JEE-Coordination-Compounds-final.pdf

Welcome to Coordination Compounds Salt and its Classification An ionic compound that contains a cation (basic radical) and an anion (acidic radical) Salt Simple salt Double salt Complex salt Simple Salt A simple salt is formed by the neutralisation of an acid and a base. NaOH + HCl NaCl + H2O Simple...

Welcome to Coordination Compounds Salt and its Classification An ionic compound that contains a cation (basic radical) and an anion (acidic radical) Salt Simple salt Double salt Complex salt Simple Salt A simple salt is formed by the neutralisation of an acid and a base. NaOH + HCl NaCl + H2O Simple salt Click to add text Double Salt It is formed by the combination of two or more stable compounds in a stoichiometric ratio. KCl + MgCl2 + 6H2O It dissociates into simple ions completely when dissolved in water. KCl.MgCl2.6H2O Carnallite Double Salt Fe2SO4 + (NH4)2SO4 + 6H2O Fe2SO4.(NH4)2SO4.6H2O Mohr’s salt + Ferrous sulphate = Ammonium sulphate Mohr’s salt Double Salt FeSO4.(NH4)2SO4.6H2O KCl.MgCl2.6H2O When dissolved in water When dissolved in water The solution shows the properties The solution shows the properties + 2+ of K , Mg , and Cl− ions. 2+ + 2- of Fe , NH4 , and SO4 ions. Double Salt K2SO4.Al2(SO4)3.24H2O When dissolved in water The solution shows the properties of K+, Al3+ and SO42− ions. Note!! On qualitative analysis, a solution of double salt gives the test for each individual constituent ions. Generally, a double salt contains same anion but different cations. Complex Salt Fe(CN)2.4KCN CuSO4.4NH3.H2O When dissolved in water When dissolved in water The solution does not show the _ 2+ properties of Fe and CN The solution shows the properties of SO42- but does not show for NH3 Complex Salt CuSO4 + 4NH3 + H2O CuSO4.4NH3.H2O Tetraamminecopper(II) sulphate monohydrate It is formed by the combination of two or more stable compounds in a stoichiometric ratio. However, it does not dissociate into simple ions when dissolved in water. [Cu(NH3)4]SO4.H2O Gives test for SO42- Complex salts are generally coordination compounds. Exists as a distinct entity Complex Salt Fe(CN)2.4KCN Fe(CN)2 + 4KCN Fe(CN)2.4KCN When dissolved in water Potassium ferrocyanide The solution does not show the properties _ of Fe2+ and CN. Complex Salt K4[Fe(CN)6] Complex ions are shown by square brackets. Exists as a distinct entity Solution of K4[Fe(CN)6] K + [Fe(CN)6] 4 _ Gives the identification Does not dissociate test of K + ions to give the identification test 2+ into Fe and CN _ Coordination Compounds The addition compounds that retain their identity (i.e., do not lose their identity) in solutions are known as coordination compounds. Metal atom or ion is attached to group of neutral molecules/ions by coordinate covalent bonds. Example: K4[Fe(CN)6] Coordination Compounds a+ General formula [M(L)n] a+ [M(L)n] Coordination entity X a− a− X The central atom/ion and the ​ligands attached to it are enclosed​ in square brackets along with electrical charge on it, collectively termed as the coordination entity. Coordination Entity 3+ − [Co(NH3)6] 3Cl M Central atom Coordination entity Generally, metals Central Atom In a coordination entity, the atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement around it, is known as the central atom or ion. Also referred to as Lewis acids [Co(NH3)6]Cl3 Central metal atom/ion Coordination Compounds General formula [M(L)n]a+ Xa− Ligand Ligand The electron-rich species bound to the central atom/ion inside the square bracket are known as ligands. In general, donates the electron to the central metal atom [Co(NH3)6]Cl3 Ligand Behaves as Lewis base Ligands Ligand Behaves as Lewis base [Co(NH 3 [Co (NH33))6]Cl 6]Cl 3 In general, donates the electron to the central metal atom Ligand Coordination Compound General formula [M(L)n] a+ Charge on the coordination entity a− X 3+ 3+ [[Co Co(NH (NH33))66]] Charge on the coordination entity − − 3Cl 3Cl Coordination Compound a+ [M(L)n] X General formula Coordination sphere 3+ − [Co(NH3)6] 3Cl Coordination sphere a− The central atom/ion and the ligands attached to it are enclosed in square brackets termed as the coordination sphere. Coordination Entity vs Coordination Sphere Coordination sphere Coordination entity Collection of a central atom and the ligands surrounding the atom Coordination entity given with the electrical charge of the coordination compound Coordination Compounds General formula a+ [M(L)n] Counter ion a− X 3Cl3Cl− [Co[Co(NH (NH33))66]]3+ 3+ Counter ion − Counter Ion as Complex Compound [Ag(NH3)2]1+[Ag(CN)2]1- When there are two coordination spheres joined together by an ionisation sphere [Cr(NH3)6]3+[Co(CN)6]3- Ionisation Sphere The molecule or ionic species present outside the coordination sphere constitute the ionisation sphere. In K4[Fe(CN)6], the + ionisation sphere is K. Coordination Compounds General formula a+ [M(L)n] X a− Coordination number 3Cl3Cl− 3+ 3+ [Co[Co(NH (NH33))66]] Coordination number − The coordination number (CN) of a metal ion in a complex can be defined as The number of ligand donor atoms to which the metal is directly bonded. Coordination Number 2 _ [Pt(Cl)6] Coordination number of Pt = 6 = 4 2+ [Ni(NH3)4] Coordination number of Ni Coordination Polyhedron The spatial arrangement of the ligand atoms that are directly attached to the central atom/ion defines a coordination polyhedron about the central atom. Octahedral The most common coordination polyhedral are octahedral, square planar, and tetrahedral. Tetrahedral Square planar Oxidation Number of Central Metal Atom [Cu(CN)4] The charge that a metal would carry if all the ligands are removed along with the electron pairs that are shared with the central atom x 3 _ -1 x + 4(-1) = -3 x = +1 Oxidation number of Cu Based upon the charge on the coordination sphere Classification of complex Based upon the type of the donor group Classification of Complex Based on the charge on the coordination sphere Positive Neutral Negative Positively Charged Coordination Complex Cationic complex Charge on the coordination sphere: Positive 2+ 22 2+ [Cu(NH3)4] Neutral Coordination Sphere Neutral complex [Ni(CO)4] Charge on the coordination sphere: Absent or no charge Negatively Charged Coordination Sphere Anionic complex Charge on the coordination sphere: Negative [Fe(CN)6] 3− 3- Classification of Complex Based on the type of the donor group Homoleptic Heteroleptic Homoleptic Complex A metal is bound to only one kind of donor group. Example: [Co(NH3)6]3+ 3+ Heteroleptic Complex A metal is bound to more than one kind of donor groups. Example: [Co(NH3)4Cl2]+ Based on the charge Classification of ligands Based on the denticity Based upon bonding interactions between the ligand and the central atom Classification of Ligands Based on their charge, ligands can be classified as Positive Examples: NO+, NO+2 Neutral Examples: H2O, NO, CO, NH3, etc. Negative Examples: Cl−, OH−, CN−, etc. Denticity The number of donations accepted by a central atom from a particular ligand is known as the denticity of the ligand. Monodentate Classification based upon denticity of ligand Bidentate Polydentate Monodentate Ligand Only one donation is accepted from the ligand. H2O, NO, CO, NH3, Cl −, etc. Bidentate Ligand Two donations are accepted from the ligand. en: Ethylenediamine Ox2– : Oxalate Polydentate Ligand More than two donations are accepted from the ligand. Tridentate Ligand dien: Diethylenetriamine Tetradentate Ligand trien: Triethylenetetraamine Pentadentate Ligand _ _ Ethylenediaminetriacetate _ Hexadentate Ligand _ _ Ethylenediaminetetraacetate (EDTA) _ _ Classification of Ligands Flexidentate ligand Types of ligands Ambidentate ligand Chelating ligand Flexidentate Ligand Didentate ligand Monodentate ligand _ A ligand that shows flexible denticity 2– Examples: S O 2 _ _ 3 , 2– CH3COO , NO3 , SO4 , etc. _ Flexidentate ligand behaving as didentate _ M _ M Flexidentate ligand behaving as monodentate Ambidentate Ligand A ligand that has two different donor atoms and either of the two ligates/donates in the complex _ : : or NO2 M or : CN _ : _ : _ : _ _ M Chelating Ligand When a ligand uses two or more of its donor atoms simultaneously to bind a single metal ion and produces a ring The complex formed is referred to as a chelate complex. The process of chelate formation is known as chelation. Chelating Ligand dmg Ni 240 pm Points to Remember! More than two donor sites All polydentate ligands are examples of chelating ligands and they can form chelate complexes. Note!! A complex is perfect or imperfect will depend upon the reagent by which that complex is reacted. Stability of the complex will depend upon the reagent by which that complex is reacted. Classification of Complex Based upon stability Imperfect complex Complex ions that cannot maintain their independent existence in solution And dissociate into their constituent ions partially Perfect complex Imperfect Complex K2[Cd(CN)4] Unstable complex K ↓CdS (Yellow precipitate) + [Cd(CN)4] + H 2S Cd 2+ 2− _ CN Classification of Complex Based upon stability Imperfect complex Perfect complex Complex ions that exhibit very less degree of dissociation in solution Perfect Complex K3[Cu(CN)4] No further dissociation Stable complex K + [Cu(CN)4] 3− + H2S Cu2S (not formed) Stability Due to Chelation Chelated complexes are more stable than similar complexes with monodentate ligands as dissociation of the complex involves breaking of more bonds rather than one. Condition for Chelation Chelating ligands prefer forming 5 or 6 membered ring complexes with metal compared to other ring sizes. Stability Due to Chelation Co3+ (aq) + 6 NH3(aq) Co3+ (aq) + 3 en(aq) Co3+ (aq) + 2 dien(aq) [Co(NH3)6]3+ 1 [Co(en)3]3+ 2 [Co(dien)2]3+ 3 Stability Order : 3 > 2 > 1 Chelate Effect The greater affinity of chelating ligands for a metal ion than that of similar nonchelating (monodentate) ligands for the same metal is known as Chelate Effect. [Co(NH3)6]3+ (aq) + 3 en(aq) [Co(en)3]3+ + 6 NH3(aq) Naming of Mononuclear Coordination Compounds Given by IUPAC The names of coordination compounds are derived by the following principles of nomenclature. Rules 1 The cation is named first in both positively and negatively charged coordination entities. Naming of Mononuclear Coordination Compounds Rules 2 The ligands are named in an alphabetical order before the name of the central atom/ion. Multiplicative prefixes are not considered. Naming of 3 Anionic ligands Neutral ligands Cationic ligands Naming of Anionic Ligands The names of anionic ligands end with ‘o’. a Ligand _ NO3 Nitrato _ CO32 _ _ NO3 Nitrate Nitrato MeCO2 _ acac 2 SO4 _ 2_ S2O3 Name Carbonato Acetato or Ethanoato Acetylacetonato Sulphato Thiosulphato Naming of Anionic Ligands b Ligand names ending with ‘ite’ are changed to ‘ito’. Ligand 2 SO3 ClO2 NO2 Name Sulphito _ _ HSO3 Ligand names ending with ‘ide’ are changed to ‘ido’. _ _ OCl c _ Chlorito Hypochlorito Ligand F _ Cl Nitrito Br Hydrogensulphito I _ _ Name Fluorido Chlorido Bromido Iodido Naming of Anionic Ligands Ligand _ CN _ NC H _ _ OH _ OD _ NH2 NH 2_ Name Cyanido Ligand 3 N _ Isocyanido N3 Hydrido O2 Hydroxido 2 O2 Duterido Amido Imido _ _ _ O2 2_ HS Nitrido Azido _ S Name _ Oxido Peroxido Superoxido Sulphido Hydrosulphido Naming of Anionic Ligands Ambidentate ligands d Serial Ligand number _ Cyanido/Cyanido-C CN 1 _ _ 2 Name NC Isocyanido/Cyanido-N NO2 Nitro/Nitrito-N _ ONO Nitrito/Nitrito-O Naming of Anionic Ligands Ambidentate ligands d Serial Ligand number _ 3 _ _ 4 _ Name OCN Cyanato/cyanato-O NCO Isocyanato/cyanato-N SCN Thiocyanato/thiocyanato-S NCS Isothiocyanato/thiocyanato-N Naming of Neutral Ligands The names of neutral ligand have no specific endings. Ligand Name H2O Aqua NH3 Ammine CO Carbonyl NO MeCONH2 Nitrosyl Acetamide Ligand Name MeNH2 Methanamine MePH2 Methylphosphine C6H6 Benzene CH3NC Methyl isocyanide PPh3 Triphenylphosphine C5H5N Pyridine Naming of Cationic Ligands Rules 3 Ligand Name + The name of cationic ligands end with ‘ium’. NO Nitrosonium + NO+2 Hydrazinium Nitronium 𝛑-Donor Ligands The 𝛑-donor ligands are the unsaturated organic compounds that can donate the electron pair to the metal atom/ion. Examples Alkenes, alkynes and benzene Naming of Anionic 𝛑 -donors For the 𝛑 -donor ligands ,, Prefix like “ηx is used. Where, η : Hapticity of the ligand x : Number of atoms attached to the metal atom/ion. Hapticity The number of carbon atoms of a ligand which are directly attached to (or closely associated with) the metal in a compound is called as the hapticity. − 𝛑 -C5H5 η5- cyclopetadienyl 𝛑 -C6H6 η6- benzene − 𝛑 -C3H5 η3- allyl Examples of structures of Hapticity − M η5 - C5H5 η3 – allyl − ≡ − _ M η6 -C6H6 ≡ ≡ _ M Naming of Mononuclear Coordination Compounds Rules 2-di 3-tri 4 If a number of same ligands are present, then a prefix is used. 4-tetra Prefixes 5-penta 6-hexa 7-hepta Naming of Mononuclear Coordination Compounds Rules 2-bis 5 When the name of a ligand already includes terms like di, tri, tetra, etc., then a second set of prefix is used. 3-tris 4-tetrakis Prefixes 5-pentakis The ligand's name is kept within brackets to avoid ambiguity. 6-hexakis 7-heptakis Naming of Mononuclear Coordination Compounds (en)2 bis-(ethylenediamine) Me2NH dimethylamine (NH2Me)2 bis-(methylamine) To make it distinct from dimethyl amine Naming of Mononuclear Coordination Compounds Rules 6 There is no deletion of vowels or use of a hyphen. Example: (NH3)4 Tetraammine Naming of Mononuclear Coordination Compounds 7 For anionic complexes, the names of metal end with -ate. Aluminium Aluminate Nickel Nickelate Platinum Platinate Cobalt Cobaltate Molybdenum Molybdate Iron Ferrate Zinc Zincate Manganese Manganate Note!! For cationic and neutral complexes, the metal is named as the element. Example: Co: Cobalt Example: Pt: Platinum Naming of Mononuclear Coordination Compounds Rules 8 The oxidation state of the metal in a cation, anion, or neutral coordination entity is indicated by a Roman numeral in brackets. Examples: (II), (-I), (V) Naming of Mononuclear Coordination Compounds Rules 9 If there is water of crystallisation, it is to be included in the name. Example: [Cr(H2O)4Cl2]Cl.2H2O Tetraaquadichloridochromium(III) chloride dihydrate 10 No space is left between the names that refer to the same coordination entity. Note!! For coordination complexes, we don’t indicate the number of cations and anions in the IUPAC nomenclature. Rules for Writing Formulas from the Names Rules for Writing Formulas R U L E R U L E 1 The ligands are then listed in an alphabetical order. 