Syllabus

· Werners Theory of Co-ordination Compounds

· Definition of some Important terms Pertaining to coordination compounds.

· Nomenclature of cordination compounds

· Bonding in coordination compounds

· Stability of coordination compounds

· Organometallic Compounds

The chemistry of co-ordination compounds and organometallics is an important and challenging area of modern inorganic chemistry. The exciting field of bio-inorganic chemistry is centred on co-ordination compounds present in living systems. Developments in these areas, over years, have :

(i) led to new concepts of chemical bonding and models of molecular structure

(ii) revolutionised the chemical industry and

(iii) provided critical insights into the functioning and structures of vital components of biological systems.

Co-ordination compounds also find extensive applications in metallurgical processes, analytical and medicinal chemistry.

Co-ordination compounds are special class of compounds in which the central metal atom is surrounded by ions or molecules beyond their valence. These are referred to as co-ordination complexes or simply complexes. The formation of co-ordination compounds is one of the remarkable characteristics of transition metals.

WERNER’S THEORY OF CO-ORDINATION COMPOUNDS

The systematic study of co-ordination compounds was started by Alfred Werner. He prepared a large number of co-ordination compounds and studied their physical , chemical and isomeric behaviour by simple experimental techniques. On the basis of these studies , Wener , in 1898 , propounded his theory of co-ordination compounds. The main postulates of Werner’s theory are :

(i) Metals possess two types of valence.

(a) The primary valence

(b) The secondary valence.

The primary or principal valence is ionisable whereas, the secondary valence is non-ionisable. In modern terminology, the primary valence corresponds to oxidation state and the secondary valence corresponds to the co-ordination number.

(ii) Every metal has a fixed number of secondary valencies

(iii) Every metal tends to satisfy both primary and secondary valencies. Primary valencies are satisfied by negative ions, whereas secondary valencies are satisfied by negative ions or neutral molecules (ligands).

(iv) The secondary valencies are directed towards fixed positions in space about the central ion. As a result, the co-ordination compound acquires a definite structure and geometry. Thus the secondary valence determines the stereochemistry of co-ordination compounds. For example, when the metal has six secondary valencies, these are arranged octahedrally around

the central metal atom. When the metal has four secondary valences these are arranged tetrahedrally or in square planar geometry. The primary valences, on the other hand, are non-directional.

Structures of Co-ordination compounds on the basis of Werner’s theory

In the light of Werner’s theory, the structure of the co-ordination compounds of cobalt(III) chloride with ammonia can be explained.

(i) CoCl₃6 NH₃ : In the complex CoCl₃ 6 NH₃, cobalt has six secondary valences holding ammonia molecules to itself. The three Cl- ions ionically bound through three primary valences, are ionisable and can be precipitated with silver nitrate. Thus the complex may be represented as :

[Co(NH₃)₆]Cl₃

By convention , the solid lines (-) represent secondary valences whereas the dotted lines ( …. ) represent primary valences. The complex ionises as :

[Co(NH₃)₆]Cl₃ ® [Co(NH₃)₆]3+ + 3 Cl-

(ii) CoCl₃ 5 NH₃ : In this complex the co-ordination number of cobalt is 6 , but now five positions are occupied by NH₃ molecules and the sixth position by one of the chloride ions. This chloride ion has a dual character as it satisfies secondary as well as a primary valence as indicated by a full line as well as a dotted line as shown in Fig.

[Co(NH₃)₅Cl]Cl₂

The two Cl-ions satisfy the two remaining primary valencies of cobalt. This satisfies 6 secondary and 3 primary valencies of cobalt. However, on ionisation, only two Cl- ions will be precipitated because one Cl- ion which also satisfied secondary valence, will not be precipitated.

[Co(NH₃)₅Cl]Cl₂ ® [Co(NH₃)₅Cl]2+ + 2 Cl-

can be precipitated

(iii) CoCl₃ 4 NH₃ : In this complex, CoCl₃ 4 NH₃ , two chloride ions exhibit dual character satisfying both primary and secondary valencies. This is shown in Fig.

