Coordination Compounds = Coordination compounds are the compounds, in which the central metal atom is linked to anumber of ions or neutral molecules by coordinate bonds, For example, nickel tetracarbony!, [Ni(CO),], in which CO molecules are linked to the central nickel atom by coordinate bonds by donating lone pairs of electrons, is a coordination compound. If the species formed by linking of a number of ions or molecules by coordinate bonds to the central metal atom (or ion) carries positive or negative charge, itis called a complex ion, e.g., [Fe(CN),]* [Cu(NH,),]** ete. The branch of inorganic chemistry dealing with the study of coordination compounds is known as coordination chemistry. ‘Types of complexes. There are following three types of complexes (i) cationic complex, eg. (Co(NH,)4]°* (ii) anionic complexe.g.[Ag(CN),]~ (ii) neutral complex e.g., [Ni(CO),] Difference between a double salt and a complex. Both doubles salts as well as complexes are formed by the combination of two or more stable compounds in stoichiometric ratio. However they differ in the fact that double salts such as Mohr’s salt, FeSO, (NH,),SO,.6H,O, potash alum, K,S0,.Aly(SO,)3.24H,O, ete. dissociate into simple substance or ions completely when dissolved in water, eg, FeSO,.(NH,),S0,.6H,0 — Fe?* + 2NH} + 2807 + 6H,O (On the other hand, in complexes, the complexes ion almost does not dissociate. For example, the complex, K,[Fe(CN)g]. The solution does not give the tests of Fe2* and CN” ions because the complex ion, [Fe(CN)gI does not dissociate into Fe?" and CN™ ions. ‘Some Important terms used in coordination compounds Ligands and central metal atom/ion. The donor atoms, molecules or anions which donate a pair of electrons to the metal atom or ion and forma coordinate bond with it are called ligands. The metal atom or ion to which these ligands are attached is called central metal atom or ion. The ligands are Lewis bases whereas central metal atom or ion is Lewis acid, The ligand may contain one or more than one donor atom. If only one donor atom is present in its molecule which can coordinate, then itis called as unidentate. Thisis also referred to as monodentate. A few examples are : NH, H,O and CN Denticity and Chelation. The ligand may contain two donor atoms positioned in such a way that a five or a six membered ring is formed with the metal ion, then itis called bidentate chelating ligand and the ring is called chelate ring, the resulting complex is called a metal chelate and this property is called chelation. The well known examples of the didentate ligands are ethane-1, 2-diamine (en), oxalate ion (0x). ‘The number of coordinating or ligating groups present in a ligand is called the denticity of that ligand. Ambidentate ligands. Unidentate ligands containing more than one coordinating atoms are called ambidentate ligands, For example NO} can coordinate through either nitrogen or oxygen. Similarly, CN” can coordinate through CN” are either carbon or nitrogen and S all ambidentate ligands. can coordinate through either sulphur or nitrogen. Hence, NO3, CN Coordination number. The total number of unidentate ligands (plus double the number of didentate ligands if any) attached to the central metal ion through coordinate bonds is called the coordination number of the metal ion. For example, in the complex ions [Ag(CN),] ~, [Cu(NH,),]?" and [Cr(H,0),}°", the coordination numbers of Ag, Cu and Crare2, 4 and 6 respectively. Similarly, in the complex ion, [Fe(C,O,);]°~, the coordination number of Fe is 6 because C,07° is a didentate ligand. Chemistry by : Khaleel Sir Mob. No. 9219710253Coordination Compounds Jf Coordination sphere or coordination entity & counter ions. The central atom and the ligands which are directly attached to it are enclosed in square brackets and are collectively termed as the coordination sphere. The ionizable groups are written outside the brackets are called counter ions. For example, in the coordination compound, [Cu(NH;),] SO, the complex ion, (Cu(NH,)4]"*, in which Cu?* is the central metal ion and four NH; molecules are ligands, forms the coordination sphere and $03” ions are the counter ions. Coordination polyhedron. The spatial arrangement of the ligand atoms which are directly attached to the central atomy ion is called coordination polyhedron around the central atom/ion Oxidation number or oxidation state. (self) Homoleptic and Heteroleptic complexes : Complexes in which the metal atom or ion is linked to only one type of ligands are called homoleptic complexes, e.g., [Co(NH,)g]°*. The complexes in which the metal atom or ion is linked to more than one kind of ligands are called heteroleptic complexes, ¢.g., (Co(NH,),Cl]"- (Formula Writing and Nomenclature of Coordination Compounds) Naming of Ligand: (i) Negative ligands (organic or inorganic) end in-o, e.g., CN” (cyano), Cl” (chloro), Br (bromo), F~ (fluoro), NO} (nitrito-N), OH” (hydroxo), 0; (OXO.), H” Hydrido) SO} (sulphato), C,07° (Oxalato), NH (amido), NH?" (imido), ONO™ (nitrito-O), NO3 (nitrato), SCN“ (thiocyanato), NCS~ (isothiocyanato), CH,(NH,)COO™ (glycinato). According to 2004 IUPAC recommendations as Cl”, Br”, I”, F~ end in -ide, they should be named as chlorido, bromido ete. in place of chloro, bromo, etc. similarly, CN~ should be named as cyanide though the name cyano is still used, Hydrogen as ligand is always taken as anionic (hydride) and is named as hydrido. (ii) Neutral ligands have no special ending, e.g.,NH, (ammine), H,O (aqua), CO (carbonyl), NO (nitrosyl) (iii) Positive ligands (which are very few) and in- ium, e.g, NO" (nitrosonium), NO} (nitronium). (iv) Organic ligands. eg, CH, (methyl), CH; (ethyl), CgH5 (phenyl), CsHs (cyclopentadienyl). CHsNH) (methylamine), P(C,Hs), (triphenylphosphine), CsHsN (Pyridine or py), NH,CH,CHNH, (ethylenediamine i.e, ethane-1, 2- diamine or en). Q. Name the following coordination compounds using LU.P.A.C. system () (Cr (NH) Gi) [Mn(HO),]°* Gil) [Fe(CN) It (iv) INi(NH,)ICl, () [Co(CN) > (vi) Cap [Fe (CN)] (vii) [Co (NH) Cl (vill) [Cr (H,0),Cl,)] NO; (&) [Co(NH,),C1 (NO,)] NO, (%) KIP(NH)ChI. Q. Write down the formulae of the following coordination compounds : (@) hexaaquairon (Il) sulphate Gi) potassium tetracyanonickelate (Ill) (iii) chloronitrodiammineplatinum (II) (iv) potassium hexacyanoferrate (III) (v) chlorodiammineplatinum (Il) ion (vi) dichlorotetraamminecobalt (II) ion. (vii)_ potassium pentacyanonitrosyleobaltate (III) ISOMERISM ‘Two or more substances having the same molecular formula but different structural or spatial arrangements are called isomers and the phenomenon is called isomerism. In inorganic compounds, coordination compounds often show various types of isomerism. (A) Structural isomerism and (B) Stereo isomerism or space isomerism. Structural Isomerism. This type of isomerism arises due to the difference in structures of coordination compounds. Itmay be further subdivided into different types as follows (2) onisation isomerism. Compounds which give different ions in solution although they have same composition are called ionization isomers. Examples : [Co(NH)s NO31SO, and [Co(NHs)5(S0,)] NOs Chemistry by : Khaleel Sir Mob. No. 9219710253V ‘coordination Compounds 2) Solvate or Hydrate isomerism. Compounds which have the same composition but differ in the number of solvent molecules present.as ligands in the crystal lattice are called solvate isomers. If water is the solvent, these are called hydrate isomers.e.g., [Co(NH), (H,0)CI] Cl, and [Co(NH), Cl,] CLH,O (3) Linkage isomerism. Isomerism of this type occurs in compounds containing ambidentate ligands. Example . [Cr(H,0),(SCN)P and (Cr(H,0); (NCS)P* (4) Coordination isomers. This type of isomers is possible when both positive and negative ions of a salt are complex ions and the two isomers differ in the distribution of ligands in the cation (positive ion) and the anion {negative ion). Examples. [Co(NH3)¢] [Cr(CN)¢] and [Cr(NH3)¢] [Co(CN)¢} This type of isomerism is caused by interchange of ligands between the two complex ions. Stereo isomerism or space isomerism. Space isomerism arises on account of the different positions and arrangements of ligand in space around the metal ion. Itis of two types: (1) Geometrical isomerism (2) Optical isomerism Geometrical or Cis-trans isomerism. This type of isomerism occurs in heteroleptic complexes due to different possible geometric arrangements of the ligands. When two identical groups (ligands) occupy adjacent positions, the isomer is called cis and when arranged opposite to one another, the isomer is called trans. This isomerism is not possible for complexes with coordination number 2 and 3 and tetrahedral complexes with coordination number 4 because in this case, all the four positions are equivalent. However, cis-trans isomerism is quite common in square planer and octahedral complexes. (a) Square planar complexes. Representing the central metal atom by M and the unidentate ligands by A, B, C, D, etc,, the square planar complexes may be classified into the following different types. () MA,B, type. eg, [Pt (NH) Ch] Gi) MA,BC type. eg, [Pt (NH) Cl (py)] ii). M(AB), type. (in which AB represents an unsymmetrical didentate ligand) e.g, [Pt (gly)3] (iv) MABCD type. Square planar complexes of the type MABCD form three isomers. W Note: sessssssssesiisiiieeee Mes osssssat si llvccscscsssensissueesssssseensisiussaniiiieeseeeesees + Square planar complexes of the type MA,, MAB and MAB, do not show geometrical isomerism because in any of these cases, the possible spatial arrangements are equivalent. (b) Octahedral complexes. These may be classified into the following different types (i) MAB, or MA,B, type. [Cr(NH,),Cl* Gi) [M(AA),B,] or [M(AA),BC] type. [CoCl, (en),]+ Git) MA,B, type. [RhCI, (py)ah 1 Note % Octahedral complexes of the type [MA,] or [MA;B] would not show geometrical isomerism for obvious reasons. Octahedral complex of the type [MABCDEF], forms 15 different geometrical isomers. Optical Isomerism. Optical isomers rotate the plane of polarised light in