2 The central atom is listed first. The placement of a ligand in the list does not depend on its charge. Rules for Writing Formulas R U L E 3 Polydentate ligands are also listed alphabetically. In case of an abbreviated ligand, the first letter of the abbreviation is used to determine the position of the ligand in the alphabetical order. R U L E 4 Formula for the entire coordination entity Charged Uncharged Enclosed in square brackets Rules for Writing Formulas R U L E 5 When ligands are polyatomic, their formulas are enclosed in parentheses. Examples: (CN–), (H2O) Ligand abbreviations are also enclosed in parentheses. Examples: (en), (ox) R U L E 6 There should be no space between the ligands and the metal within a coordination sphere. Rules for Writing Formulas R U L E 7 When the formula of a charged coordination entity is to be written without that of the counter ion The charge is indicated outside the square brackets as a right superscript with the number before the sign. Examples: [Co(CN)6]3–, [Cr(H2O)6]3+, etc. R U L E 8 The charge of the cation(s) is balanced by the charge of the anion(s). Rules for Writing Formulas Dichloridobis(ethane-1,2-diamine)cobalt(III) chloride en [CoCl2(en)2]Cl Central atom Hexaamminechromium(III)hexacyanido cobaltate(III) [Cr(NH3)6][Co(CN)6] Werner’s theory Bonding in coordination compounds Valence bond theory (VBT) Crystal field theory (CFT) Werner’s Theory 1893 Werner’s coordination theory was the first attempt to explain the bonding in coordination compounds. While adding excess of AgNO3 solution in a series of compounds of Co(III) chloride with NH3, the following results were found: Werner’s Theory Experiment CoCl3.6NH3 A series of compounds formed by CoCl3 and NH3 + CoCl3 NH3 CoCl3.4NH3 exist in two isomers(cis and trans), hence, showing two colours. CoCl3.5NH3 CoCl3.4NH3 Werner’s Theory CoCl3.6NH3 1 mol + AgNO3 3AgCl Excess 3 mol Why this gives 3 mol AgCl Because it has an ionisable valency of 3 The series of compounds formed by CoCl3 and NH3 after reacting with excess AgNO3 gives different number of AgCl. Werner’s Theory CoCl3.5NH3 1 mol + AgNO3 2AgCl Excess 2 mol Why this gives 2 mol AgCl Because it has an ionisable valency of 2 Werner’s Theory CoCl3.4NH3 1 mol, (green) + AgNO3 AgCl Excess 1 mol Why this gives 1 mol AgCl Because it has an ionisable valency of 1 Werner’s Theory CoCl3.4NH3 1 mol, (violet) + AgNO3 AgCl Excess 1 mol Why this gives 1 mol AgCl Because it has an ionisable valency of 1 Conductance A series of compounds formed by CoCl3 and NH3 show different conductance values which confirms the presence of ionisable and nonionisable valencies. CoCl3.6NH3 CoCl3.5NH3 CoCl3.4NH3 CoCl3.4NH3 Conclusion of Werner’s Experiment The atoms within the square brackets form a single entity and do not dissociate under the reaction conditions. Solution conductivity corresponds to Colour Formula Yellow [Co(NH3)6]3+3Cl _ 2+ Purple [CoCl(NH3)5] 2Cl Green [CoCl2(NH3)4] Cl Violet [CoCl2(NH3)4]+Cl Isomers + 1:3 electrolyte _ _ _ 1:2 electrolyte 1:1 electrolyte 1:1 electrolyte Have identical molecular formula, but distinct properties Werner’s Theory Metals in coordination compounds show two types of valencies/linkages Primary or ionisable Secondary or non-ionisable Primary Valency [ Mn The primary valencies are normally ionisable and are satisfied by negative ions. (H2O)6 ] Cl2 Primary valency Secondary Valency The secondary valencies are nonionisable. These are satisfied by neutral molecules or negative ions. [ Mn (H ] 2O)6 Cl2 Secondary valency A secondary valency is equal to the coordination number and is fixed for a metal. Secondary Valency Primary valency Secondary valency It is equal to the oxidation number of central atom. It is equal to the coordination number. It is non-directional. It is directional in nature. It is ionisable. It is non-ionisable. Limitations of Werner’s Theory Werner could not explain : 1 Why only certain elements possess the remarkable property of forming coordination compounds? 2 Why the bonds in coordination compounds have directional properties? 3 Why coordination compounds have characteristic magnetic and optical properties? Valence Bond Theory The valence bond theory, VBT, was extended to coordination compounds by Linus Pauling. In 1931 Valence Bond Theory Postulate 1 The formation of complex involves a reaction between a Lewis base (ligand) and a Lewis acid (metal or metal ion) With the formation of a coordinate covalent (or dative) bonds. Postulate 2 In order to form a bond, VBT utilises the concept of hybridisation, in which (n-1)d, ns, np or ns, np, nd orbitals of a metal atom or ion are hybridised to yield a set of equivalent orbitals of definite geometry. Note!! Sometimes, the unpaired (n-1)d electrons pair up as fully as possible prior to hybridisation, thus making some (n-1)d orbitals vacant for hybridisation. Valence Bond Theory Postulate 3 These hybrid orbitals are allowed to overlap with ligand orbitals that can donate electron pairs for bonding. Consequently, these bonds are of equal strength and are directional in nature. Postulate 4 The number of unpaired electrons measured by the magnetic moment of the compounds determines which d-orbitals are used. Valence Bond Theory The hybridisation and shape of the complexes can be predicted with the help of some known properties. Magnetic moment Hybridisation of Coordination Compounds C.N. of metal 2 3 C.N. of metal 4 4 4 Type of hybridisation sp sp2 Type of hybridisation sp3 dsp2 d 3s Shape of complex Linear Trigonal planar Shape of complex Tetrahedral Square planar Tetrahedral Type of d-orbital - - Type of d-orbital - d x 2 – y2 dxy, dyz, dxz Hybridisation of Coordination Compounds C.N. of metal 5 5 C.N. of metal 6 6 Type of hybridisation sp3d dsp3 Type of hybridisation sp3d2 d2sp3 Shape of complex Trigonal bipyramidal Square pyramidal Shape of complex Octahedral Octahedral Type of d-orbital d x2 – y2 , dz2 Type of d-orbital d z2 d x 2 – y2 d x2 – y2 , d z2 2 3 Coordination number 4 5 6 Coordination Number = 2 + [Ag(NH3)2] μex = 0 No unpaired electrons Diamagnetic Coordination Number = 2 + Ag 4d10 = 5s0 4d10 5s0 5p0 5p0 sp-hybridised orbitals : : NH3 NH3 + [Ag(NH3)2] Hybridisation = sp Shape = Linear Coordination Number = 3 _ [Hg(CN)3] μex = 0 No unpaired electron Diamagnetic [Hg(CN)3] _ Hybridisation = sp2 Shape = Trigonal planar Coordination Number = 3 2+ Hg 5d10 = 5d10 6s0 6p0 6s0 6p0 sp2-hybridised orbitals CN _ _ _ CN CN Coordination Number = 4 [Ni(Cl)4] μex = 2.5 BM 2 unpaired electrons 2 _ 2 Paramagnetic _ [Ni(Cl)4] Hybridisation = sp3 Shape = Tetrahedral Coordination Number = 4 Ni2+ ↿⇂ ↿⇂ ↿⇂ 3d8 ↿ ↿ 4s0 4p0 ↿⇂ ↿⇂ sp3-hybridised orbitals Cl 3d8 4s0 4p0 = _ Cl _ ↿⇂ Cl ↿⇂ _ Cl _ Coordination Number = 4 [Ni(CN)4] 2 _ Hybridisation μex = 0 No unpaired electron Diamagnetic Shape = dsp2 Square planar Coordination Number = 4 Ni 2+ 3d8 4s0 4p0 = ↿⇂ ↿⇂ ↿⇂ ↿⇂ 3d8 dsp2 -hybridised orbitals _ CN 4p0 4s0 ↿⇂ ↿⇂ CN _ ↿⇂ Inner orbital or low-spin complex ↿⇂ _ _ CN CN Inner Orbital and Outer Orbital Complex In the complex formation, the inner d-orbitals are used in the hybridisation. In the complex formation, the outer d-orbitals are used in the hybridisation. Coordination Number = 4 3 VO4 _ Hybridisation = μex = 0 No unpaired electron Diamagnetic Shape Another example = d3s Tetrahedral = KMnO4 & K2Cr2O7 Coordination Number = 4 3 VO4 _ 5+ V 3d0 4s0 4p0 = 4s0 3d0 4p0 Inner orbital or low-spin complex d3s-hybridised orbitals 2 O _ 2 O _ 2 O _ 2 O _ Coordination Number = 5 Fe(CO)5 μex = 0 Diamagnetic No unpaired electron Hybridisation = dsp3 Shape = Trigonal bipyramidal Coordination Number = 5 0 Fe 3d6 = 4s0 3d6 4s2 4p0 4p0 dsp3-hybridised orbitals Inner orbital or low-spin complex Electron pairs from CO Coordination Number = 6 [Co(CN)6] 3– Hybridisation = d2sp3 = Octahedral μex = 0 Shape Diamagnetic No unpaired electron Coordination Number = 6 Co3+ = ↿⇂ ↿⇂ ↿⇂ 3d6 d2sp3 hybridised orbitals 3d6 4s0 4p0 4s0 ↿⇂ ↿⇂ ↿⇂ 4p0 ↿⇂ ↿⇂ ↿⇂ Electron pairs from CN– Inner orbital or low-spin complex Coordination Number = 6 [CoF6] 3 _ μex = 4.9 BM 4 unpaired electrons Paramagnetic Hybridisation = sp3d2 Shape = Octahedral Coordination Number = 6 Co ↿⇂ ↿ ↿ 3d6 ↿ ↿ 4s0 sp3d2 hybridised orbitals 3+ = 3d64s04p0 4d0 4p0 ↿⇂ ↿⇂ ↿⇂ ↿⇂ ↿⇂ ↿⇂ _ Electron pairs from F Outer orbital complex Limitations of Valence Bond Theory It does not give a quantitative 1 It involves a number of assumptions. 