Chemistry Notes for NEET Chapter 20 PDF

Chapt 2 (2) Complex ion : It is an electrically charged radical which is formed by the combination of a simple cation with one or more neutral molecules or simple anions or in some cases positive groups also. (3) Ligands : Neutral molecules or ions that attach to central metal ion are called ligands...

Chapt 2 (2) Complex ion : It is an electrically charged radical which is formed by the combination of a simple cation with one or more neutral molecules or simple anions or in some cases positive groups also. (3) Ligands : Neutral molecules or ions that attach to central metal ion are called ligands. The donor atom associated with the ligands supplies lone pair of electrons to the central metal atom (forming dative bond) may be one or two more. Monodentate (one donor atom), bidentate (two donor atom), tridentate (three donor atom) etc. Monodentate Ligands (with one donor site) Table : 20.1 Anionic Ligands (Negative legands) ID All transition elements exhibit a characteristic property of complex ion formation. The compounds containing complex ions are called coordination compounds. The branch of chemistry in which coordination compounds are studied that is called co-ordination chemistry. E3 60 Co-ordination Chemistry U Double Salts and Co-ordination Compounds. D YG When solutions of two or more stable compounds are mixed in stoichiometric (simple molecular) proportions new crystalline compounds called molecular or addition compounds are formed. These are of two types, (1) Double salts, (2) Co-ordination or Complex compounds (1) Double salts : Addition compounds, stable in solid state. Dissociate into ions in aqueous solution as such give test for each constituent ion. Examples: Double Salt Carnalite : Responds test for the ions K  , Mg 2  , Cl  Formula Name Formula Name X Halo O 22  Peroxo : OH  Hydroxo CH 3 COO  Acetato CN  Cyano NO 3 Nitrato O 2 Oxo S 2 O 32  Thiosulphato NH 2 Amido NO 2 Nitrito S 2 Sulphido CO 32  Carbonato Thiocyanato SO 42  Sulphato KCl. MgCl2 6 H 2 O Potash alum : K  , Al 3  , SO 42  K 2 SO 4. Al2 (SO 4 )3 24 H 2 O U (2) Co-ordination or Complex compounds : Addition compound, stable in solid state. Retain their identity even in solution. Central metal ion form dative or coordinate bond with the species surrounding it (ligands). Examples : ST Complex compound [Cu(NH 3 )4 ]SO 4 K 2 [PtF6 ] Cation Anion [Cu (NH 3 )4 ]2 2K  SO 42  CNS  [PtF6 ]2  2 Table : 20.2 Neutral Ligands 3 Name Carbonyl Formula CO : NH 3 Name Amminato Terminology of co-ordination compounds PH 3 Phosphine H 2O Aqua (1) Central metal atom or ion : A complex ion contains a metal atom or ion known as the central metal atom or ion. It is sometimes also called a nuclear atom. NO Nitrosyl C5 H 5 N : Pyridine (py) [Co (NH 3 )6 ][Cr (CN 3 )6 ] [Co (NH 3 )6 ] [Cr (CN )6 ] Formula Table : 20.3 Cationic Ligand (Positive) Formula NO 2 Nitronium H 2 NNH 3 Hydrazinium Name Nitrosonium NO  Polydentate ligands (with two or more donor site) Table : 20.4 Bidentate (Two donor sites) Formula Name Formula H 2 NCH 2 CH 2 NH 2 Ethylenediamin e (en) Name Dimethylglyoximeto (dmg) Me  C  NO  | Me  C  NOH  Oxalato (ox) O O || NH 2  CH 2  COO  || O C  C  O  Glycinate ion (gly) Table : 20.5 Formula   Name Diethylene triaminediamine (dien)  H 2 N (CH 2 )2  N H  (CH 2 )2 N H 2 Hexadentate    Triethylenetetramine (trien)  H 2 N (CH 2 )2  N H  (CH 2 )2 N H (CH 2 )2 N H 2   : OOCH 2 C CH 2 COO :   N  CH 2  CH 2  N CH 2 COO  : : OOCH 2 C Ethylenediamine tetra acetic acid (EDTA)4– D YG  Chelating Ligand : When polydentate ligands bind to the