opposite directions-the two isomers are structurally non-superimposable mirror image of each other. The molecules or ions which are non-superimposable mirror image of each other are called chiral and this property is called chirality. The optical isomers are called dextro and laevo (d and I) depending upon the direction in which plane of the polarized light is rotated (ie., towards right or left). The d and T isomers of a compound are called enantiomers or enantiomorphs. Optical isomerism is common in octahedral complexes with coordination number six involving 1, 2 or 3 symmetrical didentate ligands. () [M(AA)3] type ie, containing three symmetrical didentate ligands , ¢.g,, [Co(en)3}* [Cr(ox),]* (ii) (M(AA)3B.] oF (M(AA)2BC] type, ie, containing two symmetrical didentate ligands, e.g,, [CoCl (en)3|*. (iii) (M(AA)B,C] type, ie, containing one symmetrical didentate ligand e.g., (Co(en)(NH),Cl]. (iv) Octahedral complexes containing hexadentate ligands, eg.[Co(edta)|-, i.e, ethylene diaminetetraacetato cobaltate (III) ion, also show optical isomerism. Chemistry by : Khaleel Sir Mob. No. 9219710253al coordination compounds I (ii) In case of complexes with coordination number 4, square planar complexes do not show optical isomerism. because they contain a plane of symmetry but tetrahedral complexes containing unsymmetrical bidentate ligands, eg, [Ni (CH,NH,COO))], ic, bis (glycinato) nickel (11), shows optical isomerism. (BONDING IN COORDINATION COMPOUNDS) Werner's theory of Coordination Compounds: Alfred Werner, a Swiss chemist was the first to study the bonding, in coordination compounds, in 1892. He proposed a theory about the nature of bonding in the coordination compounds (complexes). This theory is known as Werner's theory of coordination compounds. ‘The main postulates of this theory are as follows 1. Metals posses two types of valencies called (a) primary or principal or ionizable valency (b) secondary or non-ionizable valency (@) Primary valencies are those which a metal exhibits in the formation of its simple salts. Thus, in the formation of PtCly CoCly, CuSO, and AgCl, the primary valencies of Pt, Co, Cu and Ag are 4, 3, 2 and 1 respectively. These days primary valency is referred to as oxidation state. (b) Secondary valencies are those which a metal atom or cation exercise towards neutral molecules or negative groups (ligands) in the formation of its complex ions. Every metal cation has a fixed number of secondary valencies, The secondary valency is also termed as the coordination number (C.N) of the metal cation. In [Pt (NH,),) Cly (Co(NHs),Cl [Cu(NHs), SO, and [Ag(NH,),] Cl, the secondary valencies are 6 for Pt’, 6 for Co™, 4 for Cu and 2 for Agt cations respectively. 2. Every metal atom has a tendency to satisfy both its primary and secondary valencies. Primary valencies are satisfied by negative ions whereas secondary valencies are satisfied by negative ions or neutral molecules (ligands). 3. The ligands satisfying secondary valencies are always directed towards fixed positions in space thereby giving a definite geometry to the complex but the primary valencies are non-directional. Limitations of Wemer’s theory. Though Werner's theory was able to explain a number of properties of the coordination compounds, it could not answer the following basic questions (i) Why only certain elements form coordination compounds and not others? (ii) Why the coordination sphere/entity has a definite geometry? (iii) Why do these compounds possess definite magnetic and optical properties? Valence Bond Theory. This theory was extended to the coordination compounds by Pauling in 1931. The following are the main postulates of this theory. (i) In this approach, the basic assumption made is that the metal-ligand bond arises by the donation of pairs of electrons by ligands to the metal atomyion. (ii) In order to accommodate these electrons, the metal ion must possess requisite number of vacant orbitals of equal energy. These orbitals are obtained by hybridization of the orbitals of the metal atom to give a set of hybrid orbitals of equal energy. (iii) Sometimes, the unpaired (7-1) d electrons pair up as fully as possible prior to hybridization thus making some (r-1) d orbitals vacant. The central metal atom then makes available the number of empty orbitals equal to its coordination number for the formation of coordinate bonds with suitable ligand orbitals. (iv) With the approach of the ligands, metal-ligand bonds are then formed by the overlap of these orbitals with those of the ligands, ie, by donation of electron pairs by the ligands to the empty hybridized orbitals. Consequently, these bonds are of equal strength and directional in nature. (v) Octahedral, square planar and tetrahedral complexes are formed as a result of @2sp? (or sp), dsp? and sp? hybridization respectively of the central metal atom or the ion. (A) Octahedral complexes (Coordination Number = 6): Chemistry by : Khaleel Sir Mob. No. 9219710253V ‘coordination Compounds [Cr (NH,),P*: The oxidation