2 It does not give quantitative interpretation of magnetic data. 4 It does not make exact predictions 5 3 It does not explain the colour exhibited by the coordination compounds. interpretation of the thermodynamic or kinetic stabilities of coordination compounds. 6 regarding the tetrahedral and square planar structures of 4-coordinate complexes. It does not distinguish between weak and strong ligands. Crystal Field Theory The crystal field theory (CFT) is an electrostatic model. Assumption 1 According to the electrostatic model, the metal-ligand bond has to be ionic and the interactions are considered as purely electrostatic. Crystal Field Theory Assumption 3 Assumption 2 When the ligands are The transition metal that is present as the central atom in the complex is considered as a positive point charge. Anions Point charges Neutral molecules Point dipoles Crystal Field Theory The arrangement of ligands around the central metal ion is such that the repulsions between these points are minimum. Assumption 4 It considers the effect of ligands on the relative energies of the d-orbitals of the central metal atom/ion. Crystal Field Theory Assumption 5 Orbitals that are identical in energy In an isolated gaseous metal ion Five d-orbitals are degenerate. Degenerate d-orbitals (Free metal atom/ion) Crystal Field Theory Assumption 6 This degeneracy is still maintained. If a spherically symmetrical field of ligands surrounds the metal atom/ion Average energy Degenerate d-orbitals (Free metal atom/ion) Overall energy increased due to spherical field of ligands Crystal Field Theory Either anions or the negative ends of dipolar molecules like NH3 and H2O When there is a negative field due to ligands in a complex. It becomes asymmetrical and the degeneracy of the d-orbitals is lost Results in splitting of d-orbitals Crystal Field Theory The pattern of the splitting depends upon the nature of the crystal field exerting its influence on the central metal atom/ion. For both t2g and eg orbitals, the d-electrons occupy the orbitals singly in accordance with Hund’s rule. dz dx 2 - y 2 Free metal atom/ion 2 Barycentre line Average energy (in spherical crystal field of ligand) dxy dyz dzx Splitting of d-orbitals Ligands exerting octahedral field Crystal Field Theory The crystal field splitting depends upon the field produced by CFT gives an idea about the electronic distribution of the central metal ion in the given complex. Ligand Properties like electronic spectra, colour, and magnetic moment can be determined. Metal ion Crystal Field Theory Ligands are able to produce Strong field Splitting will be larger Generally, ligands can be arranged in a series of increasing field strength. Weak field Splitting will be smaller Such a series is termed as spectrochemical series. Spectrochemical Series Experimentally determined series I − − − < Br < S 2− − < Cl − < NO3 < − F < OH < EtOH < Oxalate < H2O < EDTA < NH3 and pyridine < Ethylenediamine < Dipyridyl < − o-phenanthroline < NO2− < CN < CO Spectrochemical Series I− − − < Br 2− < < S Cl− − < NO3 < F− < OH < EtOH < Oxalate < H2O < EDTA < NH3 and pyridine < Ethylenediamine < Dipyridyl < − o-phenanthroline < NO2− < CN < CO Generally, strong field ligands Spectrochemical Series Generally, weak field ligands I − − − < Br < S 2− < Cl− − < NO3 < − F < OH < EtOH < Oxalate < H2O < EDTA < NH3 and pyridine < Ethylenediamine < Dipyridyl < − o-phenanthroline < NO2− < CN < CO Approach of Ligands in Octahedral Complex L4 Ligands approach from/towards the axes. Z L3 Y L5 X X L2 Y L6 Z L1 The orbitals present towards the axes are repelled greater than the orbitals present between the axes. As a result, the energies of d x2-y2 and dz2 are increased, and dxy, dyz, and dxz are decreased relative to the average energy in the spherical crystal field. Splitting in Octahedral Complex Results Degeneracy of the d-orbitals is lost due to Ligand's electron-metal electron repulsions in the octahedral complex Yield three orbitals of lower energy (t2g) and two orbitals of higher energy (eg) Splitting in Octahedral Complex d-orbitals i.e., l = 2 dxy dyz dzx dx2 - y2 dz eg t2g Triply degenerate 2 Gerade Orbitals with centre of symmetry Doubly degenerate Gerade Orbitals with centre of symmetry d-orbital splitting in an octahedral crystal dx2 - y2 dz2 d-orbital splitting in an octahedral crystal eg Average energy 3 Δ o 5 Δo 2 Δ o 5 Free metal atom/ion Due to spherically symmetrical field t2g dxy dyz dzx Due to octahedral ligand field Splitting in an Octahedral Complex The energy of separation between eg and t2g is known as crystal field splitting energy (Δo) in an octahedral complex. eg Barycentre line Free metal atom/ion Average energy (in spherical crystal field) t2g Trick to Remember Pattern of increasing sigma donation Halide donors < O donors < N donors < C donors Splitting in Octahedral Complex In an octahedral coordination entity with six ligands surrounding the metal atom/ion, there will be repulsion between 1 Electrons in metal’s d-orbitals 2 Electrons (or negative charge) of the ligands To understand the shift of energy of orbitals, let’s first see the structures of d-orbitals. Directed between the axes dxy dzx dyz Directed along the axes dx 2 - y 2 dz 2 Splitting in an Octahedral Complex For t2g orbitals, the d-electrons occupy the orbitals singly in accordance with Hund’s rule. For d4-configuration, two possible e¯ distributions are 4th electron will singly occupy eg orbital 4th electron will undergo pairing in t2g orbital Splitting in an Octahedral Complex Two possibilities Δo < P Δo > P Pairing Energy(P) The energy required for electron pairing in a single orbital Splitting in an Octahedral Complex The 4th electron enters one of the eg orbitals Giving the 1 3 configuration t2g eg eg Barycentre line Free metal atom/ion Average energy (in spherical crystal field) t2g Splitting in an Octahedral Complex Ligands with ∆o < P Weak field ligands Generally, ligands form a high-spin complex. Ligands with ∆o > P The 4th electron becomes energetically more favourable to occupy t2g orbital with 4 configuration t2g eg0 eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g Pairing of electrons Splitting in an Octahedral Complex Ligands with ∆o > P Similarly, calculations are done for d5, d6, and d7 configurations. Strong field ligands Generally, ligands form a low-spin complex. For d8, d9, d10 configurations, the d-electrons occupy the eg orbitals in accordance with Hund's rule. Crystal Field Stabilisation Energy CFSE The energy released by which a complex is stabilised (compared to the free ion) due to the splitting of the dorbitals = [-0.4(n) t2g + 0.6(n’) eg]Δo + n*P Where, n: Number of e– occupying t2g orbital n': Number of e– occupying eg orbital Δo: Ligand field splitting energy n*: Number of extra electron pairs P: Spin-pairing energy For d4, weak field ligand n’ = 1 eg Average energy Free metal atom/ion Barycentre line t2g n=3 Pairing does not