Coordination Compounds PDF

28. C O - O R D I N AT I O N COMPOUNDS 1. INTRODUCTION The concept of coordination compounds originates from the tendency for complex formation of transition elements. Molecular or Addition Compounds When solutions of two or more simple...

28. C O - O R D I N AT I O N COMPOUNDS 1. INTRODUCTION The concept of coordination compounds originates from the tendency for complex formation of transition elements. Molecular or Addition Compounds When solutions of two or more simple stable compounds in molecular proportion are allowed to evaporate, crystals of new substances, called molecular or addition compounds, are obtained. Some common examples are given below: CuSO4 + 4NH3 → CuSO4·4NH3 AgCN + KCN → KCN·AgCN 4 KCN + Fe (CN)2 → Fe (CN)2.4KCN K 2SO 4 + Al2 ( SO 4 ) + 24H2O → K 2SO 4.Al2 ( SO 4 ).24H2O 3 3 Alum Simple stable Addition or Compounds Molecular compounds There are two types of molecular or addition compounds: (1) double salts or lattice compounds and (2) coordination or complex compounds. Double Salts or Lattice Compounds The addition compounds which are stable only in solid state but are broken down into their individual constituents when dissolved in water are called double salts or lattice compounds. Nomenclature of Double Salts (a) A  hyphen (-) is used between the compounds while writing the names of double salts with the number of molecules appearing inside brackets at the end. Example: KCl·MgCl2·6H2O – potassium chloride-magnesium chloride-water (1/1/6) K2SO4·Cr2 (SO4)3·24H2O – potassium sulphate-chromium sulphate-water (1/1/24) (a) C  ations and anions with the same oxidation number are represented with English alphabets (cation first). If more than one type of cations are present, they are arranged in the increasing order (ascending order) of oxidation numbers. Anions of different oxidation numbers are arranged in the following order: Oxide (O2–), hydroxide (OH–), inorganic anion, organic anion, hydride (H–). The above rules are further clarified by following examples: (i) KNaCO3 – Potassium sodium carbonate (mixed salt) (ii) KHCO3 – Potassium hydrogen carbonate (mixed salt) 2 8. 2 | Co-ordination Compounds Note: Hydrogen is an exception and is written at the end. (iii) NH4·MgPO4·6H2O — Ammonium magnesium phosphate-6-hydrate or water (iv) NaCl·NaF·2Na2SO4 — (Hexa) sodium chloride fluoride (bis) sulphate Coordination or Complex Compounds: Those molecular compounds which retain their identity in solid as well as in solution, are known as coordination or complex compounds. A part (or whole compound) of these compounds is not dissociated in solution and its behavior is different from its constituents. 4− Example: K 4 Fe ( CN)  → 4K + + Fe ( CN)   6  6 Complex ion Illustration 1: Aqueous solution of potassium ferrocyanide does not give the test of Fe (II) and it is not poisonous like potassium cyanide. Why? (JEE MAIN) Sol: Being a complex salt, it ionizes to K+ and [Fe (CN)6]4– ions. Due to absence of Fe (II) it does not give the test of Fe (II). Absence of free CN– makes it nonpoisonous. Note: An imperfect complex compound may be too unstable to exist and may be completely dissociated in solution, and then becomes a double salt. 2. TYPES OF COORDINATION COMPOUNDS (a) Based on the nature of cations and anions, coordination compounds are classified into four types: (i) Simple cation and complex anion: K4 [Fe (CN)6], K2 [PtCl6], etc. (ii) Complex cation and simple anion: [Cu (NH3)4] SO4, [Ni (NH3)6] Cl2, etc. (iii) Complex cation and complex anion: [Pt (NH3)4] [PtCl4], etc. (iv) Neutral complex compounds: Ni (CO)4, Fe (CO)5, etc. (b) Based on stability, coordination compounds are of two types: (i) P  erfect or penetrating complexes: These are compounds in which the complex ion is feebly dissociated in solution and is highly stable. E.g. K4 [Fe (CN)6] → 4K+ + [Fe (CN)6]4– [Fe (CN)6]4– Fe2+ + 6CN– (Feebly dissociated) (ii) Imperfect or normal complexes: These are compounds in which the complex ion is appreciably dissociated and is less stable. E.g. K4 [Cd (CN)4] → 2K+ + [Cd (CN)4]2– [Cd (CN)4]2– Cd2+ + 4CN– (Reversible dissociation) 3. TERMINOLOGIES IN COORDINATION COMPOUNDS 3.1 Central Metal Atom/Ion Metal atom of complex ion which coordinates with atoms or group of atoms/ions by accepting NH3 NH3 +2 their lone pair of electrons. Here Cu is the central metal atom. Some complex ions may also Cu have more than one metal atom. NH3 NH3  Chem i str y | 28.3 3.2 Ligands The anions, cations or neutral molecules, which form coordinate bonds with the central metal atom by donating an electron pair (lone pair) are ligands. These electron pair donors are also known as Lewis bases. Thus, a complex ion is formed as: Mn+ + xL → [MLx] n+ In the above example, NH3 is the ligand. Classification of Ligands Ligands are classified as follows: (a) On the basis of the charge of ligand: (i)  nionic ligands: These are negatively charged and are the most common type of ligand, such as, F–, Cl–, A Br–, OH–, CN–, SO32– , S2–, SO42–, etc.  OH, (ii) Neutral ligands: These are uncharged and are the electron pair donor species such as H2 O,R  NH  3      RNH2 , R 2 NH, R3 N etc. (iii) Cationic ligands: They are positively charged and are rare such as NO+, etc. (b) On the basis of denticity: The number of donations accepted by a central atom from a particular ligand is known as the denticity of the ligand. Based on this, ligands are classified as follows: (i) Monodentate: when only one donation is accepted from the ligand. For example, H2O, NO, CO, NH3, CO2–, Cl–, etc. (ii) Bidentate: when two donations are accepted from the ligand. For example, (1) En: Ethylenediamine (2) Pn: Propylenediamine CH2 CH2 CH3 CH CH2 H2N NH2 NH2 NH2 (iii) Tridentate: when three donations are accepted from the ligand. For example, (1) Dien: Diethylenetriamine (2) Imda: Iminodiacetate CH2 CH2 NH CH2 CH2 H2C NH2 CH2 O=C C=O H2N NH2 O - O- (iv) Tetradentate: when four donations are accepted from the ligand. For example, (1) Trien: Trie thylenetetraamine (2) NTA: Nitrilotriacetate CH2 CH2 CH2 O CH2 CH2 NH CH2 C N C=O CH2 - H2N NH O O H2N CH2 C O O CH2 2 8. 4 | Co-ordination Compounds (v) Pentadentate: when five donations are accepted from the ligand. For example, EDTA: ethylenediamine triacetate CH2 CH2 CH2 CH2 O=C N N C=O CH2 H - O - O C O O - (vi) H  exadentate: when six donations are accepted from the ligand. For example, EDTA: Ethylenediamine tetraacetate CH2 CH2 CH2 CH2 O=C N N C=O - CH2 CH2 - O O C C O O - O O - Some other types of ligands: (a) Flexidentate: Some ligands exhibit variable denticity and are called flexidentate ligands. For example, (i) O O O O S S and O O O O Bidentate ion Monodentate ion (ii) [Co (NH3)4CO3] Br and [Co (NH3)4CO3] Br CO32 − is bidentate CO32– is monodentate (b) Chelating: A ligand which forms a ring structure with the central atom is called a chelating ligand. All polydentate ligands are chelating ligands. Chelated complex compounds are more stable than similar complexes with monodentate ligands because dissociation of the complex involves breaking down two bonds rather than one. However, it must be noted that NH2NH2 and H2N(−CH2−CH2−)NH2 cannot act as chelating ligands due to their three-member ring and locked structure, respectively. (c) Ambidentate ligand: A ligand with more than one kind of donor sites but only one site is utilized at a time is called an ambidentate ligand. Ambidentate ligands are of two types: (i) Monodentate and ambidentate: C N O—C N S—C N or or or C N S—C N or O—C N (ii) Bidentate and ambidentate: S O- S - O C C or C C - S O- S O dithicoxalate  Chem i