AP Chem Unit 9: Covalent Bonding Orbitals

AP Chem Unit 9: Covalent Bonding Orbitals

Guided Notes

Main Ideas and Sections:

· Hybridization and the Localized Electron Model

· The Molecular Orbital Model

· Bonding in Homonuclear Diatomic Molecules

· Bonding in Heteronuclear Diatomic Molecules

· Combining the Localized Electron and MO models

Hybridization and the Localized Electron Model

Localized Electron Model

The arrangement of valence electrons is represented by the Lewis structure or structures, and the molecular geometry can be predicted from the VSEPR model.

Hybridization

In general we assume that bonding involves only the valence orbitals.

o Carbon is said to undergo sp3 hybridization or is sp3 hybridized because is uses one s orbital and three p orbitals to form four identical bonding orbitals.

o The four sp3 orbitals are identical in shape each one having a large lobe and a small lobe. The four orbitals are oriented in space so that the large lobes form a tetrahedral arrangement.

sp3 hybridized

sp2 Hybridization

Example: Ethylene (C2H4) is commonly used in plastics and has a C=C double bond. Each carbon uses sp2 hybridization in this molecule because a double bond acts as one effective pair.

Example: sp2 Hybridization: Ethylene

Multiple Bonds

· Single bonds are sigma bonds (σ) and the electron pair is shared in an area centered on a line running between the atoms. These are hybridized bonding orbitals.

Example: sp2 Hybridization: Ethylene

sp Hybridization

· CO2 is sp hybridized

Example: sp Hybridization: CO2

Example: sp Hybridization: N2

sp3d Hybridization

· PCl5 has sp3d hybridization and is trigonal bipyramidal.

Example: sp3d Hybridization: PCl5

· Each chlorine atom displays a tetrahedral arrangement around the atom.

sp3d2 Hybridization

· SF6 has sp3d2 hybridization.

· Example Problem: How is the xenon atom in XeF4 hybridized?

Localized Electron Summary

· Draw the Lewis Structure

· Determine the arrangement of electron pairs using the VSEPR model.

· Specify the hybrid orbitals needed to accommodate the electron pairs.

Effective Pairs and Their Spatial Arrangement

The Molecular Orbital Model

Molecular Orbital Model

The localized electron model works very well with the prediction of structure and bonding of molecules, but the electron correlation problem still exists.

· The Molecular Orbital model helps us to deal with the molecular problem.

Molecular Orbitals

Molecular orbitals (MOs) have many of the same characteristics as atomic orbitals. Two of the most important are:

· The square of the MO’s wave function indicates electron probability.

For simplicity we will first look at the H2 molecule.

· The wave phases of the atomic orbitals combine/overlap. Since electrons move in wave functions, this causes constructive and destructive interference in the wave pattern.

MOs

Two MOs exist for H2:

MOs

The type of electron distribution described in these MOs is called sigma as in the localized electron model. MO1 and MO2 are sigma (σ) molecular orbitals.

The energy level of the bonding MO is lower and more stable than that of the antibonding MO. Since molecule formation favors the lowest energy state, this provides the driving force for molecule formation of H2. This is called probonding.

If two electrons were forced to occupy the higher-energy MO2 this would be anti-bonding and the lower energy of the separated atoms would be favored.

Bonding and Antibonding

Labels are given to MOs indicate their symmetry (shape), the parent atomic orbitals, and whether they are bonding or antibonding.

Molecular electron configurations can be written in much the same way as atomic (electron) configurations. Since the H2 molecule has two electrons in the σ1s molecular orbital, the electron configuration is: σ1s2

From this molecular electron configuration, we can determine a molecules stability.

Example: Would H2- be stable?

o (σ1s )2 ( σ1s*) 1

o The key idea is that H2- would exist if it were a lower energy than its separated parts. Two electrons are in bonding and one is in antibonding. Since more electrons favor bonding H2- is formed.

o This also is a good indicator of bond strength. H2 has a stronger bond than H2-. The net lowering of the bonding electrons by one is a direct relationship to bond strength. H2 is twice as strong.

Bond Order

To indicate bond strength, we use the concept of bond order.

