How many unpaired electrons are present in the complex ion [FeCl_6]^{4-}?

• A. 0
• B. 6
• C. 2
• D. 4

Answer: (D)

Verified step by step guidance

1. Identify the oxidation state of the iron (Fe) in the complex ion [FeCl_6]^{4-}. Since each chloride ion (Cl^-) has a charge of -1, and the overall charge of the complex is -4, set up the equation: \(x + 6(-1) = -4\), where \(x\) is the oxidation state of Fe.
2. Solve for \(x\) to find the oxidation state of Fe. This will help determine the electron configuration of the metal center in the complex.
3. Determine the electron configuration of the Fe ion in this oxidation state. Recall that neutral Fe has the configuration \([Ar] 3d^6 4s^2\). Remove electrons according to the oxidation state found, starting from the 4s orbital.
4. Consider the ligand field strength of Cl^- ligands. Chloride is a weak field ligand, so the complex is likely to be high spin. This affects how the \(d\) electrons are distributed among the \(t_{2g}\) and \(e_g\) orbitals.
5. Draw the \(d\) orbital splitting diagram for an octahedral field with weak field ligands and fill the electrons according to Hund's rule. Count the number of unpaired electrons in the \(d\) orbitals to find the answer.