occur CFSE = CFSE = [-0.4(3) + 0.6(1) ]Δo + 0 × P -0.6Δo n* = 0 For d4, strong field ligand n’ = 0 eg Average energy Free metal atom/ion Barycentre line t2g n=4 CFSE = [-0.4(4) + 0.6(0) ]Δo + 1 × P CFSE = -1.6Δo + P One extra paired e- n* = 1 d5 configuration – High spin eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [ 0.6(n’)eg – 0.4(n)t2g]Δo + n*P = = [0.6 × 2 – 0.4 × 3] Δo 0 d5 configuration – Low spin eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = [– 0.4 × 5] Δo + 2P = 2(P – Δo) d6 configuration – High spin eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = [0.6 × 2 – 0.4 × 4] Δo = -0.4 Δo d6 configuration – Low spin eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = [– 0.4 × 6] Δo + 2P = − 2.4 Δo + 2P For d7, weak field ligand Average energy n’ = 2 Free metal atom/ion eg Barycentre line t2g n=5 n* CFSE = [-0.4(5) + 0.6(2) ]Δo + 0 × P CFSE = -0.8Δo Pairing does not occur = 0 d7 configuration – Low spin eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = = [0.6 × 1 – 0.4 × 6] Δo + P -1.8 Δo + P d8 configuration eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = [0.6 × 2 – 0.4 × 6] Δo = -1.2 Δo d9 configuration eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = = [0.6 × 3 – 0.4 × 6] Δo -0.6 Δo d10 configuration eg Free metal atom/ion Barycentre line Average energy (in spherical crystal field) t2g CFSE = [0.6(n’)eg – 0.4(n)t2g]Δo + n*P = [0.6 × 4 – 0.4 × 6] Δo = 0 Calculation of CFSE CFSE of d4, d5, d6, and d7 electronic configurations in the presence of WFL and SFL Electronic CFSE of WFL CFSE of SFL configuration d4 -0.6Δo -1.6Δo d5 0 -2.0Δo d6 -0.4Δo -2.4Δo d7 -0.8Δo -1.8Δo Calculations show that d4 to d7 coordination entities are more stable for a strong field as compared to a weak field. Relation of a Tetrahedron with a Cube z x y Tetrahedral geometry in a cube Orientation of d-Orbitals Relative to a Cube dyz dxz d z2 dxy d x2-y2 Splitting in Tetrahedral Complexes In the splitting of tetrahedral complexes, the direction of approach of the ligands does not coincide exactly with either the eg or the t2g orbitals. t2-Orbitals e-Orbitals Nearer to the ligand’s approach direction Away from the ligand’s approach direction Raised in energy in the crystal field of ligands Energy raised but lower than t2 orbitals Example: dxy, dyz, dxz Example: dx2-y2, d z2 Note!! The crystal field splitting in a tetrahedral complex is in the opposite way to that in an octahedral complex. dxy d-orbital splitting in a tetrahedral crystal Average energy dyz dzx t2 2 Δ t 5 Δt 3 Δ 5 t Free metal atom/ion e dx2 - y2 dz2 The crystal field splitting in a tetrahedral complex is in the opposite way to that in an octahedral complex. Splitting in Tetrahedral Complex The ‘g’ subscript is used for the octahedral complexes that have centre of symmetry. Since tetrahedral complexes lack symmetry, ‘g’ subscript is not used with energy levels. The magnitude of the crystal field splitting (Δt) in tetrahedral complexes is considerably less than in octahedral fields. Δt = 2 2 4 × ΔO ≈ ΔO 3 3 9 Splitting in Tetrahedral Complexes Δt = 2 2 4 × ΔO = ΔO 3 3 9 Δt = 2 2 4 × ΔO = Δ 3 3 9 O There are only four ligands instead of six. So, the ligand field is only two-third the size. The direction of the orbitals does not coincide with the direction of the ligands. Hence, the ligand field splitting is also two-thirds the size. This reduces the crystal field splitting by roughly a further two-third. Splitting in Tetrahedral Complexes Results Strong field ligands cause a bigger energy difference between t2g and eg than weak field ligands. Tetrahedral splitting Δt is always much smaller than octahedral splitting Δo. It is never energetically favourable to pair electrons. All tetrahedral complexes are high-spin. CFSE In Tetrahedral Complexes CFSE = [ 0.4(n) t2 – 0.6(n’) e ]Δt Where, n: Number of electrons occupying t2 orbital n’: Number of electrons occupying e orbital Δt: Ligand field splitting energy (n = 3) Average energy t2 Free metal atom/ion Barycentre line d5 configuration (n’ = 2) e CFSE in Tetrahedral Complexes CFSE = [0.4(n) t2 − 0.6(n’) e]Δt = [0.4 × 3 − 0.6 × 2]Δt = [1.2 − 1.2]Δt = 0 (n = 4) Average energy t2 Free metal atom/ion Barycentre line d8 configuration (n’ = 4) e CFSE in Tetrahedral Complexes CFSE = [ 0.4(n) t2 − 0.6(n’) e ]Δt = [ 0.4 x 4 − 0.6 x 4 ]Δt = [ 1.6 − 2.4 ]Δt = -0.8Δt Square Planar Complexes The square planar arrangement of ligands is derived from octahedral ligand field by removing two trans ligands located along the z-axis. d x2 - y2 Crystal field splitting in square planar complexes d x 2 - y2 eg Δ1 d x2 - y2 d z2 Average energy d z2 Δo= 10Dq dxy Δ2 Barycentre d z2 dxy t2g dxy Δsp= Δ1 + Δ2 + Δ3 dyz Δ3 dzx dyz dzx dyz dzx Δsp More repulsion Along the axis Greater energy dxy Between the axes Energy raised, but lower than d x2-y2 d z2 Along the axis Lower energy than dxy Between the axes Least energy d x 2 - y2 d-orbitals dxz & dyz CFSE in Square Planar Complexes Δsp > Δo Because dxz and dyz orbitals interact with only two ligands in square planar complexes, but in octahedral complexes, interaction takes place with four ligands Δsp = 1.3Δo Factors Affecting Stability of Complexes Stability of a complex 1 ∝ CFSE Oxidation state of central atom CFSE ∝ For the same metal ion, same ligand, and fixed coordination number Oxidation state ∝ CFS Factors Affecting Stability of Complexes [Fe(CN)6]4− [Fe(CN)6]3− Fe2+ Fe3+ Low CFSE value High CFSE value Due to high ∆o Factors Affecting Stability of Complexes 2 Nature of ligand Strong field ligand (SFL) Weak field ligand (WFL) More CFSE value Less CFSE value Because the strength of ligand can vary the ∆ of the complex [Co(NH3)6] 3+ [CoF6] 3− SFL WFL More CFSE value Less CFSE value Factors Affecting Stability of Complexes 3 Order of CFSE/CFS Geometry of complex Square planar 4− [Ni(CN)6] Octahedral 2− [Ni(Cl)4] Tetrahedral Order of CFSE > Octahedral Tetrahedral > 2− [Ni(CN)4] Square planar 2− [Ni(CN)4] > 4− [Ni(CN)6] > 2− [Ni(Cl)4] Factors Affecting Stability of Complexes 4 Order of CFSE Nature of metal cation CFSE ∝ For the same ligand and oxidation state, i.e., 3d < 4d < 5d Size of metal cation Co3+ < Rh3+ < Ir3+ Ni2+ < Pd2+ < Pt2+ Cu+ < Ag+ < Au+ Down the group Factors Affecting Stability of Complexes 5 Formation constant (Kf) or stability constant (Ks) Stability ∝ Kf or Ks For the same metal ion, same ligand, and fixed coordination number Factors Affecting Stability of Complexes General reaction M(H2O)n + nL Ks = MLn + nH2O Kf = Ks Cu2+ + 4NH3 [Cu(NH3)4 ]2+ [MLn] [M(H2O)n] [L]n = Kf = [Cu(NH3)4]2+ [Cu2+] [NH3]4 Note!! Greater the value of Kf/Ks, more is the thermodynamic stability of a complex. Factors Affecting Stability of Complex 6 Stability Chelation ∝ Number of rings in a chelate complex A chelating complex would be generally more stable. Magnetic Moment of a Complex Generally, Coordination compounds have partially filled d-orbitals. Formula Spin-only magnetic moment μ Due to this, they show a characteristic magnetic property. = n(n + 2) B.M. Where, n is the number of unpaired electrons. Magnetic Moment of a Complex [FeF6] 3− μ = 5(5 + 2)B.M. μ = 35 B.M. μ = Fe3+ 3d5 5 unpaired electrons 4s 0 5.91 B.M. Limitations of CFT Limitation 1 It considers only the metal ion d-orbitals. It gives no consideration at all to other metal orbitals (s, px, py, and pz orbitals). Limitation 2 It is unable to explain the relative strengths of ligands. For example, it gives no explanations regarding why − H2O is a stronger ligand