str y | 28.5 (d) Classification of ligands on the basis of bonding interaction between the central metal atom and ligand: (i) Classical or simple donor ligand: These ligands donate their lone pair of electrons to the central atom. For example, O2–, OH–, F–, NH2− , NH3, N3–, etc. (ii) Nonclassical or π-acceptor ligand: These ligands donate the lone pair of electrons to the central atom and accept the electron cloud from the central atom in their low-lying vacant orbital. This kind of back donation is known as ‘synergic effect’ or ‘synergic bonding’. For example, CO, CN–, NO+, PF3, PR3 (R = H, Et, Ph…), C2H4, C2H2, CO2, etc. In the case of CO, the back donation to the π* orbital of the central atom can be depicted as: * M :C  O: M = C =O * Figure 28.1: Back bonding in metal carbonyl As per valence bond or molecular orbital theory, it is implicit that the bond order of C–O bond decreases but the C–O bond length increases due to synergic effect. Similarly, since CN– and NO+ are isoelectronic with CO, back donation takes place in these species also in the π* orbitals and the same conclusion can be drawn for the bond order and bond lengths. In PR3, the back donation can be depicted as: vacant 3d orbital accepts the back donation M :P R3 Figure 28.2: Back bonding in case of phosphine ligand In C2H4, the back donation is clearly depicted with the example of Zeise’s salt. H Cl Cl * H P1 C Cl C H * H Figure 28.3: Back donation in case of ethylene ligand Here, back donation is received in the p* orbital of the C–C bond. Hence, the bond order of C–C bond decreases and the bond length increases as compared to the free C2H4 molecule. Due to backbonding, the C2H4 molecule loses its planarity and likewise the C2H2 molecule loses its linearity. 2 8. 6 | Co-ordination Compounds Illustration 2: What is meant by the denticity of a ligand? Give examples of a unidentate and a bidentate ligand. (JEE MAIN) Sol: Denticity indicates the number of donor sites in a ligand. It can be 1, 2, 3, 4 and 6 for unidentate, bidentate, terdentate, tetradentate and hexadentate ligands respectively. Unidentate ligands: Cl, H2O, NH3, Bidentate ligands: Ethylene diamine, Propylene, diamine. Illustration 3: Although NH2·NH2 possesses two electron pairs for donation, it does not act as a chelating agent. Why? (JEE MAIN) Sol: The coordination by NH2·NH2 leads to a three-member strained ring which is highly unstable and hence it does not act as chelating agent. 3.3 Coordination Number (a) T  he number of atoms in a ligand that directly bond to the central metal atom or ion by coordinate bonds is called the coordination number of the metal atom or ion. (b) In other words, it is the number of coordinate covalent bond which the ligands form with the central metal atom or ion. (c)  ome common coordination numbers exhibited by metal ions are 2, 4 and 6. The light transition metals S exhibit 4 and 6 coordination numbers while heavy transition metals exhibit coordination numbers above 6. (d) F or 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. 3.4 Coordination Sphere Coordination Sphere The central metal atom/ion and the ligands directly attached to it Central metal ion Ligand are collectively termed as the coordination sphere. Coordination 2+ [Ni (NH3)4 ] Cl2 - lonization Sphere sphere is represented inside square brackets, e.g. Ni (NH3 )   4 Coordination number 3.5 Oxidation Number The actual charge that a metal atom experiences in a complex is known as its oxidation number. In other words, oxidation number of a metal atom will be equal to the total charge on this atom if all the ligands are removed without their electron pair. Calculation of oxidation number: Algebraic sum of oxidation numbers of all the atoms of a molecule/ion is equal to the charge on it. For example, for Ma [M’b (L)x] [a× (O.N. of M)] + [b × O.N. of M’] + [x × (O.N. of L)] = 0 For ion [Ma (L) x]y+ a × (O.N. of M) + x × (O.N. of L) = y e.g:- Oxidation number of Co (let it be x) in ion [Co (CN) (H2O) (en)2]2+ can be calculated as: x + (O.N. of CN) + (O.N. of H2O) + 2 (O.N. of en) = +2 x + (–1) + (0) + (2×0) = +2 ∴ x = +3  Chem i str y | 28.7 MASTERJEE CONCEPTS (a) Transition metals exhibit variable oxidation states. (b) Oxidation number of different species: (i) Alkali metals (Li, Na, K, Rb, Cs) = +1 (ii) Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) = +2 (iii) Oxidation number of ion = charge on ion (iv) Oxidation number of neutral molecule = 0 Vaibhav Krishnan (JEE 2009, AIR 22) Illustration 4: Specify the oxidation number of the central metals in the following coordination entities: (a) [Co (CN) (H2O) (en)2]2+ (b) [PtCl4]2– (c) [CrCl3 (NH3)3] (d) [CoBr2 (en)2]+ (e) K2 [Fe (CN)6] (JEE MAIN) Sol: (a) +3 (b) +2 (c) +3 (d) +3 (e) +3 3.6 Effective Atomic Number (a) Effective atomic number was first proposed by Sidgwick in order to explain the stability of the complex. (b) It can be defined as the resultant number of electrons present in the metal atom or ion after accepting electrons from the donor atoms of the ligands. (c) In some cases, the effective atomic number coincides with the atomic number of the next noble gas. (d) Effective atomic number is calculated as follows: EAN = atomic number of the metal – number of electrons lost in ion formation + number of electrons gained from the donor atoms of the ligands. MASTERJEE CONCEPTS Calculation of EAN: Effective atomic number = Atomic number (Z) – Electron donated (Equal to O.N.) + Electrons accepted from ligands (2 × No. of coordinate bonds formed) Or EAN = Z–O.N. + 2 × (C.N.) Note: EAN and stability: An ion with central metal atom possessing EAN equal to next inert gas will be more stable. Vaibhav Krishnan (JEE 2009, AIR 22) 2 8. 8 | Co-ordination Compounds Illustration 5: Metal carbonyls having formula M (CO)x where the number of carbonyl units coordinated to metal M are formed by Fe, Cr and Ni. If effective atomic number of each metal is 36, write the formulas of these metal carbonyls. (JEE MAIN) Sol: M (CO)x, In Fe (CO)x EAN = At. No. of Fe + 2 × No. of ligands. ( O.N. of Metals in Metal carbonyls is zero as these compounds are neutral species formed by neutral ligand CO). i.e. CO ∴ 36 = 26+2x ; ∴ x=5 ∴ Formula of iron carbonyl is Fe (CO)5, Similarly, we get Cr (CO)6 and Ni (CO)4. 4. FORMULA AND IUPAC NOMENCLATURE OF COORDINATION COMPOUNDS 4.1 Formula of a Complex (a) In formulas of both simple and complex salts, cation precedes the anion. Nonionic compounds are written as single units. (b) Complex ions are written inside square brackets without any space between the ions. (c) Metal atom and ligands are written in the following order: (i) In the complex part, the metal atom is written first followed by ligands in the order, anionic → neutral → cationic. (ii) If more than one ligand of one type (anionic, neutral or cationic) are present, then they are arranged in English alphabetical order, e.g. between H2O and NH2, H2O should be written first. Similarly, order of NO2− , SO32 − and OH– will be NO2− (iii) W  hen ligands of the same type have similar name for the first atom, then the ligand with less number of such atoms is written first. Sometimes the second atom may be used to decide the order. When number  , NH of atoms are also same e.g., Out of NO − , NH− will be written first. In H and N  will be written first 2 2 3 2 3 as it contains only one N-atom. (iv) P  olyatomic ligands and abbreviations for ligands are always written in lower case letters. e.g. (en), (py), etc. (v) Charge of a complex ion is represented as over script or square bracket. E.g. K4 [Fe (CN)6] — First cation and then anion [Rule 1 and 2] [CrCl2 (H2O)4] Br—Cl– (negative ligand) before H2O (neutral ligand [Rule 3-(i)] 4.2 Nomenclature of Coordination Compounds Mononuclear coordination compounds are named by following these rules: (a) In both the positively and negatively charged coordination compounds, the cation is named first followed by the anion. (b) T  he ligands are named in alphabetical order before the name of the central atom/ion. (This procedure is reversed in writing its formula).  