Coordination Chemistry
[CO(NH3)6]Cl3
hexaamminecobalt (III) chloride
[CO(NH3)5Cl]Cl2
pentaamminechloridocobalt (III) chloride
K3[Fe(CN)6]
potassium hexacyanoferrate (III)
[K3[Fe(C2O4)3]
potassium trioxalatoferrate (III)
K2[PdCl4]
potassium tetrachloridoplatinum (II)
[Pt(NH3)2ClNH2CH3]Cl
diamminechlorido (methylamine) platinum(II) chloride
[CO(NH3)4(H2O)2]Cl3
Tetraamminediaquacobalt(IlI) chloride
K2[Ni(CN)4]
Potassium tetracyanidonickelate(II)
[Cr(en)3]Cl3
Tris(ethane-1,2-diamine) chromium(III) chloride
[Pt (NH 3) Br Cl (N0 3)]–
Amminebromidochloridonitrito-N- platinatc(II) ion
[PtCl2(en)2](N03)2
Dichloridobis(ethane-l ,2-diamine) platinum (IV) nitrate
Fe4[Fe(CN)6]3
Iron(III)hexacyanidoferrate(II)
[Zn(OH)4]2-
tetrahydroxozincate(II) ion
[Pt(NH3)6]4+
hexaammineplatinum (IV) ion
K2[PdCl4]
potassiumtetrachloridopalladate(II)
[Cu(Br)4]2-
tetrabromidocuprate (II)ion
[CO(NH3)6]2 (SO4)3
hexaaminecobalt(III) sulphate
K2[Ni(CN)4]
potassiumtetracyanonicklate (II)
K3[Cr(OX)3]
potassiumtrioxalatochromate(III)
[CO(NH3)5ONO]2+
pentaamminenitrito-O-cobalt(III)ion
[Pt(NH3)2Cl2]
diamminedichloridoplatinum(II)
[CO(NH3)5NO2]2+
pentaamminenitrito-N-cobalt (III)ion
[Pt(NH3)2CI (NH2CH3)]Cl
Diammine chlorido (methylamine) platinum (II) chloride
[Ti(H2O)6]3+
Hexaaquotitanium(III)ion
[Ni(CO)4]
Tetra carbonyl nickel (0)
|Mn(H2O)6]2+
Hexaaquamanganese (II) ion
K3[Fe(OX)3]
Potassium tris(oxalate)ferrate(III)
Ca2[Fe(CN)6]
Calcium hexacyanoferrate(III)
[Co(NH3)5Cl]+2
Chloropentammine cobalt(III)ion
[Co(NH3)3(NO2)Cl2]
Triamminedichloronitrocobalt(III)
K4[Fe(CN)6]
Potassium hexacyanoferrate(II)
K3[Fe(C2O4)3]
Potassium tris(oxalato)ferrate(III)
[Co(en)2(NO2)Cl]Cl
Chloronitrodiethylenediamminecobaltichloride
[Cr(NH3)3(ONO)Cl2]H2O
Triamminechloronitritochromium(III)monohydrate
K[PtBrCl(NH3)(NO2)]
Potassium amminebromochloronitroplatinate(II)
[Rh(en)3(NO3)3]
Tris(ethylenediammine)rhodium(III)nitrate
K3[Fe(CN)6]
Potassium hexacyanoferrate(III)
[Co(NH3)4(H2O)2]Cl3
Tetraamminediaquocobalt(III)chloride
[Ag(NH3)2]Cl
Diamine silverchloride
[Ni(Hdmg)2]
Bis(dimethylglyoximato)nickel(II)
[Cu(H2O)4]SO4.H2O
Tetraaquocopper(II)sulphatemonohydrate
[Co(NH3)4(H2O)Cl]Cl2
Tetraammineaquochlorocobalt(III)chloride
[Co(ONO)(NH3)5]SO4
Pentaamiinenitritocobalt(III)sulphate
[Cr(NH3)6]Cl
Hexaammine chromium(III)chloride
[Co(NCS)4]-2
Tetrathiocyanatocobaltate(II)ion
Na[Co(CO)4]
Sodiumtetracarbonylcobaltate
K[Cr(NH3)2(NCS)4]
Potassium diamminetetrathiocyanatochromate(III)
FeF6-3
Hexafluroferrate(III)ion
[Co(H2O)5Cl]Cl
Pentaaquochlorocobalt(II)chloride
K2[PtCl4]
Potassium tetrachloroplatinate(II)
K[BF4]
Potassium tetrafluroborate
hexaamminecobalt (III) chloride
[CO(NH3)5Cl]Cl2
pentaamminechloridocobalt (III) chloride
K3[Fe(CN)6]
potassium hexacyanoferrate (III)
[K3[Fe(C2O4)3]
potassium trioxalatoferrate (III)
K2[PdCl4]
potassium tetrachloridoplatinum (II)
