Coordination Compounds 1 h472b

Coordination Compounds Coordination Compounds − Complex compounds in which transition metal atoms are bound to a number of anions or neutral molecules (ligands). [Co(NH3)6]Cl3 Postulates of Werner’s Theory of Coordination Compounds  In coordination compounds, there are two types of linkages (valences) − primary and secondary.  The primary valences are ionisable, and are satisfied by negative ions.  The secondary valences are non-ionisable, and are satisfied by negative ions or neutral molecules. The secondary valence is equal to the coordination number of a metal, and remains fixed for a metal.  Primary valence is non-directional. Secondary valence are directional. Ligands have different spatial arrangement depending on coordination no. Difference between a Double Salt and a Complex Double salt Complex ion 1. A double salt dissociates completely to give 1.Complex ions do not dissociate into simpler ions in simpler ions in water. water. 2. It loses its identity in solution. 2. It retains its identity in solution. Eg: carnallite (KCl.MgCl2.6H2O), Mohr’s salt Eg: [Co(NH3)6]3+, [Fe(CN)6]4− [FeSO4.(NH4)2 SO4.6H2O] Definition of Important in Coordination Compounds  Coordination Entity Constitutes a central metal atom/ ion bonded to a fixed number of ligands [ions or molecules].Eg:[Co(NH3)6]3+  Central Atom or Ion The atom or ion to which a fixed number of ligands are attached in a definite geometrical arrangement around it in a coordination entity. Eg: Co3+ in [Co(NH3)6]3+ , Fe3+ in [Fe(CN)6]3−  Ligands Ions or molecules bound to the central metal atom or ion in a coordination entity Types of ligands 1. Unidentate. When a ligand is bound to a metal ion through a single donor atom., the ligand is said to be unidentate. Eg: Cl–, H2O or NH3 2. Didentate When a ligand is bind to a metal through two donor atoms, the ligand is said to be didentate. Eg: H2NCH2CH2NH2 (ethane-1,2-diamine) or C2O42– (oxalate) 3. Polydentate. When a ligand is bound to a metal ion through a several donor atoms., the ligand is said to be polydentate. Eg: Ethylenediaminetetraacetate ion (EDTA4–) is an important hexadentate ligand. It can bind through two nitrogen and four oxygen atoms to a central metal ion. 4. Ambidentate ligand Ligand which can ligate to a metal ion through two different atoms is called ambidentate ligand. Eg: NO2– and SCN– ions. NO2– ion can coordinate either through nitrogen or through oxygen to a central metal atom/ion.

*Similarly, SCN– ion can coordinate through the sulphur or nitrogen atom. 5. Chelate ligand and chelate complex When a di- or polydendate ligand uses its 2 or more donor atoms to bind a single metal ion so that a 5 or 6 membered ring are formed, it is said to be a chelate ligand. Such complexes are called chelate complexes.

Eg: [Pt Cl2(en)]  Dendicity of a ligand The no.of ligating group present in a ligand. Eg: Dendicity of Cl–, H2O or NH3 is 1,dendicity of oxalate,en is 2  Coordination Number [CN] Number of ligand-donor atoms bonded directly to the metal. Eg: The CN of Pt and Ni in [PtCl6]2− and [Ni(NH3)4]2+ are 6 and 4 respectively.  Coordination Sphere The central atom or ion and the ligands attached to it are enclosed in square brackets, which are collectively known as the coordination sphere.Eg: In K4[Fe(CN)6], the coordination sphere is [Fe(CN)6]4−, counter ion-K+.  Counter ion The ionisable group present outside the coordination sphere.Eg: In [Co(NH3)6]Cl3, counter ion is Cl Coordination Polyhedron The spatial arrangement of the ligand atoms which are directly attached to the central atom or ion o Example: Octahedral, square planar, tetrahedral

 Oxidation Number of Central Atom The charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom. Eg: Oxidation number of copper in [Cu(CN)4]3− is 1, Co in [Co(NH3)6]3+ is 3  Homoleptic and Heteroleptic Complexes Homoleptic − Complexes in which the metal is bound to only one kind of ligand [donor group]. Eg: [Co(NH3)6]3+ Heteroleptic − Complexes in which the metal is bound to more than one kind of ligands [donor groups]. Eg: [Co(NH3)4Cl2]+ Ligand

