Chapter 9 - Covalent Bonding: Orbitals
The arrangement of valence electrons is represented by the Lewis structure and the molecular geometry can be predicted from the VSEPR model
The valence orbitals are the orbitals associated with the highest principal quantum level that contains electrons on a given atom
Hybridization is a modification of the localized electron model to account for the observation that atoms often seem to use special atomic orbitals in forming molecules
sp3 hybridization gives a tetrahedral set of orbitals
Whenever a set of equivalent tetrahedral atomic orbitals is required by an atom, this model assumes that the atom adopts a set of sp3 orbitals; the atom becomes sp3 hybridized.
A double bond acts as one effective electron pair
Whenever an atom is surrounded by three effective pairs, a set of sp2 hybrid orbitals is required.
Two effective pairs around an atom will always require sp hybridization of that atom
More rigorous theoretical models of CO2 indicate that each of the oxygen atoms uses two p orbitals simultaneously to form the pi bonds to the carbon atom, thus leading to unusually strong CPO bonds.
A set of five effective pairs around a given atom always requires a trigonal bipyramidal arrangement, which in turn requires d2 sp3 hybridization of that atom
D2 sp3 hybridization gives six orbitals arranged octahedrally.
In applying the localized electron model, we must remember not to overemphasize the characteristics of the separate atoms
In the same vein, it is not the orbitals in the isolated atom that matter, but which orbitals the molecule requires for minimum energy
Molecular orbitals have many of the same characteristics as atomic orbitals
Two of the most important are that they can hold two electrons with opposite spins and that the square of the molecular orbital wave function indicates electron probability
The electron probability of both molecular orbitals is centered along the line passing through the two nuclei
In the molecule, only the molecular orbitals are available for occupation by electrons
MO1 is lower in energy than the 1s orbitals of free hydrogen atoms, while MO2 is higher in energy than the 1s orbitals
Bonding molecular orbital: Lower in energy than the atomic orbitals of which it is composed
Anti bonding molecular orbital: Higher in energy than the atomic orbitals of which it is composed
Bonding will result if the molecule has lower energy than the separated atoms
To indicate bond strength, we use the concept of bond order. Bond order is the difference between the number of bonding electrons and the number of anti-bonding electrons divided by 2.
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
Larger bond order means greater bond strength
9.3 Bonding in Homonuclear Diatomic Molecules
To participate in molecular orbitals, atomic orbitals must overlap in space
Paramagnetism causes the substance to be attracted into the inducing magnetic field
Studies have shown that paramagnetism is associated with unpaired electrons and diamagnetism is associated with paired electrons
Diamagnetism causes the substance to be repelled from the inducing magnetic field.
There are definite correlations between bond order, bond energy, and bond length
As the bond order predicted by the molecular orbital model increases, the bond energy increases, and the bond length decreases
Comparison of the bond energies of the B2 and F2 molecules indicates that bond order cannot automatically be associated with a particular bond energy
The molecular orbital model correctly predicts oxygen’s paramagnetism, while the localized electron model predicts a diamagnetic molecule
****When the two atoms of a diatomic molecule are very different, the energy level diagram for homonuclear molecules can no longer be used
A new diagram must be devised for each molecule.
The molecular orbital containing the bonding electron pair shows greater electron probability close to the fluorine
This causes the fluorine atom to have a slight excess of negative charge and leaves the hydrogen atom partially positive.
Thus, the molecular orbital model accounts in a straightforward way for the different electronegativities of hydrogen and fluorine and the resulting unequal charge distribution
Even with resonance included, the localized electron model does not describe molecules and ions
It is really the bond that has different locations in the various resonance structures.
In molecules that require resonance, it is the bonding that is most clearly delocalized.
Molecules that require the concept of resonance in the localized electron model can be more accurately described by combining the localized electron and molecular orbital models
The arrangement of valence electrons is represented by the Lewis structure and the molecular geometry can be predicted from the VSEPR model
The valence orbitals are the orbitals associated with the highest principal quantum level that contains electrons on a given atom
Hybridization is a modification of the localized electron model to account for the observation that atoms often seem to use special atomic orbitals in forming molecules
sp3 hybridization gives a tetrahedral set of orbitals
Whenever a set of equivalent tetrahedral atomic orbitals is required by an atom, this model assumes that the atom adopts a set of sp3 orbitals; the atom becomes sp3 hybridized.
A double bond acts as one effective electron pair
Whenever an atom is surrounded by three effective pairs, a set of sp2 hybrid orbitals is required.
Two effective pairs around an atom will always require sp hybridization of that atom
More rigorous theoretical models of CO2 indicate that each of the oxygen atoms uses two p orbitals simultaneously to form the pi bonds to the carbon atom, thus leading to unusually strong CPO bonds.
A set of five effective pairs around a given atom always requires a trigonal bipyramidal arrangement, which in turn requires d2 sp3 hybridization of that atom
D2 sp3 hybridization gives six orbitals arranged octahedrally.
In applying the localized electron model, we must remember not to overemphasize the characteristics of the separate atoms
In the same vein, it is not the orbitals in the isolated atom that matter, but which orbitals the molecule requires for minimum energy
Molecular orbitals have many of the same characteristics as atomic orbitals
Two of the most important are that they can hold two electrons with opposite spins and that the square of the molecular orbital wave function indicates electron probability
The electron probability of both molecular orbitals is centered along the line passing through the two nuclei
In the molecule, only the molecular orbitals are available for occupation by electrons
MO1 is lower in energy than the 1s orbitals of free hydrogen atoms, while MO2 is higher in energy than the 1s orbitals
Bonding molecular orbital: Lower in energy than the atomic orbitals of which it is composed
Anti bonding molecular orbital: Higher in energy than the atomic orbitals of which it is composed
Bonding will result if the molecule has lower energy than the separated atoms
To indicate bond strength, we use the concept of bond order. Bond order is the difference between the number of bonding electrons and the number of anti-bonding electrons divided by 2.
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
Larger bond order means greater bond strength
9.3 Bonding in Homonuclear Diatomic Molecules
To participate in molecular orbitals, atomic orbitals must overlap in space
Paramagnetism causes the substance to be attracted into the inducing magnetic field
Studies have shown that paramagnetism is associated with unpaired electrons and diamagnetism is associated with paired electrons
Diamagnetism causes the substance to be repelled from the inducing magnetic field.
There are definite correlations between bond order, bond energy, and bond length
As the bond order predicted by the molecular orbital model increases, the bond energy increases, and the bond length decreases
Comparison of the bond energies of the B2 and F2 molecules indicates that bond order cannot automatically be associated with a particular bond energy
The molecular orbital model correctly predicts oxygen’s paramagnetism, while the localized electron model predicts a diamagnetic molecule
****When the two atoms of a diatomic molecule are very different, the energy level diagram for homonuclear molecules can no longer be used
A new diagram must be devised for each molecule.
The molecular orbital containing the bonding electron pair shows greater electron probability close to the fluorine
This causes the fluorine atom to have a slight excess of negative charge and leaves the hydrogen atom partially positive.
Thus, the molecular orbital model accounts in a straightforward way for the different electronegativities of hydrogen and fluorine and the resulting unequal charge distribution
Even with resonance included, the localized electron model does not describe molecules and ions
It is really the bond that has different locations in the various resonance structures.
In molecules that require resonance, it is the bonding that is most clearly delocalized.
Molecules that require the concept of resonance in the localized electron model can be more accurately described by combining the localized electron and molecular orbital models