[Pt(NH3)2ClNH2CH3]Cl
diamminechlorido (methylamine) platinum(II) chloride
[CO(NH3)4(H2O)2]Cl3
Tetraamminediaquacobalt(IlI) chloride
K2[Ni(CN)4]
Potassium tetracyanidonickelate(II)
[Cr(en)3]Cl3
Tris(ethane-1,2-diamine) chromium(III) chloride
[Pt (NH 3) Br Cl (N0 3)]–
Amminebromidochloridonitrito-N- platinatc(II) ion
[PtCl2(en)2](N03)2
Dichloridobis(ethane-l ,2-diamine) platinum (IV) nitrate
Fe4[Fe(CN)6]3
Iron(III)hexacyanidoferrate(II)
[Zn(OH)4]2-
tetrahydroxozincate(II) ion
[Pt(NH3)6]4+
hexaammineplatinum (IV) ion
K2[PdCl4]
potassiumtetrachloridopalladate(II)
[Cu(Br)4]2-
tetrabromidocuprate (II)ion
[CO(NH3)6]2 (SO4)3
hexaaminecobalt(III) sulphate
K2[Ni(CN)4]
potassiumtetracyanonicklate (II)
K3[Cr(OX)3]
potassiumtrioxalatochromate(III)
[CO(NH3)5ONO]2+
pentaamminenitrito-O-cobalt(III)ion
[Pt(NH3)2Cl2]
diamminedichloridoplatinum(II)
[CO(NH3)5NO2]2+
pentaamminenitrito-N-cobalt (III)ion
[Pt(NH3)2CI (NH2CH3)]Cl
Diammine chlorido (methylamine) platinum (II) chloride
[Ti(H2O)6]3+
Hexaaquotitanium(III)ion
[Ni(CO)4]
Tetra carbonyl nickel (0)
|Mn(H2O)6]2+
Hexaaquamanganese (II) ion
K3[Fe(OX)3]
Potassium tris(oxalate)ferrate(III)
Ca2[Fe(CN)6]
Calcium hexacyanoferrate(III)
[Co(NH3)5Cl]+2
Chloropentammine cobalt(III)ion
[Co(NH3)3(NO2)Cl2]
Triamminedichloronitrocobalt(III)
K4[Fe(CN)6]
Potassium hexacyanoferrate(II)
K3[Fe(C2O4)3]
Potassium tris(oxalato)ferrate(III)
[Co(en)2(NO2)Cl]Cl
Chloronitrodiethylenediamminecobaltichloride
[Cr(NH3)3(ONO)Cl2]H2O
Triamminechloronitritochromium(III)monohydrate
K[PtBrCl(NH3)(NO2)]
Potassium amminebromochloronitroplatinate(II)
[Rh(en)3(NO3)3]
Tris(ethylenediammine)rhodium(III)nitrate
K3[Fe(CN)6]
Potassium hexacyanoferrate(III)
[Co(NH3)4(H2O)2]Cl3
Tetraamminediaquocobalt(III)chloride
[Ag(NH3)2]Cl
Diamine silverchloride
[Ni(Hdmg)2]
Bis(dimethylglyoximato)nickel(II)
[Cu(H2O)4]SO4.H2O
Tetraaquocopper(II)sulphatemonohydrate
[Co(NH3)4(H2O)Cl]Cl2
Tetraammineaquochlorocobalt(III)chloride
[Co(ONO)(NH3)5]SO4
Pentaamiinenitritocobalt(III)sulphate
[Cr(NH3)6]Cl
Hexaammine chromium(III)chloride
[Co(NCS)4]-2
Tetrathiocyanatocobaltate(II)ion
Na[Co(CO)4]
Sodiumtetracarbonylcobaltate
K[Cr(NH3)2(NCS)4]
Potassium diamminetetrathiocyanatochromate(III)
FeF6-3
Hexafluroferrate(III)ion
[Co(H2O)5Cl]Cl
Pentaaquochlorocobalt(II)chloride
K2[PtCl4]
Potassium tetrachloroplatinate(II)
K[BF4]
Potassium tetrafluroborate
Effective Atomic Number (EAN) Rule
EAN rule is a theory given by Sidgwick. This theory gives an idea about the stability of the coordination compound that forms. EAN represents the total number of electrons surrounding the nucleus of a metal atom in a complex compound or coordination sphere. It is composed of the metal atom's electrons and the bonding electrons from the surrounding electron-donating atoms and molecules.This rule is also called 18 –electron rule.