Charge

Normal name

IUPAC name

Cl Br F I CN OH NO3 NO2 ONO SCN NCS CO3 SO4 C2O4 ClO4 ClO3 NH3 H2O CO en CH3NH2 CH3CH2NH2 NO PPh3

-1 -1 -1 -1 -1 -1 -1 -1 -1 -1 -1 -2 -2 -2 -1 -1 0 0 0 0 0 0 0 0

Chloride Bromide Fluoride Iodide Cyanide Hydroxide Nitrate Nitrite Nitrite Thiocyanate Isothiocyanate Carbonate Sulphate oxalate perchlorate chlorate

Chlorido Bromido Fluorido Iodo Cyano/cyanido Hydroxo/ hydroxido Nitrato Nitrito-N Nitrito-O Thiocyanato Isothiocyanato Carbonato Sulphato Oxalato Perchlorato chlorato Ammine Aqua Carbonyl Ethane-1,2-diamine Methyl amine Ethyl amine Nitrosyl Triphenyl phosphine

Naming of Mononuclear Coordination Compounds The following rules are followed while naming coordination compounds.  The cation is named first in both positively and negatively charged coordination entities.  The ligands are named in alphabetical order before the name of the central atom/ion.  Names of the anionic ligands end in −o and those of neutral and cation ligands are the same. [Exceptions: aqua (H2O), ammine (NH3), carbonyl (CO), nitrosyl (NO)]  To indicate the number of the individual ligands, the prefixes mono−, di−, tri−, etc., are used. If these prefixes are present in the names of ligands, then the −bis, −tris, −tetrakis, etc., are used. For example, [NiCl2(PPh3)2] is named as Dichlorobis(triphenylphosphine)nickel(II).  Oxidation state of the metal is indicated by a Roman numeral in parentheses.  If the complex ion is cation, then the metal is named as the element. For example, Fe in a complex cation is called iron and Pt is called platinum.  If the complex ion is anion, then the metal is named with ‘−ate’ ending. For example, Co in a complex anion, [Co(SCN)4]2− is called cobaltate.  The neutral complex molecule is named as the complex cation. Formulas of Mononuclear Coordination Compounds The following are the rules for writing the formulas.  The central atom is listed first.  The ligands are then listed in the alphabetical order.  Polydentate ligands are also listed in the alphabetical order. In the case of abbreviated ligands, the first letter of the abbreviation is used for determining the position of the ligands in the alphabetical order.  The formula of the entire coordination entity is enclosed in square brackets. Ligand abbreviations and formulas for polyatomic ligands are enclosed in parentheses.  There should be no space between the ligands and the metal within a coordination sphere.  For the charged coordination entity, the charge is indicated outside the square brackets, as a right superscript, with the number before the sign. Example: [Fe(CN)6]3−, [Cr(H2O)6]3+, etc.  The charge of the cation(s) is balanced by the charge of the anion(s). IUPAC name of some coordination compounds are as follows: [Ni(CO)4] − Tetracarbonlynickel(0) [Co(NH3)3(H2O)3]Cl3 − Triamminetriaquachromium(III) chloride [Pt(NH3)2Cl(NO2)] − Diamminechloridonitrito-N-platinum(II) [CoCl2(en)2]Cl − Dichloridobis(ethane-1,2-diammine)cobalt(III) chloride K2[Zn(OH)4] − Potassium tetrahydroxozincate(II) Q) Write the formulas for the following coordination compounds: (i) Tetraamineaquachloridocobalt(III) chloride (ii) Potassium tetrahydroxozincate(II) (iii) Potassium trioxalatoaluminate(III (iv) Dichloridobis(ethane-1,2-diamine)cobalt(III) (v) Tetracarbonylnickel(0) A) (i) [Co(NH3)4(H2O)Cl]Cl2 (ii) K2[Zn(OH)4] (iii) K3[Al(C2O4)3 (iv) [CoCl2(en)2]+ (v) [Ni(CO)4] Q) Write the IUPAC names of the following coordination compounds: (i) [Pt(NH3)2Cl(NO2)] (ii) K3[Cr(C2O4)3] (iii) [CoCl2(en)2]Cl (iv) [Co(NH3)5(CO3)]Cl (v) Hg[Co(SCN)4] A) (i) Diamminechloridonitrito-N-platinum(II) (ii) Potassium trioxalatochromate(III) (iii)Dichloridobis(ethane-1,2-diamine)cobalt(III) chloride (iv) Pentaamminecarbonatocobalt(III) chloride (v) Mercury tetrathiocyanatocobaltate(III) Q)Write the formulas for the following coordination compounds: (i) Tetraamminediaquacobalt(III) chloride (ii) Potassium tetracyanonickelate(II) (iii) Tris(ethane–1,2– diamine) chromium(III) chloride (iv) Amminebromidochloridonitrito-N-platinate(II) (v) Dichloridobis(ethane–1,2–diamine)platinum(IV) nitrate (vi) Iron(III) hexacyanoferrate(II) A) (i) [Co(NH3)4(H2O)2]Cl3 (ii) K2[Ni(CN)4] (iii) [Cr(en)3]Cl3 (iv) [Pt(NH3)BrCl(NO2)] (v) [PtCl2(en)2](NO3)2 (vi) Fe4[Fe(CN)6]3 Q) Write the IUPAC names of the following coordination compounds: (i) [Co(NH3)6]Cl3 (ii) [Co(NH3)5Cl]Cl2 (iii) K3[Fe(CN)6] (iv) K3[Fe(C2O4)3] (v) K2[PdCl4] (vi) [Pt(NH3)2Cl(NH2CH3)]Cl (i)Hexaamminecobalt(III) chloride (ii)Pentaamminechoridocobalt(III) chloride