central metal ion they form a ring called chelate and the ligand is referred as chelating ligand. Ambidentate ligands : A ligand which possesses two donor atom but in forming complex it utilizes only one atom depending upon the condition and type of complex. NO 2 (nitro) , ONO (nitrito), CN (cyano), NC (isocyano), SCN (thiocyanide), NCS (isothiocyanide) - acid ligand : Ligands which are capable of accepting an appreciable U amount of - e  density from the metal atom into emptying  or  * orbital or their own called   acceptor or   acid ligands eg. CO. ST (4) Co-ordination Sphere : Ligand with central metal ion is kept in square bracket [ ] retains its identity in the same form is called co-ordination sphere (non-ionisable) (5) Co-ordination Number : Number of monodentate ligands attached to central atom/ion are called coordination number of the central metal atom/ion. Ligand Central metal Co 6 NH Coordination 3 Fe Ferrate Cu Cuperate Ag Argentate Au Aurate Sn Stannate Pb Plumbate (5) Point of attachment in case unidentate ligands with more than co-ordinating atoms (ambidentate ligands) : The point of attachment in case of unidentate ligands with more than one co-ordinating atoms is either indicated by using different names for the ligands (e.g, thiocyanato and isothiocyanato) or by placing the symbol of the donor atom attached, the name of the ligand separated by a hypen. U Tetradentate (4) Naming of ligands : The different types of ligands i.e. neutral, negative or positive are named differently in a complex compound. When a complex species has negative charge, the name of the central metal ends in – ate. For some elements, the ion name is based on the Latin name of the metal (for example, argentate for silver). Some such latin names used (with the suffix – ate) are given below : ID Tridentate In order to name complex compounds certain rules have been framed by IUPAC. These are as follows : (1) The positive part of a coordination compound is named first and is followed by the name of negative part. (2) The ligands are named first followed by the central metal. The prefixes di-, tri-, tetra-, etc., are used to indicate the number of each kind of ligand present. The prefixes bis (two ligands), tris (three ligands), etc., are used when the ligands includes a number e.g., dipyridyl, bis (ethylenediamine). (3) In polynuclear complexes, the bridging group is indicated in the formula of the complex by separating it from the rest of the complex by hyphens. In polynuclear complexes (a complex with two or more metal atoms), bridging ligand (which links two metal atoms) is denoted by the prefix  before its name. 60 Name E3 Formula 3+ Coordin ation sphere (6) Ionisation Sphere : The part present out side of the square bracket is called ionization sphere (ionisable). IUPAC Nomenclature of complex compounds ( NH 4 ) 2 [Pt( NCS ) 6 ] ( NH 4 )3 [Cr (SCN )6 ] Ammonium hexathioxy anato - S - chromate (III) Ammonium hexathiocyanato - N - platinate (IV) or or Ammonium hexathiocyanatochrom ate (III) Ammonium hexaisothi ocyanatopl atinate (IV) (6) Name of the bridging groups : If a complex contains two or more central metal atoms or ions, it is termed as polynuclear. In certain polynuclear complexes. ligands may link the two metal atoms or ions. Such ligands which link the two metal atoms or ions in polynuclear complexes are termed as bridge ligands. These bridge ligands are separated from the rest of the complex by hyphens and denoted by the prefix . If there are two or more bridging groups of the same kind, this is indicated by di-  , tri  , etc. [(NH 3 )]5 Co  NH 2  Co (NH 3 )5 ](NO 3 )5  - amidobis [pentaamminecobalt (III)]nitrate [(NH 3 )5 Co  OH  Co (NH 3 )5 ]Cl5  - hydroxobis [pentaamminecobalt (III)] chloride (7) If any lattice component such as water or solvent of crystallisation are present, these follow the name and are preceded by the number of these groups (molecules of solvent of crystallisation) in Arabic numerals. For example, [Cu (H 2 O)4 ]SO 4.H 2 O Tetraaquacopper(II)sulphate1  water [Cr(H 2 O)4 Cl 2 ]Cl.2 H 2 O tetraaquad ichlorochromium (III(chloride 2- water (8) Following punctuation rules should also be followed while writing the name of the complex compounds. (i) The name of the complete compound should not start a capital [Cu ( NH 3 )4 ]SO 4 letter, e.g., tetraamminecopper (II) sulphate (Correct) Tetraamminecopper (II) sulphate (Incorrect) Compounds having the same molecular formula but different structures or spatial arrangements are called isomers and the phenomenon is referred as isomerism. Isomerism Structural isomerism Ionisation isomerism Co-ordination isomerism Stereoisomerism Linkage isomerism Hydrate isomerism Geometrical isomerism Optical isomerism Example : Pentaaminebromo Sulphate cobalt (III) Pentaaminesulphato bromide cobalt (III) Br  present in ionization sphere Gives white BaCl2 Gives light yellow precipitate with AgNO3 D YG SO 42  present in ionisation sphere precipitate with (ii) Co-ordination isomerism : In this case compound is made up of cation and anion and the isomerism arises due to interchange of ligands between complex cation and complex anion. Example : [Co (NH 3 )6 ] [Cr (CN )6 ] [Cr (NH 3 )6 ][Co (CN )6 ] (c) [Cr(H 2 O)4 Cl]Cl 2.2 H 2 O tetraaquadichloro chromium (III) chloride (2) Stereo isomerism or space isomerism : Here the isomers differ only in the spatial arrangement of atoms of groups about the central metal atom. It is of two types : (i) Geometrical or Cis-trans isomerism : This isomerism arises due to the difference in geometrical arrangement of the ligands around the central atom. When identical ligands occupy positions near to each other called cisisomer. When identical ligands occupy positions opposite to each other called trans –isomer. It is very common in disubstituted complexes with coordination number of 4 and 6.  Complexes of co-ordination number 4 M 1 ST contains 2 Tetrahedral geometry : In this case all the four ligands are symmetrically arranged with respect to one another as such geometrical isomerism is not possible. Square planar geometry : The four ligands occupy position at the four corners and the metal atom or ion is at the center and lie in the same plane. Type : I [Ma2 b 2 ] , M  Pt, a  Cl, b  NH 3 Example : C l C P t [Pt Cl (NH 3 )(Py)2 ] N H3 C l NH cisl –isomer N H3 P t H3 C l N –isomer trans (dark- yellow) Complexes of co-ordination number 6 Octahedral geometry : Here the metal atom or ion lies at the center and 1 to 6 position are occupied by the ligands. Cis–Positions : 1–2, 2–3, 3–4, 4–5 Trans – position : 1–4, 2–5, 3–6 hexaamine 6 1 2 M complex cation contains NH 3 ligand (with cobalt) complex anion contains NH 3 ligand (with chromium) anion 3 4 U hexaamine cobalt (III) hexacyano chromate (III) chromium (III) hexacyanocobalt (III) complex (III) dihydrate (green) U [Co SO 4 (NH 3 )5 ] Br [Co Br (NH 3 )5 ] SO 4 pentaaquachlorochromium (b) [Cr (H 2 O)5 Cl ] Cl 2. H 2 O chloride monohydrate (blue green) ID (1) Structural isomerism : Here the isomers have different arrangement of ligands around the central metal atom. It is of the following types : (i) Ionisation isomerism : The co-ordination compound