Generally, the EAN of the central metal is numerically equal to the atomic number of the noble gas element found in the same period in which the central metal atom is located.
Effective Atomic Number Formula
EAN = (Z – X) + (C.N × 2)EAN = (Z – X) + (L × D × 2)
Z is the atomic number of the central metal ion.
X is the oxidation number of the central metal ion.
L is the total number of ligands bound to the central metal atom.
D is the denticity of the ligand.
For example-
EAN of Co in [Co(NH3)6]Cl3 complex is 36.
Co is in +3 oxidation state. So it has 24 electrons.
6 NH3 donates 12 electrons (2 electrons by each NH3)
So, total electrons around the nucleus of a Co atom in [Co(NH3)6]Cl3 complex is-
EAN = 24 + 12 = 36.
Complexes Which Do Not Obey Effective Atomic Number Rule
There are many exceptions to the effective atomic number rule. Many stable complexes are known in which the effective atomic number rule is not obeyed.EAN for [Ni(CN)4]-2 complex is 34
EAN for [Ni(NH3)6]+2 complex is 38
EAN for [Cr(NH3)6]+3 complex is 33
EAN for [CoCl4]-2 complex is 33
EAN for [Cu(NH3)4]+2 complex is 35
EAN for [Fe(CN)6]-3 complex is 35
Significance of Effective Atomic Number
a. Effective atomic numbers help understand why electrons are weakly bound to the nucleus when located farther from it.b. It explains the stability of coordination compounds.
c. The EAN rule is more valid for non-classical complexes, especially carbonyl compounds; hence, this rule is used to determine stability, oxidizing and reducing character of carbonyl compounds.
d. The EAN rule is generally found to be invalid in most complexes (i.e. [Fe(CN)6]−3 is 35 and [Co(CO)4]−2 is 37), but in the case of metal carbonyls, this rule is found to be valid in all cases.
Q. Which of the following complexes do not follow EAN rule
a. [Fe(CN)6]4-b. [Fe(CN)6]3-
c. Ni(CO)4
d. [PdCl4]2-
Q. Whta is the EAN for [Mo(CO)6] complex
a. 34
b. 36
c. 54
d. none of these
Werner’s Theory
In 1893, Werner produced a theory to explain the structures, formation and nature of bonding in the coordination compounds. This theory is known as Werner's theory of coordination compounds.Some important postulates of this theory are given below-
1. The central metal atoms in coordination compounds show two types of valency. first one is the primary valency (Principal or Ionizable) and the second one is the secondary valency (Auxiliary or non-Ionizable).
2. The primary valency relates to the oxidation state and the secondary valency relates to the coordinate number.
3. The number of secondary valences is fixed for every metal atom. that means the coordination number is fixed.
4. Central metal atom satisfy both its primary and secondary valencies. Primary valency is satisfied by negative ion whereas secondary valancies are satisfied by negative ion or by neutral molecules.
5. The secondary valancies are always directed towards fixed position in space and this cause definite geometry of the coordinate compound. For examples: If a metal ion has six secondary valencies, these are arranged octahedrally around the central metal ion. If the metal ion has four secondary valencies, these are arranged in either tetrahedral or square planar arrangement around the central metal ion. The secondary valency thus determines the stereochemistry of the complex ion. While the primary valency is non-directional.
6. The secondary valencies are generally represented by solid lines while the primary valencies are represented by dashed lines and the ions which satisfy both primary and secondary valencies will be drawn with both solid and dashed lines.
Limitations of Werners Theory
1. It could not explain the inability of all elements to form coordination compounds.2. The Werners theory could not explain the directional properties of bonds in various coordination compounds.
