|MadSci Network: Chemistry|
Admin note:Please see additional information provided by James Griepenburg at the bottom of this page!
Sorry for taking so long to get back on this question but I wanted to make sure that I did indeed give the correct answer. According to a book by Turro, “the ground state of molecular oxygen is predicted (Hund’s rule) and found to be a triplet” (p. 584). In this triplet state, oxygen only has a single bond. The two electrons, which we might think would form the second bond, are actually found in separate orbitals, and thus do not contribute as a bonding pair.
However, for the junior high level the explanation for why a single bond instead of a double bond is found might not be easy to grasp. Hund’s rule states that an atom or molecule in the ground state will adopt a configuration with the greatest numbers of unpaired electrons (electrons in different orbitals). For oxygen, this results in a single bond. The fact that these unpaired electrons need not have paired spins results in the triplet state.
Lewis dot structures predict a double bond in an oxygen molecule (the octet rule is satisfied), and this might be easier for students to understand.
I consulted the following book, which includes excellent diagrams, though the explanations might be too complicated for junior high students(!).
Turro, Nicholas J. Modern Molecular Photochemistry. Benjamin Cummings Publishing. Menlo Park, CA: 1978.
James Greipenburg adds the following:
The answer given is erroneous. Molecular orbital theory gives a picture that is much closer to reality. Using sp hybrid orbitals there are in order of energy 3 bonding orbitals [sp sigma], [p pi], [p pi] each of which contains 2 electrons making an effective triple bond The 2 remaining sp orbitals are non bonding and hold 2 electrons each. The remaining 2 electrons each go into one of the [p pi] antibonding orbitals [Hund's rule] negating one bond giving a molecule with diradical[triplet] structure and double bond strength.
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