than OH in the spectrochemical series. Limitations of CFT Limitation 3 Limitation 4 The CFT could not explain the 𝝅bonding in complexes. Partial covalent character of bonding According to this theory, the bonds between the metal and ligands are purely ionic. It gives no account on the partially covalent nature of the metal-ligand bonds. Colour in Coordination Compounds Colour of complex mainly explained by d-d transition Charge transfer spectra Colour of Coordination Compounds When an electron from a lower energy d-orbital is excited to a higher energy d-orbital Average energy Free metal atom/ion h𝛎 Note!! A metal complex must contain an unpaired electron for a d-d transition to occur. Colour in Coordination Compounds Due to the absorption of light at a specific wavelength in the visible region 400 to 700 nm And the transmission or reflection of the light with rest of the wavelengths An object that absorbs all visible light appears black. Absorption and Transmission Process in Complexes [Ti(H2O)6]3+ eg h𝛎 t2g Ground state After absorption a photon of energy exactly equal to energy difference between t2g and eg, the electron does a transition from t2g to eg. Excited state ∆o or Eabsorbed d-d transition Colour in Coordination Compounds Due to the absorption of light at a specific wavelength in the visible region And transmission or reflection of the light with rest of the wavelengths Corresponding wavelength gives complementary colour Complementary Colour The complementary colour is generated from the wavelength left over. R V O B Y G Colour in Coordination Compounds In the absence of a ligand No d-d transition Crystal field splitting does not occur as all five d-orbitals are degenerate. Hence, the substance is colourless. Colour in Coordination Compounds [Ti(H2O)6]Cl3 Violet Δ TiCl3 + 6H2O Colourless No d-d transition Degenerate orbitals Charge Transfer Spectra The transition in which an electron is transferred from one atom or group to another one Charge transfer transitions are also considered as internal redox reactions. Charge Transfer Transitions Ligand to metal charge transfer (LMCT) Charge transfer transitions are mainly of three types Metal to ligand charge transfer (MLCT) Metal to metal charge transfer (MMCT) Ligand to metal charge transfer eMn+ ElectronMetal acceptor L Liganddonor Electron Ligand To Metal Charge Transfer Charge transfer occurs from the ligand's filled molecular orbitals to the metal’s empty or partially filled d-orbitals. Generally, metals having high oxidation state (+5,+6,+7) with 𝜋 donor ligands show LMCT. Ligand To Metal Charge Transfer − MnO4 2− MnO4 Deep purple Green 2− Orange 2− Yellow Cr2O7 CrO4 Electron donor Electron acceptor Mn+ e- L Metal to ligand charge transfer Charge transfer occurs from the metal’s filled orbitals to the empty orbital of ligand. Note!! HgO Generally, metals having low oxidation state (0,+1,+2) with 𝜋 acceptor ligands show LMCT. [Fe(phen)3]3+ [Fe(acac)3] Metal To Metal Charge Transfer Intervalence charge transfer The excitation and subsequent transfer of electrons from a lower O.S cation to neighbouring cation of a higher O.S Oxidation state Prussian blue complex III II 4FeCl3 + 3K4[Fe(CN)6] Fe4[Fe(CN)6]3 + 12KCl Fe4[Fe(CN)6]3 Metal To Metal Charge Transfer Turnbull’s blue complex 3FeCl2 + 2K3[Fe(CN)6] Fe3[Fe(CN)6]2 + 6KCl Turnbull’s blue Oxidation state II III Fe3[Fe(CN)6]2 Now, Let’s See the Different Types of Charge Transfer Transitions in Gemstones! Gemstones Gemstones Ruby Emerald Ruby is aluminium oxide (Al2O3) Containing about 0.5-1 % Cr3+ ions (d3). The colour shown by ruby is Randomly distributed in Al3+ positions Due to the d-d transitions Emerald Emerald Green It transmits light in the green region. Here also the colour is due to impurities in crystal. Organometallic Compounds Organometallic compounds are the compounds that contain carbon-metal bond. Metal carbonyls have simple, welldefined structures. Metal carbonyls are the transition metal complexes of carbon monoxide containing metalcarbon (M-C) bonds. Organometallic Compounds Example Organolithium Compound δ_ δ+ Organomagnesium compound δ_ δ+ δ_ Metal Carbonyls Fe(CO)5 Bonding in Metal Carbonyls Metal–carbon bond in metal carbonyls possesses 𝞂-character M–C 𝞂–bond is formed by 𝛑-character The donation of lone pair of electrons of carbonyl carbon into vacant d-orbital of metal 𝞂-Bonding in Metal Carbonyls Empty d-orbital of metal M 𝞂 NBMO M–C 𝛑–bond is formed by : The donation of lone pair of electrons from filled d-orbital of metal to vacant 𝛑* orbital of CO 𝛑–Bonding in Metal Carbonyls Synergic bond Back donation M 𝛑 𝞂 𝛑 Vacant 𝛑* Bonding in Metal Carbonyls Metal–carbon bond in metal carbonyls possesses 𝞂-character 𝛑-character M–C 𝞂–bond is formed by The donation of lone pair of electrons of carbonyl carbon into vacant d-orbital of metal 𝞂-Bonding in Metal Carbonyls Empty d-orbital of metal M 𝞂 NBMO : 𝛑–Bonding in Metal Carbonyls Back donation M–C 𝛑–bond is formed by The donation of lone pair of electrons from filled d-orbital of metal to vacant 𝛑* orbital of CO M Synergic bond 𝛑 𝞂 𝛑 Vacant 𝛑* Synergic Bond Synergic effect is due to a selfstrengthening bond, which involves transference of electrons from The 𝛑-bond formed between a metal atom/ion and a ligand by back donation of a metal atom/ion is known as a synergic bond. 1 2 Ligand to metal Metal’s filled orbitals to vacant orbitals of ligand Synergic Effect Back bond/ Synergic bond M 𝛑 𝞂 𝛑 Vacant 𝛑* CO is a nonclassical ligand. Those ligands which not only donate the lone pair of electrons but also accept the electron cloud from the central metal atom/ion in their low-lying vacant orbitals. Synergic Effect The length of the internal bond of a ligand and the length of the metal–ligand bond vary due to the synergic bonding. This effect is known as synergic effect. Because electrons are coming in 𝛑* orbitals of CO Bond order of the C–O bond decreases BUT The length of C–O bond increases Synergic Bonding More electron rich the metal, i.e., lower the oxidation state of metal More the donation of electrons from metal to ligands More the increase in C–O bond length Synergic Bonding Metal nitrosyls Synergic bonding in case of Metal phosphenes Alkenes Example: [Fe(H2O)5NO]2+ 𝛑-Bonding in Metal Nitrosyls Back bond/ Synergic bond 𝛑 M Vacant 𝛑* 𝞂 𝛑 Formation of M NO+ bond by back donation in metal nitrosyls Synergic Bonding in Metal Phosphenes Vacant 3d orbital accepts the back donation 𝛑 M 𝞂 :PR3 𝛑 Wilkinson’s catalyst Synergic Bonding in Alkenes The filled d-orbitals of metal overlap with vacant 𝛑* molecular orbital of ethylene molecule and forms 𝛑bond/backbond/ synergic bond. _ Cl C2H4 Zeise’s salt K+ K[PtCl3(𝝶2-C2H4) Pt Cl Cl Synergic Bonding in Alkenes Synergic bonds 𝛑* 𝞂-bond 𝛑* Stability of Organometallic Compounds Rules 18 electron rule Proposed by Irving Langmuir EAN rule Proposed by Sidgwick 18 Electron Rule When the central metal atom or ion of a complex acquires an outer shell electronic configuration of (n-1)d10ns2np6 There will be 18 electrons in the valence shell And the electronic configuration will be closed and stable Total valence electron = Contribution of metal ion + Contribution of ligand ± Charge on coordination sphere Counting the Number of Electrons Neutral atom (Covalent model) Two conventions for counting the number of electrons in a complex We will study only this Scandium Cobalt Oxidation state (Ionic model) = Group number of metal Group number =3 Electron count = 3 