3. It does not explain the colour, the magnetic and optical properties shown by coordination compounds.
Q. According to the postulates of Werner's theory for coordination compounds
a. primary valency is ionizable
b. secondary valency is ionizable
c. primary and secondary valencies are non-ionizable
d. only primary valency is non-ionizable
Q. Which of the following postulates of Werner's theory is incorrect
a. Secondary valence is equal to the coordination number and it depends upon the nature of ligand attached to metal.
b. The ions/ groups bound by the secondary linkages to the metal have charecteristic spatial arrangements
c. Primary valencies are satisfied by negative ions
d. None of these
Q. When AgNO3 is added to a solution of Co(NH3)3Cl3, the precipitate of AgCl shows two ionisable chloride ions. This means
a. Two chlorine atoms satisfy primary valency and one secondary valency
b. One chlorine atom satisfies primary as well as secondary valency
c. Three chlorine atoms satisfy primary valency
d. Three chlorine atoms satisfy secondary valency
Valence Bond Theory For Bonding In Coordination Compounds
This theory is exclusively used to explain the stereochemistry and magnetochemistry of complexes. Some main points of this theory are given below-1. It concern itself with the oxidation number of the central metal atom in the complex compounds.
2. The electronic configuration of the central metal in complex compound is then written in their oxidation state.
3. The outer orbital of the central metal is reperesented by a box. The electrons of the inner orbitals does not participate in the bonding.
4. The central metal electron in outer orbitals is shown by upward (↑) and downward (↓) arrow.
5. The electrons of the ligan is shown by cross (x).
6. Each ligand donates 2 electrons to the cental metal atom for the formation of M ← L coordinate bond.
7. The metal - ligan bond is formed by the overlapping of orbitals. Greater the overlapping, stronger the bond.
8. A σ bond is formed by the overlapping of a vacant metal orbital and a filled ligand orbital.
9. A π bond is formed by the overlapping of a filled metal orbital and a vacant ligand orbital.
10. The hybridization and geometry of complexes are related to the number of ligands(i.e. coordination number).
sp hybridisation – Linear
sp2 hybridisation – Triangular
sp3 hybridisation – Tetrahedral
dsp2 hybridisation – Square planar
dsp3 hybridisation – Trigonal bypyramidal
d2sp3 hybridisation – Octahedral
d3sp3 hybridisation – Pentagonal bypyramidal
11. Ligands donating an electron pair easily to the central metal atom are called strong ligands (e.g. CN−,CO etc.) and those donating with difficulty are called weak ligands (e.g. halides, water etc.).
12. Under the influence of strong ligands, metal electrons are forced to pair up even contrary to Hund's rule.
13. The magnetic propertis of complexes are governed by the number of unpaired electrons present in electronic configuration of complexes.
μs = n(n+2)1/2 B.M.
where μs is spin only magnetic moment, n is the number of unpaired electrons and B.M.(Bohr Magneton) is unit of magnetic moment.
Complexes having unpaired electrons are paramagnetic while having no unpaired electrons are diamagnetic.
Example
Co(NH3)6]+3
Co is in +3 oxidation state
Structure of this complex is octahedral in which inner 'd' orbital of central metal atom is used due to strong ligand (NH3).
μs = 0 as this complex does not have any unpaired electron.
Merits of Valence bond theory
1. Valence bond theory explains the geometrical shape and magnetic properties of complexes satisfactorily.2. It explains the formation of inner complexes in the presence of strong ligands and outer complexes in the presence of weak ligands.
3. It explains the back donation of electrons from metal ions to ligands through dπ pπ overlapping.
Demerits of valence bond theory
1. Valence bond theory does not give information regarding magnetic moment due to orbital contribution of electrons.2. It can not explain the spectral properties of complexes.
3. It does not explain the relative stability of complex compounds.
Inner Orbital Complexes
If the complex is formed by the use of inner d-orbitals for hybridisation (written as dnsp3), it is called inner orbital complex. In the formation of inner orbital complex, the electrons of the metal are forced to pair up and hence the complex will be either diamagnetic or will have lesser number of unpaired electrons. Such a complex is also called low spin complex.For example, [Fe(CN)6]-3 and [Co(NH3)6]+3 are inner orbital complexes.
[Co(NH3)6]+3-
Outer Orbital Complexes
If the complex is formed by the use of outer d-orbitals for hybridisation (written as sp3dn), it is called an outer orbital complex. The outer orbital complex will have larger number of unpaired electrons since the configuration of the metal ion remains undisturbed. Such a complex is also called high spin complex.For example, [Fe(H2O)6]+3, [CoF6]- 3 and [Co(NH3)6]+2 are outer orbital complexes.
