Using the molecular orbital theory, predict how many unpaired electrons are in the O2+ ion.

Question

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The question asks us to predict the number of unpaired electrons in the O2+ ion using molecular orbital theory.
Option 1: 0. This would imply all electrons are paired in all occupied MOs. In O2+ there are electrons occupying both bonding and antibonding MOs with at least one degenerate set that remains partially filled, so having zero unpaired electrons is not correct.
Option 2: 2. For neutral O2, there are two unpaired electrons in the antibonding pi* set. When you remove one electron to form O2+, you reduce the number of unpaired electrons by one, not by two, so this option overcounts the unpaired electrons for the cation.
Option 3: 4. Four unpaired electrons would require four singly occupied MOs with unpaired spins, which is not consistent with the MO filling pattern for O2 or O2+. The valence MO diagram for O2-type species does not provide enough electrons to yield four separate unpaired electrons in the ground state.
Option 4: 1. This matches the standard molecular orbital treatment: neutral O2 has two unpaired electrons in the degenerate π*2p orbitals. Removing one electron to form O2+ leaves only one unpaired electron, since one of the previously unpaired electrons is removed from the antibonding π* set, leaving a single unpaired electron overall.
Option 5: 3. THREE unpaired electrons would require three singly occupied MOs in the valence set, which is not supported by the electron configuration of O2+ under MO theory; the total electron count and orbital filling do not permit three unpaired electrons in the ground state.
In summary, considering O2 neutral has two unpaired electrons in the π*2p orbitals and removing one electron to form O2+ reduces the count by one, yielding exactly one unpaired electron in O2+.

Using the molecular orbital theory, predict how many unpaired electrons are in the O2+ ion.

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