having the same composition or molecular formula but gives different ions in solution are called ionization isomers. There is exchange of anions between the co-ordination sphere and ionization sphere. Examples :(a) [Cr (H 2 O)6 ] Cl 3 hexaaquachromium (III) chloride (violet) 60 Isomerism in co-ordination compounds (iv) Hydrate isomerism : Hydrate isomers have the same composition but differ in the number of water molecules present as ligands and the phenomenon is called hydrate isomerism. E3 (ii) The full name of the complex ion should be written as one word without any gap. (iii) There should be a gap between the cation and anion in case of ionic complexes. (iv) The full name of non-ionic complexes should be written as one word without any gap. 5 4 3 CN  ligand (with chromium) complex anion contains CN  ligand (with cobalt) (iii) Linkage isomerism : In this case isomers differ in the mode of attachment of ligand to central metal ion and the phenomenon is called linkage isomerism. Example : [Co ONO (NH 3 )5 ]Cl2 ; [Co NO 2 (NH 3 )5 ]Cl2 Pentaamminenitritocobalt (III) Pentaaminenitrocobalt (III) chloride : O  NO  oxygen atom donates lone pair of electrons (nitrito) NO 2 nitrogen atom donates lone pair of electrons (nitro) Type –I Ma4 b 2 , M  Co , a  NH 3 , and b  Cl Example : [CoCl 2 ( NH 3 )4 ] ion Type –II [Ma3 b3 ] , M  Rh, a  Cl, and b  Py Example : [Rh Cl 3 (Py)3 Type –III [M (aa)2 (en )2 ]  , M  Co , a  a  CH 2 NH 2 | CH 2 NH 2 (bidentate ) b  Cl (monodentate) + C C l l C o H3 N N H3 N Cis-isomer H3 (blue H3 N violet) H3 N H3 N C l C N H3 N H3 Trans-isomer l (Green) [M ( AA)2 ab]n  ; [Co (en )2 NH 3 Cl]2  (ii) Optical isomerism (a) Optical isomers are mirror images of each other and have chiral centers. py Pt C N H N CisH -d-3 H 2 Cl + p y Pt C l N CisH-l3- Mir [M a b c d e f ] ; Pt (py) NH 3 NO 2 Cl Br] Br Pt Pt H N H I C l I d- isomer Mirr D YG C l p y O NO l- isomer [M ( AA)3 ]n  ; [CO(en )3 ]3  3 e e Co e Co e e e l- isomer U d- form ST Mirr e n Co en en 'Meso' or optically inactive form [M ( AA)2 a2 ]n  ; [Co (en )2 Cl 2 ] + e C e Cis-d- e C C C C C Mir e Cis-l- Cis-d- + N e Mir Cis-l- Bonding in co-ordination compounds (Werner's Coordination theory) Werner was able to explain the bonding in complex. Primary valency (Pv) : This is non- directional and ionizable. In fact it is the positive charge on the metal ion. Secondary valency (Sv) : This is directional and non- ionizable. It is equal to the number of ligand atoms co-ordinated to the metal (co-ordination number). Example : [Co (NH 3 )6 ]Cl 3 or Co (NH 3 )6 ]3  3Cl  Pv 3Cl  Sv 6 NH 3 (6) [Co (NH 3 )5 Cl]Cl 2 or [Co (NH 3 )5 Cl]2  2Cl  Pv 2Cl  (2) Sv 5 NH 3  1Cl  (6) [Co (NH 3 )4 Cl2 ] Cl or [Co (NH 3 )4 Cl2 ] Cl  Pv Cl  (1) Sv 4 NH 3  2Cl  (6) U p y e E3 2+ py C C ID Pt 2 e 60  Example : [Ma2 b 2 c 2 ]n  ; [Pt(Py)2 (NH 3 )2 Cl 2 ]2  2 C H (b) Mirror images are not super imposable and do and have the plane of symmetry. (c) Optical isomers have similar physical and chemical properties but differ in rotating the plane of plane polarized light. (d) Isomer which rotates the plane polarized light to the right is called dextro rotatory (d-form) and the isomer which rotates the plane polarized light to the left is called laevorotatory (l–form) C l e C Nature of the complex can be understood by treating the above complexes with excess of AgNO3. CoCl 3. 6 NH 3 3 AgCl, [Co (NH 3 )6 Cl 3 (three chloride ion) CoCl 3. 5 NH 3 2 AgCl, [Co (NH 3 )5 Cl 2 (two chloride ion) CoCl 3. 4 NH 3 1 AgCl, [Co (NH 3 )4 Cl 2 (one chloride ion) CoCl 3. 