Structure & Reactivity in Chemistry
MO7. Experimental Evidence for Molecular Orbital
The molecular orbital picture of dioxygen differs from the Lewis picture.
Both models predict an oxygen-oxygen double bond, but one model suggests
unpaired electrons whereas the other indicates an electron-paired system.
Often, there is experimental evidence available to check the reliability of
predictions about structure. These data include measurements of bond lengths
and bond strengths as well as magnetic properties.
Bond dissociation energy data tell us how difficult it is to separate one
atom from another in a molecule. Bond order is one of the factors that
influences bond strength. Thus, measuring a bond dissociation energy is one
way to confirm that dioxygen really does contain an oxygen-oxygen double
First, we need something to compare it to. Peroxides (such as hydrogen
peroxide, H2O2, or sodium peroxide, Na2O2) probably contain oxygen-oxygen
single bonds, according to their Lewis structures. These bonds are
relatively weak, costing about 35 kcal/mol to break. In contrast, the bond
in dioxygen costs about 70 kcal/mol to break. Its bond is about twice as
strong; it is a double bond.
- Bond dissociation energies can be used to determine how many bonds there
are between two atoms.
Bond dissociation energies can be complicated to measure. They require a
comparison of energy changes in numerous chemical reactions so that the
energy change resulting from cleavage of a specific bond can be inferred. In
contrast, infrared absorption frequencies are easy to measure. They simply
require shining infrared light through a sample and measuring what
frequencies of the light are absorbed by the material. (A related technique,
Raman spectroscopy, gives similar information by measuring subtle changes in
the frequency of laser light that is scattered off a sample). The
frequencies absorbed depend on what bonds are present in the material. These
frequencies vary according to two basic factors: the weights of the atoms at
the ends of the bond, and the strength of the bond between them. The
stronger the bond, the higher the absorption frequency.
Peroxides absorb infrared light at around 800 cm-1 (this unusual frequency
unit is usually pronounced "wavenumbers"). Dioxygen absorbs infrared light
around 1300 cm-1 . Since the atoms at the ends of the bond in both peroxide
and dioxygen are oxygens, we can be sure that this difference in frequency
is not due to a difference in mass. It is due to a difference in bond
strength. The bond in dioxygen is much stronger than the O-O bond in
peroxide, because the former is a double bond and the latter is a single
- Vibrational spectroscopies (IR and Raman spectroscopy) can give
information about the bond order between two atoms.
A third measure of bond order is found in bond length measurements. The more
strongly bound two atoms are, the closer they are together. An O=O bond
should be shorter than an O-O bond. Bond lengths can be measured by
microwave spectroscopy (usually for gas-phase molecules), in which
frequencies absorbed depend on the distance between the molecules.
Alternatively, bond lengths can be measured by x-ray crystallography. X-rays
can be diffracted through crystals of solid materials. The interference
pattern that is produced can be mathematically decoded to produce a
three-dimensional map of where all the atoms are in the material. The
distances between these atoms can be measured very accurately.
The O-O bond in peroxides are about 1.49 Angstroms long (an Angstrom is
10-10 m; this unit is often used for bond lengths because it is a convenient
size for this task. Covalent bonds are generally one to three Angstroms
long). The O-O bond in dioxygen is about 1.21 A long. The O-O bond in
dioxygen is shorter and stronger than in a peroxide.
- Bond length data provides insight into the bond order.
In addition to bond order, there is the question of electron pairing in
dioxygen. The Lewis structure suggests electrons are paired in dioxygen. The
molecular orbital picture suggests two unpaired electrons.
Compounds with paired electrons are referred to as diamagnetic. Those with
unpaired electrons are called paramagnetic. Paramagnetic substances interact
strongly with magnetic fields.
It turns out that oxygen does interact with a magnetic fields. A sample of
liquid-phase oxygen can be held between the poles of a magnet. Oxygen has
unpaired electrons. This finding is consistent with molecular orbital
theory, but not with simple Lewis structures. Thus, MO theory tells us
something that the Lewis picture cannot.
- Magnetic information, and measurements of magnetism, give us experimental
evidence of spin states.
- We can tell if electrons are paired, unpaired, and how many unpaired spins
A final important source of experimental data is photoelectron spectroscopy.
Photoelectron spectroscopy gives information about the electron energy
levels in an atom or compound. In this technique, gas-phase molecules are
subjected to high-energy electromagnetic radiation, such as ultraviolet
light or X-rays. Electrons are ejected from various energy levels in the
molecule, and the binding energies of electrons in those levels is
determined. Thus, photoelectron spectroscopy provides verification for
exactly the sort of information that quantitative molecular orbital
calculations are designed to deliver.
- Photoelectron spectroscopy tells how much energy is needed to remove
electrons from various energy levels in a molecule.
- This technique gives us an accurate experimental picture of the energy
levels that we predict with molecular orbital calculations.
Problem MO 7.1
In the previous section you were asked to draw MO
diagrams for some molecules. Determine whether the molecules in problems
MO6.3-6.8 are paramagnetic or diamagnetic.
Use a MO diagram to determine which species in the
following pairs will have the longer bond. Give an explanation for your choice.
a. N2 or N2+
b. N2 or N2-
Use a MO diagram to determine which species in the
following pairs will have the stronger bond. Give an explanation for your
a. O2 or O2+
b. O2 or O2-
This site is written and maintained by Chris P. Schaller, Ph.D., College of Saint Benedict / Saint John's
University (with contributions from other authors as noted). It is freely
available for educational use.
Structure & Reactivity in Organic, Biological and Inorganic Chemistry by Chris Schaller is licensed under a Creative Commons Attribution-NonCommercial 3.0 Unported License.
Send corrections to firstname.lastname@example.org
Back to Molecular Orbital
Back to Structure & Reactivity in Chemistry