Chem i str y | 28.9 (c) Names of the anionic ligands end in –o. E.g. Symbol Name as ligand Symbol Name as ligand N 3– Azido OH – Hydroxo Cl– Chloro CO32– Carbonato O– Peroxo Oxalato C2 O −42 Br– Bromo Sulphato SO −42 O2H– Perhydroxo Nitrato NO3− CN– Cyano SO32 − Sulphito S2– Sulphido CH3COO– Acetato O2– Oxo (Bonded through oxygen) nitrite NO2− NH2− Amido (Bonded through nitrogen) nitro (d) N  ames of neutral and cationic ligands are the same except for aqua for H2O, ammine for NH3, carbonyl for CO and nitrosyl for NO. These are placed within parentheses ( ). Symbol Name as ligand Symbol Name as ligand H 2O Aqua NO Nitrosyl NH3 Ammine CS Thiocarbonyl CO Carbonyl (e) Positive ligands are named as: Symbol Name as ligand NH+4 NO+ Nitrosylium NH2NH3+ Hydrazinium (f) P  refixes mono, di, tri, etc. are used to indicate the number of the individual ligands in a coordination compound. When the names of the ligands include a numerical prefix, then the terms, bis, tris, tetrakis are used, and the ligand to which they refer is placed in parentheses. For example, [NiCl2 (PPh3)2] is named as dichlorobis (triphenylphosphine) nickel (II). (g)  xidation state of the metal in a cation, anion or a neutral coordination compound is indicated by a Roman O numeral in parenthesis. (h) When the complex ion is a cation, the metal is named same as the element. For example, Co in a complex cation is called cobalt and Pt is called platinum. In an anion, Co is called cobaltate. For some metals, their Latin names are used in the complex anions, e.g. ferrate for Fe. (i) Nomenclature of a neutral complex molecule is done in the similar way as that of a complex cation. 2 8. 10 | Co-ordination Compounds The following examples illustrate the nomenclature for coordination compounds: [Cr (NH3)3 (H2O)3] Cl3 is named as: Triamminetriaquachromium (III) chloride [Co (H2NCH2CH2NH2)3]2 SO4 is named as: Tris (ethane-1, 2-diammine) cobalt (III) sulphate [Ag (NH3)2] [Ag (CN)2] is named as: Diamminesilver (I) dicyanoargentate (I) ( j) Ligands which join two metals are known as ‘Bridge ligands’ and they are prefixed by ‘µ’ (mu). NH2 E.g. (NH3)4Co (NO33))44, in this complex Co (NH3)4 (NO NO2 Here, NH2 and NO2 are bridge ligands and they are named µ-amido and µ-nitro, respectively. Illustration 6: Write the formulas for the following coordination compounds: (JEE MAIN) (A) Tetraammineaquachloridocobalt (III) chloride (B) Potassium tetrahydroxidozincate (II) (C) Potassium trioxalatoaluminate (III) (D) Dichloridobis (ethane-1, 2-diamine) cobalt (III) Sol: (A) [Co (NH3)4 (H2O) Cl] Cl2 (B) K2 [Zn (OH)4] (C) K3 [Al (C2O4)3] (D) [CoCl2 (en)2]+ Illustration 7: Write the IUPAC names of the following coordination compounds: (JEE MAIN) (A) [Pt (NH3)2Cl (NO2)] (B) K3 [Cr (C2O4)3] (C) [CoCl2 (en)2] Cl (D) [Co (NH3)5 (CO3)] Cl (E) Hg [Co (SCN)4] (F) [Ni (CO)4] Sol: (A) Diamminechloridonitrito-N-platinum (II) (B) Potassium trioxalatochromate (III) (C) Dichloridobis (ethane-1, 2-diamine) cobalt (III) chloride (D) Pentaamminecarbonatocobalt (III) chloride (E) Mercury tetrathiocyanatocobaltate (III) (F) Tetracarbonylnickel (0) 5. ISOMERISM IN COORDINATION COMPOUNDS Compounds having the same molecular formula but a different arrangement of atoms and different properties are called isomers and the phenomenon is called isomerism. Types of isomerism exhibited by complex compounds are summarized below: Isomerism Structural isomerism Stereoisomerism Ionization Salt or Linkage Coordination Geometrical Optical Isomerism Isomerism Isomerism Isomerism Isomerism Hydrate Polymerization Coordination Isomerism Isomerism Position Isomerism  Chem i str y | 28.11 5.1 Structural Isomerism Structural isomerism occurs due to the difference in chemical linkages and distribution of ligands within and outside the coordination sphere. In Structural isomerism, isomers possess dissimilar bonding pattern. Different types of isomers are discussed below: (a) Ionization Isomerism: Ionization isomerism is the result of the exchange of groups or ions between the coordinating sphere and the ionization sphere. This isomerism occurs only in compounds where counter ions act as potential ligands. Ionization isomers exhibit different physical as well as chemical properties. Ionisation [CoBr(NH3 )5 ]SO 4  →[CoBr(NH3 )5 ]2 + + SO2– 4 ( A ) Red violet Pentaamminebromocobalt (III) sulphate Ionisation [Co(SO 4 )(NH3 )5 ]Br  →[Co(SO 4 )(NH3 )5 ]+ + Br – (B) Red Pentaamminesulphatocobalt (III) bromide Here (A) and (B) are ionization isomers. (A) forms white precipitate (BaSO4) with BaCl2 whereas (B) does not react with BaCl2. Similarly (B) gives yellowish white precipitate (AgBr) with AgNO3 while (A) does not react with AgNO3. Other examples of ionization isomers are: (i) [PtCl2 (NH3)4] SO4 and [Pt (SO4) (NH3)4] Cl2 (ii) [CoCl2 (NH3)4] NO2 and [CoCl (NO2) (NH3)4] Cl (iii) [Pt (OH)2· (NH3)4] SO4 and [Pt (SO4) (NH3)4] (OH)2 (b) Hydrate Isomerism (Solvate Isomerism): In a complex compound, water molecules behave in two ways: (i) W  ater molecules which behave as ligands are coordinated with the metal atom and are part of the complex ion, e.g. [M (H2O)x]. (ii) W  ater molecules act as water of crystallization and these appear outside the coordination sphere, e.g. [MLx].nH2O. Isomerism which occurs due to dissimilar number of water molecules as ligands (inside the sphere) and as water of crystallization (outside the sphere), is known as hydrate isomerism. This isomerism is analogous to ionization isomerism, in which water molecules inside and outside the sphere are exchanged. For example, Cr (H2O)6Cl3 has three possible structures: [Cr (H2O)6] Cl3 (violet) [Cr (H2O)5Cl] Cl2H2O (green) [Cr (H2O)4Cl2] Cl.2H2O (dark green) These complex compounds differ from one another with respect to the number of water molecules acting as ligands. Other hydrate isomers are: [Co (NH3)4 (H2O)Cl]Cl2. [Co (NH3)4Cl2] Cl H2O 2 8. 12 | Co-ordination Compounds (c) Linkage or Salt Isomerism: (i) Linkage isomerism occurs in complex compounds having ambidentate ligands like —CN, —NC, —NO2, —ONO, —CNO, —NCO, —CNS, —NCS, —SCN, etc. (ii) In this isomerism, an ambidentate ligand coordinates with different atoms. (iii) These isomers can be differentiated by IR spectroscopy. For example, [Co (NO2) (NH3)5] Cl2 and [Co (ONO) (NH3)5] Cl2 (A) (B) Pentaamminenitrocobalt (III) chloride Pentaammine nitritocobalt (II) chloride (Yellow-red) (Red) (A) is not decomposed by the action of acids whereas (B) liberates HNO3 by the action of acid. Other examples of linkage isomers are: (i) [Cr (SCN) (H2O)5]2+ and [Cr (NCS) (H2O)5]2+ (ii) [Co (NO2) (py)2 (NH3)2] NO3 and [Co (ONO) (py)2 (NH3)2] NO3 (d) Polymerization Isomerism: When two compounds possess stoichiometric composition but different molecular formulas, they are known as polymerization isomers of each other. Molecular formula of one isomer will be the integral multiple of the other one. Example: [PtCl2 (NH3)2] and [Pt (NH3)4] [PtCl4] (e) Coordination Isomerism: (i) This isomerism occurs only in those complexes in which both cation and anion are complex. (ii) It occurs as a result of the exchange of ligands between the cation and anion. (iii) It may occur in those complexes also in which both cation and anion have the same metal atoms. Example: (i) [Cr (NH3)6] [Cr (SCN)6] and [Cr (SCN)2 (NH3)4] [Cr (SCN)4 (NH3)2] (ii) [Co (NH3)6] [Cr (C2O4)3] and [Cr (NH3)6] [Co (C2O4)3] (f) Coordination Position Isomerism: It occurs in complexes containing bridge ligands and is the result of OH dissimilar arrangement of metal atoms forming bridge, e.g. (NH3)4Co Co (NH3)2Cl2 SO4 and Cl OH Cl (NH3)4Co Co (NH3)3Cl SO4 Cl 5.2 Stereoisomerism Stereoisomerism occurs as a result of the different arrangements of ligands around the central metal atom. It may be of two types: (1) Geometrical isomerism and (2) Optical isomerism.  