1. Title: Coordination Chemistry Dr
Title: Coordination Chemistry Dr. Sujata Kundan Assistant Professor Department of Chemistry and Chemical Sciences

2. Coordination Compounds Werner’s Theory Chelating ligands and Chelates
Contents:
Coordination Compounds
Werner’s Theory
Chelating ligands and Chelates
Effective Atomic Number
Isomerism in Coordination
Compounds

3. Coordination chemistry is the study of compounds formed between metal ions and other neutral or negatively charged molecules. What is a Coordination Compound? It is a compound which has one or more co-ordinate bonds. In a coordination bond there is a species which donates a lone pair of electrons (ligand) and a species which receives the lone pair of electrons. Most of the time the receiver of the lone pair of electrons is a metal ion (or a metal atom in rare cases) in which case we call it a metal ligand complex.
Some terms used in Explaining Coordinate Compounds
Complex Ion : An electrically charged ion which consist of a central metal atom/ion surrounded by a group of ions or neutral atom is termed as complex ion.
Example: Ni(NH3)6]+2  is a complex ion in which Ni+2 is the

4. Example: [Ag(CN)2]- is a coordinate sphere.
Central Ion and Ligand: The cation to which one or more neutral molecule are coordinated is called the Central Ion and the attached molecules are called Ligands.
Example: So in [Ni(NH3)6]+2 Ni+2 is the central ion  surrounded by  ammonia ligands.
Coordinate Number: The total number of ligands attached to a central metal ion/atom is called the Coordinate Number of that particular ion.
Example: In complex ion [Cu(NH3)4]+2, 4 ammonia ligand are attached to the central metal ion Cu. Hence the coordination number of Cu+2 ion in above complex ion is 4.
Coordinate Sphere: The central metal ion together with molecules or ion coordinated to it is termed as Coordinate Sphere. It is written inside a “[ ]” square bracket.
Example: [Ag(CN)2]- is a coordinate sphere.

5. Ion and the Ligands Coordinated to it.
Example: In complex ion [Cu(NH3)4]+2 , Cu+2 carries a charge of +2 and ammonia molecule is neutral ,hence the total algebraic sum of these two are +2 . Thus [Cu(NH3)4]+2 carries a charge of +2.
Characteristic feature of the coordination compounds :
is their ability to retain their identity in solution (which distinguishes them from double salts like carnallite-KCl.MgCl2 .6H2O ).
Coordination number- number of ligands bound to the central metal ion (or atom).
Coordination sphere- the group comprising the metal ion and the ligands.
Polynuclear complex- complex containing more than one central metal atom

6. Ligands: The molecule or ions that bond to the metal ion are known as ligands. The central metal and its ligands constitute the coordination sphere. A common feature shared by the ligands in complexes is the ability to donate electrons pairs to central metal atoms or ions act as Lewis bases. Types of Ligands: A ligand that uses one pair of electrons to form one point of attachment to the central metal atom or ion is called monodentate ligand. Some ligands are capable of donating more than a single electron pair from different atoms in the ligand and to different sites in the geometry structure of a complex are called Polydentate ligands.

7. Some Examples of The Ligands:
EDTA is a polydentate ligand with six donor atoms
Ethylenediamine (en) is a bidentate ligand

8. He is considered at "Father of coordination chemistry".
Werner’s Coordination Theory:
Alfred Werner (12 December 1866 – 15 November 1919) was a Swiss chemist who was a student at ETH Zurich and a professor at the University of Zurich. He won the Nobel Prize in Chemistry in 1913 for proposing the octahedral configuration of transition metal complexes. Werner developed the basis for modern coordination chemistry. He was the first inorganic chemist to win the Nobel prize, and the only one prior to 1973.
He is considered at "Father of coordination chemistry".