[Co(NH₃)₄Cl₂]Cl

It will give precipitate with silver nitrate corresponding to only one Cl^- ion and the number of ions in this case is 2. It may be formulated as [Co(NH₃)₄Cl₂]Cl and ionises as

[Co(NH₃)₄Cl₂]Cl ® [Co(NH₃)₄Cl₂]+ + Cl-

can be precipitated

(iv) CoCl₃ 3 NH₃ : In the compound CoCl₃ 3 NH₃, three chloride ions satisfy primary and secondary valencies as shown in Fig.

[Co(NH₃)₃Cl₃]

It is clear from the figure that all the chloride ions are non-ionisable and will not be precipitated by the addition of AgNO₃. Therefore, the complex behaves as neutral non-conducting molecule. It may be formulated as [Co(NH₃)₃Cl₃] and does not ionise .

In all the above complexes , the six secondary valencies of cobalt are directed in space and are arranged octahedrally about the metal ion.

Chemical evidence to support Werner’s formulation came not only from the fraction of chlorine which was precipitated but also from conductance measurements. The following TABLE shows the number of ions produced by some complexes and the co-ordination numbers of the metal ions.

Empirical formula	Werner’s formula	Number of ions	CN
CoCl₃ 6 NH₃	[Co(NH₃)₆]Cl₃	4	6
CoCl₃ 5 NH₃	[Co(NH₃)5Cl]Cl₂	3	6
CoCl₃4 NH₃	[Co(NH₃)4Cl₂]Cl	2	6
CoCl₃3 NH₃	[Co(NH₃)₃Cl₃]	0	6
AgCl 2 NH₃	[Ag(NH₃)₂]Cl	2	2
PtCl₄ 6 NH₃	[Pt(NH₃)₆]Cl₄	5	6
PtCl₄ 3 NH₃	[Pt(NH₃)₃Cl₃]Cl	2	6
PtCl₄ 2 NH₃	[Pt(NH₃)₂Cl₄]	0	6

TERMS INVOLVED IN COORDINATION CHEMISTRY

Some impotant terms required for the description of a coordination are :

Coordination entity, central atom, ligand, coordination number, coordination polyhedron, denticity, chelation and oxidation number (of central atom). Their definitions and meanings are discussed below :

Coordination entity (complex)

A coordination entity constitutes a central atom/ion, usually a metal , to which attached a fixed number of other atoms or groups each of which is called a ligand. It may be neutral or charged. Examples :

Cationic complexes : [Co(NH₃)₄]2+, [Ni(NH₃)₆]2+

Anionic complexes : [Ag(CN)₂]-, [Fe(C₂O₄)₃]3-, [Fe(CN)₆]4-

Neutral complexes : [Co(NH₃)₃Cl₃] , [Ni(CO)₄]

Problem

01. Designate the coordination entities and counter ions in the coordination compounds : [Cr(NH₃)₆]Cl₃ ; K₄[Fe(CN)₆], K₂[PtCl₄] ; [Ni(CO)₄] ; K₂[Ni(CN)₄]

Central atom/ion

In a cordination entity , the atom/ion to which are bound a fixed number of ligands in a definte geometric arrangement around it, is called the central atom or ion. For example, the central atom/ion in the coordination entities ;

[NiCl₂ (H₂O)₄], [CoCl(NH₃)₅]²⁺, [Fe(CN)₆]³^-, are Ni²⁺, Co³⁺ and Fe³⁺ respectively.

Ligands

The ligands are the ions or molecules bound to the central atom/ion in the coordination entity. Thus, in the complex ion, [Ni(NH₃)₆]²⁺ the Ni²⁺ ion is the central ion and the molecules of ammonia are the ligands. Similarly in the complex ion [Co(NH₃)₅Cl]²⁺, the Co³⁺ ion is the central ion while ammonia molecules and chloride ions are the ligands.

The ligands are attached to the central metal atom or ion through co-ordinate bonds. It is essential therefore that the central ion should have vacant orbitals. The ligands should have lone pairs of electrons in their outermost orbitals which have been donated to the central ion. The atom in the ligand, which can donate the electron pair is called donor atom or co-ordinating atom. For example, in ammonia, nitrogen is the donor atom or co-ordinating atom, while in water, oxygen is the donor atom.

