JEE Chemistry Coordination Compounds Notes Set 05

Basic Terms of Coordination Chemistry

A coordination compound contains a central metal atom or ion bonded to a fixed number of ions or molecules called ligands. The central atom and its ligands together form the coordination sphere and are written inside square brackets, for example \( [Co(NH_3)_6]Cl_3 \). The ions outside the brackets are the counter ions (ionisation sphere). A complex ion shows the properties of a single species and does not release its metal ion or ligands in solution as the free ions would, which separates complex salts from double salts (alum, Mohr's salt) that break up completely into their simple ions.
The coordination number is the number of donor atoms attached to the metal, not the number of ligand molecules. The oxidation number of the metal is found from the overall charge of the complex minus the charges of the ligands. Ligands are Lewis bases (lone pair donors) and the metal is a Lewis acid. A complex made from one kind of ligand is homoleptic and one with more than one kind is heteroleptic.

Types of ligand. Monodentate: \( H_2O, NH_3, Cl^-, CN^-, CO \). Bidentate (chelating, two donor atoms): ethylenediamine (en), oxalate (ox), acetylacetonate (acac), glycinate (gly), dimethylglyoximate (dmg), bipyridyl, o-phenanthroline. Tridentate: diethylenetriamine (dien). Hexadentate: EDTA, with two N and four O donors. Ambidentate ligands have two different donor atoms but use only one at a time: \( NO_2^- \) (nitro-N or nitrito-O), \( SCN^- \) (thiocyanato-S or isothiocyanato-N), \( CN^- \) (cyano or isocyano).
A chelate ring of five or six atoms is especially stable, which is the chelate effect: replacing monodentate ligands by a polydentate one increases the number of free particles and the entropy, so \( [Ni(en)_3]^{2+} \) is far more stable than \( [Ni(NH_3)_6]^{2+} \).

IUPAC Nomenclature

1. In the name the cation is written first and then the anion. 2. Within the coordination sphere the ligands are named first, in alphabetical order (ignoring the numerical prefixes di, tri, tetra), and then the metal. 3. Anionic ligands end in -o (chlorido, cyanido, hydroxido, oxalato, nitrito-O, nitro for NO2 bound through N). Neutral ligands keep their names, except aqua (H2O), ammine (NH3), carbonyl (CO), nitrosyl (NO). 4. Use di, tri, tetra for simple ligands and bis, tris, tetrakis for ligands that already contain a prefix or are complex (bis(ethylenediamine)). 5. The oxidation state of the metal is given in Roman numerals in brackets. 6. If the complex is an anion, the metal name ends in -ate (ferrate for iron, cuprate for copper, argentate for silver, aurate for gold, plumbate for lead, stannate for tin). 7. Bridging ligands get the prefix mu.
Examples. \( K_4[Fe(CN)_6] \): potassium hexacyanidoferrate(II). \( [Co(NH_3)_4Cl(ONO)]Cl \): tetraamminechloridonitrito-O-cobalt(III) chloride. \( [Cr(en)_2Br_2]Br \): dibromidobis(ethylenediamine)chromium(III) bromide. \( Na_2[Fe(CN)_5NO] \): sodium pentacyanidonitrosylferrate(II) (sodium nitroprusside, with NO acting as \( NO^+ \)). \( [Fe(CO)_5] \): pentacarbonyliron(0).

Werner's Theory and Counting Ions

Werner proposed that a metal has two kinds of valence: the primary valence (oxidation state, ionisable, satisfied by anions) and the secondary valence (coordination number, non-ionisable, satisfied by ligands and fixed in direction in space). In \( CoCl_3\cdot 6NH_3 \) all three chlorides are primary and ionisable (AgNO3 precipitates all three), in \( CoCl_3\cdot 5NH_3 \) two are ionisable, in \( CoCl_3\cdot 4NH_3 \) one is ionisable (green and violet isomers), and in \( CoCl_3\cdot 3NH_3 \) none is ionisable. The number of ions per formula unit seen in conductivity measurements (4, 3, 2 and 0) matches the formulae \( [Co(NH_3)_6]Cl_3, [Co(NH_3)_5Cl]Cl_2, [Co(NH_3)_4Cl_2]Cl, [Co(NH_3)_3Cl_3] \).
Example. A solution of 0.319 g of \( CrCl_3\cdot 6H_2O \) (M = 266.5 g/mol) passed through a cation exchanger releases acid needing 19 mL of 0.125 N NaOH. Moles of complex = \( 0.319/266.5 = 1.2\times10^{-3} \) and milliequivalents of acid = 2.375, so 2 H⁺ per mole: the complex is \( [Cr(H_2O)_5Cl]Cl_2\cdot H_2O \).