9. The important postulates of Werner’s Coordination Theory are:
In Coordination Compounds the metal atom exhibit two types of valency, viz. Primary Valency, Secondary Valency.
The Primary Valency is ionisable and nondirectional whereas the secondary valency is nonionisable and directional. The primary valency corresponds to oxidation state and the secondary valency corresponds to the coordination number.
Every metal atom has a fixed number of secondary valency, i.e, it has a fixed coordination number.
The metal atom tends to satisfy both its primary and its secondary valencies.
Primary valencies are satisfied by negative ions whereas secondary valencies are satisfied by either negative or neutral or positive ligands.
The secondary valencies are always directed towards fixed position in space and this leads to definite geometry.
Example: If a metal ion has six secondary valencies, these are arranged in octahedral manner around the central metal ion. If the metal ion has four secondary valency, these are arranged in either tetrahedral or square planar arrangement around the central metal ion. The secondary valencies thus determine the stereo chemistry of the compound.

10. Explanation of Coordinate Compounds on the Basis of WernEr’s Theory:
The following formulas describe a series of three coordination compounds:
CoCl3 .6NH3 (a)
CoCl3 .5NH3 (b)
CoCl3 .4NH3 (c)
In 1893, Swiss chemist Alfred Werner proposed that certain metal atoms, primarily those of the transition metals, have two types of valence.
1. The primary valence, is based on the number of electrons the atom loses in forming the metal ion.
2. Secondary or auxiliary, valence is responsible for the bonding of other groups, called ligands, to the central metal ion.

11. Explanation of Coordinate Compounds on the Basis of Werner’s theory:
[Co(NH3)6]Cl (a)
[CoCl2(NH3)4]Cl (b)
[CoCl(NH3)5]Cl (c)
Ionization of coordination compound (a) can be represented as:
[Co(NH3)6]Cl3 (s) H2O [Co(NH3)6]3+(aq) + 3Cl(aq)
In the presence of an excess of AgNO3 (aq)
(a) yield three moles of AgCl(s) per mole of compound
(b) yield only two moles of AgCl(s) per mole of compound
(c) and yield only one moles of AgCl(s) per mole of compound

12. Alfred Werner developed a model of coordination complexes which explains the following observations
At least three different cobalt(III) complexes can be isolated when CoCl2 is dissolved in aqueous ammonia and then oxidized by air to the +3 oxidation state. A fourth complex can be made by slightly different techniques. These complexes have different colors and different empirical formulas.
CoCl3 6 NH3  orange-yellow
CoCl3 5 NH3 H2O  red
CoCl3 5 NH3  purple
CoCl3 4 NH3  green
The reactivity of the ammonia in these complexes has been drastically reduced. By itself, ammonia reacts rapidly with hydrochloric acid to form ammonium chloride
NH3(aq) + HCl(aq)   NH4+(aq) + Cl-(aq)

13. CoCl3 6 NH3(aq) + HCl(aq) (no reaction)
These complexes don't react with hydrochloric acid, even at 100oC.
CoCl3 6 NH3(aq) + HCl(aq) (no reaction)
Solutions of the Cl- ion react with Ag+ ion to form a white precipitate of AgCl.
Ag+(aq) + Cl-(aq)    AgCl(s)
When excess Ag+ ion is added to solutions of the CoCl3.6NH3 and CoCl3.5NH3 H2O complexes, three moles of AgCl are formed for each mole of complex in solution, as might be expected. However, only two of the Cl- ions in the CoCl3.5NH3complex and only one of the Cl- ions in CoCl3.4 NH3 can be precipitated with Ag+ ions.
Measurements of the conductivity of aqueous solutions of these complexes suggest that the CoCl3.6NH3 and CoCl3.5 NH3 H2O complexes dissociate in water to give a total of four ions. CoCl3.5NH3 dissociates to give three ions, and CoCl3.4NH3 dissociates to give only two ions.