(iii)Potassium hexacyanoferrate(III) (iv)Potassium trioxalatoferrate(III) (v)Potassium tetrachloridopalladate(II) (vi)Diamminechlorido(methylamine)platinum(II) chloride Q)Using IUPAC norms write the formulas for the following: (i) Tetrahydroxozincate(II) (vi) Hexaamminecobalt(III) sulphate (ii) Potassium tetrachloridopalladate(II) (vii) Potassium tri(oxalato)chromate(III) (iii) Diamminedichloridoplatinum(II) (viii) Hexaammineplatinum(IV) (iv) Potassium tetracyanonickelate(II) (ix) Tetrabromidocuprate(II) (v) Pentaamminenitrito-O-cobalt(III) (x) Pentaamminenitrito-N-cobalt(III) Q)Using IUPAC norms write the systematic names of the following: (i) [Co(NH3)6]Cl3 (iv) [Co(NH3)4Cl(NO2)]Cl (vii) [Ni(NH3)6]Cl2 (ii) [Pt(NH3)2Cl(NH2CH3)]Cl (v) [Mn(H2O)6]2+ (viii) [Co(en)3]3+ 3+ 2– (iii) [Ti(H2O)6] (vi) [NiCl4] (ix) [Ni(CO)4] ISOMERISM

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Stereoisomers have same chemical formula and chemical bonds, but have different spatial arrangement.  Structural isomers have different chemical bonds. Geometric Isomerism This type of isomerism arises in heteroleptic complexes due to different possible geometric arrangements of the ligands. Generally found with compounds having coordination numbers 4 and 6  Square planar complex of formula [MX2L2] (X and L are uni-dentate) exhibits geometrical isomerism − cis-isomer and trans-isomer.  Cis isomer - 2 identical ligands are adjacent to each other.  Tras isomer - 2 identical ligands are opposite to each other. o Example − Geometrical isomers of Pt(NH3)2Cl2

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Octahedral complex of formula [MX2L4] also exhibits geometrical isomerism. + o Example − Geometrical isomers of [Co(NH3)4)Cl2]

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Octahedral complex of formula [MX2 (L − L)2] exhibits geometrical isomerism, where L−L is a bidentate ligand. [e.g., NH2 CH2 CH2 NH2(en)] o Example − Geometrical isomers of [CoCl2 (en)2]

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Octahedral complexes of type [Ma3b3] exhibit another type of geometrical isomerism − facial (fac) isomer and meridional (mer) isomer. Example − [Co(NH3)3(NO2)3] o Facial (fac) isomer – 3 same ligands occupy adjacent positions at the corners of an octahedral face. o Meridional (mer) isomer − When 3 identical ligands are around the meridian of the octahedron

1) Why is geometrical isomerism not possible in tetrahedral complexes having two different types of unidentate ligands coordinated with the central metal ion ? A)Tetrahedral complexes do not show geometrical isomerism because the relative positions of the unidentate ligands attached to the central metal atom are the same with respect to each other. Optical Isomerism Optical isomers are mirror images that cannot be superimposed on one another. These are called as enantiomers The molecules or ions that cannot be superimposed on its mirror image are called chiral. The two forms are called dextro (d) and laevo (l) depending upon the direction they rotate the plane of polarised light in a polarimeter. * Optical isomerism is common in octahedral complexes involving didentate ligands 3+  Example − ‘d’ and ‘l’ isomer of [Co(en)3]

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→ Only cis isomer shows optical activity. Trans isomer does not show optical isomerism as they form superimposable mirror image(optically inactive-achiral).

