valence bond theory pptx

Technology

valence bond theory

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overlap region

orbital hybridization11

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