Coordination Compounds

Coordination Compounds INTRODUCTION (a) The concept of coordination compounds arises from the complex formation tendency of transition elements. (b) These compounds play a vital role in our lives, as chlorophyll of plants and haemoglobin of animal blood are also coordination compounds of Mg and Fe respectively (b) Some common examples are as follows. CuSO4 + 4NH3  CuSO4. 4NH3 AgCN + KCN  KCN. AgCN simple stable Addition or compounds molecular compounds 4 KCN + Fe(CN)2  Fe(CN 2) . 4KCN K2SO4 + Al2 (SO4)3 + 24H2O  K2SO4. Al 2(SO 4)3. 24 H2O (c) One of the earliest known coordination compound Alum is prussian blue which was accidently prepared (c) The molecular or addition compounds are of two by a Berlin color maker, Diebach in 1704. type Double salts lattice compound and coordination (d) The field of such compound is expanding due or complex compound to their utilization in metallurgy, analytical chemistry 1. Double salts or lattice compounds and polymerisation. (a) The addition compounds which are stable in solid MOLECULAR OR ADDITION state only but are broken down into individual constituents when dissolved in water are called COMPOUNDS double salts or lattice compounds. (a) When solutions containing two or more simple (b) Their solution have the same properties as the stable compounds in molecular proportions are mixture of individual compounds allowed to evaporate, crystals of new substances (c) For example when carnallite (KCl. MgCl2. 6H2O) are obtained. These substances are termed molecular is dissolved in water it exhibits the properties or addition compounds. of KCl and MgCl2 107 Coordination Compounds 2. Coordination or complex compounds (a) The addition compounds in which some of the Ex.1 The number of ions formed in aqueous solution constituent ions or molecules lose their identity by the compound [Co(NH3)4Cl2]Cl is : and when dissolved in water they do not break (A) 2 (B) 3 up completely into individual ions are called (C) 4 (4) 7 coordination compounds Sol. (A) (b) The properties of their solutions are different than The dissociation of given compound is as follows those of their constituents. [Co(NH3)4Cl 2]Cl  [Co(NH3)4Cl 2]+ + Cl¯ (c) In such compounds there is complex ion which Ex.2 Why NH3 forms co-ordinate complexes, while is a central metal ion with lewis bases attached the isoelectronic species CH4 does not ? to it through coordinate covalent bonds. Sol. NH3 has lone pair of electrons on N unlike CH4 (d) On the basis of stability of complex ion, complex ions are further divided as follows. VARIOUS TERMS USED IN (i) Perfect complexes : Those in which complex COORDINATION COMPOUNDS ion is fairly stable and is either not dissociated or feebly dissociated in solution state, e.g. To understand concepts related to coordination + 4– compounds and their nature it is essential to discuss K4[Fe(CN)6]  4K + [Fe(CN)6] some important terms. These are as follows.  Fe2+ + 6CN¯ (feebly dissociated) The ferrocyanide ion [Fe(CN) 6]4– is so insignificantly dissociated that it can be considered as practically undissociated and does not give the test of Fe2+ or CN¯ ions 1. Central ion : (Centre of coordination) (a) The cation to which one or more neutral molecules or anions are attached is called as centre of coordination (b) Since the central ion acts as an acceptor and thus has to accommodate electron pairs donated by the donor atom of the ligands, it must have empty orbitals (ii) Imperfect complexes : Those in which complex ion is less stable and is reversibly dissociated (c) This explains why the transition metals having empty d-orbitals form coordination compounds very to give enough simple ions and thus imparts readily. their tests, e.g. K2[Cd(CN)4]  2K+ + [Cd(CN)4]2–  Cd2+ + 4CN¯ (appreciably dissociated) Note : (d) In the complexes [Ni(NH3)6]2+ and [Fe(CN)6]3– , Ni2+ and Fe3+ respectively are the central ions. 2. Ligand : (a) The neutral molecules, anions or cations which are directly linked with the central metal atom or ion in a complex ion are called ligands. An imperfect complex may be too unstable to (b) Ligand donate a pair of electrons to the central atom exist and may be completely dissociated in solution, (c) Ligands are attached to the central metal ion or it then becomes a double salts atom through coordinate bonds or dative linkage 108 Coordination Compounds (C) Tridentate ligands : The ligands having three donor atoms are called tridentate ligands. Examples are: (d) Ligands are normally polar molecules like H3; H2 or anions such as Cl¯ , OH¯, CN¯ etc. which contain at least one unshared pair of valency electrons. (e) With few exceptions, free ligands have lone pair of electron that is not engaged in bonding for example. :CN:¯ , ¯ 2.1 Types of Ligands on the basis of number of donor atoms present in them : Diethylene triamine (Dien) (A) Mono-or unidentate ligands : They have one donor atom, i.e. they supply only one electron pair to central metal atom or ion. Examples : F¯, Cl¯, Br¯, H2O, NH3, CN¯, NO2¯, ON¯,CO 2,2’,2”-Terpyridine (terpy) (B) Bidentate