Electron count of metal ion Group number = 9 Electron count = 9 Electron Count of Ligand Electron count of anionic ligand = Charge on anion Anionic Ligand classification Neutral H– Electron count = 1 Organic Cl – 2– O Electron count = 1 Electron count = 2 Neutral Ligand Electron count of neutral ligand Number of electron donated = Ethylene diamine CO Electron count = 4 Electron count = 2 𝞂 M NH3 Electron count = 2 M 𝞂 NH3 𝞂 M 𝞂 Organic Ligand Allyl ion Electron count = Heptacity = 3 Organic ligands are ligands having unsaturated bond. M Electron count of organic ligand = Hepacity of ligand Electron count = Heptacity = 5 cyclopentadiene M Counting of Number of Electrons Covalent model Fe 8 electrons 2C5H5– 10 electrons Ferrocene Total electrons = 18 electrons Counting of Number of Electrons Covalent model (CO)5Mn CH3 Mn 7 electrons 5CO 10 electrons CH3– 1 electron Total electrons = 18 electrons Limitations of 18-Electron Rule 1 The 18-Electron rule is not a hard and fast rule and is not always strictly obeyed. There are number of stable complexes in which the electron count is not equal to 18. Examples in which the number of electrons ≠ 18 [Cr(H2O)6]2+ [Zn(NH3)6]2+ 16 electrons 22 electrons Both the complexes exist and are stable. Limitations of 18-Electron Rule 2 The transition metal ions on the left side of the periodic table have a lower number of valence electrons. [Ti(F)6]2- [V(Cl)6]3- Such transition metals cannot acquire 18 electrons in their valence shell. Complexes of these transition metals have less than 18 electrons in their valence shell Both the complexes exist and are stable. 12 electrons 14 electrons Effective Atomic Number Rule (EAN) The sum of electrons on the central metal atom or ion and electrons donated from the ligand is known as effective atomic number. Proposed by Sidgwick Generally, EAN is equal to the atomic number of the nearest noble gas to the central metal atom. Krypton (Kr) 36 Xenon (Xe) 54 Radon (Rn) 86 Effective Atomic Number Rule Mathematically EAN = Z − O.S Number of electrons in the central metal atom or ion EAN = Number of electrons in the central metal atom or ion + + 2(C.N) Number of electrons donated by ligands Number of electrons donated by ligands Effective Atomic Number Rule [Cr(CO)6] O.S of Cr = 0 Electrons in Cr atom = Atomic number of Cr 24 Number of electrons donated by 6CO 2 X 6 = 12 36 Obeys EAN rule Note!! Some complexes can be stabilised by following EAN rule. Applications of the EAN Rule 1 The number of CO molecules attached in mononuclear carbonyls can be predicted. Fe(CO)x Ni(CO)y 2 Some complexes can act as reducing agents by losing electron. Cr(CO)z To attain the noble gas configuration 5 4 6 Applications of the EAN Rule [Mn(CO)6] EAN = 37 ─ e─ Oxidation [Mn(CO)6]+ EAN = 36 Obey Sidgwick's EAN rule Applications of the EAN Rule 3 Some complexes can act as oxidising agents by gaining electrons. [V(CO)6] EAN = 35 To attain the noble gas configuration + e─ Reduction [V(CO)6]─ EAN = 36 Obey Sidgwick's EAN rule Applications of the EAN Rule 4 Some complexes undergo dimerisation. To attain the noble gas configuration 2[Mn(CO)5] [Mn2(CO)10] EAN = 35 EAN = 36 Obey Sidgwick's EAN rule Isomerism in Coordination Compounds Structural isomerism Structural isomerism Isomerism Stereo isomerism An isomer in which molecules with the same molecular formula have different bonding patterns Structural Isomerism Linkage Structural Isomerism Coordination Ionisation Solvate Linkage Isomerism CN─, NO2─, SCN─, etc. The isomers formed when ambidentate ligands use different coordinating atoms/ sites for binding metal ions Ligands that can ligate through two different sites present in them. Linkage Isomerism [Co(NH3)5NO2]Cl2 2+ Nitropentamminecobalt(III) ion 2+ Nitritopentamminecobalt(III) ion Linkage Isomerism Jorgensen observed the difference in colour of the two isomers. [Co(NH3)5NO2]Cl2 2+ 2+ Nitropentamminecobalt(III) ion Nitritopentamminecobalt(III) ion Yellow Red Linkage Isomerism 2─ [PtCl3(SCN)] 2– 2– Coordination Isomerism [Pt(NH3)4]2+ [CuCl4]2─ It arises due to an interchange of ligands between cationic and anionic entities of different metal ions present in the complex. [PtCl4]2─ [Cu(NH3)4]2+ Coordination isomer Coordination Isomerism _ 2+ 2 + _ Coordination Isomerism 0 _ 0 + Coordination Isomerism 2 _ Total coordination isomers = 4 2+ Coordination Isomerism 0 This is not considered as a coordination isomers. _ 0 + Points to Remember! If the coordination entity becomes neutral, then it is not considered as a coordination complex and it exists as a free entity in the solution. Coordination Isomerism 3+ 3– 3+ 3– [Co(NH3)6]3+ [Cr(CN)6]3─ [Cr(NH3)6]3+ [Co(CN)6]3─ Coordination isomer Ionisation Isomerism It arises when a counter ion in a complex salt itself is a potential ligand. [Pt(NH3)3Br] NO2 So, it can displace an actual ligand, which can then become a counter ion. [Pt(NH3)3NO2] Br Counter ion Note!! Ionisation isomers produce different ions in aqueous solutions. [Co(NH3)5(SO4)]Br H2O [Co(NH3)5(SO4)]+ + Br ─ Red [Co(NH3)5Br]SO4 H2O [Co(NH3)5Br]2+ + SO42─ Violet Solvate Isomerism Isomers that differ with respect to the number of solvent ligand molecules as well as counter ions 1 Ammonia (NH3) 2 Water (H2O) Hydrate Isomerism Produces 3Cl─ CrCl3.6H2O [Cr(H2O)6]Cl3 Violet When water is present as a solvent, the solvate isomerism is known as hydrate isomerism. Produces 2Cl─ [Cr(H2O)5Cl]Cl2.H2O Bright blue green [Cr(H2O)4Cl2]Cl.2H2O Dark green Produces ─ 1Cl Note!! Solvate and hydrate isomerism are types of ionisation isomerism. Polymerisation Isomerism A special case of coordination isomerism but not a true isomerism In which compounds differ from each other in molecular formula & molecular weight. [Pt(NH3)2Cl2] and [Pt(NH3)4] [PtCl4] [Pt(NH3)4][Pt(NH3)2Cl3]2 and [Pt(NH3)3Cl]2 [PtCl4] Stereoisomerism Geometrical Isomerism due to different spatial arrangements of atoms or groups (ligands) about the central metal atom or ion Stereoisomerism Optical Geometrical and Optical Isomerism Coordination number Stereoisomerism in tetrahedral complexes 4 6 Geometrical Tetrahedral Square planar Octahedral Optical Geometrical Isomerism in Tetrahedral Complexes Generally, GI is not possible in tetrahedral complexes. [Ma4], [Ma2b2], or [Mabcd] The arrangement of every ligand around the central metal atom/ion in space is equivalent in every respect. Optical Isomerism in Tetrahedral Complexes A complex can show optical [Be(acac)2] isomerism only when the attached 4 ligands are different. But POS and COS must be absent acetylacetonato ─ ─ Optically inactive due to POS Geometrical and Optical Isomerism Coordination number 4 Tetrahedral 6 Square planar Octahedral Geometrical and Optical Isomerism Stereoisomerism in square planar complexes Geometrical Optical [Ma4]n± n± [M(AB)a2] [Ma3b]n± [Ma2b2]n± [M(AA)2]n± [Ma2bc]n± [M(AA)a2]n± [Mabcd]n± GI in Square Planar Complexes Where a, b, c, d ─ ─ CN , Br , NH3, etc. Simple monodentate ligands AA Symmetrical bidentate ligand AB Unsymmetrical bidentate ligand en, oxalato, etc. Glycinato [Ma4]n± [M(AB)a2]n± [Ma3b]n± [Ma2b2]n± [M(AA)2]n± [Ma2bc]n± [M(AA)a2]n± [Mabcd]n± GI in Square Planar Complexes Does not show geometrical isomerism [Ma4]n± [M(AB)a2]n± [Ma3b]n± [Ma2b2]n± [M(AA)2]n± [Ma2bc]n± [M(AA)a2]n± [Mabcd]n± Reason Simple monodentate ligand Since only one geometry is possible Symmetrical bidentate ligand Unsymmetrical bidentate ligand GI in Square Planar Complexes [Ma4]n± n± [M(AB)a2] [Ma3b]n± [Ma2b2]n± [M(AA)2]n± [Ma2bc]n± [M(AA)a2]n± [Mabcd]n± Shows