Chapter 11
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Chapter 11. Theories of Covalent Bonding. If you are doing this lecture “online” then print the lecture notes available as a word document, go through this ppt lecture, and do all the example and practice assignments for discussion time. VALENCE BOND THEORY:.

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Chapter 11

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Chapter 11

Chapter 11

Theories of Covalent Bonding

If you are doing this lecture “online” then print the lecture notes available as a word document, go through this ppt lecture, and do all the example and practice assignments for discussion time.


Valence bond theory

VALENCE BOND THEORY:

  • DEVELOPED BY LINUS PAULING, who received the Nobel Prize in 1954 for his work

  • A view of chemical bonding in which bonds arise from the overlap of atomic orbitals on two atoms to give a bonding orbital of electrons localized between the bonded atoms

  • RULE: Realize that Valence Bond Theory and all the others don't explain everything


Chapter 11

The Central Themes of VB Theory

Basic Principle

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.

(The two wave functions are in phase so the amplitude increases

between the nuclei.)


Chapter 11

The Central Themes of VB Theory

Themes

A set of overlapping orbitals has a maximum of two electrons that must have opposite spins.

The greater the orbital overlap, the stronger (more stable) the bond.

The valence atomic orbitals in a molecule are different from those in isolated atoms.

There is a hybridization of atomic orbitals to form molecular

orbitals.


Chapter 11

Orbital overlap and spin pairing in three diatomic molecules.

Figure 11.1

Hydrogen, H2

Hydrogen fluoride, HF

Regular atomic orbital overlap can explain these bonds.

Fluorine, F2


Valence bond theory1

VALENCE BOND THEORY:

Ha to Hb: 1sa to 1sb overlap radius = 74 pm

As overlap increases, strength of bond increases - both electrons are mutually attracted to both atomic nuclei.

At optimum distance between nuclei with maximum overlap, a sigma s bond (strong primary bond) forms. Max electron density is along the axis of the bond

Ha to Fb: 1sa to 2pb direct overlap or s bond


Valence bond theory2

VALENCE BOND THEORY:

F to F: the picture looks like a 2p orbital on one F is overlapping with a 2p orbital on the other F atom, but actually each F is sp3 hybridized & electrons are localized between two atomic nuclei


Valence bond theory3

VALENCE BOND THEORY:

We cannot use this direct overlap picture for CH4’s bonding. The 2s and the three 2p orbitals on each C do not fit into the CH4 molecule's 109o bond angles, since the 2p orbitals are at 90° to each other

Valence Bond Theory states that HYBRID orbitals of the outermost orbitals on an atom are formed from the atoms’ atomic orbitals


Chapter 11

Key Points

Types of Hybrid Orbitals

sp

sp2

sp3

sp3d

sp3d2

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.


Chapter 11

The sp hybrid orbitals in gaseous BeCl2.

Figure 11.2

atomic orbitals on Be

hybrid orbitals

You have to know how to draw this energy hybrid formation.

orbital box diagrams


Chapter 11

Figure 11.2

The sp hybrid orbitals in gaseous BeCl2(continued).

orbital box diagrams with orbital contours


Chapter 11

The sp2 hybrid orbitals in BF3.

Figure 11.3

You have to know how to draw this energy hybrid formation.

Note the three sigma bonds formed between B and each F.


Chapter 11

The sp3 hybrid orbitals in CH4.

Figure 11.4

You have to know how to draw this energy hybrid formation.


Chapter 11

Figure 11.5

The sp3 hybrid orbitals in NH3.

You have to know how to draw this energy hybrid formation.


Chapter 11

Figure 11.5 continued

The sp3 hybrid orbitals in H2O.

You have to know how to draw this energy hybrid formation.


Valence bond theory4

VALENCE BOND THEORY

Expanded Valence Shells have hybrid orbitals using s, p & d atomic orbitals. Example: PCl5 P: [Ne]3s23p3

dsp3 hybridization results in 5 s bonds and trigonal bipyramidal geometry

(You can write these as dsp3 or sp3d)


Chapter 11

Figure 11.6

The sp3d hybrid orbitals in PCl5.

You have to know how to draw this energy hybrid formation.


Chapter 11

The sp3d2hybrid orbitals in SF6.

Figure 11.7

You have to know how to draw this energy hybrid formation.


Chapter 11

Step 1

Step 2

Step 3

Figure 10.1

Figure 10.12

Table 11.1

Figure 11.8

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

Molecular shape and e- group arrangement

Molecular formula

Lewis structure

Hybrid orbitals


Chapter 11

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:

SAMPLE PROBLEM 11.1

Postulating Hybrid Orbitals in a Molecule

(a) Methanol, CH3OH

(b) Sulfur tetrafluoride, SF4

SOLUTION:

The groups around C are arranged as a tetrahedron.