Structural isomerism Structural isomers the same chemical formula but they have different bonds. Types of Structural isomerism (i) Linkage isomerism (ii) Coordination isomerism

(iii) Ionisation isomerism (iv) Solvate isomerism Linkage Isomerism Linkage isomerism arises in a coordination compound containing ambidentate ligand(which can ligate through two different atoms) Eg: [Co(NH3)5(NO2)]Cl2 and [Co(NH3)5(ONO)]Cl2 [Co(H2O)5(SCN)]Cl2 and [Co(H2O)5(NCS)]Cl2 Coordination Isomerism This type of isomerism arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex. Example − Ionisation Isomerism Compounds which give different ions in solution although they have same composition. (or) This type of isomerism arises when the counter ion in the complex salt is itself a potential ligand and can displace a ligand, which can then become the counter ion o

Example −

Solvate Isomerism Compounds with same molecular formula, but differ in no. of solvent molecules inside and outside the coordination sphere. (or) Solvate isomers differ by whether or not a solvent molecule is directly bonded to the metal ion or merely present as free solvent molecules in the crystal lattice. Example − (violet) and (grey green  Known as ‘hydrate isomerism’ when water is involved as a solvent

Q. Draw structures of geometrical isomers of [Fe(NH3)2(CN)4]–

Q. Out of the following two coordination entities which is chiral (optically active)? (a) cis-[CrCl2(ox)2]3– (b) trans-[CrCl2(ox)2]3–

Out of the two, (a) cis – [CrCl2(ox)2] is chiral (optically active). Trans is optically inactive, form superimposable mirror image. Q. Indicate the types of isomerism exhibited and draw the structures for these isomers (i) K[Cr(H2O)2(C2O4)2 (ii) [Co(en)3]Cl3 (iii) [Co(NH3)5(NO2)](NO3)2 (iv) [Pt(NH3)(H2O)Cl2] (i) Both geometrical (cis-, trans-) and optical isomers for cis can exist. (ii) Two optical isomers can exist. (iii) There are 10 possible isomers. (Hint: There are geometrical, ionisation and linkage isomers possible). (iv) Geometrical (cis-, trans-) isomers can exist.

Q. Give evidence that [Co(NH3)5Cl]SO4 and [Co(NH3)5SO4]Cl are ionisation isomers.(or) Give chemical test to distinguish [Co(NH3)5Cl]SO4 and [Co(NH3)5SO4]Cl Reagent [Co(NH3)5Cl]SO4 [Co(NH3)5SO4]Cl No reaction White ppt of AgCl AgNO3 solution White ppt of BaSO4 No reaction BaCl2 solution Valence Bond Theory  The metal atom or ion under the influence of ligands can use its (n−1)d, ns, np or ns, np, nd orbitals for hybridisation, to yield a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral, square planar, and so on.  These hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron pairs for bonding. Coordination number

Type of hybridisation

Distribution of hybrid orbitals in space

4 4 5 6 6

sp3 dsp2 sp3d sp3d2 d2sp3

Tetrahedral Square planar Trigonal bipyramidal Octahedral Octahedral

Octahedral Complexes 2 3 3 2 o The hybridisation involved can be d sp or sp d . Inner-orbital or low-spin or spin-paired complexes: Complexes that use inner d-orbitals in hybridization. 3+ 2 3 o Eg: [Co(NH3)6] . d sp hybridisation,octahedral, diamagnetic, inner orbital complex.

Outer-orbital or high-spin or spin-free complexes: Complexes that use outer d-orbitals in hybridisation; Example, [CoF6]3−.sp3d2 hybridisation,octahedral, paramagnetic, outer orbital complex.

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Tetrahedral Complexes 3 o The hybridisation involved is sp , tetrahedral, paramagnetic 2− o Example: [NiCl4] o The hybridisation scheme is shown in the following diagram.

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Square planar Complexes 2− 2 o Example: [Ni(CN)4] : dsp hybridisation, square planar, paramagnetic.

