Statement I: K3[Fe(CN)6] is a low spin complex.
Statement II: In the presence of the strong field ligand CN−, the 3d electrons of Fe3+ pair up, leaving only one unpaired electron.
Answer & explanation
Correct answer: option 1
The correct answer is Option (1) → Both Statement I and statement II are correct and statement II is the correct explanation of statement I
A low spin complex is one in which the d electrons of the metal ion pair up as far as possible, leaving the fewest unpaired electrons.
Oxidation state of Fe: in K3[Fe(CN)6], 3(+1) + x + 6(−1) = 0, so x = +3.
Configuration: Fe (Z = 26) is [Ar]3d64s2; losing the two 4s electrons and one 3d electron gives \( \ce{Fe^3+} \): 3d5.
Ligand: \( \ce{CN^-} \) is a strong field ligand, so Δo > P and the electrons pair in the lower t2g set: t2g5eg0.
Unpaired electrons: n = 1 (a high spin d5 ion, as in [FeF6]3−, has 5).
Magnetic moment: \( \mu=\sqrt{n(n+2)}=\sqrt{1(1+2)}=\sqrt{3}=1.73\,\mathrm{BM} \).
Statement I: K3[Fe(CN)6] is a low spin complex. — True. [Fe(CN)6]3− has only one unpaired electron. In valence bond terms it is an inner orbital complex (d2sp3), and an inner orbital complex is a low spin complex.
Statement II: In the presence of the strong field ligand CN−, the 3d electrons of Fe3+ pair up, leaving only one unpaired electron. — True. \( \ce{CN^-} \) stands near the top of the spectrochemical series. It splits the d orbitals so widely that pairing in t2g costs less energy than going up to eg, so the five 3d electrons of \( \ce{Fe^3+} \) all enter t2g as two pairs and one single electron.
Why Statement II explains Statement I?
A low spin complex is one with the fewest unpaired electrons, and Statement II gives the reason this complex has only one.
The strong field of \( \ce{CN^-} \) forces four of the five 3d electrons into pairs. So Statement II is the cause of Statement I.
NCERT (2026-27 edition):
5.5.2 Magnetic Properties of Coordination Compounds – “… [Fe(CN)6]3– has magnetic moment of a single unpaired electron while [FeF6]3– has a paramagnetic moment of five unpaired electrons. … [Mn(CN)6]3–, [Fe(CN)6]3– and [Co(C2O4)3]3– are inner orbital complexes involving d2sp3 hybridisation, the former two complexes are paramagnetic and the latter diamagnetic.”
5.5.1 Valence Bond Theory – “… In the formation of this complex, since the inner d orbital (3d) is used in hybridisation, the complex, [Co(NH3)6]3+ is called an inner orbital or low spin or spin paired complex.”
5.5.4 Crystal Field Theory – “… In general, ligands can be arranged in a series in the order of increasing field strength as given below: I– < Br– < SCN– < Cl– < S2– < F– < OH– < C2O42– < H2O < NCS– < edta4– < NH3 < en < CN– < CO Such a series is termed as spectrochemical series. … (ii) If Δo > P, it becomes more energetically favourable for the fourth electron to occupy a t2g orbital with configuration t2g4eg0. Ligands which produce this effect are known as strong field ligands and form low spin complexes.”