Types of Ligands

The ligands may contain one or more than one donor atoms for co-ordination with the central atom. The number of such ligating groups indicate the denticity of the ligand. Accordingly, the ligands are classified as follows:

(i) Unidentate or monodentate ligands : Ligands which can co-ordinate to the central ion through only one donor atom are known as unidentate or monodentate ligands. The examples of unidentate ligands are :

NH₃, H₂O, Cl-, CN^-, OH^- etc.

Some common unidentate ligands and their donor atoms are listed in the TABLE.

Some common unidentate ligands

Name of ligand	Formula	Donor atom
Water	H₂O	O
Ammonia	NH₃	N
Pyridine(Py)	C₅H₅N	N
Carbonyl	CO	O
Nitrosyl	NO	N
Triphenyl phosphine	(C₆H₆)₃P	P
Halide ion	X- ( X = F, Cl, Br, I )	X
Hydroxide ion	OH-	O
Cyanide ion	CN-	N or C
Nitro	NO₂-	N
Nitrito	ONO-	O
Oxide ion	O₂-	O
Thiocyanate	SCN-	S or N
Acetate	CH₃COO-	O

(ii) Bidentate ligands : Ligands which have two donor atoms and therefore, can co-ordinate to the central ion at two positions, are called bidentate ligands. The examples of bidentate ligands are :

Oxalate ion is abbreviated as ox and ethylenediamine is abbreviated en.

(iii) Polydentate ligands : Ligands having more than two donor atoms in the molecule are called polydentate ligands. These may be called tridentate (three), tetradentate (four), pentadendate (five) and hexadentate (six) ligands depending upon the number of donor atoms present in their molecules. Ethylene diamine tetracetate (EDTA), for example is an important hexadentate ligand. It binds through two nitrogen and four oxygen atoms to a central metal ion.

The various multidentate ligands which generally take part in the formation of complexes are given in TABLE.

Chelating Ligands

When a bidentate or polydentate ligand is attached by two or more donor atoms to the same metal ion forming a ring structure , the ligand is called chelating ligand. The complex is called chelate. For example, when a bidentate ligand such as ethylene diamine attaches to Cu2+ ion, through two amino groups and forms a ring structure, it is called chelating ligand. The resulting complex ion [Cu(NH₂CH₂CH₂NH₂)₂]2+ or [Cu(en)₂]2+ is called a chelate.

Ambidentate ligands

Ligands which can ligate through different atoms present in it are called amidentate ligands. Examples of such ligands are the NO₂^- ion can coordinate through either the nitrogen or the oxygen atoms to a central metal atom/ion. Similarly SCN ^- ion can coordinate through the sulphur or nitrogen atom. Such possibilities give rise to linkage isomerism in coordination compounds.

Coordination Number

The coordination number of a central atom/ion is determined by the number of sigma bonds between the ligands and the central atom/ion. Pi-bonds , if any between ligating atom and the central atom/ion are not considered for the determination of coordination number. The sigma bonding electrons may be indicated by a pair of dots (:) preceding the door atom in the ligand formula as in :

[Co(: NH₃)₆]³⁺, [Fe(: CN)₆]³^-, [Ni(:CO)₄], [Co(:Cl)₄)]²^-.

Coordination polyhydron

The spatial arrangement of the ligand atoms which are directly attached to chemical atom/ion defines a coordination polyhydron about a central atom. The figure shows the shapes of tetrahedral, square planar, octahedral, square pyramidal and trigonal bipyramidal coordination polyhedra.

Shapes of tetrahedral, square planar, octahedral, square pyramidal and trigonal bipyramidal coordination polyhedron. M represents the central atom/ion and L represents a unidentate ligand

Oxidation Number of central atom

The oxidation number is the residual charge that appears on the central atom when all atoms/ions are removed from it. Similarly the charge on the complex is the sum of the charges on the central metal ion and its surrounding ligands. It must be noted that :

(i) the sum of the charges of constituents is equal to zero if the complex is neutral and

(ii) for anionic or cationic complex, the sum of the charges of the constituents is equal to charge on the co-ordination sphere.