Magnetic Properties of Coordination Compounds  Complexes with unpaired electron(s) in the orbitals are paramagnetic.  Complexes with no unpaired electron(s) in the orbitals (i.e., all paired electrons) are diamagnetic. Limitations of Valence Bond Theory (i) It involves a number of assumptions. (ii) It does not give quantitative interpretation of magnetic data. (iii) It does not explain the colour exhibited by coordination compounds. (iv) It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds. (v) It does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes. (vi) It does not distinguish between weak and strong ligands. Crystal-Field Theory  An electrostatic model which considers the metal−ligand bond to be ionic, which arises from the electrostatic interaction between the metal ion and the ligand  Ligands are treated as point charges in the case of anions, or dipoles in the case of neutral molecules.  The five d-orbitals in an isolated gaseous metal atom/ion are degenerate (i.e., have the same energy).  Due to the negative fields of the ligands (either anions or the negative ends of dipolar molecules), the degeneracy of the d-orbitals is lifted, resulting in the splitting of the d-orbitals to give 2 sets of orbitals- t2g and eg. Crystal-Field Splitting in Octahedral Coordination Entities Crystal-Field Splitting: Crystal-field splitting is the splitting of the degenerate d orbitals into t2g and eg due to the presence of ligands.

Crystal field splitting energy Δo: Energy separation between t2g and eg orbitals  The splitting of d-orbitals in an octahedral crystal-field is shown in the given figure.

The energy separation is denoted by Δo (the subscript ‘o’ is for octahedral) Spectrochemical series The arrangement of ligands in the order of increasing field strength is called spectrochemical series. Ligands are arranged in a series in the increasing order of the field strength as follows:

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Strong field and weak field ligands Ligands for which Δo (crystal-field splitting) < P (pairing energy), are called weak-field ligands, and form high-spin complexes. d4 → t2g3 eg1 Ligands for which Δo (crystal-field splitting) > P (pairing energy), are called strong-field ligands, and form low-spin complexes. d4 → t2g4 eg0

Q. How does Δo decides the d orbital configuration? 4 3 1  Weak ligands for which Δo < P form high-spin complexes. d → t2g eg 4 4 0  Strong ligands for which Δo > P, form low-spin complexes. d → t2g eg Crystal-Field Splitting in Tetrahedral Coordination Entities  The splitting of d-orbitals in a tetrahedral crystal field is shown in the given figure.

 Δt = (4/9) Δo Colour in Coordination Compounds  The colour of the coordination compounds is attributed to d−d transition of electrons. o For example, the complex [Ti(H2O)6]3+ is violet in colour, which is due to d−d transition

o

The energy of yellow-green region is absorbed by the complex, and then, the electron gets excited from t2g level to the eg level, i.e.,

In the absence of ligand, crystal field splitting does not occur and hence the substance is colourless. For example, removal of water from [Ti(H2O)6]Cl3 on heating renders it colourless. Similarly, anhydrous CuSO4 is white, but CuSO4.5H2O is blue in colour.

Metal Carbonyls The homoleptic carbonyls (compounds containing carbonyl ligands only) are formed by most of the transition metals. Eg Ni(CO)4 , Fe(CO)5 , Cr(CO)6 , [Mn (CO)10 ], [Co2 (CO)8 ]

Bonding in metal carbonyls  In metal carbonyls, the metal−carbon bonds have both σ and π characters.  A σ bond is formed when the carbonyl carbon donates a lone pair of electrons to the vacant d orbital of the metal.  A π bond is formed by the donation of electron pair from the filled metal d- orbital to the vacant antibonding π* orbital (also known as back bonding of the carbonyl group).  Thus, a synergic effect is created due to this metal-to-ligand bonding. This synergic effect strengthens the bond between CO and the metal.

Stability of Coordination Compounds  The stability of a complex in a solution refers to the degree of association between the two species involved in the state of equilibrium.  Stability can be expressed quantitatively in of stability constant or formation constant.

