Question

Magnetic Moment of \( \text{Mn}^{2+} \) is:

Hint:

The magnetic moment depends on the number of unpaired electrons, calculated using \( \mu = \sqrt{n(n+2)} \, \text{BM} \), where \( n \) is the number of unpaired electrons.

Answer 1

Step 1: Ascertain the quantity of unpaired electrons. The electronic configuration for \( \text{Mn}^{2+} \) is: \[ \text{[Ar]} \, 3d^5 \] This configuration shows 5 unpaired electrons within the \( 3d \) subshell.

Step 2: Apply the formula for magnetic moment. The magnetic moment, denoted by \( \mu \), is calculated using: \[ \mu = \sqrt{n(n+2)} \, \text{BM}, \] where \( n \) represents the count of unpaired electrons. 

\[ \mu = \sqrt{5(5+2)} = \sqrt{35} = 5.9 \, \text{BM}. \] Consequently, the magnetic moment of \( \text{Mn}^{2+} \) is established as 5.9 BM.