14. Werner explained these observations by suggesting that transition-metal ions such as the Co3+ ion have a primary valence and a secondary valence. The primary valence is the number of negative ions needed to satisfy the charge on the metal ion. In each of the cobalt(III) complexes previously described, three Cl- ions are needed to satisfy the primary valence of the Co3+ ion.
The secondary valence is the number of ions of molecules that are coordinated to the metal ion. Werner assumed that the secondary valence of the transition metal in these cobalt(III) complexes is six. The formulas of these compounds can therefore be written as follows.
[Co(NH3)63+][Cl-]3  orange-yellow
[Co(NH3)5(H2O)3+][Cl-]  red
[Co(NH3)5Cl2+][Cl-]2  purple
[Co(NH3)4Cl2+][Cl-]  green

15. The cobalt ion is coordinated to a total of six ligands in each complex, which satisfies the secondary valence of this ion. Each complex also has a total of three chloride ions that satisfy the primary valence. Some of the Cl- ions are free to dissociate when the complex dissolves in water. Others are bound to the Co3+ ion and neither dissociate nor react with Ag+.
The [Co(NH3)6]Cl3 complex dissociates in water to give a total of four ions, and all three Cl- ions are free to react with Ag+ion.
 [Co(NH3)6]Cl3(s) H2O Co(NH3)63+(aq) + 3 Cl-(aq)
One of the chloride ions is bound to the cobalt in the [Co(NH3)5Cl]Cl2 complex. Only three ions are formed when this compound dissolves in water, and only two Cl- ions are free to precipitate with Ag+ ions.
[Co(NH3)5Cl][Cl]2(s) H2O  Co(NH3)5Cl2+(aq) + 2 Cl-(aq)

16. [Co(NH3)5(H2O)]Cl3(s) H2O Co(NH3)5(H2O)3+(aq) + 3 Cl-(aq)
Once again, the three Cl- ions are free to dissociate when [Co(NH3)5(H2O)]Cl3 dissolves in water, and they precipitate when Ag+ ions are added to the solution.
 [Co(NH3)5(H2O)]Cl3(s) H2O Co(NH3)5(H2O)3+(aq) + 3 Cl-(aq)
Two of the chloride ions are bound to the cobalt in [Co(NH3)4Cl2]Cl. Only two ions are formed when this compound dissolves in water, and only one Cl- ion is free to precipitate with Ag+ ions.
 [Co(NH3)4Cl2][Cl](s) H2O Co(NH3)4Cl2+(aq) + Cl-(aq)

17. Werner assumed that transition-metal complexes had definite shapes
Werner assumed that transition-metal complexes had definite shapes. According to his theory, the ligands in six-coordinate cobalt(III) complexes are oriented toward the corners of an octahedron, as shown in the figure below.

18. No. of CL- ions precipitated
Experimental Verification for Werner’s Coordination Theory:
Conductivities of CoCl3·6NH3 were measured when the compounds were dissolved in water.
S.No.
Formula
Conductivity
No. of CL- ions precipitated 1.
CoCl3.6NH3
High 3 2.
CoCl3.5NH3
Medium 2 3.
CoCl3.4NH3
Low 1 4.
IrCl3.3NH3
Zero
AgNO3(aq) + Cl−(aq) −→ AgCl(s) + NO3− (aq)
The no. of Cl– ions precipitated supported the Werner’s theory

19. Chelating Ligands and Chelates:
When a bidentate or a polydentate ligand is attached by two or
more donor atoms to the same central ion forming a ring structure, the ligand is called chelating lignad. The complex is called Chelate and the process is called Chelation. Chelating agent is a polydentate ligand. It simultaneously attaches to two or , more position in the coordination sphere of the central atom of a complex ion. Usually these ligands are organic compounds, and are called chelants, chelators, chelating agents, or sequestering agents.
Some of the Chelating Agents are:
Ethylenediamine (en)
ethylenediaminetetraacetato(edta)
Oxalato (ox)

20. Factors Affecting Stability of Complexes
The term stability may mean either thermodynamic or kinetic stability. Inshort, if a complex is thermodynamically stable means it has large and positive free energy of reaction ∆G and if a complex is kinetically stable means it has large, and positive free energy of activation ∆G#.
Main Factors Affecting Stability of Complexes are:
Charge on the metal ion
Principal quantum number
Nature of ligands
Chelation
Macro cyclic ligands
Hardness and softness
Surrounding conditions

