🤓 Based on our data, we think this question is relevant for Professor Altomare's class at UCF.

Explain the following.

The O_{2}^{2+} ion has a stronger O-O bond than O_{2} itself.

We’re being asked why O_{2}^{2+} ion has a stronger O-O bond than O_{2} itself.

For this, we need to determine the bond order for each species.

The ** bond order** tells us the stability of a bond:

* Step 1:* Calculate the total number of valence electrons present.

* Step 2:* Draw the molecular orbital diagram.

**Step 3:**** **Calculate the bond order of the molecule/ion.

Recall that the formula for ** bond order** is:

$\overline{){\mathbf{Bond}}{\mathbf{}}{\mathbf{Order}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}\frac{\mathbf{1}}{\mathbf{2}}{\mathbf{[}}{\mathbf{\#}}{\mathbf{}}{\mathbf{of}}{\mathbf{}}{{\mathbf{e}}}^{{\mathbf{-}}}{\mathbf{}}{\mathbf{in}}{\mathbf{}}{\mathbf{bonding}}{\mathbf{}}{\mathbf{MO}}{\mathbf{}}{\mathbf{-}}{\mathbf{}}{\mathbf{\#}}{\mathbf{}}{\mathbf{of}}{\mathbf{}}{{\mathbf{e}}}^{{\mathbf{-}}}{\mathbf{}}{\mathbf{in}}{\mathbf{}}{\mathbf{antibonding}}{\mathbf{}}{\mathbf{MO}}{\mathbf{]}}}$

Bond Order