Question

Manganate ion is paramagnetic due to the presence of :

• A. one unpaired electron
• B. two unpaired electrons
• C. three unpaired electrons
• D. fully paired electrons

Hint:

Manganate (\( MnO_{4}^{2-} \)) is Green and Paramagnetic (\( d^{1} \)).
Permanganate (\( MnO_{4}^{-} \)) is Purple and Diamagnetic (\( d^{0} \)).
Always distinguish carefully between these two closely related ions.

Answer 1

The Correct Option is A

Concept:

• Paramagnetism is a property of substances that are weakly attracted by a magnetic field, occurring due to the presence of unpaired electrons.

• To determine the magnetic nature of an ion, we must identify the oxidation state of the central metal and write its electronic configuration.

• Manganate ion is represented by the chemical formula \( MnO_{4}^{2-} \).
Step 1: Calculate the oxidation state of Manganese (\( Mn \))
In the \( MnO_{4}^{2-} \) ion, let the oxidation state of \( Mn \) be \( x \).
Oxygen generally has an oxidation state of \( -2 \).
The sum of oxidation states must equal the net charge of the ion (\( -2 \)):
\[ x + 4(-2) = -2 \] \[ x - 8 = -2 \] \[ x = +6 \] Thus, Manganese is in the \( +6 \) oxidation state.

Step 2: Determine the electronic configuration of \( Mn^{6+} \)
The atomic number of Manganese (\( Mn \)) is 25.
The ground state electronic configuration of neutral \( Mn \) is:
\[ Mn: [Ar] 3d^{5} 4s^{2} \]
To form the \( Mn^{6+} \) ion, we must remove 6 electrons (starting from the outermost shell):
- Remove 2 electrons from the \( 4s \) orbital.
- Remove 4 electrons from the \( 3d \) orbital.

The configuration of \( Mn^{6+} \) is:
\[ Mn^{6+}: [Ar] 3d^{1} 4s^{0} \]

Step 3: Analyze the unpaired electrons
The \( 3d^{1} \) configuration indicates that there is exactly one electron in the d-subshell.
Since this single electron occupies one of the five d-orbitals alone, it is an unpaired electron.
The presence of this single unpaired electron causes the manganate ion to be paramagnetic.

Step 4: Calculate the magnetic moment (Extra detail)
The spin-only magnetic moment (\( \mu \)) can be calculated using:
\[ \mu = \sqrt{n(n+2)} \text{ B.M.} \] Where \( n = 1 \):
\[ \mu = \sqrt{1(1+2)} = \sqrt{3} = 1.73 \text{ B.M.} \]