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The greater the value of the stability constant, the greater is the stability of the complex. Free metal ions are usually surrounded by solvent molecules in solution, which will compete with the ligands molecules (L) and be successively replaced by them as follows:

Where, K1, K2, …= Stepwise stability constants

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Alternatively, M +4L → ML4 β4 = [ML4]/[M][L]4 Where, β1, β2, …. = Overall stability constant The stepwise and overall stability constants are therefore related as follows:

or more generally,  The reciprocal of the formation constant is called instability constant or dissociation constant. Q. Calculate the overall complex dissociation equilibrium constant for the Cu(NH3)42+ion, given that β4 for this complex is 2.1 × 1013 . The overall dissociation constant is the reciprocal of overall stability constant i.e. 1/ β4 = 4.7 × 10−14 Applications of Coordination Compounds  Analytical chemistry o In many qualitative and quantitative chemical analyses: Colour reactions are given by metal ions with a number of ligands (especially chelating ligands), which help in detection and estimation. o Hardness of water can be estimated by simple titration with Na2EDTA, which forms stable complexes with Ca2+ and Mg2+ ions.  Metallurgy o Some important extraction processes of metals such as gold and silver make use of complex formation. . Gold, for example, combines with cyanide in the presence of oxygen and water to form the coordination entity [Au(CN)2]− in aqueous solution. Gold can be separated in metallic form from this solution by the addition of zinc . o Similarly, purification of metals can be achieved through formation and subsequent decomposition of their coordination compounds. For example, impure nickel is converted to [Ni(CO)4], which is decomposed to yield pure nickel.  Medicinal chemistry o Chelate therapy is used in medicinal chemistry. For example, EDTA is used in the treatment of lead poisoning. Cis-platin inhibit the growth of tumours. o Excess metal ions in plant and animal system can be removed by chelating ligands by forming complex.  Biological system o Coordination compounds are of great importance in biological systems. For example, chlorophyll is a coordination compound of magnesium, haemoglobin is a coordination compound of iron and vitamin B12 is a coordination compound of cobalt.   

Coordination compounds are used as catalysts. For example, Wilkinson catalyst, [(Ph3P)3RhCl], is used for the hydrogenation of alkenes Articles can be electroplated with silver and gold much more smoothly and evenly from solutions of the complexes, [Ag(CN)2]– and [Au(CN)2]− than from a solution of simple metal ions. In black and white photography, the developed film is fixed by washing with hypo solution which dissolves the undecomposed AgBr to form a complex ion, [Ag(S2O3)2]3− .

Intext Questions

9.5 Explain on the basis of valence bond theory that [Ni(CN) 4]2− ion with square planar structure is diamagnetic and the [NiCl4]2− ion with tetrahedral geometry is paramagnetic. 9.6 [NiCl4]2− is paramagnetic while [Ni(CO)4] is diamagnetic though both are tetrahedral. Why? 9.7 [Fe(H2O)6]3+ is strongly paramagnetic whereas [Fe(CN)6]3− is weakly paramagnetic. Explain. 9.8 Explain [Co(NH3)6]3+ is an inner orbital complex whereas [Ni(NH3)6]2+ is an outer orbital complex. 9.9 Predict the number of unpaired electrons in the square planar [Pt(CN) 4]2−ion. 9.10 The hexaquo manganese(II) ion contains five unpaired electrons, while the hexacyanoion contains only one unpaired electron. Explain using Crystal Field Theory.

Exercises 9.1 Explain the bonding in coordination compounds in of Werner’s postulates. 9.2 FeSO4 solution mixed with (NH4)2SO4 solution in 1:1 molar ratio gives the test of Fe2+ ion but CuSO4 solution mixed with aqueous ammonia in 1:4 molar ratio does not give the test of Cu2+ ion. Explain y?

