To determine the most stable oxidation state of iron, we can follow these steps:
### Step 1: Understand the Electronic Configuration of Iron
The electronic configuration of iron (Fe) is:
\[ \text{Fe: } 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^6 \, 4s^2 \]
### Step 2: Identify Possible Oxidation States
Iron can exhibit multiple oxidation states, commonly +2 and +3. The oxidation states arise from the loss of electrons from the outermost s and d orbitals.
### Step 3: Analyze the +2 Oxidation State (Fe²⁺)
When iron loses two electrons to form Fe²⁺, the electronic configuration becomes:
\[ \text{Fe}^{2+}: 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^6 \]
This indicates that two electrons from the 4s orbital are removed, resulting in a configuration of 3d⁶.
### Step 4: Analyze the +3 Oxidation State (Fe³⁺)
When iron loses three electrons to form Fe³⁺, the electronic configuration becomes:
\[ \text{Fe}^{3+}: 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^5 \]
This indicates that one electron is removed from the 4s orbital and two from the 3d orbital, resulting in a configuration of 3d⁵.
### Step 5: Stability of the Oxidation States
The stability of oxidation states is influenced by the electron configuration:
- A half-filled (3d⁵) or fully filled (3d¹⁰) d subshell is particularly stable due to symmetry and exchange energy.
- The +3 oxidation state (3d⁵) is half-filled, making it more stable than the +2 oxidation state (3d⁶).
### Conclusion
The most stable oxidation state of iron is:
\[ \text{Fe}^{3+} \text{ (or +3)} \]