Effective Atomic Number (EAN) Rule

Sidgwick's rule says that a metal tends to accept electrons from ligands until its total electron count equals that of the next noble gas. \( EAN = Z - (\text{oxidation number}) + 2\times(\text{number of ligands}) \), counting each donor pair as two electrons. The rule works best for metal carbonyls and related low-valent compounds: \( [Cr(CO)_6] \) has 24 + 12 = 36; \( [Fe(CO)_5] \) has 26 + 10 = 36; \( [Ni(CO)_4] \) has 28 + 8 = 36. \( [Fe(CN)_6]^{4-} \) has 24 + 12 = 36 and \( [PtCl_6]^{2-} \) has 78 - 4 + 12 = 86. A metal-metal bond adds one electron per metal, which is why dimers such as \( Co_2(CO)_8 \) (27 + 8 + 1 = 36) and \( Mn_2(CO)_{10} \) are stable, whereas mononuclear \( Co(CO)_4 \) (35) and \( Mn(CO)_5 \) (35) dimerise or are reduced. In \( [Fe(CO)_2(NO)_2] \) the NO ligands are treated as \( NO^+ \), so iron is in the -2 state and EAN = 26 + 2 + 8 = 36.

Valence Bond Theory

Vacant orbitals of the metal ion hybridise and each accepts a lone pair from a ligand. Coordination number 2: sp, linear (\( [Ag(NH_3)_2]^+ \)). Coordination number 4: sp³, tetrahedral (\( [NiCl_4]^{2-}, [Zn(NH_3)_4]^{2+}, Ni(CO)_4 \)) or dsp², square planar (\( [Ni(CN)_4]^{2-}, [PtCl_4]^{2-}, [Cu(NH_3)_4]^{2+} \)). Coordination number 5: dsp³, trigonal bipyramidal (\( Fe(CO)_5 \)). Coordination number 6: d²sp³, octahedral, inner orbital (low spin) when strong field ligands pair up the (n-1)d electrons, or sp³d², outer orbital (high spin), when weak ligands leave the d electrons unpaired.
Magnetic moment. \( \mu = \sqrt{n(n+2)}\ \text{BM} \) where n is the number of unpaired electrons: n = 1, 2, 3, 4, 5 give 1.73, 2.83, 3.87, 4.90 and 5.92 BM. Examples: \( [Fe(CN)_6]^{4-} \) (d6, low spin) is diamagnetic and d2sp3; \( [FeF_6]^{3-} \) (d5, high spin) has 5 unpaired electrons (5.92 BM) and sp3d2; \( [Co(NH_3)_6]^{3+} \) is diamagnetic whereas \( [CoF_6]^{3-} \) has 4 unpaired electrons (4.90 BM). VBT cannot explain colour, the difference between strong and weak ligands or the tetrahedral/square planar choice quantitatively, which is why crystal field theory is used.

Crystal Field Theory

Ligands are treated as point charges that raise the energy of the metal d orbitals. In an octahedral field the five d orbitals split into the lower t2g set (\( d_{xy}, d_{yz}, d_{zx} \), energy \( -0.4\Delta_o \)) and the higher eg set (\( d_{x^2-y^2}, d_{z^2} \), energy \( +0.6\Delta_o \)), because the eg orbitals point directly at the ligands. In a tetrahedral field the order is reversed: e (lower, \( -0.6\Delta_t \)) and t2 (higher, \( +0.4\Delta_t \)), with \( \Delta_t = \tfrac{4}{9}\Delta_o \). Tetrahedral complexes are almost always high spin because \( \Delta_t \) is small. In a square planar field the d orbitals split into four levels, with \( d_{x^2-y^2} \) highest.
The spectrochemical series orders ligands by field strength: \( I^- < Br^- < SCN^- < Cl^- < F^- < OH^- < C_2O_4^{2-} < H_2O < NCS^- < NH_3 < en < NO_2^- < CN^- < CO \). Delta_o also grows with the charge on the metal ion and when moving down a group (3d < 4d < 5d), so 4d and 5d metals are nearly always low spin. For d4 to d7 in an octahedral field the electrons pair (low spin) if the pairing energy P is smaller than \( \Delta_o \) (strong field) and stay unpaired (high spin) if \( \Delta_o < P \) (weak field).