Example: H2 has a bond order of 1

H2- has a bond order of ½

Bond order is an indication of bond strength because it reflects the difference between the number of bonding electrons and the number of antibonding electrons.

Bonding in Homonuclear Diatomic Molecules

Homonuclear Bonding

When looking at bonding beyond energy level 1, we need to consider what orbitals are overlapping and therefore bonding.

· Li2 has electrons in the 1s and 2s orbitals; the 2s orbitals are much larger and overlap, but the 1s orbitals are smaller and do not overlap.

· To participate in molecular orbitals, atomic orbitals must overlap in space. This means that only the valence orbitals of the atoms contribute significantly to the molecular orbitals of a particular molecule.

Example: What is the molecular electron configuration and bond order of Li2?

Example: What is the molecular electron configuration and bond order of Be2?

MOs from p orbitals

· p orbitals must overlap in such a way that the wave patterns produce constructive interference. As with the s orbitals, the destructive interference produces a node in the wave pattern and decreases the probability of bonding.

· When the parallel p orbitals are combined with the positive and negative phases matched, constructive interference occurs, giving a bonding π orbital. When orbitals have opposite phases, destructive interference results in an antibonding π orbital.

MOs from p orbitals continued

· Since the electron probability lies above and below the line between the nuclei (with parallel p orbitals), the stability of a π molecular bonding orbital is less than that of a σ bonding orbital. Also, the antibonding π MO is not as unstable as the antibonding σ MO. The energies associated with the orbitals reflect this stability.

Example: B2

Exceptions

· B2, C2, and N2 molecules use the same set of molecular orbitals that we expect but some mixing of orbital energies occurs. The s and p atomic orbitals mix or hybridize in a way that changes some MO energy states. This affects filling order and pairing of electrons.

Paramagnetism

Most materials have no magnetism until they are placed in a magnetic field. However, in the presence of such a field, magnetism of two types can be induced:

· Paramagnetism –

· Diamagnetism –

Paramagnetism

· Any substance that has both paired and unpaired electrons will exhibit a net paramagnetism since the effect of paramagnetism is much stronger than that of diamagnetism.

Summary

· Comparison of bond orders between different molecules cannot predict bond energies of different molecules.

o B2 and F2 both have bond order of 1 but bond energies are very different. B-B is a much stronger bond.

· N2 has a bond order of 3 and has a very large bond energy. N2 is a very stable molecule and is used to drive powerful reactions.

Example Problem: For O2, O2+, and O2-, give the MO electron configuration and the bond order for each. Which has the strongest bond?

MO order with mixing / MO order without mixing
σ2p* / σ2p*
π2p* / π2p*
σ2p / π2p
π2p / σ2p
σ2s* / σ2s*
σ2s / σ2s
σ1s* / σ1s*
σ1s / σ1s
MO order with mixing / MO order without mixing
σ2p* / σ2p*
π2p* / π2p*
σ2p / π2p
π2p / σ2p
σ2s* / σ2s*
σ2s / σ2s
σ1s* / σ1s*
σ1s / σ1s
MO order with mixing / MO order without mixing
σ2p* / σ2p*
π2p* / π2p*
σ2p / π2p
π2p / σ2p
σ2s* / σ2s*
σ2s / σ2s
σ1s* / σ1s*
σ1s / σ1s

Example Problem: Use the molecular orbital model to predict the bond order and magnetism of each of the following molecules: Ne2 and P2

Bonding in Heteronuclear Diatomic Molecules

Heteronuclear Molecules

When dealing with different atoms within diatomic molecules we can still use the MO model to determine bond order and magnetism

Example: NO

Example Problem: Use the MO model to predict the magnetism and bond order of the NO+ and CN- ions.

Heteronuclear Diatomics

What happens with the diatomic molecules are very different?

Combining the Localized Electron and MO models

Resonance

When a molecule has resonance. It is usually a double bond that can have different positions around the molecule.

· Benzene: All C-C bonds are known to be equivalent and the molecule has resonance

· Benzene:

Resonance continued

· Benzene: The p orbitals are perpendicular to the plane and form π bonds above and below the plane. The electrons in the π bonds delocalize and give six equivalent C-C bonds that give the structure true resonance.

NO3-