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:

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

Explain the following.

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

Frequently Asked Questions

What scientific concept do you need to know in order to solve this problem?

Our tutors have indicated that to solve this problem you will need to apply the Bond Order concept. If you need more Bond Order practice, you can also practice Bond Order practice problems.

What professor is this problem relevant for?

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