The above points are illustrated by taking the following examples.

(i) The complex [Co(en)₂Cl(ONO)] the charge on the complex is zero, whereas the charge on the central metal ion may be calculated as under :

(ii) Similarly in the complex ion [Co(NH₃)₂Cl₄]2- the charge on the complex is - 2 and charge on the central metal ion is :

(iii) In the complex [Fe(CN)₆]4- the charge on the complex is - 4 and the oxidation number of the central metal is

It must be noted that the total electrovalent charge on the complex ion changes with number of charged ions which may enter into co-ordination sphere to replace neutral molecules.

Co-ordination sphere

In the complexes, the central metal ion along with the non-ionisable ligands attached to it is always written in square bracket [ ]. The part of the complex is called co-ordination sphere. For example, in a complex [Cu(NH₃)₄]SO₄ , the part [Cu(NH₃)₄]2+ is called co-ordination sphere and SO₄2- constitutes ionisation spheres. It may be noted that :

(i) Species present in the co-ordination sphere are non-ionisable.

(ii) Species present in the ionisation sphere are ionisable.

It may be noted that only some metals like Cr(III) and Co(III) etc exhibit a constant co-ordination number, whereas in most of the ions the co-ordinatrion number varies with the ligand. However, the most common oxidation numbers are four and six.

In the formation of a complex of CoCl₃ with NH₃, the complex compound is suitably labelled to illustrate the various terms discussed above.

Nomenclature of Co-ordination Compounds

The nomenclature system given by IUPAC (International Union of Pure and Applied Chemistry) has been given below:

(i) Order of naming ions : In ionic complexes, cations are named first and then the anion. Non-ionic or molecular complexes are given one word name.

(ii) Naming the co-ordination spheres : The ligands are named first and then the name of the the central metal ion.

(iii) Naming of the ligands : The names of the negative (anionic) ligands end in ‘o’ and names of positive (cationic) ligands end in ‘ium’. Neutral ligands are named as such. Examples :

(a) Negative ligands : names end in ‘o’. Normally ‘’ide’’ ending are changed into ‘o’ ; ‘ite’ endings are changed to ‘ato’.

Anions	Name of ligand	Anions	Name of ligand
F-	Fluoro	ONO-	Nitrito*
Cl-	Chloro	NO₂-	Nitro*
Br-	Bromo	NO₃-	Nitrato
I-	Iodo	SO₄2-	sulphato
OH-	Hydroxo	SCN-	thiocyanato**
O2-	Oxo	NCS-	Isothiocyanato**
CN-	Cyano	-OOC-COO-	oxalato
CO₃2-	carbonato	O₂2-	Peroxo
S2-	sulphido	N3-	Nitrido

· If nitrite ion is attached through nitrogen atom (-NO₂), the name ‘nitro’ is used ; if attached through O atom (-ONO), the name is ‘nitrito’.

** If the thiocyanate ion attached through S atom

(- SCN) , the name ‘thiocyanato’ is used ; if attached through N atom (-NCS), the name is ‘isothiocyanato’.

(b) Positive ligands : The names end in ‘ium’.

Cation	Name of ligand
NO+	Nitrosonium ion
NO₂+	Nitronium ion
NH₂NH₃+	Hydrazinium ion

( c) Neutral Ligands: they are named as such.