[Co(NH3)6]+2-
8-Hydroxyquinoline or Oxine
8-Hydroxyquinoline is a chelating agent which has been used for the quantitative determination of metal ions. It reacts with metal ions, losing the proton and forming 8-hydroxyquinolinato-chelate complexes. Its metal chelates have the formula, M(C9H6NO)n, where 'n' is 2 for metals having coordination number 4 (e.g. Mg, Cu, Zn, Cd etc.), 3 for metals having coordination number 6 (e.g. Al, Fe, Bi etc.) and 4 for metals having coordination number 8 (e.g. Zr, Hf etc.).
This reagent separates Mg+2 ion from other alkaline earth metals in ammonical buffers. When Mg(II) solution is treated with oxine solution in the presence of NH3, a pale yellow precipitate is formed.
It is a buff colored compound and insoluble in water. Its 2% ethanolic acid solution is used.
Dimethylglyoxime
Dimethylglyoxime is a white powder chemical compound with the molecular formula C4H8N2O2.
Dimethylglyoxime can be prepared from butanone first by reaction with ethyl nitrite to give biacetyl monoxime. The second oxime is installed using sodium hydroxylamine monosulfonate
Its abbreviation is dmgH2 for neutral form, and dmgH for anionic form, where H stands for hydrogen. This compound is the dioxime derivative of the diketone butane-2,3-dione. dmgH2 is used in the analysis of palladium or nickel.
Dimethylglyoximato is an example of asymmetric ligand. Charge of dimethylglyoximato is -1 and so it is an anionic ligand. The number of donor atoms in dimethylglyoximato ligand is 2, So, it is bidentate ligand. It co-ordinates through 2 N atoms.
Dimethylglyoxime is an extraordinarily sensitive and specific reagent for nickel. Nickel cation reacts with dimethylglyoxime forms an insoluble red precipitate of nickel dimethylglyoxime.
Ni2+ + 2C4H8N2O2 → Ni(C4H7N2O2)2↓(red precipitate) + 2H+
Cupferron
Cupferron is an important analytical reagent having molecular formula NH4[C6H5N(O)NO], as Ammonium salt of N-nitroso-N-phenylhydroxylamine, commonly known as cupferron. It is usually white or light yellow bright flake crystal, with sweet odour. It is soluble in water, benzene, alcohol, ether. The reagent is thermally decomposed to prepare nitrobenzene.
The anion binds to metal cations through the two oxygen atoms, forming five-membered chelate rings.
Cupferron is prepared from phenylhydroxylamine and an NO+
C6H5NHOH + C4H9ONO + NH3 → NH4[C6H5N(O)NO] + C4H9OH
Cupferron is used for quantitative analysis for aluminum, zinc, copper, iron, gallium, mercury, manganese, niobium, tin, tantalum, thorium, titanium, vanadium, zirconium and other elements.
It is also used as polymerization inhibitor, due to unique polymerization inhibition characteristics of cupferron, also the amount is small, can be used as an alternative of phenol polymerization inhibitor BHT, which is current largely applicated.
Cupferron is used for colorimetric determination for the weight of aluminum, bismuth, copper, iron, mercury, zinc, manganese, niobium, gallium, tantalum, thorium, titanium, vanadium, tin and other elements.
It is used for quantitative determination of iron in the strong acid solution, quantitative analysis of vanadate, separation of copper and iron together with the other metals, as a masking agent for measuring rare earths.
Isomerism
Compounds having the same molecular formula but different arrangement of atoms are called isomers and the phenomenon is called isomerism. Due to different arrangement of atoms, isomers differ in their physical or chemical properties.Isomers can be classified into two major categories
1. Structural isomers
2. Stereoisomers
Structural isomerism
The isomers which have same molecular formula but different structural arrangement of atoms or groups of atoms around the central metal ion are called structural isomers. It is of five types and they are given below-1. Ionisation isomerism
2. Solvate Isomerism
3. Coordination isomerism
4. Linkage isomerism
5. Polymerization isomerism
1. Ionisation Isomerism
This type of isomerism arises due to the exchange of groups between the complex ion and ligands outside it.[Co(NH3)5(SO4)]Br and [Co(NH3)5Br]SO4 are Ionisation isomerism. In this complexes Br is replaced by SO4
2. Solvate Isomerism
This form of isomerism is known as hydrate isomerism in case where water is involved as a solvent. This is similar to ionisation isomerism. 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.[Cr(H2O)6]Cl3 and its solvate isomer [Cr(H2O)5Cl]Cl2.H2O and [Cr(H2O)4Cl]Cl2.2H2O are hydrate isomers.