ligands : Ligands which have two donor atoms and have the ability to link with (D) Tetradentate ligands : These ligands possess four donor atoms. Examples are: central metal ion at two positions are called bidentate ligands. Some examples are : (Nitriloacetato) Ethylenediamine (en) Oxalate (ox) 1,10-Phenanthroline (o-phen) Triethylene tetramine (Trien) (E) Pentadentate ligands : They have five donor atoms. For example, ethylenediamine triacetate ion. Glycinato (Gly) 2,2’-Dipyridyl (Dipy) Carbonato Ethylenediamine triacetato ion 109 Coordination Compounds (F) Hexadentate ligands : They have six donor 2.2 Classification of ligands on the basis of donor atoms. The most important example is and acceptor properties of the ligands ethylenediamine tetraacetate ion. (A) Ligands having one (or more) lone pair (or pairs) of electrons (a) Ligands which contain vacant -type orbitals that can receive back donated e¯ from M-ion in low oxidation state. e.g.: CO, NO, CN, NC, R–N=C, R 3P, R3As (G) (a) (b) (c) (d) (e) (f) Ethylenediamine tetraacetato ion (EDTA)4– Chelating ligands. Polydentate ligands whose structures permit the attachment of two or more donor sites to the same metal ion simultaneously, thus closing one or more rings are called chelating ligands and the compounds formed are known as chelate compounds. A chelate may be defined as a ring structure formed by the combination of a polydentate ligand having two or more donor atoms with a metal ion forming part of the ring. The process of formation of chelates is called chelation. Chelate complexes are more stable than ordinary complexes in which the ligand is a monodentate This increased stability of the compound due to chelation is called the chelate effect In the complex ion given below, 5 membered rings are formed. So all these are called chelate complexes –dipyridyl, o-phenanthroline (b) All these ligands also have filled donor orbital in addition to vacant -type acceptor orbitals. (c) Thus in there complexes both metal and ligand function as donors and acceptors     L) (M    (d) Ligands which do not have vacant orbitals to receive back donated electron from metals eq. H2O, NH3 F. (B) Ligands having no lone pairs of electrons but have p bonding electron eq. Ethylene, benzene, cyclopentadenyl-ion Note: (1) Polydentate ligands have Flexidentate character it is not necessary that all the donor atoms present in the polydentate ligands should form coordinate bonds with central metal atom or ion. (2) There are certain ligands which have two or more donor atoms but during formation of complexes only one donor atom is attached to metal ion. Such ligands are called ambidentate ligands such as CN, CNS (g) Generally the chelate complexes with 5 or 6 3 membered rings are more stable. (h) Out of these, 5 membered rings are very stable (a) when they involve saturated ligands. (i) On the other hand 6-membered ring structures acquire maximum stability when they involve unsaturated ligands containing conjugate double (b) bond. This is due to the resonance effects involving metal d-orbitals and ligand p-orbital electrons. Coordination number : The number of atoms of the ligands that are directly bond to the central metal atom or ion by coordinate bonds is known as the coordination number of the metal atom or ion. It is actually the number of coordinate covalent which the ligands form with the central metal atom or ion 110 Coordination Compounds (c) Some common coordination numbers exhibited by metal ions are 2,4,6. The light transition metals exhibit 4 and 6 coordination numbers while heavy transition metals may exhibit coordination number more than 6 (d) For example, the coordination number of Ni in the complex [Ni(NH3)4]Cl2 is 4 and that of Pt in the complex K2[PtCl6] is 6 4. Coordination sphere : (a) The central metal atom and the ligands directly attached to it are collectively termed as the coordination sphere (b) Coordination sphere is written inside square bracket, for example [Co(NH3)6]3+ (c) The central metal atom and the ligands inside the square bracket behave as a single entity (d) The part outside the bracket is called ionisation sphere. (e) The species present in the coordination sphere are nonionizable (f) The species present in the ionization sphere are ionisable 5. Oxidation state (a) It is a number which represents the electric charge on the central metal atom of a complex ion (b) For example, the O.N. of Fe, Co and Ni in [Fe(CN)6]4–, [Co(NH3)6]3+ and Ni(CO)4 is +2, +3 and zero respectively (c) The charge of the complex is the sum of the charges of the constituent parts. 