3 NH 3 no AgCl, [Co (NH 3 )3 Cl 3 (no chloride ion) The nature of bonding between central metal atom and ligands in the coordination sphere has been explained by the three well-known theories. These are : (1) Valence Bond theory of coordination compounds (i) The suitable number of atomic orbitals of central metal ion (s, p, d) hybridise to provide empty hybrid orbitals. (ii) These hybrid orbitals accept lone pair of electrons from the ligands and are directed towards the ligand positions according to the geometry of the complex. (iii) When inner d-orbitals i.e. (n-1) d orbitals are used in hybridization, the complex is called – inner orbital or spin or hyperligated complex. (iv) A substance which do not contain any unpaired electron is not attracted by 2 magnet. It is said to be diamagnetic. On the other hand, a substance which contains one or more unpaired electrons in the electrons in the d-orbitals, is attracted by a magnetic field [exception O 2 and NO]. It is said to be paramagnetic. Paramagnetism can be calculated by the expression, μ s  n(n  2), where   magnetic moment. weak field ligand is incapable of making the electrons of the metal atom/ ion to pair up. Strong field ligands : CN  , CO, en, NH 3 , H 2 O, NO  , Py. s= spin only value and n= number of unpaired electrons. Hence, if n  1, s  1(1  2)  1.73 B.M. , if  3(3  2)  3.87 B.M. and so on Weak field ligands : n  3,  s I  , Br  , Cl  , F  , NO 3 , OH  , C 2 O 42  , NH 3 , H 2 O. Limitations of valence bond theory The valence bond theory was fairly successful in explaining qualitatively the geometry and magnetic behaviour of the complexes. But, it could not explain the following : (i) The origin of their absorption spectra could not be explained. (ii) Why did different complexes of the same metal show different colours. (iii) Relative stabilities of different complexes could not be explained. (iv) Why should certain ligands form high spin, while others low spin complexes. Table : 20.6 Geometry and magnetic nature of some complexes +2 2– U sp3 2+ D YG [Ni(CN) ] 4 Magnetic nature (7) 2 Paramagnetic Ni Ni(CO) 4 Rearrange ment 4s 4p Tetrahedral 2 Paramagnetic 2+ sp +2 (outer) sp d2 ST 4s Square planar 0 Diamagnetic 2 Paramagnetic 3 Tetrahedral 0 Diamagnetic sp d 3 p 5 Octahedral 2 Paramagnetic 5 Paramagnetic 1 Paramagnetic +2 d sp 2 +2 4– 6] Rearrange ment 2 4 Mn (d ) [Mn(CN) 2 0 3 3d dsp 0 4d U 3d 3 sp3 [Ni(NH ) ] 6 +2 sp ds p2 Rearrange ment 2+ No. of unpaired electrons (6) Geometry shape (5) +2 4 3 Type of hybridizat ion (4) ID 4p 8 [NiCl ] Oxidatio n state of metal (3) Configuration (2) Ni (d ) 2+ E3 Atom/ion/ complex (1) 60 On the basis of value of magnetic moment, we can predict the number of unpaired electrons present in the complex. If we know the number of unpaired electrons in the metal complex, then it is possible to predict the geometry of the complex species. (v) There are two types of ligands namely strong field and weak field ligands. A strong field ligand is capable of forcing the electrons of the metal atom/ion to pair up (if required). Pairing is done only to the extent which is required to cause the hybridization possible for that co-ordination number. A d2s p3 sp3 3 (Inner) Octahedral sp +2 2 4 Cu (d ) +2 [CuCl ] +2 sp +2 dsp 2+ 9 2– 4 [Cu(NH ) ] 3 2+ 4 Tetrahedral 3 Tetrahedral 3 2 dsp One electron is shifted from 3d-to 4p-orbital (1) (2) Cr (d ) 3d 3 4 4d 6 sp3 d2 CO (d ) 3– 6 [Co(NH ) ] 3+ 3 6 D YG 6 [CoF ] sp3 d2 Rearrange ment CO (d ) 2+ d2s p3 U 7 sp3 d2 6 [Fe(CN) ] [Fe(H O) ] 2 sp d d2s p3 d sp Paramagnetic Octahedral 3 Paramagnetic Octahedral 3 Paramagnetic 4 Paramagnetic Octahedral 4 Paramagnetic