Neutal molecule	Name of ligand
NH₂CH₂CH₂NH₂	Ethylenediamine
CH₃NH₂	Methyl amine
(C₆H₅)₃P	Triphenylphosphine
C₅H₅N	Pyridine

However, there are some exceptions as listed below :

Species	Name of ligand
H₂O	Aqua
CS	Thiocarbonyl
CO	Carbonyl
NO	Nitrosyl
NH₃	Ammine

(iv) Number of ligands : The number of ligands of a given type is indicated by a Greek prefix ; mono = 1 ; di = 2 ; tri = 3 ; tetra = 4 ; penta = 5 ; hexa = 6 etc. For example dichloro signifies two Cl- ions as ligands, pentaammine denotes five NH₃ molecules, tetraaqua denotes four H₂O molecules and so on. However, in the case of complex ligands, which have already prefixes, di, tri , tetra etc in their names , prefixes bis, tris, tetrakis, etc are used to indicate the number of ligands. For example, to denote two en molecules as ligands, we write bis(ethylenediamine). Similarly, terakis(ethylenediamine) indicates that there are four en molecules as ligands.

(v) Order of naming ligands : When more than one type of ligands are present, they are named in alphabetical order of preference without separation by hyphen. For example in the complex [Co(NH₃)₄Cl(NO₂)]+, the ligands are named in the order : ammine, chloro and nitro.

(vi) Ending of ligands

(a) For cationic and neutral complexes, the name of the central metal ion is unchanged, but we denote its oxidation state by a Roman numeral. For example, the complex ion [Co(NH₃)₆]3+ , is named as hexaamminecobalt(III) ion. The complex [Co(NH₃)₆]Cl₃ is named as hexaamminecobalt(III) chloride. More examples :

[Co(H₂O)₄Cl₂]Cl : Tetraaquadichlorocobalt(III)chloride

[Co(NH₃)₃ (NO₃)₃] : Triamminetrinitratocobalt(III)

[Co(en)₃]Cl₃ : Tris(ethhylenediamine)cobalt(III)chloride

[Pt(NH₃)₄(NO₂)Cl]SO₄:

Tetraamminechloronitroplatinum(IV)sulphate

[Co(NH₃)₃ (NO₂)Cl₂] : Triamminedichloronitrocobalt(III)

(b) In complex anions the oxidation state of the central metal ion is denoted by a Roman numeral and the name of the central metal ion modifies to end in ‘ate’.

For example, the complex [Fe(CN)₆]4- , is named as hexacyanoferrate(II) ion. The complex K₄[Fe(CN)₆] is named as potassiumhexacyanoferrate(II). More examples :

K[Pt(NH₃)Cl₅] : Potassiumamminepentachloroplatinate(IV).

K₂[HgCl₄] : Potassiumtetrachloromercurate(II)

K[Au(CN)₂] : Potassiumdicyanoaurate(I)

[Al(H₂O)₂ (OH)₄]- : Diaquatetrahydroxoaluminate(III) ion

(iv) Point of attachment

When a ligand can be co-ordinated through more than one atom, then the point of attachment of ligands to the central atom is indicated by placing the symbol of the donor atom after the name of the ligand. In some cases different names may be used for alternative modes of attachment. For example, SCN group may be bonded in two ways : (i) M-SCN and (ii) M-NCS. M-SCN(through S) is called thiocyanato or systematically as thiocyanato-S. M-NCS (through N) is called isothiocyanato or systematically as isothiocyanato-N . Thus K₃[Co(NCS)₆] is named as potassiumhexaisothiocyanato-N cobaltate(III).

Problems

02. Give the IPAC names of following compounds :

i) [Co(NH₃)₃(NO₂)₃]

ii) [Pt(NH₃)₄(NO₂)Cl]SO₄

iii) K[Ag(CN)₂]

iv) [Co(NH₃) ₃ (NO₂)Cl₂]

v) Na[Au(CN)₂]

vi) [Cu(H₂O) 2(NH₃)₄]SO₄

vii) Na₃ [Fe(CN)₅NO]

viii) Na₃ [Co(NO ₂)₆]

ix) K₄[Ni(CN)₄]

x) [Co(NH₃)₆]3+

xi) [Pt(NH₃)₄][PtCl₄]

xii) [Co(en)₂Cl ₂]₂SO₄

xiii) K₃ [Cr(C₂O₄)₃]

xiv) [Ni(CO)₄]

xv) [Pt(NH₃) ₆]Cl₄

xvi) [Pt(NH₃) ₂Cl ₂]

xvii) [Co(NH₃)₅(ONO)]SO₄

03. Give the IUPAC names of following complexes.

(i) [Cr(NH₃)₆]3+ vi) [Ni(NH₃)₆]Cl ₂ [Mn(H ₂O)₆] 2+ vii) [Co(en)₃]3+

(ii) [Fe(CN)₆]4- viii) [NiCl₄]2-

(iii) [Co(en)₃]3+ ix) [Fe(CN)₆]3-

(iv) [Co(en)₂Cl(ONO)]+

4. Give the formulae of the following complexes.

(i) Ammineaquadibromocopper(II)