6. Effective atomic number (EAN) (a) In order to explain the stability of the complex, Sidgwick proposed effective atomic number (b) It can be defined as the resultant number of electrons with the metal atom or ion after gaining electrons from the donor atoms of the ligands. (c) The EAN generally coincides with the atomic number of next noble gas in some case. (d) EAN is calculated by the following relation EAN = atomic number of the metal–number of electrons lost in ion formation + number of electrons gained from the donor atoms of the ligands (e) The EAN values of different metal in their respective complexes are tabulated as follows Complex K4[Fe(CN)6] [Cu(NH3)4]SO4 [Co(CH3)6]Cl3 Ni(CO)4 K2[Ni(CN)4] K2[PtCl6] K3[Cr(C2O4)3] K3[Fe(CN)6] K2[Hg4] [Ag(NH3)2]Cl K2[PdCl4] Metal (Oxid. state) At. No. of metal Coordination number +2 +2 +3 0 +2 +4 +3 +3 +2 +1 +2 26 29 27 28 28 78 24 26 80 47 46 6 4 6 4 4 6 6 6 4 2 4 Effective atomic number (26 (29 (27 (28 (28 (78 (24 (26 (80 (47 (46 –2) –2) –3) –0) –2) –4) –3) –3) –2) –1) –2) + + + + + + + + + + + (6 (4 (6 (4 (4 (6 (6 (6 (4 (2 (4 x x x x x x x x x x x 2) 2) 2) 2) 2) 2) 2) 2) 2) 2) 2) = = = = = = = = = = = 36 35 36 36 34 86 33 35 86 50 52 [Kr] [Kr] [Kr] [Rn] [Rn] 111 1. Ex.3 The EAN of nickel in Ni(CO)4 is (A) 36 (B) 38 (C) 28 (D) 54 Sol. (A) Effective atomic no. (EAN) = at. no. of central atom – oxidation state +2 x (no. of ligands) = 28 - 0 + 2 x 4 = 36 Ex.4 The charge on cobalt in [Co(CN)6]3– is (A) –6 (B) +3 (C) –3 (D) +6 Sol. (B) The algebraic sum of oxidation no. of all atoms in a complex ion is equal to charge on that ion, i.e., a + 6 x (–1) = –3,  a = + 3 Ex.5 Co-ordination number of Co in [Co (NH 3)6]2+ is (A) 4 (B) 5 (C) 6 (D) 8 Sol. (C) Co-ordination number is equal to total number of ligands in a complex. Ex.6 The neutral ligand is (A) Chloro (B) Hydroxo (C) Ammine (D) Oxalato Sol. (C) :NH3 has no charge but acts as ligand. Ex.7 NH2.NH2 serves as : (A) Monodentate ligand (B) Chelating ligand (C) Bridging ligand (D) (a) and (c) both Sol. (D) NH2.NH2 serves as monodenate as well as bridging ligand because a 3-membered ring will be too strained to be stable. Coordination Compounds The order of listing the ions (a) In common salts cation is named first and then the anion. (b) In the complex ion (Cation or anion) ligands are named first followed by the name of central atom (c) The oxidation state of the central metal is indicated by Roman numeral in brack ets immediately after its name ; (0) in dicates zero o.s. (d) In case the complex is non ionic, it is amed as one word e.g. [Ni(CO)4] is called tetracarbonyl nickel (0) (e) The suffix - ate is added to the name of central metal forming anionic complex ion. In cationic complex ion, the name of metal (usual name) is fallowed by the oxidation number in bracket. However in some metals Latin names are preferred in place of English names e.g. ion as ferrate, lead as plumbate and silver as argentate. 2. Name of the ligand : (a) If there are two or more different kinds of ligands, they are named in alphabatical order without separation by hyphen. (b) When there are several ligands of same kind, they are listed alphabetically (c) Anionic ligands ending with ‘ide’ are named by replacing ‘ide’ by suffix ‘O’. e.g. Symbol Due to wide variety of coordination compounds it is essential to use a unified system of nomenclature. Though a lot of compounds are known by their trivial names only, but I.U.P.A.C. has introduced rules for the naming of coordination compounds as follows N 3– Nitrido Peroxo ligand Cl¯ Chloro O 2 2– Br¯ Bromo O2H¯ Perhydroxo CN¯ Cyano S 2– O I.U.P.A.C. NAMING OF COORDINATION COMPOUNDS Name as 2– OH¯ Oxo NH Sulphido 2– Amido Hydroxo (d) Ligands whose names end in ‘ite’ or ‘ate’ become ‘ito’ or ‘ato’ i.e., by replacing the ending ‘e’ with ‘o’ as follows Symbol Name as ligand CO 32– Carbonato C 2O 4 2– Oxalato 112 Coordination Compounds SO 4 2– Sulphato NO 3¯ Niitrato SO 32– Sulphito CH 3COO¯ Acetato NO 2¯ (bonded through oxygen) nitrite (bonded through nitrogen) nitro (e) Neutral ligands are given the same names at the neutral molecules. For example. Ethylene diamine as a ligand is named ethylene diamine in the complex. However some ex ceptions to this rule are H2O Aquo NH3 Ammine CO Carbonyl NO Nitrosyl CS Thiocarbonyl (f) Names of positive ligands ends in ‘ium’ e.g. NH4+ Ammonium NO+ Nitrosylium NH2NH3+ Hydrazinium (g) If the number of a particular ligand is more than one in the complex ion, the number is indicated by using Greek numbers such as di,tri, tetra, penta, hexa, etc. However, when the name of the ligand includes a number, e.g. dipyridyl, ethylene diamine, then bis, tris, tetrakis are used in place of di, tri, tetra, etc (h) In case of chelating ligands or ligands having di, tri, tetra, etc, in their name the prefixes bis, tris, tetrakis are used before ligands placed in paranthesis (i) In poly nuclear complexes the bridging group is indicated in the formula of the complex by separating it from the rest of complex by hyphens and adding before its name or in poly nuclear complex (a complex with two or more metal atoms) bridging ligands (which links two metal atoms) is denoted by the prefix  before its name. Ex.8 The correct IUPAC name of [CoF2 (en)2] ClO4 is (A) Difluorobis (ethylene diamine) Cobalt (III) per chlorate (B) Bis (ethylene diamine) difluoro cobalt (III) per chlorate (C) Bis (ethylene diamine) difluoro cobalt (II) per chlorate (D) None of these Sol. (B) Follow IUPAC rules. Ex.9 The formula of dichlorobis (urea) copper (II) is (A) [CuO =C(NH2)2]Cl2 (B) [CuCl2{O=C(NH2)2}] (C) [Cu{O=C(NH2)2}Cl]Cl (D) [CuCl2][O=C(NH2)2]H2 Sol. (B) Follow IUPAC rules. Ex.10 Write the formulae of the following coordination compounds (I) Potassium hexanitritoferrate (III) (ii) Dichloroplatinum (IV) tetrachloropl-atinate (II) (II) Bis (acetylacetonato) oxovanadium (IV) (iv) Potassium tetrahydroxozincate (II) (v) Sodium pentacyanonitrosylsulphidofe - rrate (III) Sol. (i) K3[Fe(ONO)6] (ii) [PtCl2][PtCl4] (iii) [VO(acac)2] (iv) K2[Zn(OH)4] (v) Na4[Fe(CN)5NOS] Ex.11 Which of the following names are not correct? Point out the mistakes. (i) [Cu(H2O) (NH3)]Br2 ; Amminoaquodibromocopper (I) (ii) Na3[Al(C 2O4)3] Trisodium trioxalatoaluminate (III) (iii) Na2 [Ni(EDTA)] ; Sodium ethylenediaminetetracetato nickel (III) (iv) [Co(NH 3 ) 5 ONO]SO 4 ; Pentaamminenitrocobalt (III) sulphate Sol. (i) Ammine aqua copper (II) bromide (ii) Sodium trioxalato aluminate (III) (iii) It is correct (iv) Pentammine nitrito cobalt (III) sulphate 113 BONDING IN COORDINATION COMPOUNDS 1. (a) (b) (c) (d) (e) (f) (g) (h) Werner’s coordination theory Alfred Werner (considered as the father of coordination chemistry) studied the structure of coordination complexes such as CoCl3. 6NH3 and CuSO4. 4NH3 in 1893. According to himEach metal in coordination compound possesses two types of valencies : (i) primary valency or principal valencies or ionisable valencies. (ii) Secondary valency or nonionisable valencies Primary valencies are satisfied by anions only. The number of primary valencies depends upon the oxidation state of the central metal. It may change from one compound to other. These are represented by dotted lines between central metal atom and anion. Secondary valencies are satisfied only by electron pair donor, the ions or the neutral species. These are represented by thick lines. Each metal has a fixed number of secondary valencies also referred as coordination number. The coordination number depends mainly on the size and the charge on the central atom. The maximum number of ions or molecules that the central atom can hold by secondary valencies is known as coordination number. The ions attached to primary valencies possess ionising nature whereas the ions attached to secondary valencies do not ionise when the complex is dissolved in a solvent. Every central ion tends to satisfy its primary as well as secondary valencies. The secondary valencies are directional and are directed in space about the central metal ion. The primary valencies are non-directional. The presence of secondary valencies gives rise to stereoisomerism in complexes. Initially, Werner had pointed out coordination number of a metal atom to be four or six. Coordination Compounds (i) The six valencies were regarded to be directed to the corners of a regular octahedron circumscribed about the metal ion. For metals having four coordination number, the four valencies are either arranged in a planar or tetrahedral nature. (j) Thus on the basis of werner theory, the CoCl3. 6 NH3 is called hexamminecobalt (III) chloride because there are six ammonia ligands and the cobalt is in the +3 oxidation state, i.e. cobalt has three primary valencies and six secondary valencies. The complex can therefore be represented as shown. Note: Now, it has been proposed that coordination number of a metal may be any whole number between 2 and 9 2. Sidgwick Theory and EAN rule Sidgwick suggested that the metal ion will continue accepting electron pairs till the total number of electrons in the metal ion and those donated by ligands is equal to that of the next noble gas. This total number of electrons is called EAN of the metal which we have already discussed. 3. Valence Bond Theory The salient features of the theory are summarised below. (a) The central metal ion has a number of empty orbitals for accommodating electrons donated by the ligands. The number of empty orbitals is equal to the coordination number of the metal ion for the particular complex. 114 (b) (c) (d) (e) (f) (g) Coordination Compounds Draw back of V.B. Theory The atomic orbitals (s, p or d) of the metal ion hybridize to form hybrid orbitals with definite (i) It describes bonding in coordination compounds only qualitatively directional properties. These hybrid orbitals now overlap with the ligand orbitals to form strong (ii) It does not offer any explanation for the optical absorption spectia of complex chemical bonds. (iii) It does not describe the detailed magnetic The d-orbitals involved in the hybridization may properties of coordination compounds. be either inner (n –1) d orbitals or outer n dorbitals. The complexes formed in these two ways 4. Crystal field Theory This is the more advanced theory and satisfactorily are referred to as low spin and high spin complexes, explains the optical absorption spectra and respectively. magnetic properties of complexes which were not Each ligand contains a lone pair of electrons. explained by V.B.T. But its details are beyond the A covalent bond is formed by the overlap of scope of this text. a vacant hybridized metal orbital and a filled