Samacheer Class 12 Chemistry - Coordination Chemistry

i) EDTA is the ethylenediaminetetraacetato ligand (EDTA4−); Na[Ni(EDTA)] is named sodium nickel(II) ethylenediaminetetraacetate. ii) [Ag(CN)2]−: CN− is named cyanido (or cyanido), so [Ag(CN)2]− is dicyanidoargentate(I). iii) Interpreted as a complex in which sulfate coordinates to Co together with two en ligands: name = sulfato-bis(ethylenediamine)cobalt(III) (coordination-site indicated as sulfato- if O‑bound). iv) ONO denotes nitrito bound through O; [Co(NH3)5(ONO)]2+ is named pentaamminenitrito-Ocobalt(III) ion. v) Interpreted as a platinum complex containing NH3, Cl and NO2; when NO2 is N‑bound it is named nitro; an appropriate IUPAC style name is amminechloridonitroplatinum(II).

a) Fe(II) + 6 CN− → complex [Fe(CN)6]4−; potassium salt K4[Fe(CN)6]. b) Pentacarbonyliron(0) is neutral Fe(CO)5. c) Nitrito-κN (NO2− bound through N) is an anionic ligand; Co(III) with five neutral NH3 and one NO2− gives [Co(NH3)5(NO2)]2+. d) [Co(NH3)6]3+ cation; to balance sulfate (SO42−) need overall formula [Co(NH3)6]2(SO4)3. e) Cr(III) with 4 F− and 2 OH−: complex ion [CrF4(OH)2]3−; sodium salt Na3[CrF4(OH)2].

The molar conductivity of a coordination compound in solution depends on how many ions it dissociates into (i.e. the total number of ions per formula unit). Arrange from least to greatest dissociation (fewest → most ions). (Because the OCRed formulas are unclear, apply this rule to the given complexes: the species that gives the smallest ionic count in solution has the lowest molar conductivity and the one giving the largest ionic count has the highest.)

Cisplatin (cis‑Pt(NH3)2Cl2) is an anticancer drug. Important biological coordination compounds include haemoglobin (Fe2+/Fe3+ in a porphyrin ring), chlorophyll (Mg2+ in a porphyrin ring) and vitamin B12 (corrin ring with Co).

VB theory: Cr3+ (d3) in [Cr(NH3)6]3+ with weak‑field NH3 does not pair electrons; three unpaired electrons give paramagnetism. Ni2+ (d8) in presence of strong‑field CN− pairs electrons and forms a square‑planar dsp2 hybrid (as in [Ni(CN)4]2−), resulting in all electrons paired and diamagnetism.

In the octahedral [Co(en)2Cl2]+ (two bidentate en and two Cl−): - trans: the two Cl− are opposite → achiral. - cis: the two Cl− adjacent → the arrangement of the two bidentate ligands gives non‑superposable mirror images (Δ and Λ) → cis isomer is optically active (enantiomeric pair).

Colour in many transition complexes arises from d–d electronic transitions. Ti3+ (3d1) has one d electron that can be excited (giving colour). Sc3+ (3d0) has no d electrons, so there are no d–d transitions and the complex is colourless (except for possible weak charge‑transfer bands).

For an octahedral complex of type MA2B2C2 (three different pairs of monodentate ligands) there are several distinct geometric isomers. One important classification: the facial (fac) isomer (each pair of identical ligands adjacent on a face) and meridional (mer) isomer (identical ligands lie in a plane through the centre). Specific counts and chirality depend on which ligands are trans/cis; e.g. [Co(NH3)2Cl2Br2] shows these distinct arrangements.

Add aqueous AgNO3 to the solutions. If Cl− is present as a free (ionic) counter‑ion (as in [Co(NH3)4]Cl with external Cl−) a white precipitate of AgCl forms. If Cl− is coordinated to Co (inside the coordination sphere) it is not available as a free ion and no AgCl precipitate is obtained. (Conversely, BaCl2 can be used to test for free sulfate: BaSO4 precipitates if SO4 2− is outside the coordination sphere.)

In an octahedral field the five d orbitals split into two sets: the lower‑energy t2g set (dxy, dxz, dyz) and the higher‑energy eg set (dz2, dx2−y2). The t2g orbitals are stabilized by −0.4Δo each and the eg orbitals are destabilized by +0.6Δo each; the energy gap between the sets is Δo.

Ambidentate ligands (e.g. NO2−, SCN−) can coordinate through alternative atoms, giving linkage isomers that differ in which atom is bonded to the metal. The chemical and spectroscopic properties of the isomers are different.

Key differences: (1) Bonding — double salts are held by ionic interactions between simple ions; coordination compounds have coordinate (dative covalent) bonds between metal and ligands. (2) Dissociation — double salts dissociate into original ions in solution; coordination compounds retain the complex ion. (3) Properties — coordination compounds often show distinct stereochemistry and characteristic colours, whereas double salts behave as mixtures of their component salts in solution.