CFSE for octahedral complexes: \( \text{CFSE} = (-0.4\,n_{t_{2g}} + 0.6\,n_{e_g})\Delta_o + (\text{pairing energy terms}) \). d3: -1.2 \( \Delta_o \); d4 high spin: -0.6, low spin: -1.6; d5 high spin: 0, low spin: -2.0; d6 low spin: -2.4; d7 high spin: -0.8. High spin d5 (\( [Fe(H_2O)_6]^{3+}, [Mn(H_2O)_6]^{2+} \)) and d10 have zero CFSE. Colour arises from d-d transitions in which the complex absorbs light of energy equal to \( \Delta_o \) and shows the complementary colour. The wavelength is \( \lambda = hc/\Delta_o \), so a stronger ligand shifts absorption to shorter wavelength. Complexes of d0 and d10 ions (\( Ti^{4+}, Zn^{2+}, Cu^+, Sc^{3+} \)) are normally colourless (intensely coloured d0 ions such as \( MnO_4^- \) owe their colour to charge transfer, not d-d transitions). Limitations of CFT: it treats ligands as point charges, cannot explain the position of neutral ligands such as CO in the series, and ignores pi bonding and covalency.

Examples. \( [Fe(H_2O)_6]^{2+} \): d6 weak field, t2g4 eg2, 4 unpaired, 4.90 BM, sp3d2. \( [Co(NH_3)_6]^{3+} \): d6 strong field, t2g6, 0 unpaired, d2sp3. \( [Cu(H_2O)_6]^{2+} \): d9, one unpaired electron, 1.73 BM, blue. \( [CoCl_4]^{2-} \): tetrahedral Co2+ (d7), e4 t2^3, 3 unpaired, 3.87 BM. \( [Ni(CN)_4]^{2-} \): square planar d8, diamagnetic. \( [NiCl_4]^{2-} \): tetrahedral, 2 unpaired, 2.83 BM. Co(II) is stable in water, but with strong field ligands such as ammonia it is easily oxidised to Co(III), because the extra CFSE (t2g6) of the d6 low spin ion compensates for the ionisation energy.

Isomerism in Coordination Compounds

Structural isomers. (a) Ionisation isomers exchange an anion between the sphere and the counter ion: \( [Co(NH_3)_5Br]SO_4 \) and \( [Co(NH_3)_5SO_4]Br \) (they give different precipitates with BaCl2 and AgNO3). (b) Hydrate (solvate) isomers differ in the number of water molecules inside the sphere: \( [Cr(H_2O)_6]Cl_3 \) (violet), \( [Cr(H_2O)_5Cl]Cl_2\cdot H_2O \) (blue-green) and \( [Cr(H_2O)_4Cl_2]Cl\cdot 2H_2O \) (dark green). They lose different amounts of water on drying with conc. H2SO4. (c) Linkage isomers arise from ambidentate ligands: \( [Co(NH_3)_5NO_2]^{2+} \) and \( [Co(NH_3)_5ONO]^{2+} \). (d) Coordination isomers exchange ligands between a complex cation and a complex anion: \( [Co(NH_3)_6][Cr(CN)_6] \) and \( [Cr(NH_3)_6][Co(CN)_6] \).
Geometrical isomers. Square planar \( Ma_2b_2 \) has cis and trans forms, and \( Mabcd \) has 3 isomers. Octahedral \( Ma_4b_2 \) has cis and trans; \( Ma_3b_3 \) has facial (fac) and meridional (mer) forms; \( M(AA)_2b_2 \) has cis and trans. Tetrahedral complexes and complexes with coordination number 2 or 3 show none. Counts (verified by enumeration): \( Ma_2b_2c_2 \) has 5 geometrical forms (6 stereoisomers including one enantiomeric pair); \( Mabcdef \) has 15 geometrical forms and 30 stereoisomers.
Optical isomers are non-superimposable mirror images, found in octahedral complexes without a plane of symmetry: \( [Co(en)_3]^{3+} \), cis-\( [Co(en)_2Cl_2]^+ \), \( [Cr(ox)_3]^{3-} \). The trans-isomers of \( [M(AA)_2b_2] \) are optically inactive (plane of symmetry). Square planar and tetrahedral \( Ma_2b_2 \) complexes are achiral.

