Chapter 7 Theories of Chemical Bonding 7.1 Localized Bonds 7.2 Hybridization of Atomic Orbitals 7.3...

Chapter 7 Theories of Chemical Bonding

Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.

Orbitals and Bonding Theories

Review:

Electrons are attracted to the nuclei.

Electrons act as waves. Orbitals are wavefunctions.

Electrons are responsible for bonds in a molecule.


7.1 Localized Bonds

Learning objective:

Use the orbital overlap model to explain the bonding in simple molecules


7.1 Localized Bonds

Two distinct methods for characterizing bonding:

Localized Bonding: electrons are localized in the bonds between two atoms, usually in pairs.

Delocalized Bonding: bonds delocalized over several atoms, which explains some chemical properties.


Orbital Overlap

Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms.


Orbital Overlap: As two hydrogen atoms approach each other, the overlap of their 1s atomic orbitals increases. The wave amplitudes add, generating a new orbital with high electron density between the nuclei.

What Else is Going on?


Conventions of the Orbital Overlap Model

1. Each electron in a molecule is assigned to a specific orbital.

2. No two electrons in a molecule have identical descriptions, because Pauli exclusion principle applies to electrons in molecules as well as in atoms.

3. The electrons in molecules obey the aufbau principle, meaning that they occupy the most stable orbitals available to them.

4. Even though every atom has an unlimited number of atomic orbitals, the valence orbitals are all that are needed to describe bonding.


Diatomic Molecules: HF and F2


Bonding in H2S

What is the Lewis structure of H2S?


Example 7 - 1

Phosphine is a colourless, highly toxic gas with bond angles of 93.6°. Describe the bonding in PH3.


7.2 Hybridization of Atomic Orbitals

Learning objective:

Assign the correct hybrid orbitals used by each inner atom in a molecule, and the molecular geometry that results.


7.2 Hybridization of Atomic Orbitals

Remember Mendel? He made hybrids of pea plants by mixing purebreds. We will apply a similar method to atomic orbitals, first described by Linus Pauling.

Atomic orbitals can be hybridized to generate a new set of directional orbitals.

These mixed orbitals match the orbital geometry of the compounds.

Remember, all electrons around the central atom must be in orbitals --- whether they are nonbonding electrons or bonding electrons.


s and p Hybridization


Methane Hybridization


A single sp hybrid orbital

Four sp hybrid orbitals on the carbon atom

1s atomic orbitals on the H atoms

The s and all the p orbitals are needed for directional bonding, therefore, the s and the px, py, and pz hybridize.

The new orbitals are called sp3. These overlap with the 1s atomic orbitals of the hydrogen atoms to make CH4.


Describe the bonding of the hydronium ion, H3O+.


Example 7 - 2

Describe the bonding of methanol, CH3OH. Sketch an orbital overlap picture of the molecule.


Example 7 - 3

General Features of Hybridization

1. The number of valence orbitals generated by the hybridization process equals the number of valence atomic orbitals participating in hybridization.

2. The steric number of an inner atom uniquely determines the number and type of hybrid orbitals.

3. Hybrid orbitals form localized bonds by overlap with atomic orbitals or with other hybrid orbitals.

4. There is no need to hybridize orbitals on outer atoms, because atoms do not have limiting geometries. Hydrogen always forms localized bonds with its 1s orbital. The bonds formed by all other outer atoms can be described using valence p orbitals.


sp2 Hybrid Orbitals

Mixes an s orbital with two p orbitals (s+p+p) (the third p-orbital is unchanged!)

Required by central atoms with steric number of 3 (trigonal planar electron group geometry)


sp Hybrid Orbitals

Mixes an s orbital with a p orbital (s+p) (the other two p-orbitals are unchanged!)

Required by central atoms with steric number of 2 (linear electron group geometry)


sp3d Hybrid Orbitals

Mixes an s orbital with three p orbitals and a d orbital (s+p+p+p+d)

Required by central atoms with steric number of 5 (trigonal bipyramidal electron group geometry)


sp3d2 Hybrid Orbitals

Mixes an s orbital with three p orbitals and two d orbitals (s+p+p+p+d+d)

Required by central atoms with steric number of 6 (octahedral electron group geometry)


Describe the bonding of chlorine trifluoride and xenon tetrafluoride


Example 7 - 4

Summary of Valence Orbital Hybridization


Summary of Valence Orbital Hybridization – cont’d


For each of the following Lewis structures, name the electron group geometry and the hybrid orbital used by the inner atoms.


Example 7 - 5

7.3 Multiple Bonds

Learning objective:

Describe the and bonding systems in multiple bonds.