Werner explained coordination compounds by introducing primary valence (oxidation state) and secondary valence (coordination number/geometry). He used these to rationalize the composition and isomerism of coordination complexes, predicting octahedral and tetrahedral arrangements corresponding to different coordination numbers.

Because a tetrahedron has no pair of positions that are uniquely trans or cis relative to others (every ligand sees three equivalent neighbours), arrangements that would be distinct in an octahedron are equivalent in a tetrahedron. Thus geometrical isomers (cis/trans) are not observed for most tetrahedral ML4 complexes.

Chiral coordination complexes lack mirror planes or inversion centres; their mirror images are non‑superposable (enantiomers). In tris(bidentate) complexes like [Co(en)3]3+, the arrangement of three bidentate en ligands can be right‑ or left‑handed (Δ and Λ), giving optical activity.

Hydrate (or solvate) isomerism arises when water molecules interchange between being coordinated to the metal and being present as crystal‑water. For example, Cr(III) chloride hydrates include [Cr(H2O)6]Cl3 (all six waters coordinated) and [Cr(H2O)5Cl]Cl2·H2O (one Cl− coordinated, one water uncoordinated as lattice water).

Ligand electric fields remove the degeneracy of metal d orbitals; the resulting energy difference between the higher‑energy and lower‑energy sets of d orbitals is the crystal field splitting energy (Δ). In an octahedral field Δo separates eg (higher) and t2g (lower) levels.

When d electrons occupy split orbital levels (t2g and eg) the lower‑lying electrons lower the energy of the complex. CFSE quantifies this stabilization (in units of Δ). For example, a d3 ion in octahedral field has CFSE = 3(−0.4Δo) = −1.2Δo.

[Ni(H2O)6]2+ is an octahedral, weak‑field (water) complex of Ni2+ (d8). Weak‑field ligands give moderate d–d splitting so d–d transitions absorb visible light, producing a green colour. [Ni(CN)4]2− has CN−, a strong‑field ligand, and the complex adopts square‑planar geometry for d8 Ni2+; electrons are paired and d–d transitions are suppressed or shifted out of the visible region, so the complex appears colourless.

Bonding in metal carbonyls is described by synergic (two-way) interaction: (1) σ-donation — the filled lone pair on carbon (CO HOMO, 5σ) donates electron density into an empty metal orbital; (2) π-backbonding — filled metal d orbitals donate electron density into the empty π* antibonding orbitals of CO (LUMO, 2π*). The net result strengthens M–C bonding and weakens the C≡O bond (lower C–O stretching frequency). This synergic bonding often leads to complexes that obey the 18‑electron rule, explains stability of terminal vs. bridging CO, and accounts qualitatively for spectroscopic (IR) shifts and observed geometries. MO descriptions combine these interactions into bonding and antibonding molecular orbitals delocalized over metal and CO.

When excess NH3 is added to CuSO4·5H2O, ammonia replaces water ligands to give the tetraammine complex: Cu^{2+} + 4 NH3 → [Cu(NH3)4]^{2+}. In aqueous medium two water molecules may remain axially: [Cu(NH3)4(H2O)2]^{2+}. The salt is commonly written as [Cu(NH3)4]SO4 (or with waters coordinated); the complex is deep blue.

Co in [Co(CN)6]3− is in the +3 oxidation state (Co3+, d6). CN− is a strong‑field ligand and pairs the 3d electrons. The paired d electrons occupy three 3d orbitals leaving two 3d orbitals (along with 4s and 4p) available for hybridization to form six d2sp3 hybrid orbitals directed octahedrally. Six CN− ligands coordinate by donating lone pairs into these hybrids to form coordinate covalent bonds. All 3d electrons are paired (low‑spin d6), so the complex is diamagnetic.

Main limitations of valence‑bond (VB) theory for coordination compounds: (1) It gives only a qualitative picture and cannot predict energy levels or spectral lines (electronic spectra). (2) It fails to account for the variable magnetic moments and the existence of both high‑spin and low‑spin states in some cases quantitatively. (3) It does not treat metal–ligand π‑bonding (π‑backbonding) adequately — important in carbonyls and π‑acceptor ligands. (4) It cannot explain the bonding in complexes with extensive delocalization or metal–metal bonding. (5) It sometimes gives incorrect predictions of geometry (e.g. complexes of d10 metals) and does not account well for thermodynamic stabilities and reaction kinetics. (6) It lacks the molecular orbital treatment needed for quantitative description.

Oxidation state of Mn: +3 (Mn3+, d4). Coordination number: 6. Nature of ligand: CN− (monodentate, strong‑field). In an octahedral crystal field CN− causes pairing → low‑spin d4 with electronic configuration t2g4 eg0. Magnetic property: paramagnetic with two unpaired electrons.