Chem i str y | 28.13 5.2.1 Geometrical Isomerism Isomerism which occurs due to different relative arrangements of ligands around the central metal atom is known as geometrical isomerism. Geometrical isomers are of two types: (a) Cis-isomer: In a disubstituted complex molecule/ion, when two similar ligands are at right angle (90º), the geometrical isomer is known as Cis-isomer. (b) T  rans-isomer: When two ligands are positioned in opposite directions, i.e. at 180º to each other, the isomer formed is trans-isomer. Cis- and Trans- positions are indicated in figures: 5 4 1 4 1 M M 3 2 3 2 6 Square planar Octahedral Cis- positions: (1, 2), (2, 3), (3, 4), (1, 4) (1, 2), (2, 3), (3, 4), (1, 4), (1, 5), (4, 5), (3, 5), (2, 5), (1, 6), (2, 6), (3, 6) and (4, 6) Trans- positions: (1, 3) and (2, 4) Geometrical Isomerism and Coordination Numbers Geometrical Isomerism with Coordination Number 4: Tetrahedral complexes do not show geometrical isomerism as all the four valences are identical. Square–planar complexes: (a) C  omplexes of type MA4, MA3B and MAB3 do not show geometrical isomerism, where A and B are monodentate ligands. (b) C  omplexes of formula MA2B2 and MA2BC types have two geometrical isomers, where A and B are monodentate ligands. Example: (i) [PtCl2 (NH3)2] resembles MA2B2 in formula and exists in two isomeric forms: Cl NH3 Cl NH3 Pt Pt Cl NH3 NH3 Cl Cis-isomer Trans-isomer (light yellow) (Dark yellow) (ii) [PtCl (NH3) (py)2] resembles MA2BC and exists in two isomeric forms: Py NH3 NH3 Py Pt Pt Py Cl Py Cl Cis Trans 2 8. 14 | Co-ordination Compounds (c) Complexes of formula MABCD exist in three isomeric forms: A B A C A C A B M M M M D C D B B D C D (I) (II) (III) (IV) (III) and (IV) are similar. e.g. [Pt (NO2) (NH2OH) (NH3) (py)] + exists in 3 isomeric forms. A = NO2, B = NH2OH, C = NH3, D = py (d) If A is an unsymmetrical bidentate ligand, then compounds having formula MA2 tend to exhibit geometrical isomerism, e.g. [Pt (gly)2]  —CH —COO–) gly = glycinate ( NH 2 2 CH2—NH2 NH2—CH2 OC O NH2—CH2 Pt Pt and OC O O CO CH2—NH2 O CO Cis Trans (e) Bridged dinuclear complexes of formula M2A2B4 also exhibit geometrical isomerism, e.g. PtCl2 P ( C6H5 )3   2 ( ) (C6H5)3P Cl Cl Cl Cl Cl Pt Pt Pt Pt Cl Cl P(C6H5)3 and (C6H5)3P Cl P(C6H5)3 Trans Cis Geometrical Isomerism with Coordination Number 6: (a) Complexes of type MA6 and MA5B type do not show geometrical isomerism. (b) Complexes of type MA4B2 or MA4BC exist in two isomeric forms, e.g. [CoCl2· (NH3)4]+ Cl Cl NH3 Cl NH3 NH3 + + Co Co NH3 NH3 NH3 NH3 NH3 Cl Cis Trans (c)  omplexes of type MA3B3 exist in two geometrical forms which are named as facial (fac–) and meridonial C (mer–) isomers. When three ligands of the same type are arranged in one triangular face, then isomer is facial. fac- and mer- isomers of complex MA3B3 are as follows: A B B A A A M M B A B A B B fac-isomer mer-isomer  Chem i str y | 28.15 E.g. [Co (NO2)3) (NH3)3] can be represented in fac- and mer- isomeric forms as follows: NO2 NH3 NH3 NO2 NO2 NO2 Co Co NH3 NO2 NH3 NO2 NH3 NH3 fac-isomer mer-isomer Similarly, [RhCl3 (py)3] also exists in fac- and mer- forms. (d) C  omplex compound of formula MABCDEF may exist in 15 isomeric forms and only one compound of this type is identified so far [Pt (Br) (Cl) (I) (NO2) (NH3) (py)]. (e) C  omplexes of formula M(AA)2B2 and M(AA)2BC also exhibit geometrical isomerism, where A is the symmetrical bidentate ligand, e.g. ethylenediamine (en), oxalate (ox), etc. [CoCl2 (en)2]+ Cl Cl Cl en CO+ en CO + en en Cl Cis-isomer Trans-isomer (f) Complex of type M(AA')3 also exists in Cis- and Trans- forms. Where AA’ is unsymmetrical bidentate ligand, e.g. [Cr (gly)3], gly: glycinate (NH2CH2COO–) O O CH2 C O CH2 C O NH2 NH2 NH2 NH2 CH2 CH2 + + CO CO C= O C= O O O NH2 O O=C NH2 CH2 C O CH2 = O Cis-isomer Trans-isomer Illustration 8: Draw the structure of geometrical isomers of [Pt (gly)2] where gly is NH2CH2COO–. (JEE ADVANCED) Sol: CH2 NH2 NH2 CH2 and CH2 CO NH2 O Pt Pt OC O O CO OC O H2N CH2 Cis-isomer trans-isomer 5.2.2 Optical Isomerism Optical activity: Compounds which rotate on the plane of polarized light are optically active. If the plane rotates clockwise, then the isomer is said to be dextro rotator (d or +) and if the plane rotates anticlockwise then the isomer is said to be laevo rotator (l or -). Equimolar mixture of d– and isomer is optically inactive and is called racemic mixture. Optical isomers differ in optical properties. 2 8. 16 | Co-ordination Compounds (a) Optical isomerism in complexes with coordination number 4: (i) T  etrahedral complexes: Like carbon compounds, complex MABCD must be optically active but due to their labile nature, such complex cannot be resolved in d or l form. However, tetrahedral complexes with unsymmetrical bidentate ligand are optically active. In optically active tetrahedral compounds, the ligand must be unsymmetrical. It is not necessary whether it is chiral (asymmetric) or not, e.g. bis (benzoyl acetonato) beryllium (II) H5C6 C6H5 H5C6 C6H5 C= O O=C C= O O=C CH Be CH CH Be CH and C O O C C O O C CH3 CH3 CH3 CH3 Dextro Leavo Another example of this type is [Ni (CH2NH2COO)2]—bis (glycinato) nickel (II) O O O O Ni Ni N N N N NO2 Illustration 9: Draw all the optical isomers for [(en)2Co Co(en)2]4+ (JEE MAIN) NO2 Sol: Complex compound shows optical isomerism and exists in d l and meso forms. 4+ en en 4+ en en en 4+ NO2 NO2 NO2 (1) CO CO 2) CO CO (3) en CO CO NO2 NO2 NO2 en en en en en en I and II d and l form (mirror image of each other), III meso-form (ii) S  quare planar complexes: Generally square planar complexes are not optically active as they have all the ligands and metal atoms in one plane. That is why there is a plane of symmetry. Note: However some optically active square planar complexes are identified, e.g. isobutylenediaminemesostilben- ediaminoplatinum (II) ion. 2+ C6H5 CH NH2 NH2 CH NH2 Pt C6H5 CH NH2 NH2 CH C6H5 (a) Optical isomerism in compounds of coordination number 6 – Octahedral complexes: (i) C  omplexes of type MA4R2 exist in cis- and trans- forms and both forms are optically inactive due to plane of symmetry. (ii) Complexes of type MA3B2 exist in facial and meridonial forms but both are optically inactive. (iii) C  omplexes of type MA2B2C2 are optically active, e.g. five geometrical isomers of [PtCl2 (NH3)2 (py) 2]2+ are possible. Out of these five possible isomers, three have been prepared. Their cis- form is optically active while trans- forms are optically inactive due to symmetry.  Chem i str y | 28.17 py 2+ py 2+ Cl NH3 py Cl Pt Pt Cl NH3 NH3 Cl Py NH3 Cis-isomer Trans-isomer (iv) C  omplex MABCDEF has 15 geometrical isomers and each isomer exists as pair of enantiomers and hence total 30 optical isomers will be possible. Only one such compound has been identified so far – [Pt (Br) (Cl) (I) (NO2) NH2) (py)]. MASTERJEE CONCEPTS Number of Possible Isomers for Specific Complexes Formula Number of Stereoisomers Pairs of Enantiomers MA6 1 0 MA5B 1 0 MA4B2 2 0 MA3B3 2 0 MA4BC 2 0 MA3BCD 5 1 MA2BCDE 15 6 MABCDEF 30 15 MA2B2C2 6 1 MA3B2CD 8 2 MA3B2C 3 0 M (AA) BCDE 10 5 M (AB)2 CD 11 5 Saurabh Gupta (JEE 2010, AIR 443) 2 8. 