orbital (A) Crystal field splitting in octahedral coordination entities of the ligand. The bond is also sometimes called In an octahedral coordination entity with six ligands as a coordinate bond. surrounding the metal atom/ion, there will be If the complex contains unpaired electrons, it is repulsion between the electrons in metal d orbitals paramagnetic in nature, while if it does not contain and the electrons (or negative charges) of the unpaired electrons, it is diamagnetic in nature. ligands. The number of unpaired electrons in the complex, Such a repulsion is more when the metal d orbital is directed towards the ligand than when it is away points out the geometry of the complex as well as hybridisation of central metal ion and vice versa. In practice, the number of unpaired electrons in a complex is found from magnetic moment measurements as illustrated below.  = n (n  2) where n = no of lone pair Thus the knowledge of the magnetic moment can be of great help in ascertaining the type of complex (h) Under the influence of a strong ligand, the electrons can be forced to pair up against the Hund’s rule of maximum multiplicity. Magnetic moment 0 1.73 2.83 3.87 4.90 5.92 (Bohr magnetons) Number of unpaired electrons 0 1 2 3 4 5 from the ligand. Thus, the d x 2  y 2 and d z 2 orbitals which point towards the axes along the direction of the ligand will experience more repulsion and will be raised in energy; and the dxy, dyz and dxz orbitals which are directed between the axes will be lowered in energy relative to the average energy in the spherical crystal field. Thus, the degeneracy of the d orbitals has been removed due to ligand electron-metal electron repulsions in the octahedral complex to yield three orbitals of lower energy, t2g set and two orbitals of higher energy, eg set. This splitting of the degenerate levels due to the presence of ligands in a definite geometry is termed as crystal field splitting and the energy separation is denoted by o (the subscript o is for octahedral). Thus, the energy of the two eg orbitals will increase by (3/5) o and that of the three t2g will decrease by (2/5) o. 115 Coordination Compounds Fig.: d orbital splitting in an octahedral crystal field The crystal field splitting, o, depends upon the Which of these possibilities occurs, depends on field produced by the ligand and charge on the the relative magnitude of the crystal field splitting, metal ion. Some ligands are able to produce strong o and the pairing energy, P (P represents the energy required for electron pairing in a single fields in which case, the splitting will be large orbital). The two options are: whereas others produce weak fields and (i) If o < P, the fourth electron enters one of the eg consequently result in small splitting of d orbitals. In general, ligands can be arranged in a series in orbitals giving the configuration t 32g e1g . Ligands for the order of increasing field strength as given below: which o < P are known as weak field ligands and – – – – – – – form high spin complexes. I < Br < SCN < Cl < S2 < F < OH (ii) If o > P, it becomes more energetically favourable 2– – 4– < C2O4 < H2O < NCS < edta for the fourth electron to occupy a t2g orbital with < NH3 < en < CN– < CO configuration t2g4eg0. Ligands which produce this Such a series is termed as spectrochemical series. effect are known as strong field ligands and form It is an experimentally determined series based on low spin complexes. the absorption of light by complexes with different Calculations show that d4 to d7 coordination entities ligands. are more stable for strong field as compared to weak field cases. Let us assign electrons in the d orbitals of metal ion in octahedral coordination entities. Obviously, the (B) Crystal field splitting in tetrahedral single d electron occupies one of the lower energy coordination entities t2g orbitals. In d2 and d3 coordination entities, the d In tetrahedral coordination entity formation, the d electrons occupy the t 2g orbitals singly in orbital splitting is inverted and is smaller as accordance with the Hund’s rule. compared to the octahedral field splitting. 4 For the same metal, the same ligands and metalFor d ions, two possible patterns of electron ligand distances, it can be shown that t = (4/9) distribution arise: (i) the fourth electron could either o. Consequently, the orbital splitting energies are enter the t2g level and pair with an existing electron, not sufficiently large for forcing pairing and, or (ii) it could avoid paying the price of the pairing therefore, low spin configurations are rarely energy by occupying the eg level. observed. 