Answer
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Step 1:: Determine the total number of valence electrons in the $\mathrm{O}_{2}{ }^{-}$ anion.
The $\mathrm{O}_{2}{ }^{-}$ anion has 2 oxygen atoms, so there are a total of 2 * 6 = 12 valence electrons.
Oxygen (O) is in period 2 and group 16 of the periodic table, so it has 6 valence electrons.
Step 2:: Draw a skeletal structure for the molecule.
\mathrm{O}=\mathrm{O}
Place the two oxygen atoms next to each other, sharing a double bond:
Step 3:: Distribute the remaining valence electrons around the atoms.
\mathrm{O}:\stackrel{.}{e}:\stackrel{.}{e}:\mathrm{O}
Add the remaining valence electrons as lone pairs on the oxygen atoms. Since there are 12 valence electrons in total, and 4 are used in the double bond, there are 12 - 4 = 8 remaining electrons. Divide these 8 electrons into 4 lone pairs (2 electrons per lone pair) and add them to the oxygen atoms:
Step 4:: Calculate formal charges for each atom.
\text{Formal charge} = 6 - 4 - 4 = -2
Formal charge is calculated as: For each oxygen atom, there are 6 valence electrons, 4 non-bonding electrons (2 lone pairs), and 4 bonding electrons (2 from the double bond). Plugging these values into the formula, we get:
Step 5:: Analyze the octet-rule exception.
The octet rule is not obeyed for the oxygen atoms, as they each have 10 electrons in their outermost shell (6 valence electrons + 4 from the double bond). This is a hypervalent molecule, which is an exception to the octet rule.
Step 6:: Summarize the Lewis structure and octet-rule exception.
\mathrm{O}:\stackrel{.}{e}:\stackrel{.}{e}:\mathrm{O}^{-}
The octet-rule exception is hypervalency, with each oxygen atom having 10 electrons in its outermost shell. The formal charge on each oxygen atom is - 2. There are 4 lone pairs of electrons around the molecule.
Final Answer
\mathrm{O}:\stackrel{.}{e}:\stackrel{.}{e}:\mathrm{O}^{-}
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