(ii) Potassiumtetracyanocuperate(II)

(iii) Diamminedichloroplatinum(II)

(iv) Potassiumpentacyanonitrosylferrate(II)

(v) Hexammineplatinum(IV)chloride

(vi) Potassiumhexacyanoferrate(III)

(vii) Hexaaquamanganese(II)

(viii) Dichlorodiammineplatinum(II)

(ix) Chlorobis(ethylenediamine)cobalt(III)ion

(x) Sodiumtetracarbonylcobaltate(-I)

ISOMERISM IN CO-ORDINATION COMPOUNDS

The compounds which have same molecular formulae but different structures and therefore different physical and chemical properties are called isomers. Isomers can be broadly classified into two major categories.

(i) Structural isomers (ii) Stereoisomers.

Structural isomerism

The isomers which have same molecular formula but different structural arrangements of atoms or groups of atoms around the central metal ion is called structural isomers. These are discussed below :

(i) Ionisation isomerism : In this type of isomerism, the difference arises from the interchange of groups within or outside the co-ordination spheres. Therefore, these isomers give different ions in solution. For example, there are two isomers of the compound of formula [Co(NH₃)₅Br]SO₄

Structure Mode of ionisation

[Co(NH₃)₅Br]SO₄ ® [Co(NH₃)₅Br]2+ + SO₄2-

Pentamminebromo Gves test for SO₄2- ions with

cobalt(III)sulphate BaCl₂ solution

(red-violet)

[Co(NH₃)₅ SO₄]Br ® [Co(NH₃)₅ SO₄]+ + Br-

Pentamminesulphato Gves test for Br- ions with

cobalt(III)bromide AgNO₃ solution

(red)

Other compounds showing this type of isomerism are :

[Co(NH₃)₄Cl₂]NO₂ and [Co(NH₃)₄Cl(NO₂)]Cl

[Pt(NH₃)₄Cl₂]Br₂ and [Pt(NH₃) 4Br₂]Cl₂

(ii) Hydrate isomerism : This type of isomerism is similar to ionisation isomerism. In some cases, water can occur in more than one way inside and outside the co-ordination sphere as a co-ordinated group or water of hydration. For example, there are three isomers having molecular formula CrCl₃.6H₂O. These are :

[Cr(H₂O)₆]Cl₃ , [Cr(H₂O)₅Cl]Cl₂. H₂O and[Cr( H₂O)₄Cl₂]Cl .2 H₂O

They differ from one another as :

(i) [Cr(H₂O)₆]Cl₃ ,(violet) : It does not lose water when treated with con. Sulphuric acid and the three chloride ions are precipitated with silver nitrate.

(ii) [Cr(H₂O)₅Cl]Cl₂. H₂O (blue-green) : It loses water molecule when treated with con. Sulphuric acid and two chloride ions are precipitated with silver nitrate.

(iii) [Cr( H₂O)₄Cl₂]Cl .2 H₂O (green) : It loses two water molecules on treatment with con. Sulphuric acid and one chloride ion is precipitated with silver nitrate.

(iii) Co-ordination isomerism : The isomerism occurs in compounds containing both cationic and anionic complexes. The two isomers arise due to the interchange of ligands in the co-ordination sphere of cationic and anionic parts. The examples are

(i) [Co (NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆]

(ii) [Cu(NH₃)₄][PtCl₄] and [Pt(NH₃)₄][CuCl₄]

(iv) Linkage isomerism : In few ligands, there exists two possibilities in the mode of attachment to the metal atom or ion because there can be two atoms which can donate their lone pairs. For example, in NO₂- ion, the nitrogen atom as well as the oxygen atom can donate their lone pairs. This gives rise isomerism. If nitrogen donates its lone pair, one particular compound will be formed. If oxygen does so, a different compound (although having the same molecular formula) is obtained. If the bonding is through nitrogen, the ligand is named as nitro and if it is through O, it is named nitrito.