state of chromium in [Cr (NH,)}* ion is +3. The electronic configuration of Cris [Ar] 3d°4s!, Hence, We have Cratom (Z= 24) In ground state Cx ion Psp hybridized orbitals of Cr?" ion Formation of [Cr(NH,),)* Cr" ion provides six empty orbitals to accommodate six pairs of electrons from six molecules of ammonia (shown as crosses). The resulting complex [Cr(NH,),]°* involves dsp hybridization and it thus octahedral. The presence of these unpaired electrons in the complex explains its paramagnetic character. Inner and outer orbital complexes. When the complex formed involves the inner (7-1)d-orbitals for hybridization (sp), the complexis called inner orbital complex or hyperligated complex. In this case, the electrons of the metal are made to pair up, so the complex will be either diamagnetic or will have lesser number of unpaired electrons. This type of complex is also known as low spin complex. e.¢,, (Co(NH,),)** When the complex formed involves the use of outer nd-orbitals for hybridization, (sp'd2), the complex is called outer orbital complex or hypoligated complex. The complex will have large number of unpaired electrons as the configuration of the metal remains unchanged. This type of complex is also called high spin complex. Example-{MnCl,]* (8) Tetrahedral complexes. (Coordination number = 4) (© Square planar complexes. (Coordination Number = 4) Drawbacks of Valence Bond Theory. (@ It cannot explain why some complexes of a metal ion in a particular oxidation state are low spin, i.e, inner orbital complexes while some other complexes of the same metal ion in the same oxidation state are high spin, i.e, outer orbital complexes. For example, (Co(NH,),I* is a low spin, while [CoF,]-is a high spin, It could not give any satisfactory explanation for the colour of the complex, It does not give an exact explanation of thermodynamic or kinetic stabilities of coordination compounds. (iv) Tt does not distinguish between weak and strong ligands. Crystal Field Theory: This theory gives a much more satisfactory explanation for the bonding and the properties of the complexes. The main points of this theory are as follows (i) Whereas valence bond theory considers the metal-ligand bond to be covalent, crystal ficld theory considers the bond to be ionic arising purely from electrostatic interactions between the metal ion and the ligands. As these interactions are similar to those between the ions in a crystal, that is why it has been named as crystal field theory (CFT). Gi) It treats each ligand as a point of negative charge. The arrangement of the ligands around the central metal ion is, such that the repulsions between these negative points are minimum Gil) According to this theory, in a fee transition-metal ion, all the five d-orbitals have equal energies, ic, they are degenerate. This degeneracy is maintained if the negative charges present around the central metal atom/ion form a spherically symmetrical field, However, if the negative field is due to ligands which may be anions or polar molecules such as H,O, NHy, etc, with their negative ends towards the central metal atom/ion, the field no longer remains symmetrical. Consequently, the degeneracy is split. The pattern of splitting depends upon the nature of the crystal field exerting its influence on the central metal atom/ion. (A) Crystal Field Theory for Octahedral Complexes. As the ligands approach the metal ion, there is repulsion between the ligands and the d-orbitals, thereby raising their energy relative to that of the free ion. If the d-orbitals Chemistry by : Khaleel Sir Mob. No. 9219710253Gl coordination compounds I were present in a spherically symmetrical crystal field, they could be represented as shown in the state Il in the fig, ‘The mean value of energy of these perturbed d-orbitals is taken as zero. ‘This is sometimes called as the Bari centre. on However, asd ~dj andd? orbitals have lobes along the axes, hence they point towards the ligands whereas the lobes of dgy dg. and dy. orbitals lie between the axes, hence they lie between the ligands. As a aga, result, the repulsions between the 7 i == ligands and the di -dy and dz dowd: dy dd, Average energy Splitting of d orbitals orbitals are greater than the ay hind by gal pan erence repulsions between the ligands and Free metal ion spherical crystal field. crystal field the remaining three d-orbitals. ey dy, and d,, orbitals are lower than those of d? ~d? and d? orbitals, The former three orbitals of lower energy are called fag orbitals (read as t-two-g) which are three-fold degenerate and the latter Consequently, the energies of two of higher energy are called eg orbitals (read as e-g) which are two-fold degenerate. The splitting of the d-orbitals, Into two sets of orbitals in an octahedral complex may thus be represented as shown in Fig ‘The difference of energy between the two sets of d-orbitals is called crystal field splitting energy or crystal field stabilization energy (CFSE). It is usually represented by the symbol Ag ‘The magnitude of Ay depends upon the nature of the ligand. Some ligands produce strong fields and hence the splitting is large whereas some others produce weak fields and hence they result in small splitting of d-orbitals. The experimentally observed values of some of the ligands in increasing order are given below. I< Br <$?