18 | Co-ordination Compounds MASTERJEE CONCEPTS Number of Possible Isomers for Specific Complexes Formula Number of Stereoisomers Pairs of Enantiomers M (AB) (CD)EF 20 10 M (AB)3 4 2 M (ABA)CDE 9 3 M (ABC)2 11 5 M (ABBA)CD 7 3 M (ABCBA)D 7 3 Saurabh Gupta (JEE 2010, AIR 443) 6. PREPARATION AND IDENTIFICATION OF COMPLEX COMPOUNDS Preparation of Complex Compounds (a) By substitution reactions: Ion of a salt can be substituted by a ligand to form a complex compound, e.g. CuSO 4 + 4NH3 → [Cu(NH ) ]SO 3 4 4 tetraamminecopper(II)sulphate (b) By combination reaction: Various complexes can be formed by combination reactions:  → [Ni(NH ) ]Cl NiCl2 + 6NH 3 3 6 2  → [Ag(NH ) ]Cl AgCl + 2NH 3 3 2 (c) By redox reactions: Two important examples are:  + 2NH NO + H O → 2[Co(NO )(NH ) ](NO ) + 14H O 2[Co(H2 O)6 ](NO3 )2 + 8NH 3 4 3 2 2 3 3 5 3 2 2 2CoCl2 + 2NH4 Cl + 10NH3 + H2 O2 → 2[Co(NO3 ) 6 ]Cl3 + 2H2 O + 9H2 Identification of Complex Compound Formation (a) Change in solubility: Solubility of a complex compound changes abnormally when complex is formed, e.g. AgCN+ KCN → K[Ag(CN)2 ] (Solubility increases) Partialy soluble Soluble AgCl+ 2NH3 → [Ag(NH3 )2 ]Cl (Solubility increases) Insoluble Soluble Ni2 + + 2dmg → [Ni(dmg)2 ]2 + (dmg = dimethylglyoxime) Soluble Insoluble  Chem i str y | 28.19 (b) C  hange in conductivity: As complex formation changes the solubility, the number of ions in solution also changes and hence conductance changes suddenly. (c)  hange in chemical properties: Change in chemical properties of metal ion also indicates the formation of C complex, e.g.  , KCl does not precipitate Ag+ due to formation Ag+ is precipitated by KCl solution but in the presence of NH 3 of [Ag (NH3)2] Cl. (d) Change in color: Change in color indicates complex formation, eg. Co2 + + 4Cl− → [CoCl4 ]2– Pink Blue  → [Cu(NH ) ]2 + Cu2 + + 4NH Lightblue 3 3 4 Deepblue [Ni(H2 O)6 ]2 + + 6NH3 → [Ni(NH3 )6 ]2 + + 6H2 O Green Blue (e) Change in pH, EMF, Magnetic properties or colligative properties also indicate the complex formation. 7. THEORIES OF BONDING IN COORDINATION COMPOUNDS 7.1 Werner’s Theory In 1898, Werner propounded his theory of coordination compounds. Werner proposed the concept of a primary valence and a secondary valence for a metal ion. Main postulates of his theory are: (a) In coordination compounds metals show two types of linkages (valencies) – primary and secondary. (b) The primary valencies are normally ionizable and are satisfied by negative ions. (c)  he secondary valencies are nonionizable. These are satisfied by the neutral molecules or negative ions T (ligands). The secondary valency is equal to the coordination number and is constant for a metal. (d) T  he ion groups bounded by the secondary linkages to the metal have a characteristic spatial arrangement corresponding to their different numbers. Illustration 10: PtCl4 and NH3 may form five complexes, A (PtCl4·6NH3), B (PtCl4·5NH3), C (PtCl4·4NH3), D PtCl4·3NH3 and E (PtCl4·2NH3). One mole of each A, B, C, D and E reacts with excess of AgNO3 to yield 4, 3, 2 and 1 mole AgCl respectively, while E gives no AgCl. The conductances of their solutions are in the order A > B > C > D > E. On the basis of Werner’s theory, write their structure and give the total number of ions given by one complex. (JEE ADVANCED) Sol: Formula Structural formula Ionization No. of ions (A) PtCl4·6NH3 [Pt (NH3)6] Cl4 [Pt (NH3)6]4+ + 4Cl– 5 (B) PtCl4·5NH3 [PtCl (NH3)5] Cl3 [PtCl (NH3)5]3+ + 3Cl– 4 (C) PtCl4·4NH3 [PtCl2 (NH3)4] Cl2 [PtCl2 (NH3)4]2+ + 2Cl– 3 (D) PtCl4·3NH3 [PtCl3 (NH3)3] Cl [PtCl3 (NH3)3]+ + Cl– 2 (E) PtCl4·2NH3 [PtCl4 (NH3)2] No isonisation possible 0 2 8. 20 | Co-ordination Compounds 7.2 Valence Bond Theory The salient features of the valence bond theory are summarized below: (a) T  he 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. (b) T  he atomic orbitals (s, p or d) of the metal ion hybridize to form hybrid with definite directional properties. These hybrid orbitals tend to form strong chemical bonds with the ligand orbitals. (c)  he d-orbitals involved in the hybridization may be either inner (n – 1) d orbitals or outer n d-orbitals. The T complexes formed in these two ways are referred to as low spin and high spin complexes, respectively. (d) Each ligand possesses a lone pair of electrons. (e) A  covalent bond is formed when a vacant hybridized metal orbital and a filled orbital of the ligand overlap. The bond is also known as a coordinate bond or dative bond. (f)  hen a complex contains unpaired electrons, it is paramagnetic in nature, whereas if it does not contain W unpaired electrons, it is diamagnetic in nature. (g) The number of unpaired electrons in the complex, determines the geometry of the complex as well as  hybridization of the 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. µ = n(n + 2) where n = no. of lone pair. (h) Thus the knowledge of the magnetic moment can be of great help in ascertaining the type of complex. (i) U  nder the influence of a strong ligand, the electrons can be forced to pair up against the Hund’s rule of maximum multiplicity. Coordination Number Type of Hybridization Distribution of hybrid orbital in space 4 sp3 Tetrahedral 4 dsp2 Square planar 5 sp3d Trigonal bipyramidal 6 sp3d2 Octahedral 6 d2sp3 Octahedral Application of Valence Bond Theory on Coordination Complexes (a) C  omplex with Coordination Number 4: Tetra coordinated complexes have either tetrahedral or square planar geometry depending on the nature of orbitals involved in hybridization. If one ns and three np orbitals are involved in bonding, geometry will be tetrahedral and hybridization sp3. If (n – 1) d, ns and two np are involved in bonding, geometry will be square planar and hybridization dsp2. Tetra coordinated complexes are common with Ni (II), Cu (II), Pt (II), Pd (II), etc. (n-1)d ns np 4dsp² hybrid orbital, squar e planar geometry (n-1)d ns np 4dsp² hybrid orbital, tetrahedral geometr y Some examples of tetra coordinated complexes are given below:  Chem i str y | 28.21 (i) Tetrahedral Complexes: Ni (CO)4: In Ni (CO)4, Ni has zero oxidation state and exists as Ni (0). Four ligands (CO) are attached to central metal atom Ni and require four orbitals. The electronic configuration in Ni (CO)4 can be written as: 8 2 0 3d 4s 4p 8 2 28Ni atom = 3d , 4s When four CO ligands are present, it is a strong ligand and the electrons pair up against “Hund’s rule for maximum multiplicity”. 10 0 0 3d 4S 4p 3 sp hybridization 3 CO 4sp hybridization Ni CO CO CO CO OC CO Four 1CO pairs from four CO CO molecules Explanation: Four sp3 hybrid orbitals are arranged tetrahedrally making it a tetrahedral complex. Since all the electrons are paired, it is diamagnetic. (ii) Square planar complexes: [Ni (CN)4]2–: Here, Ni is in (II) oxidation state and the electronic arrangement is as follows: 8 2 0 3d 4s 4p Ni-[Ar]: 3d8. 4s2 NC NC 2- 2- CN CN 4dsp2 hybrid orbital, square planar geometry Ni Ni Ni2+ ion [Ar] 3d8, 4s0 NC NC CN CN 4dsp2 hybrid orbital, 3d 8 tetrahedral geometry 3 4dsp hybrid orbital Ni2+ ion in [Ni (CN) 4]2– 3d8, 4s0 CN CN CN CN Explanation: CN– is a strong ligand and so it pairs up 3d-electrons against Hund’s rule. The d-orbital thus made vacant, takes part in hybridization. Four dsp2 hybrid orbitals are arranged in this manner and hence the geometry is square planar. Complex compounds are diamagnetic because all the electrons are paired. The complex makes use of the inner d-orbital, and so it is known as inner orbital or low spin or hyper ligand or spin paired complex. 