Understanding the Superoxide Ion \(O_2^{-}\)

The question asks about the properties of the superoxide ion, represented as \(O_2^{-}\). Let's analyze each option based on the chemical characteristics of this species.

We will examine the bond length, oxidizing capability, IR activity, and basicity of the \(O_2^{-}\) ion.

\(O_2^{-}\) Bond Length Comparison

To compare the O-O bond length in \(O_2^{-}\) to that in \(O_2\), we can look at their bond orders using Molecular Orbital (MO) theory.

• The \(O_2\) molecule has 16 electrons. Its MO configuration leads to a bond order of 2.
• The \(O_2^{-}\) ion has 17 electrons (16 from \(O_2\) plus one extra electron in an antibonding orbital). Its MO configuration leads to a bond order of 1.5.

Bond order is inversely proportional to bond length. A higher bond order means a stronger and shorter bond. Since \(O_2^{-}\) has a lower bond order (1.5) than \(O_2\) (2), the O-O bond in \(O_2^{-}\) is expected to be longer than that in \(O_2\). Therefore, the statement that \(O_2^{-}\) has a shorter O-O bond length than that in \(O_2\) is incorrect.

Oxidizing Agent Strength of \(O_2^{-}\)

An oxidizing agent is a species that gains electrons (gets reduced) and causes another species to be oxidized. The strength of an oxidizing agent is related to its tendency to gain electrons. Standard reduction potentials (\(E^\circ\)) can indicate oxidizing strength; a more positive \(E^\circ\) generally means it's a stronger oxidizing agent.

• Reduction of \(O_2\): \(O_2 + e^{-} \rightarrow O_2^{-}\) (\(E^\circ \approx -0.33 \text{ V}\))
• Reduction of \(O_2^{-}\): \(O_2^{-} + e^{-} \rightarrow O_2^{2-}\) (\(E^\circ \approx +0.89 \text{ V}\) in acidic conditions, positive values in basic conditions)

Comparing the reduction potentials shows that the tendency of \(O_2^{-}\) to gain an electron and form peroxide (\(O_2^{2-}\)) is significantly higher (more positive potential) than the tendency of \(O_2\) to gain an electron and form superoxide (\(O_2^{-}\)). This indicates that \(O_2^{-}\) can act as a stronger oxidizing agent than \(O_2\) in suitable conditions, being more readily reduced itself to \(O_2^{2-}\) than \(O_2\) is to \(O_2^{-}\). Thus, the statement that \(O_2^{-}\) is a stronger oxidizing agent than \(O_2\) is correct.

IR Activity of \(O_2^{-}\)

A molecule is infrared (IR) active if its vibration causes a change in its dipole moment. Isolated diatomic molecules like \(O_2\) and \(O_2^{-}\) are linear. Homonuclear diatomic molecules like \(O_2\) have no dipole moment and their symmetric stretching vibration does not create one, making them IR inactive. The superoxide ion \(O_2^{-}\) is also a homonuclear diatomic species (O-O bond) and is linear. Its fundamental vibration is a symmetric stretch which, in an isolated ion, does not change the dipole moment (which is zero along the bond axis). Therefore, the isolated \(O_2^{-}\) ion is generally considered IR inactive. The statement that \(O_2^{-}\) is IR active is incorrect under standard assumptions for isolated species.

Proton Abstraction by \(O_2^{-}\)

\(O_2^{-}\) is the conjugate base of the hydroperoxyl radical (\(HO_2\)). The equilibrium is \(HO_2 \rightleftharpoons H^{+} + O_2^{-}\). The pKa of \(HO_2\) is approximately 4.8. This pKa value suggests that \(O_2^{-}\) is a moderate base, capable of abstracting protons from acids that are stronger than \(HO_2\) (i.e., acids with pKa less than ~4.8). Therefore, \(O_2^{-}\) is able to abstract a proton from many weak acids. The statement that \(O_2^{-}\) is unable to abstract proton from weak acids is incorrect.

Conclusion

Based on the analysis of each property, the only correct statement among the options provided is that \(O_2^{-}\) is a stronger oxidizing agent than \(O_2\).