< SCN-< Cl <0,” P, upto d® ions, pairing will occur in the thy orbitals and ¢, orbitals will remain vacant, Thus, in such cases, we get low spin complexes. Ligands for Which Ag >P are known is strong field ligands. (8) Grystal Field Theory for Tetrahedral Complexes. ; Average eneey ofthe Spliting of erbinis feeion doris inspheral inte erysal The splitting pattern for tetrahedral complexes is just the ‘ysl eld me reverse of the splitting pattern of the octahedral complexes, Spilitting of d-orbitals in tetrahedral complex i.e, the three-fold degenerate set, f5,, has higher energy than the two-fold degenerate set, eg. Moreover the splitting is much smaller than that in case of octahedral complexes. The difference of energy, represented by A = Fao This energy is so small that itis unable to force the electrons to pair up. Hence, tetrahedral complexes have high spin configuration. The splitting pattern of tetrahedral complexes may be represented as follows : Chemistry by : Khaleel Sir Mob. No. 9219710253V ‘coordination Compounds Limitation of Crystal Field Theory. The crystal field theory was successful in explaining the colour, magnetic properties, the effects of weak and strong field ligands etc. in the coordination compounds. However, it has the following limitations () As ligands are considered as point charges, the anionic ligands should exert greater splitting effect. However, actually the anionic ligands are present at the low end of the spectrochemical series. (ii) Tetreats the metal-ligand bond as purely ionic and does not take into account the covalent character of the bond. Stability of Coordination Compounds in Solution The stability of a complex in solution can be expressed in terms of equilibrium constant of the dissociation equilibrium. This constant is called the instability constant or dissociation constant. For example, for the complex, [Cu(NH,),}*" in solution, [Cu(NH,),)* © Cu? + 4NHy 2 n [Cu(NH3)4 I Ifinstead of dissociation, we talk of the formation of a complex ion in the solution, the reaction will be reverse of the above reaction, Cu + ANH, © [Cu(NH,),P* ‘The equilibrium constant for this reaction is, therefore, called the stability constant. Representing if by 6, [CuNH), Stability constant = aN) [Cu ]INH3] ‘Thus, stability constant is the reciprocal of the instability constant (f+ 1/K,). The numerical value of the stability constant is a measure of stability of the complex in solution, Greater the magnitude of the stability constant, more stable is the complex. For example, the formation of [Cu(NH;), ’* takes place in four steps as follows : Cu?" + NH ICHNHS IP") Ki = CHING] ae cunt yftae nt CuNH,),P*, Ky =_lCuSHa PY [Cu(NHs )F Is (Cu(NHs)2 "Ka <7 ny INF] [Cu(NH}5 P* [Cu(NH; ), ?*+ NH [Cu(NH3),*, Ks = [(Cu(NH3 ), * [NH] ___[Cu(NHy),P* {Cu(NHg)3}°* [NH3] [Cu(NHs)s P+ NH [Cu(NHs),P*) Ka where K,, Kz,K, and K, are called stepwise stability constants. The overall stability constant will be B=K, XK, xK3 xKy Factors affecting the stability of a complex ion. () Charge on the central metal ion: Greater the charge on the central metal ion. greater is the stability of the complex (ii) Nature of the metal ion: Metal ions are grouped into two classes ‘Class a’ acceptors are the metals belonging to groups 1 and 2, earlier members of the transitions series, i, groups 3 to 6 and inner transition elements. They form stable complexes when the donor atoms of the ligands are N, Oor F. ‘Class b’ acceptors are the transition metals mostly after group 6 which have relatively filled d-orbitals (e.g., Rh, Pd, Ag, Au, Hg, etc.) They form stable complexes when the donor atoms of the ligands are heavier members of N, O and F family. (iii) Basic nature of the ligands : Greater the basic strength of the ligand, greater is the stability of the complex. Chemistry by : Khaleel Sir Mob. No. 9219710253Coordination Compounds Jf (iv) Presence of chelate rings : Formation of chelate rings increases the stability of the complex.The stabilization due to the chelation is called chelate effect. It is found to be maximum for the 5- and 6- membered rings. Organometallics Compounds: Organometallic compounds are the compounds which contain at least one ‘metal-carbon bond. For example, metal carbonyl like [Ni (CO),] Classification of organometallic compounds. 