How many number of unpaired electrons are present in `Mn^(2+)` ?

Explanation

To determine the number of unpaired electrons in the ion \( \text{Mn}^{2+} \), we can follow these steps:

### Step 1: Determine the atomic number of manganese (Mn)
Manganese (Mn) has an atomic number of 25.

### Step 2: Write the electronic configuration of neutral manganese (Mn)
The electronic configuration of manganese is:
\[ \text{Mn}: [\text{Ar}] \, 3d^5 \, 4s^2 \]
This means that manganese has 5 electrons in the 3d subshell and 2 electrons in the 4s subshell.

### Step 3: Remove electrons to find the configuration of \( \text{Mn}^{2+} \)
When manganese loses two electrons to form \( \text{Mn}^{2+} \), the electrons are removed from the 4s subshell first (as it is higher in energy compared to 3d). Therefore, the electronic configuration of \( \text{Mn}^{2+} \) becomes:
\[ \text{Mn}^{2+}: [\text{Ar}] \, 3d^5 \, 4s^0 \]

### Step 4: Analyze the 3d subshell for unpaired electrons
In the \( 3d^5 \) configuration, the 5 electrons will occupy the 3d orbitals. According to Hund's rule, each of the five 3d orbitals will get one electron before any pairing occurs. Thus, the distribution of electrons in the 3d subshell is:
- Each of the five 3d orbitals has one electron.

### Step 5: Count the number of unpaired electrons
Since all five electrons are in separate orbitals and not paired, the number of unpaired electrons in \( \text{Mn}^{2+} \) is:
\[ \text{Number of unpaired electrons} = 5 \]

### Final Answer
The number of unpaired electrons in \( \text{Mn}^{2+} \) is **5**.

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