21. 1. Charge on the metal ion: 2. Principal quantum number:
The effect of this factor on stability of complexes can be explained on the basis of crystal field theory. For a given ligand, greater the charge on the metal ion greater is the magnitude of crystal field splitting which ultimately affects the stability of the complex.
Example,
            Ions                            Ligands                                 CFSE (∆o in cm-1)
              V2+                              6H2O                                      12600
              V3+                              6H2O                                      17700
2. Principal quantum number:
Even though the metal ions have same charge, if the principal quantum numbers are different, then the magnitude of CFSE will be different and hence stability will be different.
        Ions                            Ligands                                 CFSE(∆o in cm-1)
                                    d6-Co3+                         6H2O                                     
                                   d6-Rh3                        6H2O                                     

22. 3. Nature of ligands: 4. Chelation: 5. Macro cyclic ligands:
Properties of ligands like size, charge, dipole moment, polarizability and π-bonding capacity will affect the CFSE and stability of complexes. Smaller the size of the ligand, greater is the approach of the ligand with the metal ion and greater is the crystal field splitting. Larger the charge on the anion, greater the polarizability and greater is the magnitude of crystal field splitting.
4. Chelation:
Chelation increases stability. This is because the entropy factor is favorable in case of chelate complexes.
For example; [Cd(en)2]2+ is more stable than [Cd(MeNH2)4]2+ since in the former there is chelation. 
5. Macro cyclic ligands:
The increased stability of complexes due to macro cyclic ligands is termed as “macro cyclic effect”. The reason for this effect is mainly entropy and enthalpy factors. The macro cyclic ligands have cavities of particular size and hence selectively form strong complexes with metal ions of corresponding sizes.
For example; 18-crown-6 forms stronger complex with potassium ion than with sodium ion

23. 6. Hardness and softness:
Stability of complexes depends also on hardness and softness of the metal and the ligands. As per HSAB theory hard acids prefer hard bases and soft acids prefer soft bases.
For example; Ni2+ is a hard acid and hence it forms stable complex with NH2 and not with soft ligand PH3.But Pd2+ being soft acid forms stable complex with PH3 rather than with NH2. 
7. Surrounding conditions:
Even though the above factors outline the stability of complexes, many complexes which are stable under particular conditions may not be stable under some other conditions.
For example; [Co(NH3)6]3+  is unstable in an acidic solution but is stable in water under neutral conditions.
[Co(NH3)6]3+ + 6H3O+ → [Co(H2O)6]3+ + 6NH4+
Hence when somebody says a complex is stable one must always ask, “Under what conditions?”
The conditions may be heat, light, acidity or basicity.

24. SiDGWiCK THEORY / Effective atomic number (EAN):
THE number that represents the total number of electrons surrounding the nucleus of a metal atom in a metal complex. It is composed of the metal atom’s electrons and the bonding electrons from the surrounding electron-donating atoms and molecules. Thus the effective atomic number of the cobalt atom in the complex [Co(NH3)6]3+ is 36, the sum of the number of electrons in the trivalent cobalt ion (24) and the number of bonding electrons from six surrounding ammonia molecules, each of which contributes an electron pair (2 × 6 =12). Also known as the inert-gas rule, nine-orbital rules or the effective atomic number rules.
English chemist Nevil V. Sidgwick made the observation, since known as the EAN rule, that in a number of metal complexes the metal atom tends to surround itself with sufficient ligands that the resulting effective atomic number is numerically equal to the atomic number of the noble-gas element found in the same period in which the metal is situated. This rule seems to hold for most of the metal complexes with carbon monoxide, the metal carbonyls as well as many organometallic compounds. By using this rule it is possible to predict the number of ligands in these types of compounds and also the products of their reactions.
The EAN rule is often referred to as the “18-electron rule” since, if one counts only valence electrons (6 for Co3+ and 2 × 6 = 12 for 6 NH3), the total number is 18.