geometrical isomerism GI in Square Planar Complexes [Mabcd] n± Total possible isomers: 3 [Pt(NH3)BrCl(py)] GI in Square Planar Complexes [M(AB)2]n ± [Pt(gly)2] cis trans Optical Isomerism in Square Planar Complexes [Pd(EDTA)]2─ Acts as a tetradentate ligand 2- Mirror Optically active compounds 2- GI in Square Planar Complexes [Pt(NH3)2(Cl)2] cis trans GI in Square Planar Complexes [Ma2bc]n± Same group opposite side (trans) Same group same side (cis) Optical Isomerism in Square Planar Complexes Generally, square planar complexes do not show optical isomerism. (Isobutylenediamine) (meso-diphenylethylenediamine) Palladium(II) or Platinum(II) complex 2+ 2+ This is because all the four ligands and metal cations exist in the same plane. Mirror Have POS Where M = Pd(II) or Pt(II) Geometrical and Optical Isomerism Stereoisomerism in octahedral complexes Geometrical Optical Geometrical Isomerism in Octahedral Complexes Octahedral complexes where ligands are Monodentate Symmetrical bidentate and monodentate type Having no chiral centre Unsymmetrical bidentate and monodentate type Octahedral Complexes with Monodentate Ligands Case 1(a) [M(a6)]n± Does not show Geometrical Isomerism Case 1(b) [M(a5b)]n± Does not show Geometrical Isomerism a a a b Octahedral Complexes with Monodentate Ligands Example Case 1(c) b b a b Cis [M(a4b2)]n± Cis form [M(a4b2)]n± Trans form Trans Octahedral Complexes with Monodentate Ligands Case 1(d) a b [Ma2b2c2]n± Octahedral Complexes with Monodentate Ligands Case 1(d) a a a a [Ma2b2c2]n± [Ma2b2c2]n± Octahedral Complexes with Monodentate Ligands Case 1(d) c b c b [Ma2b2c2]n± Possible number of geometrical isomers [Ma2b2c2]n± = 5 Octahedral Complexes Having Monodentate Ligands Case 1(e) [Ma3b3] Have three identical ligands on one triangular face n± a a b b Have three identical ligands in a plane bisecting the molecule Facial (fac) Meridional (mer) Possible number of geometrical isomers = 2 Octahedral Complexes Having Monodentate Ligands Case 1(e) [Ma3b3]n± Have three identical ligands on one triangular face a a b b Have three identical ligands in a plane bisecting the molecule Facial (fac) Meridional (mer) Possible number of geometrical isomers = 2 Octahedral Complexes Having Monodentate Ligands NH3 NO2 Facial (fac) NH3 NH3 Meridional (mer) Octahedral Complexes Having Monodentate Ligands Case 1(f) [M(abcdef)] n± Possible number of geometrical isomers = 15 Geometrical Isomerism in Octahedral Complexes Octahedral complexes where ligands are Monodentate Symmetrical bidentate and monodentate type Unsymmetrical bidentate and monodentate type Symmetrical Bidentate and Monodentate Ligands Case 2(a) n± [M(AA)3] A A Case2(b) 2(b) Case n± [M(AA)2a2] a a A a cis trans Does not show GI but chiral Possible number of geometrical isomers = 2 Symmetrical Bidentate and Monodentate Ligands [Pt(en)2Cl2]2+ 2+ cis 2+ trans Symmetrical Bidentate and Monodentate Ligands n± [M(AA)a2b2] Case 2(c) a a b b a b Possible number of geometrical isomers = 3 Symmetrical Bidentate and Monodentate Ligands Case 2(d) [M(AA)a2bc]n± a a a b a c b c Possible number of geometrical isomers = 4 Optical Isomerism in Octahedral Complexes Octahedral complexes where ligands are Monodentate Symmetrical bidentate and monodentate type Unsymmetrical bidentate and monodentate type Unsymmetrical Bidentate and Monodentate Ligands Case 3(a) n± [M(AB)3] mer-isomer [Co(gly)3] fac-isomer Possible number of geometrical isomers = 2 Possible number of geometrical isomers = 2 Unsymmetrical Bidentate and Monodentate Ligands Case 3(b) n± [M(AB)2a2b2] Possible number of geometrical isomers = 4 Optical Isomerism in Octahedral Complexes Octahedral compounds where ligands are Monodentate Symmetrical bidentate and monodentate Unsymmetrical bidentate and monodentate Octahedral Complexes Having Monodentate Ligands Case 1(a) n± [M(a4b2)] Does not show optical isomerism Possible number of enantiomers = 0 Octahedral Complexes Having Monodentate Ligands [M(a)2(b)2(c)2] Case 1(b) All possible GI’s n± a a a b b a a b Optically active Optically Optically inactive inactive c c Optically Optically inactive inactive Octahedral Complexes Having Monodentate Ligands [M(a)2(b)2(c)2] Case 1(b) n± Optical isomer a a b b Possible number of enantiomeric pairs = 1 Total possible stereoisomers = 6 Symmetrical Bidentate and Monodentate Ligands Case 2(a) : [M(AA)3]n A A A A ± Possible number of Enantiomeric pairs = 1 Total number of stereoisomers = 2 Symmetrical Bidentate and Monodentate Ligands [M(AA)2ab]n± Case 2(b) Possible GI’s A a A b Optically active Optically inactive Symmetrical Bidentate and Monodentate Ligands Optical isomers Case 2(b) [M(AA)2ab]n± A A A A Possible number of enantiomeric pairs = 1 Total possible stereoisomers = 3 Unsymmetrical Bidentate and Monodentate Ligands Case 3(a) [M(AB)3]n± Possible GI’s Optically active Optically active Unsymmetrical Bidentate and Monodentate Ligands Optical isomers [Co(gly)3] Possible number of enantiomeric pairs = 2 Total possible stereoisomers = 4 Unsymmetrical Bidentate and Monodentate Ligands Case 3(b) [M(AB)2a2]n± Optical isomers A A A A A A B B Unsymmetrical Bidentate and Monodentate Ligands Optical isomers Case 3(b) A A a a Possible number of enantiomeric pairs = 3 Total possible stereoisomers = 8 Analytical chemistry Metallurgy Applications and importance of coordination chemistry in Biological systems As Catalysts Medicinal chemistry Other industrial applications Applications in Analytical Chemistry 1 Detection and estimation of a metal Due to the colour produced by the metal when it reacts with a ligand Chelating ligands such as EDTA, DMG, etc. Applications in Analytical Chemistry 2 In estimation of hardness of water Because Ca2+ and Mg2+ ions form stable complexes with EDTA logKf Ca2+ + EDTA Ca(EDTA)2─ 10.69 Mg2+ + EDTA Mg(EDTA)2─ 8.79 Done by titration with Na2EDTA Selective estimation due to the difference in stability constants Coordination Chemistry in Metallurgy In extraction of metals Examples: Au, Ag Roasted + CN─ + H2O gold ore O2 [Au(CN)2]─ + OH─ In aqueous solution Coordination Chemistry in Metallurgy Purification of metals By formation and subsequent decomposition of their coordination compounds Ni + 4CO Impure 330−350 K Ni(CO)4 450−470 K Ni + CO Pure Coodination Chemistry in Biological System Name Metal involved Function Chlorophyll Mg Photosynthesis Haemoglobin Fe Carries Oxygenated blood Vitamin B12 (Cyanocobalamin) Co Anti-pernicious anaemia factor Carboxypeptidase A Carbonic anhydrase Zn Enzymes Used as Catalysts Coordination compounds are used as catalysts in industrial process. Example Wilkinson’s Catalyst Hydrogenation of alkene (Ph3P)3RhCl H2C = CH2 + H2 (Ph3P)3RhCl H3C ─ CH3 Application in Medicinal Chemistry When metal is present in toxic proportion in plants/animals Excess metal is removed using chelating ligand which is known as Chelating therapy Cis-platin and related compounds effectively inhibit the growth of a tumour. Hence, it is used in the treatment of cancer. Application in Medicinal Chemistry Excess Metal Chelating ligand used in removal Cu D - penicillamine Fe Desferrioxamine B Pb poisoning EDTA Other Industrial Applications Electroplating Articles electroplated by silver [Ag(CN)2]─ and gold [Au(CN)2]─ Even and smoother solutions of other metal. Other Industrial Applications Black and white photography. Developed film is fixed and washed with hypo (Na2S2O3) Dissolves undecomposed AgBr to form [Ag(S2O3)2]3─