2 8. 22 | Co-ordination Compounds MASTERJEE CONCEPTS Exception: Structure of [Cu (NH3) 4]2+ ion: It is an exceptional case which involves sp2 d hybridization. Here, Cu is tetra coordinated and may exist as square planar or tetrahedral complex. Physical measurement have indicated that tetrahedral geometry for [Cu (NH3)4]2+ is not possible. If square planar geometry is supposed to be correct, then the following electronic arrangement must be followed: 3d 10 4s 1 4p Cu [Ar]: 3d10, 4s1 Cu2– ion: → [Ar]3d9, 4s0 3d 9 4s 1 0 4p For dsp2 hyb.3d-electron must be excited to 4-p with the following configuration. 2+ NH3 NH3 Cu2+ ion in [Cu (NH3)4]2+: Cu 2 NH3 NH3 dsp hybridisation NH3 NH3 NH3 NH3 Now if the above configuration is correct, the unpaired electrons present in higher energy, 4-p orbital should be expected to be easily lost and Cu2+ must be easily oxidized to Cu3+, but it never occurs, so the configuration is not satisfactory. To explain it Huggin suggested sp2d hybridization. Cu2+ ion in [Cu (NH3)4]2+ : Note: Pt (II) and Au (III) always form square planar complexes irrespective of their ligands being strong or weak. Neeraj Toshniwal (JEE 2009, AIR 21) (b) Complexes with Coordination Number 6: Hexacoordinated complexes are of two types, inner orbital complexes and outer orbital complexes. They possess octahedral geometry. (i) Inner orbital complexes: In this type of complexes the d-orbitals used are of lower quantum number, i.e. (n – 1). Some examples are given below: Complexes formed by using the inner orbitals are diamagnetic or have reduced paramagnetism. These are also known as low spin or spin paired complexes, Example 1: [Fe (CN)6]4– e– configuration of Fe26 = [Ar] 3d64s2 6 0 0 3d 4s 4p e– configuration of Fe+2 = [Ar] 3d6 = e– configuration of Fe+2 after rearrangement = 2 3 d sp The above rearrangement is due to presence of the cyanide ligand.  Chem i str y | 28.23 At this stage, Fe2+ undergoes d2sp3 hybridization to form six d2sp3 hybrid orbitals, each of which accepts an electron pair donated by CN– ions. The complex is Diamagnetic as it has no unpaired electron. Example 2: [CO(NH8)6]3+ Example 3: [Cr (NH3)6]3+ 5 3d 4s 4p Cr → 24 24 Cr Cr3+ → Cr 3+ Cr3+ in d2sp3 hybridized state 2 3 d sp bybridized state As this d2sp3 hybridization leads to octahedral geometry, the complex [Cr (NH3)6]3+ will be octahedral in shape. Since the complex ion has 3 unpaired electrons, it is paramagnetic. Other complexes of chromium with similar inner structure are [Cr (CN)6]3– and [Cr (H2O)6]3+. (ii) Outer orbital complexes In these complexes s, p and d orbitals which are involved in hybridization, belong to the highest quantum number (n). Complex compound formed by the use of outer n and d orbitals will be paramagnetic. Outer orbital complexes are also known as high-spin or spin free complexes. The outer orbital complexes have a high number of unpaired electrons, E.g. [CoF6]3– 3d 4s 4p 4d Co 27 → Co3+ ion → Co3+ ion in sp3d2 hybridized state 3 2 sp d Owing to the octahedral orientation of six sp3d2 hybridized orbitals, shape of [CoF6]3– complex ion is octahedral. As it possesses four unpaired electrons in the 3d orbital, [CoF6]3– ion is paramagnetic. Some other examples are [FeF6]3–, [Fe (NH3)6]2+, [Ni (NH3)6]2+, [Cu (NH3)6]2+, [Cr (H2O)6]3+, etc. Limitations of valence bond theory: Even though the valence bond theory explains the formation, structures and magnetic behavior of coordination compounds to a larger extent, it suffers from the following short comings: It includes a number of assumptions. It fails to provide quantitative interpretation of magnetic data. It lacks explanation to the color exhibited by coordination compounds. It does not provide a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds. It is unable to predict the tetrahedral and planar structures of 4-coordinate complexes accurately. This theory does not distinguish between weak and strong ligands in compounds. 2 8. 24 | Co-ordination Compounds 7.3 Crystal Field Splitting Theory The important terms in Crystal Field Splitting theory are as follows: (a) D  egenerate orbitals: in free state, all the d-orbitals (viz., dxy, dyz, dxz, dx2 − y2 and dz2 ) will possess the same energy and are said to be degenerate. (b) t2g and eg set of orbitals: In a d-subshell, there are five d-orbitals and on the basis of orientation of lobes of these five d-orbitals with respect to coordinates, they have been grouped into two sets. (i) eg set of orbital: dx2 − y2 , and orbitals have their lobes arranged along the axes and they constitute eg set. These orbitals are also called axial orbital. Term eg refers to ‘doubly Degenerate’, according to group theory (e = doubly degenerate set) (ii) t2g set of orbital: This set includes orbitals whose lobes lie between the axes and this set includes dxy, dyz and dxz orbitals. These orbitals are also known nonaxial orbitals. Group theory called these orbital t2g where ‘t’ refers to ‘triply degenerate’. Crystal Field Theory: The crystal field splitting theory (CFT) is an electrostatic model which considers the metal– ligand bond to be ionic occurring purely due to the electrostatic interaction between the metal ion and the ligand. Ligands are treated as point charges in case of anions and dipoles in case of neutral molecules. The five d orbitals in an isolated gaseous metal atom/ion have the same energy, i.e. they are degenerate. This degeneracy is maintained if a spherically symmetrical field of negative charges surrounds the metal atom/ion. However, when this negative field is resulted by the ligands (either anions or the negative ends of dipolar molecules like NH3 and H2O) in a complex, it becomes asymmetrical and the degeneracy of the d orbital is lifted. It results in splitting of the d orbitals. The pattern of splitting depends upon the nature of the crystal field. Let us discuss this splitting in different crystal fields in detail. (a) C  rystal field splitting in octahedral field: The orientation of d-orbital in octahedral field is represented in the diagram. The lobes of t2g, set of orbital (dxy, dyz and dxz) point in between x, y and z axes while lobes of eg set ( d 2 and z dx2 − y2 ) point along the x, y and z axes. Thus, energy of the eg set increases higher than that of the t2g set. The splitting of orbital can be represented by Fig. 4: eg, (dz2, dx2-y2) +6 Dq. CFSE (o) = 10Dq. -4 Dq. Bari centre Energy Degenerate d-orbital n+ In free ion (M ) Figure 28.4: Splitting of d-orbital is a octahedral crystal field The difference in energy of t2g and eg set is known as crystal field splitting energy or crystal field stabilization energy (CFSE), which is represented by ∆o (o stands for octahedral) or 10 Dq. The value of 10 Dq or ∆o can be measured by UV-visible spectrum. The crystal field splitting, ∆o, depends upon the field produced by the ligand and the charge on the metal ion. Some ligands are able to produce strong fields, and correspondingly, the splitting will be large whereas others produce weak fields and these consequently result in small splitting of d orbital. Ligands can be arranged according to their order of increasing field strength as follows: I– < Br– < SCN– < Cl– < S2– < F– < OH– < C2O42– < H2O < NCH– < edta4– < NH3 < en < CN– < CO  Chem i str y | 28.25 This series is known as the spectrochemical series. It is an experimentally determined series based on the absorption of light by complex compounds with various ligands. Let us assign electrons in the d orbital