1, 6 bonded organometallic compounds. The majority of organometallic compounds of the main group (s and p-block) elements are covalent compounds with M-C 6 bonds. Such type of organometallics are called s-bonded organometallic compounds. Examples-Grignard reagents, tetramethyl silane, [(CH,),Si) , Diethyl Zinc, [(CHyCHy)3 Zn]. ete 2, bonded organometallic compounds. Transition metals form organometallic compounds of this type because they have empty d-orbitals. examples - Ferrocene , Fe(n™-C,lH,) etc. Metal Carbonyls: Metal carbonyls are the organometallic compounds in which carbon monoxide (CO) acts as the ligand.A few stable metal carbonyls are given below [Ni(CO)g], [Fe(CO)s}, [Cr(CO)«L [V(CO)g}, [Mo(CO)g] Structures of metal carbonyls- (Draw) a co on] JM >Fe—co oc~ | “co oc | oo co 9 Nuco), Fe(CO), fo 0 rs Tgmapyrananl — 06% —800 oc” \¢/ Sco ca_ 7 co co [0 ° Se oe co tentcons co~| co bo & to C0), taeda inn fCOn4 Bonding in Metal Carbonyls : ‘The metal-carbon bonds in metal carbonyls have both o and x character. The first overlap takes place between the filled bonding 13, orbitals of the carbon monoxide with an empty metal 4-orbital resulting in a s-bond between the metal and carbon atom of carbon monoxide. Here, donation of lone pair of electrons on carbon into a vacant d-orbitals of the metal takes place. The second overlap takes place between the filled metal d-orbitals with an empty antibonding 12p"-orbitals of the carbon monoxide resulting in additional x bond between the metal and same carbon monoxide molecule. Here, donation of electrons from a filled metal d-orbitals into a vacant antibonding 1* -orbitals of CO occurs (back bonding). ‘The effect of 6 bond formation strengthens the x bond and vice-versa. This is called synergic effect (ie,, working together towards the same goal). Thus, as a result of synergic effect, the bond between CO and metal is strengthened Importance and applications of Coordination Compounds/Complexes- (1) In Analytical Chemistry- ‘Synergic bonding (a) In qualitative analysis. () The presence of Ni2* ion is detected by adding dimethyl glyoxime in presence of NH,OH to the salt solution when a brilliant red precipitate is formed due to the formation of a complex of Ni?* with dimethyl glyoxime. Chemistry by : Khaleel Sir Mob. No. 9219710253NATTA o| (ii) The presence of Co* is tested by adding ammonium thiocyanate solution when a blue colour is obtained due to formation of a complex, (b) In quantitative analysis (Estimations). (®) Gravimetric Analysis. The amount of metal present in a given sample can be estimated by converting a known amount of the sample into an insoluble complex which can be filtered, dried and weighed. Gi) Volumetric Analysis. A number of metal ions react completely with polydentate ligands at an appropriate PH to form complexes. Hence, the solutions of metal ions can be titrated against. the solutions of the polydentate ligands in the presence of a suitable buffer and the end point can be detected by using a suitable indicator. The most common polydentate ligand used is ethylene diammine tetra acetic acid (EDTA). Q) In metallurgy (Extraction of metals). The noble metals like silver and gold are extracted from their ores through the formation of cyanide complexes, [Ag(CN),]- and [Au(CN),I- respectively. (3) In purification of metals. Some metals are purified by formation of their metal carbonyls followed by their decomposition, e.., impure nickel is converted into nickel tetracarbonyl. Ni(CO), which on decomposition gives pure nickel. (4) In Biological Systems: (a) Chlorophyll, the green plant pigment that acts as photosensitizer in the photosynthesis in plants is a coordination compound of magnesium, (b) Haemoglobin, a red pigment of blood which acts as an oxygen cartier is a coordination compound of iron. (©) Vitamin B12 which is chemically cyanocobalamine and is the anti-pernicious anaemia factor is a coordination compound of cobalt. (5) In Industry: (a) As catalysts-include rhodium complex, {(PhyP);RhCI], a Wilkinson catalyst, is used for the hydrogenation of alkenes. (b) Inclectroplating. Articles can be electroplated with silver or gold much more smoothly and evenly from solutions of the complexes, [Ag(CN),]~ and [Au(CN),] ~, than from the solutions of simple metal ions. (©) In photography. The developed film is fixed by washing with hypo solution which dissolves the undecomposed AgBr to form a complex, (6) In medical field- (i) EDTA is quite often used for treatment of lead poisoning. (ii) The platinum complex, cis [Pt(NH,),Cl,], known as cisplatin has been found to be useful in the treatment of cancer (tumours). QI. Givea chemical test to distinguish between [Co(NH,);BrISO, and [Co(NH,),SO,]Br. Name the type of isomerism exhibited by these compounds. Ans: The first complex will give SO} ions in the solutions which give white precipitate with BaCl, solutions, ‘The second complex will give Br ions in the solution which give yellow precipitate with AgNO3 solution ‘They exhibit ionization isomerism, Q2. Which isomer of [CoCl,(en),]* does not show optical isomerism ? Ans: Trans isomer does not show optical isomerism. Q3. How many isomers are possible for the neutral complex, [Co(NH,),Cl3] ? Ans: Only two isomers are possible, viz cis and trans (or fac and mer) Q4. A coordination compound has the formula, CoCl,.4NH3, It does not liberate ammonia but precipitates chloride ions as silver chloride. Give the