(a) CH3OH

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


Chapter 11

hybridized C atom

hybridized O atom

single C atom

single O atom

hybridized S atom

S atom

SAMPLE PROBLEM 11.1

Postulating Hybrid Orbitals in a Molecule

continued

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


Valence bond theory5

VALENCE BOND THEORY

There can be more than one central atom, and each has its own hybridization and geometry

C2H6 and H2O2 and CH3COOH

C2H6: both C's are sp3 hybridized and can rotate around axis of bond.

H2O2: both O's are sp3 , etc.


Chapter 11

both C are sp3 hybridized

s-sp3 overlaps to s bonds

sp3-sp3 overlap to form a s bond

relatively even distribution of electron density over all s bonds

The s bonds in ethane(C2H6).

Figure 11.9


Valence bond theory multiple bonds

VALENCE BOND THEORY: Multiple Bonds

H2CO: the Lewis structures shows a double bond between C and O, but we know it does not have twice the bond dissociation energy of a single C-O bond

Pauling proposed that there was only one sigma s bond between any two atoms, and the other multiples were weaker pi p bonds

If there are only 3 s bonds around this carbon, it can't be sp3 hybridized - instead we have sp2 hybrid orbitals

sp2 hybridization results in only 3 s bonds, and trig planar geometry, with 120° angles

p bond is a sideways or parallel overlap of the p atomic orbitals rather than the direct overlap of s bonds


Chapter 11

overlap in one position - s

p overlap - 

electron density

The s and p bonds in ethylene (C2H4).

Figure 11.10

Proper name is ethene.


Valence bond theory6

VALENCE BOND THEORY

Look at acetylene: its geometry is linear. C is forming a triple bond to another C and a single bond to H, so that's only two s bonds

Therefore sp hybridization results in only 2 s bonds, and linear geometry

There are 2 p bonds from the parallel overlap of the 2p orbitals remaining on both C's


Chapter 11

overlap in one position - s

p overlap - 

The s and p bonds in acetylene (C2H2).

Figure 11.11


Chapter 11

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.

sp3 hybridized

sp3 hybridized

sp2 hybridized

SAMPLE PROBLEM 11.2

Describing the Bond in Molecules

SOLUTION:

bond

bonds


Chapter 11

H2CO hybrid orbitals and sigma and pi bond formation

Remember the C=C double bond has sigman and pi bonds.

The C has sigma bonds from its hybrid orbitals to the two H’s and the O. The leftover p orbitals will form the pi bond.


Chapter 11

CIS

TRANS

Figure 11.13 from 4th ed.

Restricted rotation of p-bonded molecules in C2H2Cl2.

This cis/trans arrangement will be important in chem 2, organic chem and biology!


Valence bond theory resonance

VALENCE BOND THEORY: RESONANCE

Resonance Structures and p Bonding:

p resonance structures involve an electron pair used alternately as a p bond or a LP

Ozone: O3 O==O--O or O--O==O

All are sp2, trig planar, each has 3 sp2 orbitals and a p orbital remaining.


Valence bond theory7

VALENCE BOND THEORY

Benzene: C6H6 has carbons with sp2 hybrids and 120o angles, each C has 2 s bonds to other C's, 1 s bond to H, and 1 p bond electron available

Get "ring" of delocalized p e-s

SUMMARY: draw the Lewis structure; determine arrangement of electron pairs using VSEPR, specify the hybrid orbitals to accommodate the e- pairs


Chapter 11

Benzene sigma bond formation between C’s and C-Hs

The leftover p orbitals will form alternating pi bonds as shown in sketch.


Molecular orbital theory

MOLECULAR ORBITAL THEORY:

- explains why H2 forms easily and He2 does not

- is an alternate way of viewing e- orbitals in molecules where pure s and pure p orbitals combine to produce orbitals that are delocalized over the molecule

- they can have different energies and are assigned electrons just like we do in an atom - Pauli exclusion principle and Hund's rule included

Pauling's Valence Bond Theory does not explain everything

MO Theory doesn't either, but it does correctly predict the electronic structure of certain molecules that do not follow Lewis's approach, including the paramagnetism of certain molecules, like O2


Chapter 11

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. The number of molecular orbitals produced is always = # of atomic orbitals brought by the combining atoms (only orbitals on different atoms are combined).