of the metal ion in octahedral coordination entities. Obviously, the single d electron occupies one of the lower energy t2g orbital. In d2 and d3 coordination compounds, the d electrons occupy the t2g orbital singly in accordance with the Hund’s rule. For d4 ions, two possible patterns of electron distribution arise: (a) The fourth electron could either enter the t2g level and pair with an existing electron, or (b) It could avoid paying the price of the pairing energy by occupying the eg level. Either of these two possibilities depends on the relative magnitude of the crystal field splitting, ∆o and the pairing energy, P (presents the energy required for electron pairing in a single orbital). The two options are: (i) If ∆o < P, the fourth electron enters one of the eg orbital exhibiting the configuration t32ge1g. Ligands for which ∆o < P are known as weak field ligands, form high spin complexes. (ii) If ∆o > P, it becomes more energetically favorable for the fourth electron to occupy a t2g orbital with 4 0 configuration t2geg. Such ligands are known as strong field ligands and they form low spin complexes. It is observed from calculations that d4 to d2 coordination entities are more stable for strong field cases compared to their weak counterparts. t  Energy (dz2, dx2-y2) Figure 28.5: Splitting of d-orbital is a octahedral crystal field (b) Crystal field splitting in tetrahedral complexes: The orientation of ligands in a tetrahedral complex is given in fig. 5. Although none of the d-orbitals point towards axes, the t2g set is close to the direction in which ligands are approaching so their energy is higher. The magnitude of ∆t is considerably less than that in the octahedral field, which is mainly due to two reasons: (i) In tetrahedral complex, number of the ligands is only four instead of six. (ii) In tetrahedral complexes, the direction of the orbitals does not coincide with the direction of the ligands, both the factors reduce the CFSE by 2/3 and so ∆t is roughly 4/9 times to ∆o. (c) Crystal Field Splitting in Square Planar Complexes: The square planar geometry can be considered to be derived from the octahedral by removing negative charges from the z-axis. As these negative charges are removed, dxy, dxz and dyz orbital, all of which have a Z-component become more stable as shown in Fig. below. This type of splitting can be further explained as follows: As the lobes of point towards the ligands, this orbital has the highest energy. The lobes of dxy orbital lie between the ligands but are coplanar with them, hence this orbital has the second highest energy. The lobes of dx2 orbital point out of the plane of the complex but the belt around the center of the orbital (which contains about 1/3rd of the electron density) lies in the plane. Therefore, dz2 orbital is next highest in energy. The lobes of dxz and dyz orbital point out of the plane of the complex, and so they are least affected by the electrostatic field of the ligands, they degenerate and have the lowest in energy. 2 8. 26 | Co-ordination Compounds Energy t Figure 28.6: Splitting of d-orbital in a square planar crystal Planar MASTERJEE CONCEPTS Weak ligands favor high spin complexes because they cannot pair up the electrons against Hund’s rule while strong ligands favor low spin complexes. Vaibhav Krishnan (JEE 2009, AIR 22) Illustration 11: Mn2( aq. + ) ion is light pink colored while [Mn (CN)6] is blue in color. Explain. 4– (JEE ADVANCED) Sol: In complexes, where Mn (II) is present, configuration of a metal ion is d5. There may be two types of spin arrangements in the presence of different kinds of ligands. (A) High spin complex (with weak field ligands) and (B) Low spin complex (with strong field ligands) The arrangement of electrons in these complexes can be depicted as: eg d-d transition spin allowed laportae selection rule d-d transition spin forbidden (l = l)forbidden and laport selection rule (l = l)forbidden In high spin complex compounds, it is observed that d–d transition requires reversion of spin which is against the spin selection rules and this makes them spin forbidden and the intensity of color is of only about 1/100 when the transition is allowed. In [Mn (CN)6]4– on the other hand, d–d transitions do not have any such restrictions and are spin allowed. Intense color also is observed when transition takes place.  Chem i str y | 28.27 7.3.1 Factors Affecting CFSE (a) N  ature of ligand: The value of ∆ depends upon the nature of ligands. Ligands with a small degree of crystal field splitting capacity are termed as weak field ligands and those ligands which cause large splitting are called strong field ligands. In general, ligands can be arranged in the ascending order of CFSE caused by them. This series remains practically constant for different metals and is known as spectrochemical series. It is an experimentally determined series. The order is difficult to explain due to involvement of both σ and π bonding. Some ligands in spectrochemical series are given below: I– < Br– < S2– < Cl– < N3− , F– < Urea, OH– < Oxalate, O2– < H2O < NCS– < EDTA < py, NH3 < en = SO32- < bipy, phen < NO2− < CH3− < C6H5− < CN– < CO. For strong field ligands, the order depends on the donor atom and is in the following order: C-donor > N-donor > O-donor > Halogen donor (b) G  eometry of the Complex: ∆t is approximately 4/9 times of ∆o. The lower value of ∆t is due to lesser number of ligands in tetrahedral complex. Also, in tetrahedral complexes the orbital does not point toward the axes, resulting in less interaction. (c) Oxidation state of metal ion: It is observed that the higher the charge on the central metal atom  (or oxidation state), the higher the CFSE. E.g., ∆o for [Fe (H2O)6]3+ is greater than [Fe (H2O)6]2+, ∆o for [Co (H2O)6]3+ is greater than [Co (H2O)6]2+ and ∆o for [V (H2O)6]2+ is greater than [Cr (H2O)6]3+. (d) N  ature of metal ion: The value of CFSE is also determined by the transition series to which the metal belongs and the order for this is observed to be 3d < 4d < 5d. The value of ∆ increases by 30% to 50% for 3d to 4d series and from 4d to 5d series. Hence, metals of 4d and 5d series have more tendency to form low spin complexes, e.g. CFSE for the given complexes follow the order: [Co (NH3)6]3+ < [Rh (NH3)6]3+ < [Ir (NH3)6]3+] When two metal ions possess the same charge but different number of d-electrons, the magnitude of ∆o decreases with increase in the number of d-electrons in the central metal atom. E.g. ∆o for [Co (H2O)6]2+ is greater than ∆o for [Ni (H2O)6]2+ because Co2+ possesses 3d7 configuration while Ni2+ has 3d8 configuration. 7.3.2 Applications of CFSE (a) M  agnetic character of complexes: Complexes containing unpaired electrons tend to be attracted by magnetic fields and hence known as paramagnetic. In contrast, when all the electrons are paired, the complex is slightly repelled by a magnetic field and is said to be diamagnetic. The Weak field Strong field magnetic moment of a transition metal wholly depends on the number 2+ 6 2+ 6 Fe (3d ) ion in Fe (3d ) ion in of unpaired electrons and is equal to n(n + 2) B.M., where n is number [Fe(H2O)6] 2+ [Fe(CN)6] 4- (paramagnetic) of unpaired electrons. For diamagnetic substance, the magnetic moment (diamagnetic) will be zero. Magnetic moments of coordination compounds can be experimentally determined and this data provide information to examine the nature of coordination entities further. These measurements are termed as magnetic susceptibility measurements. For example, [Fe (H2O)6]2+ is paramagnetic while [Fe (CN)6]4– is diamagnetic. This observation can be explained on the basis of the electronic configurations of Fe2+ in [Fe (H2O)6]2+ and [Fe (CN)6]4–. H2O is a weak field ligand while CN– is a strong field ligand. So [Fe (CN)6]4– is the inner orbital low spin complex whereas [Fe (H2O)6]2+ is an outer-orbital high spin complex. The configurations of Fe2+ in both the compounds are further explained in the diagram here. 2 8. 