7.3 Multiple Bonds

Double bond – has two sets of bonding electrons, therefore it must require two sets of overlapping orbitals.

Triple bond – has three sets of bonding electrons, therefore it must require three sets of overlapping orbitals.


Sigma Bonds and Pi Bonds

A single overlap of two orbitals (as in H2) is called a sigma () bond. Electron density is distributed along the internuclear axis.

A double overlap of two orbitals is called a pi () bond. Electron density is distributed above and below the bond axis


• The carbons each have sp2 hybridization.

• How do we account for the double bond, the second electron pair?

• HINT: how many p-orbitals are there? How many have been used for hybridization?

Bonding in Ethylene


Formation of an sp2 hybrid set leaves one unused valence p orbital.In ethylene, these orbitals overlap to form a second bond between thecarbon atoms. Notice: there is double overlap! This is a bond.


An sp2 hybrid orbital on each carbon overlaps with one on the other carbon atom. This is a bond. C-H bonds are formed by overlap of an sp2 hybrid with a 1s atomic orbital on the hydrogen atom.

An imine is a molecule that contains a carbon – nitrogen double bond. Describe the bonding of the simplest possible imine, H2CNH, by sketching the and bonding systems.


Example 7 - 6

Bonds Involving Oxygen Atoms


Acetic Acid

Carbon vs Silicon ( bonds)


What About Triple Bonds?

If a single bond is a sigma (and a double bond is a sigma plus a pi ( + ), then triple bonds are a sigma plus a pi plus a pi ( + + ) This requires 2 empty p-orbitals on each atom for double

orbital overlap. Let’s look at acetylene, C2H2



Acetylene, C2H2

Three views of the p bonding in acetylene. Left, molecule viewedfrom the side. Middle, molecule viewed at a 45o angle. Right, molecule viewed from one end. Notice that the p bonds are perpendicular to each other and have electron density off the internuclear axis.

Bonds: p Orbital Overlap


Composite orbital overlap view of the and p bonds of acetylene.The bonds are shown in an “exploded” view in order to make the C-C bond visible.


Acetylene, C2H2

Example 7 - 7

Hydrogen cyanide (HCN) is an extremely poisonous gas with an odour resembling that of almonds. Approximately one billion pounds of HCN are produced each year, most of which are used to prepare starting materials for polymers. Construct a complete bonding picture for HCN and sketch the various orbitals.


7.4 Molecular Orbital Theory: Diatomic Molecules

Learning objective:

Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule and explain trends in bond length and energy


7.4 Molecular Orbital Theory: Diatomic Molecules

Bonding in the diatomic molecules second row elements can be explained in two ways:

1. localized bonding theory (already discussed)2. Pure s and p atomic orbitals of the atoms in a molecule combine

to produce orbitals that are spread out, or delocalized, over several atoms, leading to molecular orbitals (MOs).

Advantages of MOs over Localized Bonding Theory:

- predicts relative bond lengths and energies- predicts magnetic properties- correctly explains electronic structures of molecules

which do not follow Lewis Dot structure.


The Principles of MO Theory

1st Principle: the total number of molecular orbitals produced by a set of interacting atomic orbitals is always equal to the number of interacting atomic orbitals.

To explain this, let’s look at the hydrogen molecule, H2


Bonding and Antibonding in H2

Each hydrogen atom contributes a 1s orbitalThese orbitals can be added or subtracted

Addition of the orbitals: high energy density between the two orbitals, leads to overlap and a bonding molecular orbital, called 1s (just like the in VB theory)

Subtraction of the orbitals: low energy density between the two orbitals, there is no overlap and is called an antibonding molecular orbital, called

1s


When two hydrogen 1s orbitals interact, they generate two new orbitals, one bonding molecular orbital (1s) and one antibonding (

1s). Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.

Principles of MO Theory

2nd Principle: the bonding molecular orbital is lower in energy than the parent orbitals and the antibonding orbital is higher in energy. The system is “stabilized” when electrons are assigned to the

bonding orbitals The system is “destabilized” when electrons are assigned to

the antibonding orbitals because the energy of the system is higher.


MO Diagram of H2



3rd Principle: electrons of the molecule are assigned to orbitals of successively higher energy according to the aufbau principle and Hund’s Rule. Electrons occupy the lowest energy orbitals first. Atoms are most stable with the highest number of unpaired

electrons.


Bond Order in MO Theory

Bond order (BO) allows us to calculate the net amount of bonding between two atoms.

1BO = # electrons in bonding MOs - # electrons in antibonding MOs

2


Example 7 - 8

Use a molecular orbital diagram to predict if it is possible to form the He2

+ cation.


MO diagram of He2. BO = 0


4th Principle: atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy. e.g. a 1s will not bond with a 2s.


