Construct the molecular orbital diagram of NO^{+ }and calculate for the** bond order** and determine if its **paramagnetic or diamagnetic**

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

*Step 2:* Draw the molecular orbital diagram.

*Step 3:* Determine if there's an unpaired MO (paramagnetic or diamagnetic)

*Step 4:* Calculate the bond order of the molecule/ion. Recall that the formula for ** bond order** is:

$\overline{){\mathbf{Bond}}{\mathbf{}}{\mathbf{order}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}\frac{\mathbf{1}}{\mathbf{2}}\left[\mathbf{\#}\mathbf{}{\mathbf{e}}^{\mathbf{-}}\mathbf{}\mathbf{in}\mathbf{}\mathbf{bonding}\mathbf{}\mathbf{MO}\mathbf{}\mathbf{-}\mathbf{}\mathbf{\#}\mathbf{}{\mathbf{e}}^{\mathbf{-}}\mathbf{}\mathbf{in}\mathbf{}\mathbf{anti}\mathbf{-}\mathbf{bonding}\mathbf{}\mathbf{MO}\mathbf{}\right]}$

Use the molecular orbital diagram shown to determine the bond order for NO ^{+}. Is NO^{+} paramagnetic or diamagnetic? In molecular orbital names, s = sigma and p = pi.

a. 3, paramagnetic

b. 2, paramagnetic

c. 2.5, paramagnetic

d. 2, diamagnetic

e. 3, diamagnetic

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 MO Theory: Bond Order concept. You can view video lessons to learn MO Theory: Bond Order. Or if you need more MO Theory: Bond Order practice, you can also practice MO Theory: Bond Order practice problems.

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