25. EAN is calculated by EAN = Z-X+Y
where Z= THE ATOMIC NUMBER OF THE METAL x= NUMBER OF ELECTRONS LOST DURING THE FORMATION OF METAL ION Y= NUMBER OF ELECTRONS DONATED BY THE LIGANDS.
The EAN rule is helpful for organometallic compounds and carbonyl complexes, which obey in most cases this rule:
[Cr(CO)6] [Fe(CO)5] [Ni(CO)4]
Cr 24 e Fe 26 e Ni 28 e-
6CO 12 e CO 10 e CO 8 e-
36 e e e-
Similarly the formation of olefin complexes and metallocenes may be explained by the EAN rule:
olefines donate electrons /double bond
ethylene
Butadiene
Benzene
cyclopentadienyl radical 5
[Fe(C5H5)2] [Mn(CO)5C2H4]+ [Cr(C6H6)2]
Fe Mn Cr 24
2 C5H5· CO 10; C2H C6H6 12

26. Limitations of Sidgwick Theory:
Examples:
[Co(NO2)6] [PtCl6] [Ag(NH3)4]+
Co3+ 24 e Pt4+ 74 e Ag+ 46 e-
6NO e Cl- 12 e NH3 8 e-
36 e e e-
But Many elements form complexes which do not obey the EAN
[Cr(NH3)6] [Ni(NH3)6] [CoCl4]2-
Cr3+ 21 e Ni2+ 26 e Co2+ 25 e-
6NH3 12 e NH3 12 e Cl-8 e-
33 e e e-
Limitations of Sidgwick Theory:
Many complexes are stable but do not follow EAN rule.
Theory doesn’t predict the geometry of the complexes.
Theory doesn’t predict the magnetic behaviour of the complexes.

27. Isomerism in Coordination Compounds:
Compounds having the same molecular formula, but differing in physical and chemical properties are known as isomers. This phenomenon is known as isomerism. The isomers can be identified and distinguished from one another because of difference in their physical and chemical properties.
Types of Isomerism:

28. Types of StructuraL Isomerism:
Ionisation Isomerism:
It arise due to the exchange of the anion within and outside the co-ordination sphere. As a result , these isomers give different ions in solution.
Example:
[Co(NH3)5Br]SO SO4- anion in solution
violet
[Co(NH3)5SO4]Br Br- anion in solution
red
Other Examples:
[PtBr(NH3)3]NO NO2- anions in solution
[Pt(NH3)3(NO2)]Br Br- anions in solution
2. Hydrate Isomerism:
It arises when different number of water molecules are present inside and outside the co-ordination sphere.
This isomerism is illustrated by the 3 isomers that have the formula CrCl3.6H2O
[CrCl2(H2O)4]Cl.2H2O Green
[CrCl(H2O)5]Cl2.H2O Blue-Green
[Cr(H2O)6]Cl Purple
These isomers have very different chemical properties and on reaction with AgNO3 to test for Cl- ions, would find 1, 2, and 3 Cl- ions in solution respectively.
These isomers can loose two, one and no water molecules on dehydration with conc. sulphuric acid respectively

29. Coordination isomerism:
It is observed in the co-ordination compounds having both cationic and anionic complex ions. The ligands are interchanged in both the cationic and anionic ions to form isomers.
For examples;
one isomer [Co(NH3)6] [Cr(C2O4)3] another isomer [Co(C2O4)3] [Cr(NH3)6]
4. Linkage Isomerism:
It Occurs in co-ordination compounds with ambidentate ligands. These ligands are capable of coordinating in more than one way. The best known cases involve the monodentate ligands SCN- / NCS- and NO2- / ONO-.
Examples:
[Co(NH3)5(ONO)]Cl the nitrito isomer -O attached
[Co(NH3)5(NO2)]Cl the nitro isomer - N attached.
[Co(NH3)5(NCS)]Cl Co-NCS isothiocyanate and
[Co(NH3)5(SCN)]Cl Co-SCN thiocyanate