Octahedral 0 Diamagnetic 3 Paramagnetic 3 Paramagnetic 4 Paramagnetic Octahedral 0 Diamagnetic Octahedral 4 Paramagnetic 3 sp d 2 (Outer) d sp Octahedral 3 (Inner) sp d 3 sp3 d2 Paramagnetic 3 +2 +2 2+ 6 1 2 (Inner) 2 Rearrange ment Paramagnetic (7) +2 +2 4– 6 1 2 (Outer) 3 6 Fe (d ) 2+ sp d (Outer) 2 +3 Paramagnetic 3 3 +3 1 (6) +3 +2 2+ ST [Co(H O) ] 2 d sp (Inner) 3 +3 Square planar (5) ID d2s p3 3+ 2 (4) 2 +3 3+ 6 [Cr(H O) ] 3+ +3 p [Cr(NH ) ] 3 4s U 3+ (3) Paramagnetic E3 2 5 60 [MnCl ] 2 (Outer) Same 2+ 3 6 +2 Fe (d ) 3+ sp d 3 [Fe(NH ) ] 2 (Outer) +3 5 [Fe(CN) 4 Paramagnetic Octahedral 5 Paramagnetic 1 Paramagnetic d sp 2 +3 3– 6 Octahedral 3 (Inner) Fe 0 Fe(CO) dsp 0 dsp 3 ⇌ Cu 2   4 NH 3 K Paramagnetic 0 Diamagnetic [Cu ( NH 3 )42  ] [Cu 2  ][ NH 3 ]4  [Cu (NH 3 )2  is given Stability of co-ordination in solution and Spectrochemical series ; 1 K, For example, the stability constant of [Fe(CN )6 ]3  is much greater than the stability constant of [Fe(CN) ]. 4– 6 Fe [Cu (NH 3 )4 ]2  ⇌ Cu 2   4 NH 3 , is given by the expression; Fe3   6 CN  ⇌ [Fe(C N )6 ]3  ; K  1.2  10 31 D YG Instability constant for the complex ion [Cu (NH 3 )4 ]2  i.e. 4 under Greater is the stability constant, stronger is the metal – ligand bond Factors affecting the stability of complex ion (1) Nature of central metal ion : The higher the charge density on the central metal ion the greater is the stability of the complex U Stronger is the metal-ligand bond, less is the dissociation in the solution and hence greater is the stability of a coordination compounds. 2 as ID (2) Ligand field theory : According to this theory when the ligands come closer to metal atom or ion, a field is created. This field tends to split the degenerate d-orbitals of the metal atom into different energy levels. The nature and number of ligands determine the extent of splitting. Energy gap between the splitted energy levels determine the magnetic (paramagnetic or diamagnetic) and spectral properties (colour etc) of the complexes. Trigonal bipyramidal 3 (Inner) 4 E3 5 Ki  60 d2s p3 [Cu ][ NH 3 ]. [Cu ( NH 3 )4 ]2  Stability constant of the above complex i.e. 2  6 CN ⇌ [Fe(C N )6 ]4  ; K  1.8  10 6  Effective atomic number (EAN) or Sidgwick theory : In order to the stability of the complexes sidgwick proposed effective atomic number. EAN generally coincides with the atomic number of next noble gas in some cases. EAN is calculated by the following relation : Complex At. No. of metal Coordination number K4 [Fe (CN )6 ] +2 26 6 (26 – 2) + (6  2) = 36 [Kr] [Cu (NH 3 )4 ] SO 4 +2 29 4 (29 – 2) + (4  2) = 35 [Co (CH 3 )6 ] Cl3 +3 27 6 (27 – 3) + (6  2) = 36 [Kr] 0 28 4 (28 – 0) + (4  2) = 36 [Kr] K2 [ Ni (CN )4 ] +2 28 4 (28 – 2) + (4  2) = 34 K3 [Cr (C2O4 )3 ] +3 24 6 (24 – 3) + (6  2) = 33 K3 [Fe (CN )6 ] +3 26 6 (26 – 3) + (6  2) = 35 [ Ag (NH 3 )2 ] Cl +1 47 2 (47 – 1) + (2  2) = 50 ST Metal oxidation state U EAN = Atomic no. of the metal – e  lost in ion formation +No. of e  gained from the donor atom of the ligands. EAN = Atomic number – Oxidation number + co-ordination no.  2 Table : 20.7 Ni (CO )4 (2) Nature of ligand : Greater the base strength is the ease with which it can donate its lone pair of electrons and therefore, greater is the stability of the complex formed by it. For example [Cu (CN )4 ]2  ; K  2.0  10 27 Effective atomic number : [Cu (NH 3 )4 ]2  ; K  4.5  10 11 ; (3) Presence of chelate ring : Chelating ligands form more stable complex as compared to monodentate ligands. For example : Ni ⇌ 2 [ Ni(NH 3 )6 ] ; K  6  10 ; 8 Ni 2 2  6 NH 3 ⇌  3en [ Ni(en )3 ] 2  ; K  4  10 8 Spectro chemical series : Ligands can be arranged in increasing order of their strength (ability to cause crystal field splitting) and the series so obtained is called as spectro chemical series.      