116 Coordination Compounds Fig. : d orbital splitting in a tetrahedral crystal field. Ex.15 Find out the hybridization, geometry and magnetic moment of the complexes : Ex.12 When the configuration is d7 in a transition metal, the paramagnetic susceptibility will be equal (i) [Co(NH3)6]3+ (ii) [Cr(CN)6]3– to Sol. (i) The oxidation state of cobalt in the complex (A) 3.87 B.M. (B) 2.68 B.M. [Co(NH3)6]3+ is +3 (C) 5.92 B.M. (D) 6.92 B.M. The electronic configuration of Co3+ ion is : Sol. The magnetic moment values in Bohr magneton for 3d 4s 4p 1 lone electron is 1.73, for 2 is 2.83, for 3 is 3.87, for 4 is 4.90, for 5 is 5.92, for 6 is 6.92 and so on. d7 configuration has three lone electrons, hence,  = 3.87 BM. [Co(NH ) ]3+ 3 6 Ex.13 The shape of the complex Ag(NH3)2+ is : (A) Octahedral (C) Tetrahedral (B) Square planar (D) Linear Sol. (D) Ag(NH3)2+ has sp- hybridisation and linear complex. Rearrangement Octahedral, zero magnetic moment (ii) The oxidation state of chromium in the complex is +3. The electronic configuration of Cr3+ ion is Ex.14 Hexafluoroferrate (III) Ion is an outer orbital complex. The number of unpaired electrons present in it is. (A) 1 (C) 4 (B) 5 (D) Unpredicatable 3– Sol. (B) Electronic configuration of [FeF6] is : 3d 4s 4p 4d 3d 4s 4p [Cr(CN6)]3+    Rearrangement d 2 sp 3hybridizat ion Octahedral, mag. moment = = 15 = 3.87 BM 3 x ( 3  2) 117 Coordination Compounds 2+ (e) In this state Fe undergoes d2sp3 hybridisation Ex.16 Complex with CN¯ ligands are usually to from six d2sp3 hybrid orbitals, each of which (A) High spin complexes (B) Low spin complexes accepts electron pair donated by CN¯ ions. (C) Both (D) None (f) The resultant complex is inner octahedral as shown Sol. (B) A strong field ligand produces low spin in figure and it should be diamagnetic as it has complexes. no unpaired electron. Formation of [Co(NH3)6]3+ takes place in the same manner Ex.17 Which of the following has least conductivity in aqueous solution (A) Co(NH3)4 Cl3 (B) Co(NH3)3Cl3 (C) Co(NH3)5Cl 3 (D) Co(NH3)6Cl3 Sol. (B) Werner’s theory [Co(NH3)3Cl3] is non electrolyte SHAPE OF COMPLEXES 1. Octahedral shape of [Co(NH3)6]3+ 1.1.2 [Fe(CN)6]3– 3d 4s 4p 26 The shape of complexes depends upon hybridization (a) Fe 3d 4s 4p state of central atom, it is described as follows: (b) Fe3 + Octahedral complexes : 3+ On the basis of hybridized orbitals it can be of (c) Fe (rearranged due to presence of CN¯) 3d 4s 4p two type as d2sp 3 ( inner orbital) or sp3d2 (outer orbital) hybridized. 1.1 Inner orbital complexes : We have already discussed that in these type of complexes the d-orbitals used are of lower quantum number i.e. (n - 1) various examples are as follow (a) Complexes formed by the use of inner orbitals are diamagnetic or have reduced paramagnetism.  d2sp3 Hybridization (d) The resulting complex is octahedral due to d2sp3 hybridization (e) due to presence of unpaired e¯ it is paramagnetic 1.1.3 [Cr (NH3)6]3+ 3d5 4s’ 4D (b) These are called as low spin or spin paired complexes (a) Cr24 1.1.1 [Fe(CN)6]4– (b) Cr3+ (a) e¯ configuration of Fe26 = [Ar] 3d64s 2 (c) Cr3+ in d2sp 3 hybridized state (b) e¯ configuration of Fe+2 = [Ar] 3d6 =    d 2sp3 bybridized state (c) e¯ configuration of Fe+2 after rearrangement = d2sp3 (d) The above rearrangement is due to presence of cyanide ligand (d) This d2sp3 hybridisation leads to octahedral geometry, the complex [Cr(NH3)6]3+ will be octachedral in shape (e) Since the complex ion has 3 unpaired electrons it must be paramagnetic which is founded to be so (f) Other complexes of chromium with similar inner structure are [Cr(CN)6]3– and [Cr(H2O)6]3+ 118 Coordination Compounds (c) Zn in sp hybridised state 2+ 3 1.2 Outer orbital complexes. (a) In these complexes s, p as well as d orbitals involved in hybridization, belong to the highest sp3 hybridised quantum number (n) 2+ (b) complexes formed by the use of outer n, d orbitals (d) Zn (NH3)4] will be paramagnetic. (c) These complexes are called high- spin or spin NH3 NH3 NH3 NH3 free complexes (e) Since the complex is formed by sp3 hybridisation, (d) The outer orbital complexes have greater number it is tetrahedral of unpaired electrons. (f) Since all electrons are paired it is diamagnetic 1.2.1 [CoF6]3– 2.2 [Ni(CO)4] (a) Co 27 3d 4s 4p 4d (a) Ni28 3d 4s 4p (b) Co3+ ion (b) Ni after rearrangement 3d (c) Co3+ ion in sp3d2 hybridised state sp3d2 4s 4p (c) The empty one 4s and three 4p orbitals mix to form four sp3 hybridised orbitals (d) Due to octahedral orientation of six sp3d2 hybridised orbitals shape [CoF6]3– complex ion is octahedral (d) each orbital accepts one electron pair from carbon monoxide molecule forming [Ni(CO)4] (e) Due to presence of four unpaired electrons is 3d orbital [CoF6]3– ion has paramagnetic character (e) The shape of nickel tetra carbonyl is tetrahedral as shown below (f) Other examples are [FeF 6] 3–, [Fe(NH 3)6] 2+, (f) It is diamagnetic in nature [Ni(NH3)6]2+, [Cu(NH3)6]2+, [Cr(H2O)6]3+ 2. Tetrahedral Complexes : These are formed by sp3 hybridisation. Complexes of Zn2+ are invariable tetrahedral because they involve sp3 hybrid orbitals 2.1 [Zn(NH3)4]2+ (a) configuration of Zn30 3d 4s 4p 3. Square planar complex. These are formed due to dsp2 hybridisation. (b) e¯ configuration of Zn2+ 3d 4s 4p These complexes tend to be formed when the central ion has only one d orbital available in the inner shell 119 3.1 [Ni(CN)4] 2– (a) e¯ configuration of Ni 3d 4s 4p Coordination Compounds (c) The rearrangement is due to presence of strong