NO₂- : Nitro ONO- : Nitrito

For example:

[Co(NH₃)₅ (NO₂)]Cl₂ : Pentaamminenitrocobalt(III)chloride (yellow brown)

[Co(NH₃)₅(ONO)]Cl₂ : Pentaamminenitritocobalt(III)chloride (red)

(ii) Stereoisomers

Stereoisomers are those isomers which have the same position of atoms or groups but they differ in the arrangements around the central atom. Stereoisomerism is of two types (1) Geometrical isomerism and (2) Optical isomerism.

1. Geometrical isomerism

Geometrical isomerism is due to ligands occupying different positions around the central ion. The ligands occupy positions either adjacent to one another or opposite to one another. These are referred to as cis-trans isomerism.

In complexes of co-ordination number-4, geometrical isomerism is not possible in tetrahedral complexes because all the positions are adjacent to one another in these complexes. However, square planar complexes show isomerism.

It may be noted that in square planar complexes, the positions 1-2, 2-3, 3-4 and 1-4 are cis with respect to each other, while the positions 1-3 and 2-4 are trans to each other. In octahedral complexes, positions 1-6, 2-4 and 3-5 are trans while the others such as 1-2, 1-3, 2-3, 6-3, 6-3, 6-4, etc are cis to each other.

Consider some examples :

(i) [Pt(NH₃)₂Cl₂] exists in cis and trans forms as :

An octahedral complex [Co(NH₃)₄Cl₂]+ exists in two isomers.

(ii) An octahedral complex [Co(en)₂Cl₂]+ exists as two isomers.

Geometrical isomers of [Co(en)₂Cl₂]+

2. Optical isomerism

There are certain substances which can rotate the plane of polarised light. These are called optically active substances. The isomers which rotate the plane polarised light equally but in opposite directions are called optically active isomers. The isomer which rotates the plane of polarised light to the left is called laevorotatory designated as (l), while that which rotate the plane polarised light towards right is called dextro rotatory(d). The d and l isomers are mirror images of each other just as left hand is mirror image of the right hand. The essential condition for a substance to show optical activity is that the substance should not have a plane of symmetry in its structure. The optical isomers have identical physical and chemical properties. They differ only in the direction in which they rotate the plane of polarised light.

The common examples showing optical isomerism are octahedral complexes having bidentate ligands.

For example,

(i) Complexes such as [Co(en)₃]3+ and [Cr(ox)₃]3- exist as optical isomers.

Optical isomers of [Co(en)₃]3+

(ii) Complexes of the type [Co(en)₂Cl₂]+ exists as optical isomers..This complex forms geometrical isomers (cis and trans forms.The trans form does not show optical isomerism. The reason is that the molecules in this case is symmetric. On the other hand , the cis isomer is asymmetrical and can be resolved into optical isomers.

(optically inactive)

Optical isomers of cis and trans [Co(en)₂Cl₂]⁺

Problem

06. Draw the structures of geometrical isomers of [CoCl₂(NH₃)₄}⁺.

07. Whivh out of the following two coordination entities is chiral ?

(a) Cis – [CrCl₂(Ox)₂]³^- (b) trans-[CrCl₂(Ox)₂]³^-

BONDING IN CO-ORDINATION COMPOUNDS

Werner’s theory was the first successful attempt to account for the nature of bonding in the case of cordination compounds. However, with the development of modern theories like (i) valence bond theory (ii)Crystal field theory (iii) Ligand field theory and (iv) Molecular orbital theory, it has been possible to explain the properties of these compounds such as colour, geometry and magnetic properties.