Second-Row Diatomic Molecules


Evidence for Antibonding Orbitals


Homonuclear Diatomic Molecules

MO diagram for O2 can be generalized for any 2nd row diatomic molecule.

But under experimentation, it turns out the B2 is an exception to theory, so the theory must be revised to explain the exceptions


Orbital Mixing

In B2, the overlap of 2s and 2pz orbitals stabilizes s and destabilizes p. This is referred to as orbital mixing

The amount of mixing depends on the energy difference between the 2s and 2p atomic orbitals.

Mixing is largest when the energies of the orbitals are nearly the same


Example 7 - 9

Use molecular orbital diagrams to explain the trend in the following bond energies:

B2 = 290 kJ/mol, C2 = 600 kJ/mol and N2 = 942 kJ/mol.


Homonuclear Diatomic Molecules

Let’s examine the MO diagram of NO


Example 7 – 10

The first ionization energy of NO is 891 kJ/mol, that of N2 is 1500 kJ/mol, and that of CO is 1350 kJ/mol. Use MO electron configurations to explain why NO ionizes more easily than either N2 or CO.


7.5 Three-Centre Orbitals

Learning objective:

Describing the bonding in three-atom systems


7.5 Three-Centre Orbitals

Three atom systems can delocalize electrons over all of three atoms.

Consider Ozone, O3:


Nonbonding Molecular Orbitals

If a lone pair is spread over both outer atoms, but not across the inner atom, it is a nonbonding MO.

A delocalized system is present whenever p orbitals on more than two adjacent atoms are in position for side-by-side overlap


A Composite Model of Bonding

1. Construct the sigma bonding framework using hybrid orbitals for inner atoms and atomic orbitals for outer atoms. Any hybrid orbital not used to form bonds contains lone pairs of electrons

2. If the molecule contains multiple bonds, construct the bonding system, using MO theory. Watch for resonance structures, which signal the presence of delocalized electrons.

3. Place one pair of valence electrons in any atomic orbital that is not used in hybridization or in the system.

4. Sum the electrons allocated in steps 1 – 3. The result must match the total number of valence electrons used in the Lewis structure. In addition, a complete bonding description must account for all the valence orbitals.


Example 7 - 11

The acetate anion (CH3CO2-) forms when acetic acid, the

acid present in vinegar is treated with hydroxide ion:

CH3CO2H + OH- → CH3CO2- + H2O

Use the four guidelines of the composite model of bonding to describe the bonding of this anion. Sketch the bonding system and the occupied orbitals.


Carbon Dioxide


7.6 Extended Systems

Learning objective:

Describe the bonding in extended systems


7.6 Extended Systems

Can be long chains or compact clusters of ringsConsider butadiene:


Example 7 - 12

Describe the bonding of methyl methacrylate. This compound is an important industrial chemical, used mainly to make plastics such as poly(methyl methacrylate) (PMMA).


The Carbonate Ion


Larger Delocalized Systems


Larger Delocalized Systems


Benzene, C6H6

7.7 Band Theory of Solids

Learning objective:

Explain such properties as electrical conductivity and colour of metals, non-metals, and metalloids in terms of band theory


7.7 Band Theory of Solids

An extension of the delocalized orbital ideasAccounts for the properties of metals, and explains the

properties of metalloids such as Silicon


Delocalized Orbitals in Lithium Metal


Electrical Conductivity

Close spacing between orbital energies result in energy bands.

When a potential is applied, electrons flow through the metal from the negative end to the positive end (through the orbitals)


Insulators vs Conductors

Band Gap (Eg) – the energy difference between the filled bonding orbitals and the empty antibonding orbitals.


Metalloids

Semiconductors – compounds that become conductors in the presence of an appropriate energy source.


Example 7 - 13

“Electrical eye” door openers use photoconductors that respond to infrared light with a wavelength of 1.5 m. Which is suitable for photoconductors operating at this wavelength, germanium (Eg = 64 kJ/mol) or silicon (Eg = 105 kJ/mol)?


Doped Semiconductors

Doped semiconductor: when a specific impurity is added deliberately to a pure substance, the result is a material with nearly the same band structure as the pure compounds, but with different electron populations. n-type semiconductor: Si or Ge doped with atoms from Group

15, more electrons than the pure substance p-type semiconductor: Si or Ge doped with atoms from Group

13, less electrons than the pure substance.



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Chapter 7 Theories of Chemical Bonding 7.1 Localized Bonds 7.2 Hybridization of Atomic Orbitals 7.3...

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Transcript of Chapter 7 Theories of Chemical Bonding 7.1 Localized Bonds 7.2 Hybridization of Atomic Orbitals 7.3...