IUPAC name of the complex and write its structural formula, Ans: Remembering that coordination number of Co is 6, the formula of the complex will be [Co(NH,),Cly]CL ‘The name will be tetraamminedichloridocobalt (Ill) chloride. Q5. Arrange the following complexes in order of increasing electrical conductivity : [Co(NH,), Cl], (Co(NH,), Cll Cl, [Co(NH,),ICly, [Co(NH,),C1,ICL Chemistry by : Khaleel Sir Mob. No. 9219710253Ans : Q6. Ans: Q7. Ans: Q8. Ans : Qo Ans: Q10. Qu. Ans: Qn Ans: Q13. Ans: Qu. Ans: Coordination Compounds Jf [Co(NH),Cls]< [Co(NH,),ChICI < [Co(NH,)sCIC1, < [Co(NH,)sICly. This is because number of ions produced from these complexes are 0, 2, 3 and 4 respectively. [Co(CN),]* and [CoF,]*" both are octahedral complexes. Then what is the difference between the two? [Co(CN),]*- is an inner orbital or low spin complex involving the hybridisation d’ sp* while [CoF,]*~ isan. outer orbital or high spin complex involving sp*d hybridisation. ‘The molar conductivity of the complex CoCl, 44NH,.2H,0 is found to be same as that of 3:1 electrolyte. What is the structural formula of the complex ? The complex should be of the type A;B or AB,. As coordination number of Co is 6, hence its structure formula will be [Co(NH,), (H;0),] Cl, Why only transition metals are known to form complexes ? Transition metals/ions have empty d-orbitals into which the electron-pairs can be donated by ligands containing electrons, ie, electrons in their molecular orbitals, e.g., CH, = CH,, CsHs, C/Hy, ete ITi(H,O),}* is coloured while [Sc(F,0),1°* is colourless. Why ? In [Ti(H,0),)°, Ti ion has one electron in the d-subshell (lower energy th, d-orbitals) which can absorb light in the visible region resulting into d-d orbitals (jump into higher energy eg orbitals). Asa result, green and yellow portions of light are absorded. The complex has the complementary colour, viz, purple. In [Sc(F,O),P", Se has no d-electron. Hence, no light isa absorded for d-d transition. Hence, itis colourless On the basis of the following observations made with aqueous solutions, assign secondary valencies to metals in the following compounds: Formula Moles of AgCI ppt per mole of the compound with excess of AgNO; (@ PAC, 4NH, 2 (i) NiCl,, 6H, 2 (iii) PtCly. 2HC1 0 (iv) CoCl;. 4NH, 1 (v) PtCh, 2NH 0 The spin only magnetic moment of [MnBr,]?~ is 5.9 BM. Predict the geometry of the complex ion. As the coordination number of Mn2*in the given complex is 4, the geometry can be either tetrahedral (sp? hybridization) or square planar (dsp? hybridization), In [MnBr,)?>, Mn is in 42 state, ie,, its configuration is [Ar] 3d. u ~ 5.9 BM means that the number of “unpaired electrons in the complex = 5 (because pt = V7i(n +2) BM). Thus, all the five d-electrons remain unpaired. Hence, the hybridization will be sp? and cannot be dsp2. Therefore, geometry is tetrahedral and cannot be square planar. INiCI,]?- is paramagnetic while [Ni(CO),] is diamagnetic though both are tetrahedral. Why? In [NiCl,]*~, Ni is +2 oxidation state with the configuration 3d%4s®. Cl- is weak ligand. It cannot pair up the electrons in 3d orbitals. Hence, it is paramagnetic. In [Ni(CO),], Ni is in zero oxidation state with the configuration 34432. In the presence of CO ligand, the 4s electrons shift to 3d to pair up 3d electrons. Thus, there is no unpaired electron present. Hence, it is diamagnetic. [Fe(F,0),)* is strongly paramagnetic whereas [Fe(CN),]> is weakly paramagnetic. Explain. In both the complexes, Fe is in +3 oxidation state with the configuration 3d8, CN” is a strong ligand. In its presence, 3d electrons pair up leaving only one unpaired electron. The hybridisation is d2sp" forming inner orbital complex. H,0 is a weak ligand. In its presence, 3d electrons do not pair up. The hybridisation is sp'd? forming an outer orbital complex containing five unpaired electrons. Hence, it is strongly paramagnetic. Explain [Co(NH,),}* is an inner orbital complex whereas [Ni(NH,),]*" is an outer orbital complex. In [Co(NH,),]°*, Co is in +3 state with the configuration 34°. In the presence of NHy, 3d electrons pair up leaving two d-orbitals empty, Hence, the hybridisation is dsp? forming an inner orbital complex. In Chemistry by : Khaleel Sir Mob. No. 9219710253V ‘coordination Compounds Q15. Ans: QW. ‘Ans: Qi. Ans: @ (ii) Q18. Ans: Q19. Ans: [Ni(NH,),)?*, Ni is in +2 state with the configuration 348. In presence of NHy, the 3d electrons do not pair up. The hybridization involved is sp3¢2 forming an outer orbital complex. Calculate the overall complex dissociation equilibrium constant for the Cu(NH,),?* ion, given that for this complex is 2.1 x10", 1 Bo 2axi0" FeSO, solution mixed with (NH,),5O, solution in 1; 1 molar ratio gives the test of Fe but CuSO, solution mixed with aqueous ammonia in 1: 4 molar ratio does not give the test of Cu?