9.3 Explain with two examples each of the following: coordination entity, ligand, coordination number, coordination polyhedron, homoleptic and heteroleptic. 9.4 What is meant by unidentate, didentate and ambidentate ligands? Give two examples for each. 9.5 Specify the oxidation numbers of the metals in the following coordination entities: (i) [Co(H2O)(CN)(en)2]2+ (ii) [CoBr2(en)2]+ (iii) [PtCl4]2− (iv) K3[Fe(CN)6] (v) [Cr(NH3)3Cl3] 9.6 Using IUPAC norms write the formulas for the following: (i) Tetrahydroxozincate(II) (ii) Potassium tetrachloridopalladate(II) (iii) Diamminedichloridoplatinum(II) (iv) Potassium tetracyanonickelate(II) (v) Pentaamminenitrito-O-cobalt(III) (vi) Hexaamminecobalt(III) sulphate (vii) Potassium tri(oxalato)chromate(III) (viii) Hexaammineplatinum(IV)(ix) Tetrabromidocuprate(II) (x) Pentaamminenitrito-N-cobalt(III) 9.7 Using IUPAC norms write the systematic names of the following: (i) [Co(NH3)6]Cl3 (ii) [Pt(NH3)2Cl(NH2CH3)]Cl (iii) [Ti(H2O)6]3+(iv) [Co(NH3)4Cl(NO2)]Cl (v) [Mn(H2O)6]2+ (vi)[NiCl4]2−(vii) [Ni(NH3)6]Cl2(viii) [Co(en)3]3+(ix) [Ni(CO)4] 9.8 List various types of isomerism possible for coordination compounds, giving an example of each. 9.9 How many geometrical isomers are possible in the following coordination entities? (i) [Cr(C2O4)3]3– (ii) [Co(NH3)3Cl3] 9.10 Draw the structures of optical isomers of: (i) [Cr(C2O4)3]3– (ii) [PtCl2(en)2]2+ (iii) [Cr(NH3)2Cl2(en)]+ 9.11 Draw all the isomers (geometrical and optical) of: (i) [CoCl2(en)2]+ (ii) [Co(NH3)Cl(en)2]2+ (iii) [Co(NH3)2Cl2(en)]+ 9.12 Write geometrical isomers of [Pt(NH3)(Br)(Cl)(py)] and how many of these will exhibit optical isomers? 9.13 Aqueous copper sulphate solution (blue in colour) gives: (i) a green precipitate with aqueous potassium fluoride and(ii) a bright green solution with aqueous potassium chloride. Explain these experimental results. 9.14 What is the coordination entity formed when excess of aq KCN is added to an aq solution ofCuSO4? Why is it that no precipitate of copper sulphide is obtained when H2S(g) is ed through this solution? 9.15 Discuss the nature of bonding in the following coordination entities on the basis of valence bond theory: (i)[Fe(CN)6]4− (ii)[FeF6]3− (iii)[Co(C2O4)3]3− (iv)[CoF6]3− 9.16 Draw figure to show the splitting of d orbitals in an octahedral crystal field. 9.17 What is spectrochemical series? Explain difference between a weak field ligand & a strong field ligand. 9.18 What is crystal field splitting energy? How does the magnitude of Δo decide the actual configuration of d orbitals in a coordination entity? 9.19 [Cr(NH3)6]3+ is paramagnetic while [Ni(CN)4]2− is diamagnetic. Explain why? 9.20 A solution of [Ni(H2O)6]2+ is green but a solution of [Ni(CN)4]2− is colorless. Explain. 9.21 [Fe(CN)6]4− and [Fe(H2O]2+ are of different colours in dilute solutions. why? 9.22 Discuss the nature of bonding in metal carbonyls. 9.23 Give the oxidation state, d orbital occupation and coordination number of the central metal ion in the following complexes: (i) K3[Co(C2O4)3] (ii)cis-[Cr(en)2Cl2]Cl (iii)(NH4)2[CoF4] (iv) [Mn(H2O)6]SO4 9.24 Write down the IUPAC name for each of the following complexes and indicate the oxidation state, electronic configuration and coordination number. Alsogive stereochemistry and magnetic moment of the complex:(i) K[Cr(H2O)2(C2O4)2].3H2O (ii) [Co(NH3)5Cl-]Cl2 (iii) CrCl3(py)3 (iv) Cs[FeCl4] (v) K4[Mn(CN)6] 9.25 What is meant by stability of a coordination compound in solution? State the factors which govern stability of complexes. 9.26 What is meant by the chelate effect? Give an example. 9.27 Discuss briefly giving an example in each case the role of coordination compounds in: (i) biological systems (ii) medicinal chemistry (iii) analytical chemistry (iv) extraction/metallurgy of metals. 9.28 How many ions are produced from the complex [Co(NH3)6]Cl2 in solution?(i) 6 (ii) 4 (iii) 3 (iv) 2 9.29 Amongst the following ions which one has the highest magnetic moment value? (i) [Cr(H2O)6]3+ (ii) [Fe(H2O)6]2+ (iii) [Zn(H2O)6]2+ 9.30 The oxidation number of cobalt in K[Co(CO)4] is (i) +1 (ii) +3 (iii) –1 (iv) –3 9.31 Amongst the following, the most stable complex is (i) [Fe(H2O)6]3+ (ii) [Fe(NH3)6]3+ (iii) [Fe(C2O4)3]3− (iv) [FeCl6]3− 9.32 What will be the correct order for the wavelengths of absorption in the visible region for the following: [Ni(NO2)6]4− , [Ni(NH3)6]2+ , [Ni(H2O)6]2+ ?