28 | Co-ordination Compounds MASTERJEE CONCEPTS Complexes possessing d0 or d10 configuration of a metal ion are always diamagnetic. Nikhil Khandelwal (JEE 2009, AIR 94) (b) C  olour of complexes: In many complexes the d-orbital split takes place in the two sets t2g and eg, which possess different energies. The difference in energies of t2g and e and eg lies in the visible region of the spectrum and this helps g E = energy Lying in transition metal complexes to absorb color. This makes them E +hv energy (E) E visible colored complementary to the color absorbed. This transition reagion. involves t2g and eg sets of d-orbital and is called as d–d transition. Thus d–d transition is responsible for the color of Ground state Excited state transition metal complexes. d–d transition can be represented diagrammatically as shown here. MASTERJEE CONCEPTS 2− 2− Complexes like CrO 4 , Cr2 O7 , and MnO 4 , etc. have d0 configuration of the metal ion but still exhibit intense − color. Here the color is caused by the charge transfer spectra (CT) and not by the d–d transition. Saurabh Gupta (JEE 2010, AIR 443) 7.3.3 Stability of Complexes Complexes normally exhibit two kinds of stabilities: (i) Thermodynamic stability and (ii) Kinetic stability. Thermodynamic stability deals with the metal–ligand bond energy, stability constants, redox potentials, etc., that affect the equilibrium. On the basis of thermodynamic stability, Blitz classified the complexes into stable or penetration complexes and unstable or normal complexes. Kinetic stability deals with the rates of reaction of complexes in a solution. On the basis of kinetic stability, Taube classified the complexes into labile and inert complexes. Ligands of labile complexes are easily replaceable while ligands of inert complexes cannot be replaced with easily. Chelate effect: Complexes containing chelate rings are more stable, e.g. [Ni (NH3)6]2+ and so is less stable than [Ni (en)3]2+. Macrocylic effect: When a multidentate ligand is cyclic without any considerable steric effect, then the complex formed is more stable than acyclic ligand. This phenomenon is called the macrocyclic effect. Illustration 12: [Cu (CN)4]2- is a more stable complex than [Cu (NH3)4]2+. Why? (JEE MAIN) Sol: The higher stability constant K = 2 × 1027 for Cu2+ + 4CN– → [Cu (CN)4]2– than for [Cu (NH3)4]2+ (which is 4.5 × 1011) explains stability. Also CN– is stronger ligand than NH3.  Chem i str y | 28.29 8. ORGANOMETALLIC COMPOUNDS Organometallic compounds are defined as compounds in which carbon forms a bond with an atom (metal/non- metal) which is less electronegative than carbon. These compounds are classified into two – covalently bonded compounds and ionic organometallic compounds. Covalently bonded compounds: In covalently bonded compounds, the metal and carbon atoms are attached to each other by a covalent bond. They can be further classified into: three groups: (i)  (sigma) bonded complexes: A σ-bonded complex consists of σ H3C CH3 CH3 a metal atom and a carbon atom of the ligand joined together with a σ bond. In another words, the ligand contributes one Al Al electron and is called one electron donor. Tetramethyltin, (CH3)4Sn H3C CH3 Ch3 and trimethyl aluminum, (CH3)3 Al are examples of σ-bonded Trimethyl aluminium organometallic compounds. The latter exists as dimmer and has a structure analogous to diborane. In this, two methyl groups bridge between two aluminum atoms. (ii) π  complexes: Organometallic compounds with H H π-bonds present in them are called π-complexes. Zeise’s salt, ferrocene and dibenzene chromium C are π-complexes. In these compounds, the π Cl C K+ 2+ Fe Cr electrons interact with the metal ion and occupy one of the coordination sites. For example, in Pt H H ferrocene and dibenzene chromium, the iron and Cl Cl chromium atoms are sandwiched between two aromatic rings. Zeise’s salt Ferrocene Dibenzene chromium K[PtCl3(2-C2H4)] Fe[2-C5H5)]2 6 Cr[ -C6H6)]2 The number of carbon atoms taking part in the formation of π-complexes is indicated by the power of ηx (pronounced as eta). For example, ferrocene is represented as [Feη5–C5H5)2] indicating that five carbon atoms or cyclopentadienyl anion are involved in the π- complication with the metal. Similarly, one can write dibenzene chromium as [Cr (η6–C6H6)2] indicating that all the six carbons of benzene are involved in π-complexation with chromium. 8.1 Bonding in Organometallic Compounds Bonding in Metal Carbonyls: The metal–carbon bond in metal carbonyls exhibits σ as well as π characteristics. (i)  -overlap: The lone pair of electron is present on the bonding orbital of carbon monoxide in a σ bonded σ complex and it interacts with the empty d-orbital of the metal to form a metal–carbon bond. + + M + C=O: M C=O Metal orbital Bonding Bonding in metal orbiital carbon (ii) π  -overlap: Besides, the antibonding orbitals of CO also overlaps with the filled d-orbital of the metal resulting in back bonding as previously explained. Thus metal carbonyls become much more stable due to this multiple bonding. It is important to note that the σ-bond is positioned in the nodal plane of the σ-electrons whereas π-overlap is perpendicular to the nodal plane. M C=O M C = O: Metal orbital Anibonding orbital Backbonding metal corbonyl of carbon mono-oxide 2 8. 30 | Co-ordination Compounds Bonding of Alkenes to a Transition Metal: There are two components in the bonding of alkenes to a transition metal to form complexes. First, the σ-electron density of the alkene overlaps with a π-type vacant orbital of the metal atom. Second is the backbonding resulting from the flow of electron density from a filled d-orbital on the metal into the vacant σ-antibonding molecular orbital on the carbon atom as depicted in the following diagram: C C M + M C C p overlap C C M + M C C MASTERJEE CONCEPTS As the electron density on metal atom increases, strength of backbonding from the metal to carbon increases and the metal–carbon bond length decreases. Likewise, when C–O bond order decreases, C–O bond length increases and vice versa. Neeraj Toshniwal (JEE 2009, AIR 21) 8.2 Synthesis of Organometallic Compounds Some important methods to generate metal–carbon bond as follows: By the direct reaction of metals: (a) n-Butyl lithium is prepared by the reaction of n-butyl bromide with lithium in ether. Ether n − C 4H9Br + 2Li  → n − C 4H9Li + 2LiBr n −Butyl bromide n −Buthyl lithium (b) Likewise, tetra ethyl lead can be prepared as follows: 4C2H5 Cl + 4NaPb → (C2H5 )4 Pb + 4NaCl + 3Pb Sodium −lead Tetra − ethyl lead alloy (c) Grignard reagents are obtained by the reaction of alkyl halide (in ether) with magnesium: R Mg + RX Mg X Grignard reagent By using an alkylating agent: Grignard reagent and alkyl lithium or reaction with most of the metal and non-metal halides in the presence of ether as solvent yield other organometallic compounds.  Chem i str y | 28.31 Ether PCl3 + 3C6H5MgCl  → P(C6H5 )3 + 3MgCl2 Triphenyl phosphine SnCl4 + 4n − C 4H9Li → (n − C 4H9 )Sn + 4LiCl Tetrabutyl tin Preparation of Metal Carbonyls: (a) Nickel carbonyl is obtained when finely divided nickel reacts with CO at room temperature. Ni + 4CO → Ni(CO)4(g)  ∆ → Ni + 4CO The nickel carbonyl so formed is in gaseous