* ion. Explain Why? FeSO, solution mixed with (NH,),SO, solution in 1 :1 molar ratio forms a double salt, FeSO4. (NH,),SO, 6H1,0 (Mohr salt) which ionizes in the solution to give Fe2+ ions. Hence, it gives the tests of Fe2* ions. ‘CuSO, solution mixed with aqueous ammonia in 1:4 molar ratio forms a complex salt, with the formula [Cu(NH,),] SO,, The complex ion, [Cu(NH,),)>* does not ionize to give Cu®* ions. Hence, it does not give the tests of Cu? ion. 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. Aqueous CuSO, solution exists as [Cu(H,0),]SO, which has blue colour due to [Cu(H,O),]?* ions. When KF is added, the weak HO ligands are replaced by F- ligands forming [CuF]?- ions which is a green precipitate [Cu(H,0),2* + 4F —> [CuF,P> + 40,0 ‘Tetrafluorocuprate (II) (Green ppt) When KCl is added, Cl- ligands replace the weak H,O ligands forming [CuCl,}? ion which has bright green colour, [CuH,O),2" + 4CF —> [CuCl,P- + 41,0 Tetrachlorocuprate (II) (Green solution) What is the coordination entity formed when excess of aqueous KCN is added to an aqueous solution of copper sulphate ? Why is that no precipitate of copper sulphide is obtained when H,S (g) is passes through this solution ? First cupric cyanide is formed which decomposes to give cuprous cyanide and cyanogen gas. Cuprous cyanide dissolves in excess of potassium cyanide to form the complex, K; [Cu(CN)4] CuSO,+2KCN —> Cu(CN), + KO] «2 2Cu(CN)) —> Cu,(CN))+ (CN), Cuy(CN)+6KCN —+ 2K; [Cu(CN),] 2CuSO, +10 KCN —> 2K, [Cu(CN)4] + 2K,S0, + (CN), Thus, coordination entity formed = [Cu(CN),]>- As CNC isa strong ligand, the complex ion is highly stable and does not dissociation/ionize to give Cu? ions. Hence, no precipitate with HS is formed. What is crystal field splitting energy? How does the magnitude of A, decide the actual configuration of d-orbitals in an coordination entity ? When ligands approach a transition metal ion, the d-orbitals split into two sets. One with lower energy and the other with higher energy. The difference of energy between the two sets of orbitals is called crystal field splitting energy (Ao for octahedral field). IfAg

P, the 4th electron pairs up in one of the #2 orbitals giving the configuration 3,¢° thereby forming low spin complexes. Such ligands for which Ay > P are called strong field ligands. ‘What will be the correct order for the wavelengths of absorption in the visible region of the following : INi(NO,)gI*, INi(NH,),}*, [Ni(H,0),P* ‘As metal ion is fixed, the increasing field strengths (CFSE values) of the ligands from the spectrochemical series are in the order : HO As E=hefh, the wavelengths absorbed will be in the opposite order. 2 Why complexes are preferred in the electrolytic bath for electroplating ? ‘They dissociate slowly and hence give a smooth and even deposit. ‘Ametal ion Mn‘ having d valence electronic configuration combines with three didentate ligands to form a complex compound. Assuming 4 > P (j) draw the diagram showing d-orbitals splitting during this complex formation. (ii) write the electronic configuration of the valence electrons of the metal Mn* ion in terms of , and ey (Gil) What type of hybridization will Mn* ion have? (iv) Name the type of isomerism exhibited by this complex. () Mo > P, pairing will occur in the #2, orbitals and ¢, orbitals will remain vacant. (ii) Hyg (Gi) «Psp? (as there are three didentate ligands to combine) (iv) [M(AA),] type complexes show optical isomerism (a) Give the electronic configuration of the d-orbitals of Ti in (Ti(H,0),}** ion in an octahedral crystal field. (b) Why is this complex coloured ? Explain on the basis of distribution of electrons in the d-orbitals, (©) How does the colour change on heating [Ti(H,O),)** ? (a) In[Ti(H,0),}*, Tiis in +3 oxidation state. The electronic configuration of Ti" ion is 3d. In an. octahedral field, itis ty, &, (&) The electron present in ty, absorbs green and yellow radiation of white light for excitation to ¢,. The complementary colour is purple. (©) On heating, H,O is lost. In the absence of the ligands, crystal field splitting does not occur and hence the substance is colourless. Thus, no heating, the purple coloured complex becomes colourless. Explain the following () Nickel does not form low spin octahedral complexes. (di) What is the bases of formation of the spectro-chemical series? (Gi) Co®* is easily oxidized to Co™ in presence of a strong ligand. Why ? (i) Ni(Z=28) in its atomic/ionic state cannot have two vacant 3d-orbitals and hence d? sp’ hybridisation is impossible. (Gi) The series is based on the absorption of light by complexes with different ligands. (iii) Electronic configurations of Co?* and Co™* are as follow Co** (Z =27): [Ar]83d7, Co?’ (Z = 27): [Ar] 83d° CIN. of Co is 6 :In the presence of a strong ligand, pairing of electrons in 3d-subshell takes place leading of fully filled f,, level with six electrons. This lead to greater stability and lesser energy. Which type of isomerism is shown by [Co(NH);ONO]** and [Co(NH,);NO,]”*? Linkage isomerism Arrange the following complex ions is increasing order of crystal field splitting energy (Ap) [CrCl P™, [CHCN) 6 I) LCANH3 Js P* [CrCl], < (CHNH5 Jo P* < [CCN PO Chemistry by : Khaleel Sir Mob. No. 9219710253