30. Types of Stereoisomerism

31. Geometrical Isomerism in square Planer Complexes
It is due to ligands occupying different positions around the central metal atom or ion. Also known as cis-trans isomerism .
Common in square planar and octahedral complexes but not in tetrahedral complexes because all tetrahedral complexes [such as Ma4 , Ma2b2 , Mabcd, where a,b,c,d represents ligands ]exist in only one geometric form in which all positions are adjacent to each other.
Geometrical Isomerism in square Planer Complexes
cis- and trans- refer to the position of 2 groups relative to each other. In the cis- isomer they are "next to each other" i.e. at 90 degrees in relation to the central metal ion, whereas in the trans- isomer they are "opposite each other", i.e. at 180 degrees relative to the central metal ion. a
M b a M b
cis trans-

32. Compound type No. of isomers
Examples of Geometrical Isomerism in Square Planar Complexes:
Compound type No. of isomers
Ma2b (cis- and trans-)
cis- and trans- isomers of [PtCl2(NH3)2]
Ma2bc (cis- and trans-)
Cis- and trans-[Pt(NH3)2ClNO2]
Mabcd (use cis- and trans- relations)
[Pt(NH3)(NH2OH)(NO2)(py)]NO2
Here a, b, c, and d refer to monodentate ligands.
M(ab) (cis- and trans-)
[Pt(gly)2] where gly is glycine
Here ab refer to unsymmetrical bidentate ligand.
A number of examples of these types have been isolated and characterised and they show very different chemical and biological properties.

33. Geometrical Isomerism in Octahedral Complexes
Compound type No. of isomers
Ma2b (cis- and trans-)
CIS- and Trans-[FeCl2(NH3)4]
Ma3b (fac- and mer-)
New labels are introduced to reflect the relative positions of the
ligands around the octahedral structure. Thus; placing the 3 groups
on one face of the octahedral gives rise to the facial isomer and
placing the 3 groups around the centre gives rise to the meridinal
isomer.
[Fe(NH3)3Cl3]

34. M(aa)2b2 3 (2cis- and trans-)
[Co(en)2Cl2]
where en is ethylene diamine bidentate ligand which exists in cis- and trans- form.
Mabcdef
Only one compound known is [Pt(py)(NH3)(NO2)ClBrI]. In this compound only three forms are obtained but no attempt has been made to isolate all the 15 isomers.
Here a,b,c,d,e,f represents monodentate ligand and aa represents bidentate ligand

35. Optical isomers Important Points:
Geometrical isomerism is not observed in:
complexes of co-ordination number 2 and 3 .
2.Complexes of tetrahedral geometry.
3.Complexes Ma3b or Mab3 or Ma4 of square planar geometry.
4.Complexes Ma6 and Ma5b of octahedral geometry.
Optical isomers
Isomers that are non-superimposable mirror images said to be “chiral” (handed) referred to as enantiomers. A substance is “chiral” if it does not have a “plane of symmetry”.

36. Properties of Optical Isomers:
Optical isomers are related as non-superimposable mirror images and differ in the direction with which they rotate plane-polarised light. These isomers are referred to as enantiomers or enantiomorphs of each other and their non-superimposable structures are described as being asymmetric.
Common in tetrahedral and octahedral complexes but not in square planar because of the presence of axis of symmetry. Very common in octahedral complexes.
Properties of Optical Isomers:
Enantiomers
possess many identical properties
solubility, melting point, boiling point, color, chemical reactivity (with nonchiral reagents)
different in:
interactions with plane polarized light
reactivity with “chiral” reagent.
Example;
d-C4H4O62-(aq) + d,l-[Co(en)3]Cl3(aq) 
d-[Co(en)3](d-C4H4O62- )Cl(s) + l-[Co(en)3]Cl3(aq) +2Cl-(aq)

37. Example:
cis-[Co(en)2Cl2]+
rotate mirror image 180°

38. Non-Superimposable
cis-[Co(en)2Cl2]+
enantiomers

39. Optically active complexes are said to exist in the following forms:
Which rotates plane of polarized light towards right side (clockwise direction) is said to be dextro-rotetory or d-form. It is also represented by (+) sign.
Which rotates plane of polarized light towards left side (anticlockwise direction) is said to be laevo-rotatory or l-form. It is also represented by (-) sign. (+) and (-) refer to sign of rotation of optical isomer.
Which is not capable of rotating the plane polarized light is called optically inactive. This isomer is call racemic-[dl, or (±)] from which is made up of 50% d & 50% l- form. In recemic form, one form rotates the plane of polarized light in one direction is balanced by other form in opposite direction.
The d- and l-form have following characteristics
Since d and l form are capable of rotating the plane of polarized light, are said to be optically active or optical isomer. This phenomenon is called optical isomerism or optical activity. Both isomers have exactly identical physical and chemical properties.
If d and l form are mirror image to each other and not superimposed on each other, they are called enantiomerism.