ligand CN¯. The four orbitals make dsp 2 hybridisation (d) The shape of resulting complex is square planar (e) due to paired electrons it is diamagnetic. (b) e¯ configuration of Ni 2+ 3d 4s 4p L = Ligands indicated by white circles (same or different) M = Central metal indicated by black circles. dsp2 Co-ordination numbers, Hybridised orbitals and geometry of some co-ordination compounds Coordination Number 2 Hybridised orbital Geometrical shape of the complex sp Linear Examples of complex [Ag(NH3)2]+ [Ag(CN)2]¯ 3 sp2 [Hg3]¯ 4. sp3 [FeCl4]¯ [Ni(CO)4]0 Zn(NH3)4+2 [ZnCl4]–2,[CuX4]–2 where X = CN¯ Cl¯,Br¯, ¯, CNS 4. dsp 2 The d-orbital involved is dx2–y2 orbital of the inner shell, i.e. it is (n – 1)dx2–y2 orbital 5. dsp 3 The d-orbital is (n - 1)dz2 orbital [Ni(CN) 4] –2 [Pt(NH3)4]+2 [CuCl5]–3 [MoCl5]0 [Fe(CO)5]0 120 Coordination Compounds sp3d The d-orbital is ndx2–y2 orbital 5. [SbF5]–2 IF5 Square pyramidal d2sp3 When d-orbitals are (n-1) d-orbitals (Inner orbital complexes) or sp3d2 When d-orbitals are nd orbital (Outer orbital complexes) In both cases p-orbitals are dz2 and dx2-y2 orbitals. 6. [Cr(NH3)6+3] [Ti(H2O)6]+3 [Fe(CN)6]–2 [Co(NH3)6]+3 [PtCl6]–2[CoF6]–3 Octahedral In the formation of [Ni(CN)4]2–, Ni2+ ion undergoes dsp2 hybridization, hence it is square planar in shape Ex.18 The structure of iron pentacarbonyl is (A) Square planar (C) Triangular (B) Trigonal bipyramid (D) None 3d 4s 4p Ni2+ Sol. (B) [Ni(CN)4]2–    dsp3 xx Electron pair donated by CO. Rearrangement  dsp 2 hybridizat ion ISOMERISM IN COMPLEXES Ex.19 How would you account for the following : Ni(CO)4 possesses tetrahedral geometry while (a) Compounds which have the same molecular formula, but differ in their properties due to the difference [Ni(CN4]2– is square planar. in structure ar called as Isomers. Sol. In the formation of Ni(CO)4, nickel undergoes sp3 hybridization, hence it is tetrahedral in shape. (b) Isomerism is commonly considered, to be the 3d 4s 4p characteristic of only organic compounds, it is also found although less frequency among inorganic Ni(0)3d8 substances. Ni(CO)4 (c) Classification of isomerism :  Rearrangement sp 3 hybridizat ion 121 Coordination Compounds (c) Such type of isomers are distinguished by infra red STRUCTURAL ISOMERISUM (I.R.) spectroscopy. Ex. It arises due to the difference in the type of chemical (i) [Co(NH3)5NO2] Cl2 and [Co(NH3)5ONO] Cl2 linkages and distribution of ligands within and outside (ii) In NO2– ligand, The coordinating sites are the coordination sphere. nitrogen (i.e., NO2– Nitrogen ligand) or through 1. Ionisation isomers : oxygen (i.e. ONO Nitrogen ligand) The type of isomerism is due to the exchange (iii) The nitro isomer is yellow and is stable to acids whereas nitrito ismoer is red and is decomposed of groups of or ion between the coordinating by acids. sphere and the ionisation sphere Ex. 4. Coordination isomers : (a) Co(NH3)4 Br2SO4 can represent (a) This type of isomerism is exhibited when the complex (i) [Co(NH3)4Br2] SO4 (red violet) has two complex ions in it -*Cationic and anionic^. (ii) [Co(NH3)4] SO4] Br2(red). (b) This type of isomerism is caused by the interchange These complexes give sulphate ion and bromide of ligands and metal ions between the two complex ion respectively ions of the same complex. Ex. (i) [Co(NH3)6] [Cr(CN)6]’ (b) [Pt(NH3)4Cl2] Br2 and [Pt(NH3)4Br2]Cl 2 [Cr(NH3)6] [Co(CN)6] (c) [Co(NH3)4(NO3)2] SO4 and (ii) [Co(NH3)6] [Cr(C2O4)3]’ [Co(NH3)4SO4](NO3)2 [Cr(NH3)6] [Co(C2O4)3] 2. Hydrate isomers : This type of isomerism is due to presence of STEREO ISOMERISM different number of water molecules inside a (a) They have same molecule formula, same constitution, they differ only with respect to the spatial orientation coordination sphere. Ex. of ligands in space around the metal ion. (a) Cr(H2O)6Cl3 has three possible structure (b) The two stereo isomers which are possible are(i) [Cr(H2O)6]Cl3 violet Geometrical and optical. (ii) [Cr(H2O)5Cl] Cl2.H2O green 1. Geometrical or cis - Trans isomers : (iii) [Cr(H2O)4Cl 2] Cl. 2H2O dark green. (a) The ligands occupy different positions around the central metal ion. These complexes differ from one another with respect to the number of water molecules acting (b) When two identical ligands are coor dinated to the metal ion from the ion same side then it is cis as ligands. isomer. (Latin, cis means same). Other hydrate isomers are – (c) If the two identical ligands are coordinated to the (i) [Co(NH3)4]H2O Cl] Cl2 metal ion from opposite side then it is Trans isomer (ii) [Co(NH3)4 Cl2] Cl. H2O (in Latin, Trans means across) (d) These geometrical isomers differ in physical as well 3. Linkage or salt isomers : as in chemical properites. (a) This type of isomerism arises due to pres ence of ambidentate ligands like NO2– CN– and SCN– (e) Geometrical isomerism is most important in compounds with coordination numbers 4 and 6. (b) These ligands have two donor atoms but at a (f) 4-coordinated complexes with tetrahedral geometry time only one atom is directly linked to the central do not exhibit cis-Trans isomerism. metal atom of the complex. (g) It is exhibited by 4-coordinated complexes with square planar geometry. 