3. 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- [Co(NH3)6][Cr(CN)6], in which the NH3 ligands are bound to Co3+ and the CN– ligands to Cr3+. In its coordination isomer [Cr(NH3)6][Co(CN)6], the NH3ligands are bound to Cr3+ and the CN–ligands to Co3+.
4. Linkage isomerism
Linkage isomerism arises in a coordination compound containing ambidentate ligand.complexes containing the thiocyanate ligand, NCS–, which may bind through the nitrogen to give M–NCS or through sulphur to give M–SCN. Similarly, complexes containing the notro ligand, NO2, which may bind through the nitrogen to give M–NO2 or through oxygen to give M–ONO.
5. Polymerization Isomerism
When simple molecule unite together to form a large molecules, the processes is called polymerization. The simple molecules are called monomer and the large molecules are called polymers. The monomer and polymer have the same emperical formula. Complexes having the same emperical formula is called Polymerization isomerism. [Pt(NH3)2Cl2] and [Pt(NH3)4][PtCl4] are Po complexes.lymerization Isomerism complexes.
Stereoisomers
1. Geometrical isomerism2. Optical isomerism
1. Geometrical isomerism
Geometrical isomerism arises in heteroleptic complexes due to ligands occupying different positions around the central ion. The ligands occupy positions either adjacent to one another (i.e. cis-form) or opposite to one another (i.e. trans-form). This type of isomerism is, therefore, also referred to as cis-trans isomerism.Square planar and Octahedral complexes exist in such isomeric forms.
Square planar complexes of the type MA2X2 , MA2XY, MABX2, MABXY can exist as geometrical isomers. Where A and B are neutral ligands whereas X and Y are anionic ligands such.
Example- [Pt(NH3)2Cl2], [PtCl(C5H5N)2 (NH3)], [Pt(NO2)(py) (NH2OH)(NH3)]+ exist in cis-trans form.
The square planar complexes containing unsymmetrical bidentate ligands such as [M(AB)2] also show geometrical isomerism. For example, the complex [Pt(gly)2] where gly = NH2CH2COO¯ exists in cis and trans form.
Geometrical isomerism is also shown by bridged binuclear complexes of the type M2A2X4.
For example: the complex [PtCl2P(C2H5)3]2 exhibits geometrical isomers.
In octahedral complexes, MA4X2, MA2X4 , MA4XY type complexes shows cis-trans isomers.
Example- [Co(NH3)4Cl2] exist in cis-trans form.
M(A-A)2X2 (e.g. [Cr(en)2Cl2]) and M(A-A)X4 (e.g. [Cr(en)(NH3)4]) and M(AA)2XY type of complexes also shows cis-trans isomers.
Octahedral complexes of the type [MA3B3] like [Co(NO2)3(NH3)3] also exist in two geometrical isomers. When the three ligands (with same donor atoms) are on the same triangular face of the octahedron, the isomer is called facial or fac isomer. When the three ligands are on the same equatorial plane of the octahedron i.e., around the meridian of the octahedron, the isomer is called meridional or mer isomer. In facial isomer, the three ligands are at the corners of a triangular face while in meridional isomer, the three ligands are at the three corners of a square plane.
Octahedral complexes having six different ligands of the type M(ABCDEF) also exhibit geometrical isomerism. These isomers may be written by fixing a ligand at one position and then placing the other ligands trans to it. For example- [Pt(Br)(Cl)()(NO2)(py)(NH3)].
The complexes containing unsymmetrical bidentate ligands also show geometrical isomerism. For example- the complex triglycinatochromium (III), [Cr(gly)3], where gly is H2NCH2COO¯, exists in cis and trans forms.
2. 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 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 (d rotates to the right, l to the left). Optical isomerism is common in octahedral complexes involving didentate ligands.In a coordination entity of the type [PtCl2(en)2]2+, only the cis-isomer shows optical activity.
Complexes of the type M(AA)3 (where AA is symmetrical bidentate ligands) such as [Co(en)3]3+ and [Cr(ox)3]3- exist as optical isomers.
Complexes of the type [M(AA)X2Y2] containing one symmetrical bidentate ligand show optical isomerism. For example [CoCl2(en)(NH3)2]+ exists in d- and l- forms.
Octahedral complexes containing hexadentate ligands such as ethylenediaminetetracetato (EDTA) also show optical isomerism. For example- [Co(edta)]¯ exists as two optical isomers.
