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

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

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