122 Coordination Compounds Geometrical isomers with coordination number (vi) Diglycinato platium (iv) complexes = 4 (Square planar complexes) (i) Complexes with general formula, Ma2b2 (where both a and b are monodentate) can have Cis-and trans isomers. 2. (ii) Complexes with general formula Ma2bc can have 3. Geometrical isomers with coordination number =6 Cis-and trans-isomers. (i) [Fe(NH3)4Cl2] (iii) [Pt (NH3)2Cl2] (ii) Facial and Meridional isomerism (Ma3b3) (iv) [Pt(NH3)2ClBr] Optically isomers : (a) Optically active complexes are those which are nonsuperimposable over the mirror image structrue. (v) Complexes with general formula, Mabcd have three isomers. can (b) An optically active complex is one which is asymmetric in nature i.e., not divisible into two identical halves. (c) The complex which rotates plane polarised light to left hand side is laevo rotatory i.e., ‘’ or ‘–’ and if the complex rotates the plane polarised light to right hand side then it is dextro rotatory ‘d’ or ‘+’. 123 (d) Thus complexes which have same physical and chemical properties but differ in their action towards plane polarised light are called as optical isomers. (e) The ‘d’ and ‘’ isomers of a compound are called as Enantiomers or Enantio morphs. Coordination Compounds (f) Only those 6-coordinated complexes in which there are chelating agents i.e. bidentate ligands, exhibit optical isom erism. This is due to the absence of elements of symmetry in the complex. (g) Optical isomerism is not found in square planar complexes on acoount of the presence of axis of symmetry. Optical isomer with Coordination number = 6 (i) [Ma2b2c2]n+  [Pt(py)2(NH3)2Cl 2]2+ (ii) [Mabcedf]  [Pt(py)NH3NO2ClBrI] (iii) [M (AA)3]n+  [Co(en)3]3+ 124 Coordination Compounds n+ (iv) [M(AA)2ab]  [Co(en)2NH3Cl] 2+ (v) [M(AB)3]  [Cr(gly)3] (vi) [M (AA)B2C2][Fe(NH3)2(en)Cl2] (vii) [M(AA)2B2][Fe(en)2Cl2] (i) Trans- [Fe(en)2Cl2] do not show optical activity due to superimposition of their mirror image. (ii) Some more examples are optically active: [Cr(ox)3]3–;[Fe(dipy)3]2+; [Cr(ox)2(H2O)2; [Pt(en)3]4+ etc. 125 Ex.20 Which of the following sets is/are example(s) of co-ordination isomerism in complexes ? (A) [Co(NH3)6] [Cr(CN)6] and [Co(CN6] [Cr(NH3]6] (B) [Cr(H2O)5Cl]Cl2.H2O and [Cr(H2O)4Cl 2]Cl.2H2O (C) [Co(NH3)5Br]SO3 and [Cr(NH3)5 SO4]Br (D) [Pt(NH3)2Cl2] and [Pt(NH3)4] [PtCl4] Sol. (A) When both positive and negative ions are complex, co-ordination isomerism may occur due to the interchange of ligands within the co-ordination sphere itself. e.g. [Co(NH3)6] [Cr(CN)6] and [Co(CN)6] [Cr(NH3)6] are two co-ordination isomers, where ammonia and cyanide ligands have been changed within the co-ordination sphere. Choices (b), (c) and (d) are examples of hydrate isomers, ionisation isomers and polymerisation isomers, respectively. Ex.21 Which of the following does/do not exhibit optical isomerism (A) tetrahedral complexes (B) square planar complexes (C) octahedral complexes (D) polynuclear complexes Sol. A square planar complex does not exhibit optical activity because it has a horizontal plane of symmetry Coordination Compounds (f) The higher the value of stability constant stabler is the complex. (g) The value of stability constants for some of the complexes are given below : Complex Stability constant 2+ [Cu(NH3)4] 4.5 x 1011 [Ag(NH3)2]+ 1.6 x 107 [Co(NH3)6]2+ 1.12 x 106 [Co(NH3)6]3+ 5.0 x 1033 [AgCl2]¯ 1.11 x 105 [AgBr2]¯ 1.28 x 107 [Ag(CN)2]¯ 1.0 x 1022 [Cu(CN)4]2– 2.0 x 1027 [Fe(CN)6]3– 7.69 x 1043 1. Factors influencing the stability of complex 1.1 Nature of central ion (a) The complex will be more stable for higher values  Ch arg e  (b) (c) 1.2 (a) (b) of charge density  radius  The higher the electronegativity of the central ion, the greater is the stability of its complexes The higher the oxidation state of the metal the more stable is the compound Nature of ligand A basic ligand is likely to easily donate its electrons. Thus a more basic ligand will form more stable complex. Chelating ligands form more stable complexes as compared to monodentate ligands. STABILITY OF COMPLEX (a) A complex is formed in solution by the stepwise addition of ligands to a metal ion Ex.22 [Cu(CN)4]2– is more stable complex than (b) This can be expressed as follows M + L ML [Cu(NH3)4]2+. Explain where M = metal and L is ligand Sol. The higher stability constant (K = 2 x 10 27) for (c) The stability constant K for this reaction is as Cu2+ + 4CN¯  [Cu(CN)4]2– than for [Cu(NH3)4]2+ ML which is 4.5 x 1011 explains stability. shown K = [ M] [ L] Ex.23 [Ti(H2O)6]3+ is coloured while [Sc(H2O)6]3+ is (d) This metal can again get a ligand colourless. explain ML + L  ML2 Sol. Sc(H2O)63+ has no unpaired electron in its d subshell (e) The forthcoming stability constant K’ is and thus d-d transition is not possible whereas [ ML 2 ] Ti(H2O)63+ has one unpaired electron in its d subshell K1 = [ ML] [ L] its value is less than K which gives rise to d-d- transition, to impart colour. 126