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).

The electrons of the molecule are placed in bonding or antibonding orbitals of successively higher energy (just like Hund's rule).

Atomic orbitals combine most effectively with orbitals of the same type and similar energy (s w/s, n=2 w/ n=2)


Chapter 11

Amplitudes of wave functions subtracted.

Figure 11.13

An analogy between light waves and atomic wave functions.

Amplitudes of wave functions added


Chapter 11

Figure 11.14

Contours and energies of the bonding and antibonding molecular orbitals (MOs) in H2.

The bonding MO is lower in energy and the antibonding MO is higher in energy than the AOs that combined to form them.


Molecular orbital theory1

MOLECULAR ORBITAL THEORY

BOND ORDER: the number of bonding e- pairs shared by 2 atoms in a molecule

Fractional bond orders are possible in MO Theory!

Silberberg method:

B.O. = ½(# of e- in bonding orbitals - # of e- in antibonding orbitals)


Chapter 11

1s

1s

AO of H

AO of H

Figure 11.15

The MO diagram for H2.

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

s*1s

Energy

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

s1s

MO of H2


Chapter 11

s*1s

Energy

1s

1s

1s

1s

s1s

AO of He

AO of He+

AO of He

AO of He

Figure 11.16

MO diagram for He2+ and He2.

s*1s

Energy

s1s

MO of He+

MO of He2

He2 bond order = 0

He2+ bond order = 1/2


Chapter 11

PROBLEM:

Use MO diagrams to predict whether H2+ and H2- exist. Determine their bond orders and electron configurations.

s

s

1s

1s

1s

1s

AO of H-

AO of H

AO of H

AO of H

s

s

SAMPLE PROBLEM 11.3

Predicting Stability of Species Using MO Diagrams

SOLUTION:

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

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

H2+ does exist

H2- does exist

MO of H2-

MO of H2+


Chapter 11

s*2s

s*2s

2s

2s

2s

2s

s2s

s2s

Figure 11.17

Bonding in s-block homonuclear diatomic molecules.

Energy

Be2

Li2

Be2 bond order = 0

Li2 bond order = 1


Chapter 11

Contours and energies of s and p MOs through combinations of 2p atomic orbitals.

Figure 11.18

Or the pz orbitals


Chapter 11

Figure 11.19

Relative MO energy levels for Period 2 homonuclear

diatomic molecules.

with 2s-2p mixing

without 2s-2p mixing

Memorize this!

MO energy levels for O2, F2, and Ne2

MO energy levels for B2, C2, and N2


Chapter 11

Figure 11.20

MO occupancy and molecular properties for B2 through Ne2


Chapter 11

Figure 11.21

The paramagnetic properties of O2


Chapter 11

PROBLEM:

As the following data show, removing an electron from N2 forms an ion with a weaker, longer bond than in the parent molecules, whereas the ion formed from O2 has a stronger, shorter bond:

N2

N2+

O2

O2+

Bond energy (kJ/mol)

945

841

498

623

Bond length (pm)

110

112

121

112

SAMPLE PROBLEM 11.4

Using MO Theory to Explain Bond Properties

Explain these facts with diagrams that show the sequence and occupancy of MOs.

SOLUTION:

N2 has 10 valence electrons, so N2+ has 9.

O2 has 12 valence electrons, so O2+ has 11.


Chapter 11

2p

2p

2p

2p

s2s

s2s

SAMPLE PROBLEM 11.4

Using MO Theory to Explain Bond Properties

continued

N2

N2+

O2

O2+

2p

antibonding e- lost

bonding e- lost

2p

2p

2p

s2s

s2s

bond orders

1/2(8-2)=3

1/2(7-2)=2.5

1/2(8-4)=2

1/2(8-3)=2.5


Mo theory practice

MO Theory Practice

1. Draw the bonding and antibonding molecular orbitals for H2.

2. Do Valence Bond Theory (hybridization) and MO Theory for both O2 and O22-. Which theory works better to explain the molecule and ion?

3. For N2, N2+ and N2- compare

a. Magnetic character

b. Net number of p bonds

c. Bond Order

d. Bond length

e. Bond strength


Answers

Answers

1. See picture in text.

2. VB Theory shows O2 as sp2 hybridized with one s bond and one p bond. There are two lone pairs on each O. O22- has one s bond, and each O has three lone pairs. MO Theory shows a bond order of 2 for O2 and that it is paramagnetic. MO Theory shows a bond order of 1 for O22- and diamagnetic. But MO Theory fits the real data that O2 is paramagnetic.


Answers con t

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