I  Br  Cl  F  OH  OX 2 These are the compounds in which a metal atom or a metalloid (Ge, Sb) or a non-metal atom like B, Si, P, etc, (less electronegative than C) is directly linked to a carbon atom of a hydrocarbon radical or molecule. Organometallic compounds contain at least one. (1) Metal – Carbon bond, (2) Metalloid – Carbon bond, (3) Non metal – Carbon bond. Example :  H 2O C2 H 5 MgBr , C6 H 5 Ti(OC3 H7 )3 , (CH 3 )4 Si Preparation and Application of coordination compounds [Cu (H 2 O)4 ]2  (aq )  3 NH 3 (aq )) [Cu ( NH 3 )4 ]2 (aq ) are called - bonded organometallic compound. (2) -bonded organometallic compounds : The transition metals binds to unsaturated hydrocarbons and their derivatives using their d-orbitals. Here metal atom is bonded to ligands in such a way that donations of electrons and back acceptance by the ligand is feassible. These are called – orbitals of the ligand. These are called –complexes. Examples : (i)   cyclopentadienyl – iron complex Ferrocene [Fe ( 5  CH 5 )2 ] , Bis (cyclopentadienyl) iron (II) It is a  bonded sandwitch compound. The number of carbon atoms bonded to the metal ion is indicated by superscript on eta ( x ) i.e.  5 in this complex. U Deep blue (1) –bonded organometallic compounds : Compounds such as RMgX, R2 Zn, R3 Pb, R3 Al, R4 Sn etc, contains M  C   bond and ID (1) Preparation : Coordination compounds are generally prepared by the application of the following methods, (i) Ligand substitution reaction : A reaction involving the replacement of the ligands attached to the central metal ion in the complex by other ligands is called a ligand substitution reaction. 60 ligands which can not cause forcible pairing of electrons within 3d level and thus form outer orbital octahedral complexes. On the other hand NH 3 and all ligands lying right to it are stronger ligands which form inner orbital octahedral complexes after forcible pairing of electrons within 3d level. Classification of organometallic compounds : Organometallics have been classified as : E3 Ligands arranged left to NH 3 are generally regarded as weaker  4 H 2 O(l) (ii) Direct mixing of reagent : D YG PtCl2  2 H 2 N  CH 2  CH 2  NH 2 [Pt(en )2 Cl 2 ] (en) (Dichloro bis ethylene diam mine platinum(II) (ii) Dibenzene chromium (   complex) It is also a   bonded sandwitch compound. Its formula is [Cr(  C6 H 6 )2 ] 6 (iii) Alkene complex (   complex) (iii) Redox reactions : In these reactions, either oxidation or reduction is involved Zeise’s salt 2CO(NO 3 )2  8 NH 3  2 NH 4 NO 3  H 2 O 2 Pt Cl 3 (  C 2 H 4 )] ; 2[CO (NH 3 )NO 2 ](NO 3 )2  2 H 2 O Potassium platinate (IV). Penta ammine nitrocobalt (II) nitrate. Application (C2 H 5 )2 Zn , Organometallic bond: Mg  C , Zn  C , Ti  C , Si  C  Py  NH 3  en  NO 2  CN   CO Light blue Compounds : and Mg 2 ions U (1) Estimation of hardness in water, as Ca form complexes with EDTA.  (2) Animal and plant world e.g. chlorophyll is a complex of Mg 2  ST and haemoglobin is a complex of Fe 2  vitamin B12 is a complex of Co 2 . (3) Electroplating of metals involves the use of complex salt as electrolytes e.g. K [ Ag (CN ) 2 ] in silver plating. (4) Extraction of metals e.g. Ag and Au are extracted from ores by dissolving in NaCN to form complexes. (5) Estimation and detection of metal ions e.g. Ni 2  ion is estimated using dimethyl glyoxime. (6) Medicines e.g. cis-platin i.e. cis [PtCl2 (NH 3 )2 ] is used in treatment of cancer. Organometallic compounds K 2 It is a  bonded complex.  