Counting examples. \( [Co(NH_3)_4Br_2]Cl \) shows geometrical (cis/trans) and ionisation isomerism but not optical isomerism. \( [Co(en)_2ClBr]^+ \) has three stereoisomers (trans plus the cis pair), and \( [Rh(en)_2Cl(NO_2)]^+ \) has 3 stereoisomers and, with linkage isomerism, 6 isomers in all. A reaction that replaces one \( NH_3 \) in cis-\( [Co(NH_3)_4Br_2]^+ \) by Br^- gives both fac and mer products, but the trans reactant gives only mer.

Organometallic Compounds and Metal Carbonyls

An organometallic compound contains at least one direct metal-carbon bond: Grignard reagents RMgX, \( (C_2H_5)_4Pb \), \( (CH_3)_4Sn \), Zeise's salt \( K[PtCl_3(\eta^2\text{-}C_2H_4)] \) and ferrocene \( [Fe(\eta^5\text{-}C_5H_5)_2] \). Sodium ethoxide is not organometallic because sodium is bonded to oxygen. Wilkinson's catalyst \( [RhCl(PPh_3)_3] \) is a square planar rhodium(I) complex used for homogeneous hydrogenation of alkenes.
Metal carbonyls have the metal in zero or a low oxidation state. Structures: \( Ni(CO)_4 \) tetrahedral (sp3, diamagnetic); \( Fe(CO)_5 \) trigonal bipyramidal (dsp3, diamagnetic); \( Cr(CO)_6 \) octahedral (d2sp3); \( V(CO)_6 \) octahedral and paramagnetic (17 electrons); \( Mn_2(CO)_{10} \) two octahedral units joined by a Mn-Mn bond. CO bonds by a sigma donation from carbon to the metal and a pi back donation from a filled metal d orbital to the empty pi* orbital of CO (synergic bonding). The back donation strengthens the M-C bond (partial double bond character, shorter than a single bond) and weakens the C-O bond (bond order below 3, longer than the 1.128 angstrom of free CO and a lower stretching frequency). The more negative the metal, the stronger the back donation: C-O bond order is \( [Mn(CO)_6]^+ > [Cr(CO)_6] > [V(CO)_6]^- \).

Stability, Applications and Qualitative Tests

The stability of a complex is measured by the formation constant \( \beta \) (the larger it is, the more stable the complex). It rises with a higher charge/size ratio of the metal ion, with ligand basicity, and with chelation. Applications: EDTA titrations of Ca2+ and Mg2+ (hardness of water), removal of lead and iron poisoning by chelation, extraction of silver and gold as \( [Ag(CN)_2]^- \) and \( [Au(CN)_2]^- \), purification of nickel through volatile \( Ni(CO)_4 \) (Mond process), electroplating baths, cis-platin as an anticancer drug, haemoglobin (iron), chlorophyll (magnesium) and vitamin B12 (cobalt).
Complexes in analysis. Excess ammonia converts Cu2+ to deep blue \( [Cu(NH_3)_4]^{2+} \), and Ni2+ with dimethylglyoxime in ammoniacal solution gives the rosy-red precipitate \( [Ni(dmg)_2] \) (square planar, diamagnetic, intramolecular hydrogen bonds). Fe3+ gives blood red \( [Fe(SCN)(H_2O)_5]^{2+} \) with thiocyanate and Prussian blue \( Fe_4[Fe(CN)_6]_3 \) with ferrocyanide. The brown ring of the nitrate test is \( [Fe(H_2O)_5(NO)]^{2+} \) (3 unpaired electrons, 3.87 BM). Co2+ gives yellow \( K_3[Co(NO_2)_6] \); Hg2+ and Co2+ give the blue \( Co[Hg(SCN)_4] \); and Nessler's reagent is alkaline \( K_2[HgI_4] \). Cu2+ with excess KCN forms the colourless stable \( [Cu(CN)_4]^{3-} \) (Cu+ is d10, and CN- first reduces Cu2+).