According to Molecular Orbital Theory, bond length decreases as bond order increases . The bond order is calculated as $\frac{1}{2}(N_b - N_a)$, where $N_b$ and $N_a$ are the number of bonding and antibonding electrons, respectively. The bond orders for the given oxygen species are:

• $\text{O}_2$ (16 electrons): Bond order = $\frac{10 - 6}{2} = 2.0$ .
• $\text{O}_2^+$ (15 electrons): One electron is removed from the antibonding $\pi^*$ orbital, making the bond order = $\frac{10 - 5}{2} = 2.5$.
• $\text{O}_2^-$ (17 electrons): One electron is added to the antibonding $\pi^*$ orbital, making the bond order = $\frac{10 - 7}{2} = 1.5$.
• $\text{O}_2^{2-}$ (18 electrons): Two electrons are added to the antibonding $\pi^*$ orbitals, making the bond order = $\frac{10 - 8}{2} = 1.0$. Since $\text{O}_2^+$ has the maximum bond order (2.5), it possesses the minimum bond length.