Molecular Orbital Theory (10.2.4) Flashcards
• Lewis dot structures, VSEPR theory, and valence bond theory all rely on atomic orbitals, and consequently fail to predict some of the properties of molecules.
• Lewis dot structures, VSEPR theory, and valence bond theory all rely on atomic orbitals, and consequently fail to predict some of the properties of molecules.
• Molecular orbital theory (MO theory) can be used to explain why certain molecules do not form.
• Molecular orbital theory (MO theory) can be used to explain why certain molecules do not form.
Lewis dot structures, VSEPR theory, and valence
bond theory all rely on atomic orbitals, and
consequently fail to predict some of the properties of
molecules. Molecular orbital theory (MO theory)
provides a better picture of how bonds form.
Adding the wave functions for the two 1s orbitals
of the atoms of dihydrogen (H2) in phase shows
that the greatest electron probability density lies
between the two nuclei. This high electron
probability density is a bonding orbital (σ), and is
lower in energy than the atomic orbitals.
Adding the wave functions for the two 1s orbitals of
the atoms of H2 out of phase results in the formation
of two lobes of high electron probability density on
opposite sides of the molecule with a node of
electron probability density between the nuclei.
These two lobes are called an antibonding orbital
(σ). The antibonding orbital (σ) is higher in energy
than the individual atomic orbitals.
The two electrons of H2 reside in the low energy
bonding orbital (σ), and the antibonding orbital (σ)
is empty. Since this bonding orbital is lower in
energy than the atomic orbitals of the free atoms,
the molecule is stable.
Molecular orbital theory (MO theory) can be used to
explain why certain molecules do not form.
For example, molecular orbital theory predicts that
He2 will not be a stable molecule. The extra energy
needed to place electrons in the antibonding orbital
(σ) is greater than the energy gained by placing
electrons in the bonding orbital (σ). Consequently,
helium exists in nature as a monatomic gas.
Lewis dot structures, VSEPR theory, and valence
bond theory all rely on atomic orbitals, and
consequently fail to predict some of the properties of
molecules. Molecular orbital theory (MO theory)
provides a better picture of how bonds form.
Adding the wave functions for the two 1s orbitals
of the atoms of dihydrogen (H2) in phase shows
that the greatest electron probability density lies
between the two nuclei. This high electron
probability density is a bonding orbital (σ), and is
lower in energy than the atomic orbitals.
Adding the wave functions for the two 1s orbitals of
the atoms of H2 out of phase results in the formation
of two lobes of high electron probability density on
opposite sides of the molecule with a node of
electron probability density between the nuclei.
These two lobes are called an antibonding orbital
(σ). The antibonding orbital (σ) is higher in energy
than the individual atomic orbitals.
The two electrons of H2 reside in the low energy
bonding orbital (σ), and the antibonding orbital (σ)
is empty. Since this bonding orbital is lower in
energy than the atomic orbitals of the free atoms,
the molecule is stable.
Molecular orbital theory (MO theory) can be used to
explain why certain molecules do not form.
For example, molecular orbital theory predicts that
He2 will not be a stable molecule. The extra energy
needed to place electrons in the antibonding orbital
(σ) is greater than the energy gained by placing
electrons in the bonding orbital (σ). Consequently,
helium exists in nature as a monatomic gas.
Which statement about molecular orbitals is not true?
Electrons within a molecule are best described as being in atomic, not molecular, orbitals. (C)
We are learning that electrons within a molecule can sometimes be best described or accounted for by examining their aggregate, molecular orbital behavior.
Look at the plot for the formation of a molecular orbital for dihelium molecule (a molecule composed of two helium atoms).
Which statement about this diagram is not correct?
The electrons around each nucleus are distributed evenly. (C)
If you look at the distributions (the two circles) in the diagram, you can see that each distribution is not symmetrical around the “+.” This means that most of the electrons in the atom are distributed toward the outside of the atom (i.e., as far away from the other atom as possible).
Which statement best explains the difference between the electron distributions of a bonding orbital (e.g., H2 ) and an anti-bonding orbital (e.g., He2 )?
In a bonding orbital situation, the electron distribution is between the nuclei, because both nuclei want to share the electrons. In an anti- bonding circumstance, the electron distribution is around each atom. There is a void between the nuclei and the nuclei repel each other. (C)
The electrons in bonding orbitals tend to be distributed most heavily between the atoms. The opposite is true when anti-bonding orbitals are involved. There is a net void between the atoms.
Look at the plot for the formation of a molecular orbital for a hydrogen molecule.
Which statement about this diagram is not correct?
The greatest electron density will be around the nuclei of either atom. (C)
If you look at the atoms individually (i.e., as if they are not interacting), then this would be true. But these two atoms interact, so the electron density will be greatest between the two atoms, not around each individual atom.
We earlier investigated the probability density of a 1s orbital and discovered that there was a greater probability for electrons to be found near the nucleus. Now we are investigating molecules. Which statement about probability densities is true?
A molecule is not spherical. Therefore, we describe electron density along internuclear axes, because more than one atom is involved in a molecule. (D)
This is a simple, correct statement about how we approach investigating electron probability density in a molecule.
The energy diagram shows the energies involved in bonds of bonding and anti-bonding orbitals. The bonding levels for two unique 1s orbitals are equal and are also shown for your reference.
Which statement about this diagram is not correct?
Sites A and B represent the bonding orbitals in di-hydrogen and site A represents the bonding orbital in di-helium. (B)
The opposite is actually true. Site A represents the bonding orbital for di-hydrogen, and sites A and B represent the theoretical orbitals in di-helium.
The plot shows the atomic wave functions for two atoms involved in the formation of a molecule.
Which statement best describes this graph?
This plot shows the two atomic wave functions for two 1s orbitals involved in the formation of the H2 molecule. (B)
We are looking at two identical atomic wave functions for a 1s orbital in the hydrogen atom.
Which statement about electrons and orbitals is not correct?
From now on, we will explain electron behavior in terms of molecular orbital levels, not atomic orbital levels. (B)
There are many situations where describing electron behavior on the atomic orbital level is more appropriate. However, we now have learned that there are also situations in which it is more appropriate to describe electron orbitals on the molecular (rather than atomic) level.
Which statement about scientific models is not correct?
Models have limitations that sometimes make them inexact. This significantly limits their usefulness. (C)
By understanding their limitations, we can use models in many significant ways to predict and understand the behavior of matter.
Which statement about our standard definition for a chemical bond is not true?
In a chemical bond, the two nuclei get as close together as possible so that their electrons are completely shared. (B)
“As close together as possible” implies one atom right on top of the other atom. In this situation, internuclear repulsion occurs. This results in the atoms being highly unstable, with large potential energies.