40. Square planar complexes
Optical isomerism in 4-coordinate complexes:
 Mirror image isomerism is not possible tetrahedral and square planar complexes of type [Ma4], [Ma3b] and [Mab3].
Square planar complexes
Square planar complexes seldom show optical isomerism
Since they have all four ligands and the central metal ion in the same plane, hence contain plane of symmetry, therefore complex become optically inactive and cannot show optical isomerism even though all ligands are different.
In 1935 Mills and Quibell succeeded in resolving isobutylenediamine–meso-stilbenediamineplatinum(II)chloride (i.e. [Pt(NH2CH(C6H5)CH (C6H5)(NH2)(NH2CH2C(CH3)2NH2)]Cl2 complex into a highly stable enantiomorphs. This complex show optical isomerism. This in fact, provided a very elegant proof of the planar arrangement of four Pt(II) valences. If the structure were tetrahedral, it would have a plane of symmetry and hence it will not be optically active.

41. Example:
This structure has no plane of symmetry and hence is unsymmetrical and optically active and gives optical isomer.

42. Tetrahedral complexes:
asymmetric tetrahedral molecule (i.e. it should have no plane of symmetry) where all the ligands are different (i.e. [Mabcd] type show optical isomerism.
Example: [As(CH3)(C2H5)S(C6H4COO)]+2

43. Optical isomerism in 6-Coordinate complexes:
Octahedral complexes containing only monodentate ligands:
[Ma2b2c2]±n type
(ii) [Mabcdef]±n type
[Co(NH3)2Cl2(NO2)2]-1
[Pt(py)(NH3)(NO2)(Cl)(Br)(I)]

44. Octahedral complexes containing only symmetrical bidentate chelating ligands:
[M(AA)3]±n type
Example: [Co(en)3]+3, [Co(pn)3]+3, [Pt(en)3]+3, [Cr(C2O4)3]+3, [Cd(pn)3]+2, [Fe(C2O4)3]-3

45. Octahedral complexes containing monodentate and symmetrical bidentate chelating ligands:
[M(AA)2a2]±n type
[Co(en)2Cl2]+: This complex has two geometrical isomers (i.e. cis-trans isomers). In cis isomer there is no plane of symmetry, hence it show optical active isomer as shown.
In trans isomer there is plane of symmetry, hence it is optically inactive and show meso form as shown.

46. (ii) [M(AA)2ab]±n type These complexes also exist in three form in which two form are optically active and third form is inactive as shown below:
(iii) [M(AA)a2b2]±n type
These complexes also exist in three form in which two form are optically active and third form is inactive as shown in below figure:

47. Octahedral complexes containing optically active ligand:
[Co(en)(pn)(NO2)2]+

48. Octahedral complexes containing polydentate ligand:
The complexes having polydentate ligand like EDTA-4 i.e. [Co(EDTA)]- exists in two optical isomers (d-form and l-form) as shown below:

49. References:
1. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.).
2. “Inorganic Chemistry: Principles of Structure and Reactivity”, James E.Huheey, Ellen A.Keiter, Richard L.Keiter, Okhil K.Medhi, Pearson Education, Delhi, 2006.
3. “Inorganic Chemistry”, Shriver and Atkins, 3/e, Oxford University Press, 2002,.
4. “Concise Inorganic Chemistry”, 5/e, Blackwell Science, 2005.
5. “Concepts and Models of Inorganic Chemistry”, 3/e, John Wiley & Sons.
Concise Inorganic Chemistry by J.D.Lee, Amazon, Ed 5th , 1998.