a. What is the electron configuration of Mn in [MnO4]1+? b. What is the magnetic moment of Mn in [MnO4]1+? Use the formula µ = n(n + 2) µB. Is this diamagnetic or paramagnetic? c. Permanganate is an intense purple. What transition is responsible
The Correct Answer and Explanation is:
Electron Configuration of Mn in [MnO₄]⁺
To determine the electron configuration of Mn in [MnO₄]⁺:
- Manganese (Mn) has atomic number 25. Its ground-state electron configuration is:
Mn: [Ar] 3d⁵ 4s² - In the permanganate ion [MnO₄]⁺, the oxidation state of Mn is +6. Here’s how:
Let x be the oxidation state of Mn. x+4(−2)=+1⇒x−8=+1⇒x=+9x + 4(-2) = +1 \Rightarrow x – 8 = +1 \Rightarrow x = +9x+4(−2)=+1⇒x−8=+1⇒x=+9 This is incorrect. Let’s correct that. Actually, for [MnO₄]⁻, Mn is +7.
For [MnO₄]⁺, the total charge is +1, and there are 4 oxygen atoms (each -2). x+4(−2)=+1⇒x−8=+1⇒x=+9x + 4(-2) = +1 \Rightarrow x – 8 = +1 \Rightarrow x = +9x+4(−2)=+1⇒x−8=+1⇒x=+9 Still seems inconsistent. Let’s double-check this. Let us denote Mn oxidation state as x: x+4(−2)=+1⇒x−8=+1⇒x=+9x + 4(-2) = +1 \Rightarrow x – 8 = +1 \Rightarrow x = +9x+4(−2)=+1⇒x−8=+1⇒x=+9 But Mn cannot have an oxidation state of +9. That’s chemically not possible.
So this must be a misrepresentation. There is no stable species known as [MnO₄]⁺. However, for the sake of argument, let us assume it exists similarly to [MnO₄]⁻ where Mn is +7. So in [MnO₄]⁺, Mn would be in the +8 oxidation state, which is extremely rare and possibly hypothetical.
Assuming Mn is in +6 oxidation state (more reasonable than +8 or +9), then:
- From Mn: [Ar] 3d⁵ 4s²
- Mn⁶⁺ loses 6 electrons → electron configuration: [Ar] 3d¹
b. Magnetic Moment of Mn in [MnO₄]⁺
If Mn is in the +6 oxidation state:
- Mn⁶⁺ has one 3d electron → n = 1
Use the magnetic moment formula: μ=n(n+2)μB=1(1+2)=3≈1.73μB\mu = \sqrt{n(n + 2)} \mu_B = \sqrt{1(1 + 2)} = \sqrt{3} \approx 1.73 \mu_Bμ=n(n+2)μB=1(1+2)=3≈1.73μB
- Since it has unpaired electrons, it is paramagnetic, not diamagnetic.
c. Color of Permanganate
Permanganate ([MnO₄]⁻) is intensely purple. The color arises from charge transfer transitions, not d–d transitions. This is because Mn in +7 oxidation state has no d-electrons (3d⁰), making d–d transitions impossible.
- The intense purple color comes from ligand-to-metal charge transfer (LMCT).
- Electrons from oxygen p-orbitals are promoted to empty Mn 3d orbitals.
- This transition absorbs light in the green-yellow region, making the compound appear purple.
Summary
a. Electron configuration of Mn in [MnO₄]⁺ (assuming Mn⁶⁺): [Ar] 3d¹
b. Magnetic moment: 1.73 µB, paramagnetic
c. The purple color of permanganate arises from ligand-to-metal charge transfer transitions, not d–d transitions.
