JEE Chemistry Coordination Compounds Notes Set 04

Basic Terms

1. A coordination compound has a central metal atom or ion bonded to a fixed number of ions or molecules called ligands. Example: \( \ce{[Cu(NH3)4]SO4} \).
2. Coordination sphere: the central metal plus its ligands, written inside square brackets. The ions outside the bracket are counter ions.
3. Coordination number: the number of donor atoms attached directly to the metal through sigma bonds. Common values are 4 and 6.
4. Ligand types: monodentate (one donor atom, e.g. \( \ce{Cl^-} \), \( \ce{NH3} \), \( \ce{H2O} \)), bidentate (en, oxalate), hexadentate (EDTA). Ligands that bind through two or more atoms to form rings are called chelating ligands.
5. Ambidentate ligands can bind through either of two different atoms, for example \( \ce{NO2^-} \) (through N as nitro or through O as nitrito) and \( \ce{SCN^-} \) (through S or N).
6. Werner's theory: every metal has a primary valency (the oxidation state, satisfied by ions and ionisable) and a secondary valency (the coordination number, satisfied by ligands and directional in space).

Double Salts and Complex Salts

A double salt such as Mohr's salt or potash alum loses its identity in water and gives the tests of all its simple ions. A complex salt such as \( \ce{K4[Fe(CN)6]} \) keeps the complex ion intact, so the test for the central metal ion is not given in the usual way.

Naming Complexes

1. Name the cation first and the anion second.
2. Inside the complex, name the ligands first in alphabetical order (ignore the numerical prefixes di, tri), then the metal.
3. Use di, tri, tetra for simple ligands and bis, tris, tetrakis for complicated ones.
4. Anionic ligands end in -o (chloro, cyano, oxalato, nitro); neutral ligands keep their names except aqua (\( \ce{H2O} \)), ammine (\( \ce{NH3} \)), carbonyl (CO) and nitrosyl (NO).
5. If the complex is an anion, the metal name ends in -ate (ferrate, cuprate, argentate, aurate, plumbate, stannate). Write the oxidation state of the metal in Roman numerals in brackets.

Oxidation State and Effective Atomic Number

Oxidation state of the metal = overall charge of the complex minus the sum of ligand charges. For \( \ce{[Fe(CN)6]^{3-}} \): x + 6(-1) = -3, so x = +3.

Effective atomic number (EAN) = atomic number of the metal - oxidation state + 2 x (number of ligands, for monodentate ligands). For \( \ce{[Fe(CN)6]^{4-}} \): 26 - 2 + 12 = 36, the electron count of krypton. Stable carbonyls often follow the EAN (18 electron) rule, such as \( \ce{Ni(CO)4} \) (28 + 8 = 36) and \( \ce{Fe(CO)5} \).

Isomerism

Tetrahedral complexes cannot show geometrical isomerism. Square planar \( \ce{MA2B2} \) has two geometrical isomers; octahedral \( \ce{MA4B2} \) also has two; \( \ce{MA3B3} \) has fac and mer forms. An octahedral complex with two bidentate ligands and two monodentate ligands, cis-\( \ce{[M(AA)2B2]} \), is chiral.

Valence Bond Theory and Shapes

Strong-field ligands (CN-, CO, en, NH3) push electrons together in the lower d orbitals and give low-spin inner-orbital complexes. Weak-field ligands (F-, Cl-, H2O) leave electrons unpaired and give high-spin outer-orbital complexes. Spin-only magnetic moment: \( \mu = \sqrt{n(n+2)} \) BM, where n is the number of unpaired electrons.

Crystal Field Theory

1. The metal d orbitals split in the field of the ligands. In an octahedral field they split into the lower t₂g set (three orbitals) and the upper eg set (two orbitals), separated by \( \Delta_o \).
2. If \( \Delta_o \) is larger than the pairing energy, electrons pair in t₂g (low spin); otherwise they occupy eg too (high spin).
3. Crystal field stabilisation energy for octahedral complexes = \( (-0.4\, n_{t_{2g}} + 0.6\, n_{e_g})\Delta_o \).
4. In a tetrahedral field the splitting is \( \Delta_t \approx \tfrac{4}{9}\Delta_o \), too small to pair electrons, so tetrahedral complexes are almost always high spin.
5. Colour arises from d-d transitions. A stronger ligand gives larger splitting and absorbs shorter wavelengths. Ions with d⁰ or d¹⁰ configurations are colourless.

Spectrochemical series (weak to strong): \( I^- < Br^- < Cl^- < F^- < OH^- < H_2O < NH_3 < en < NO_2^- < CN^- < CO \).

Stability and Applications

1. Stability depends on the charge to size ratio of the metal ion and on the basic strength of the ligand. Chelating ligands form more stable complexes than monodentate ones (chelate effect); EDTA complexes are especially stable.
2. Tests: dimethylglyoxime gives a red precipitate with \( \ce{Ni^{2+}} \); potassium ferrocyanide gives Prussian blue with \( \ce{Fe^{3+}} \); excess ammonia gives deep blue \( \ce{[Cu(NH3)4]^{2+}} \).
3. Extraction: silver and gold dissolve in aerated cyanide as \( \ce{[Ag(CN)2]^-} \) and \( \ce{[Au(CN)2]^-} \) and are then displaced by zinc. Nickel is purified through volatile \( \ce{Ni(CO)4} \) (Mond process).
4. Biology and medicine: haemoglobin (Fe), chlorophyll (Mg), vitamin B12 (Co), and the anticancer drug cisplatin, cis-\( \ce{[Pt(NH3)2Cl2]} \).
5. Organometallic compounds have at least one metal-carbon bond: Grignard reagents, ferrocene, Zeise's salt, and the Ziegler-Natta catalyst (triethylaluminium with \( \ce{TiCl4} \)). Wilkinson's catalyst \( \ce{[RhCl(PPh3)3]} \) hydrogenates alkenes.