F trichloroethylene 2 indicates that two carbons of ethylene are bonded to metals. C Ferrocene Dibenzene chromium (3) Complexes containing both – and –bonding characteristics : Metal carbonyls, compounds formed between metal and carbon monoxide belong to this class. Metal carbonyls have been included in organometallics. (i) Mononuclear carbonyls : Contain one metallic atom per molecule. e.g Ni (CO )4 , Fe(CO )5 , Cr(CO )6 (ii) Polynuclear carbonyls : Contain two or more metallic atoms per molecule. e.g., Mn2 (CO )10 , Fe(CO )9 , Fe(CO )12 Applications of organometallics (1) Grignard reagent (RMgX) has been extensively used for synthesis of various organic compounds. (2) Wilkinson's catalyst [(PH3 P)3 RhCl] i.e. tris (triphenylphosphine) chlororhodium (I) is used as a homogeneous catalyst for the hydrogenation of alkenes. (3) Zeigler Natta catalyst (composed of a transition metal salt, generally TiCl4 and trialkyl aluminium) are used as heterogeneous catalysts in the polymerisation of alkenes. have flexidentate character if they do not use all its donar atoms to get coordinated to the metal ion e.g., EDTA generally act as a hexadentate ligand but it can also act as a pentadentate and tetradentate ligand.  Octahedral complex of the type [M(ABCDEF)] have 15  Badecker reaction : This reaction involves the following chemical change. different geometrical isomers with a pair of enantiomers. Although a few geometrical isomer have been prepared but none has been resolved e.g., [Pt (Py)(NH3)(NO2)(Cl)(Br)I]. Na 2 [Fe(CN )5 NO ]  Na 2 SO 3 Na 4 [Fe(CN )5 (NO. SO 3 )] salt : It is reduction of prussian blue. K 2 [Fe(CN )6 ] obtained by 60  Everitt's without physical separation of it is so transformed that is does not enter into a particular reaction e.g., masking of Cu 2  by CN  ion.  Macrocyclic effect : This term refers to the greater NH NH is more stable than with NH NH NH NH NH H2 N 2 ID thermodynamic stability of a complex with a cyclic polydentate ligand when compared to the complex formed with a non-cyclic ligand. e.g., Zn (II) complex with ligand; E3  Masking : Masking is the process in which a substance  Prussian blue and Turnbull's blue is pot. ferric D YG U ferrocyanide. However colour of Turnbull's blue is less intense than prussian blue. Decrease in colour is due to the presence in it of a white compound of the formula K 2 {Fe[Fe(CN )6 ]} named as potassium ferrous ferrocyanide.  Crystal field theory explains why certain geometries are more favoured than other by certain metals in terms of crystal field stablization energies.  The colour of complexes are explained in terms of electronic transitions between the various d orbitals of different energies.  In octahedral complexes, the energy of the d z 2 and U d x 2  y 2 orbitals increases much more than the energy of the d xy , d yz , d xz orbitals. The potential energy increases as a result ST of repulsive interactions. The greater the repulsion, greater is the increases in energy.  Ligands with larger groups from unstable rings than the ligands with smaller groups. This is ascribable to steric hinderance produced by a larger group.  When Ca2+ or Mg2+ forms complexes with EDTA, the pH of the solution decreases.  The copper sulphate solution e.g., turns deep blue when excess of ammonia is added. This is also due to the formation of [Cu(NH3)4]SO4  Co-ordination compound with a general formula MA4, MA3B or MABCD do not show any geometric isomerism.  Flexidentate character : polydentate ligand are said to