VALENCE BOND THEORY

The valence bond treatment of co-ordination compounds was developed by Pauling. The brief outline of the theory are as follows:

(i) A suitable number of vacant orbitals must be present on the central metal atom for the formation of co-ordinate bond with suitable ligand orbitals.

(ii) Depending upon the total number of bonds to be formed, the central metal ion can use appropriate number of atomic orbitals, i.e., s, p, or d for hybridisation yielding a set of equivalent orbitals called hybrid orbitals.

(iii) The hybrid orbitals are then allowed to overlap with those ligand orbitals that can donate an electron pair for bonding.

(iv) The outer orbital (high spin) or inner orbital (low spin) complexes are formed depending upon whether the d-orbitals of outer shell or the d-orbitals of inner shell are used in hybridisation scheme.

A few examples are discussed to illustrate the above points.

1. Hexaamminechromium(III) ion

In this complex, the chromium ion is in +3 oxidation state and has an electronic configuration of 3d3 as shown below:

The two 3d , one 4s and three 4p-orbitals then hybridise to yield d2sp3 hybrid orbitals pointing towards the six ends of an octahedron. The six ammonia molecules then donate a pair of electrons to each of the vacant orbitals. Consequently the structure of the complex is octahedral. The presence of three unpaired electrons in the remaining orbitals of chromium (III) makes the complex paramagnetic.

(ii) Tetracarbonylnickel (0)

The outer electronic configuration of nickel(0) is 3d84s2. Under the conditions of formation of chemical bonds, the 4s electrons are forced to pair with 3d electrons. Thus one s and three p orbitals hybridise to give four equivalent orbitals oriented tetrahedrally. Each CO group donates a pair of electrons to form tetrahedral complex. The magnetic studies reveal that the compound is diamagnetic in nature and thus supports the forced pairing of 4s electrons with 3d electrons. The hybridisation scheme and formation of complex is illustrated below

(iii) Tetracyanonickelate(II) ion

It is a square planar complex in which nickel is in the +2 oxidation state and has an electronic configuration of 3d8. In this case, the two unpaired electrons of the 3d-orbitals are forced to pair in the presence of a strong ligand CN- ion. Since four ligands are to be accomodated, therefore nickel(II) ion undergoes dsp2 hybridisation forming four equivalent dsp2 hybrid orbitals. These accommodate the four pair of electrons from the ligands.

In the formation of the complex , since the inner d-orbitals is used in the hybridisation [Ni(CN)₄]2- is called

inner orbital or spin paired or low spin or hyper ligand complex. The hybridisation scheme is as follows :

As there are no unpaired electrons, the complex is diamagnetic and because of dsp2 hybridisation of nickel , it has square planar geometry.

(v) Hexafluorocobaltate(III) ion

In this complex, the outer electronic configuration of cobalt in excited state i.e., under conditions of formation of complex is 3d6. Since F- ion provides a weak ligand field, therefore, one 4s , three 4p and two 4d-orbitals hybridise to yield six sp3d2 hybrid orbitals pointing towards the six corners of an octahedron. The 6 F- ions then donate a pair of electrons to each of these vacant orbitals to have an octahedral geometry. The presence of unpaired electrons in 3d orbitals make the complex paramagnetic. The schematic formation of complex is illustrated below:

The formation of complex involves d-orbitals of outer shell and hence is called outer orbital or high spin or hypoligated complex.

Drawbacks of VB Theory

The valence bond theory, to a large extent , explains the formation , structures and magnetic behaviour of coordination compounds, it suffers from the following short-comings.

(i) A number of assumptions are involved.

(ii) There is no quantitative interpretation of magnetic data.

(iii) It has nothing to say about the spectral properties of coordination compounds.

(iv) It does not give quantitative interpretation of 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 strong and weak ligands

CRYSTAL FIELD THEORY

The Crystal Field Theory (CFT) was originally proposed for explaining the optical properties of crystalline solids. It was applied to study of coordination compounds in 1950s. CFT assumes the ligands to be point charges and the interaction between them and the electrons of the central metal to be electrostatic in nature. The five d-orbitals in an isolated gaseous metal atom/ion have the same energy, i.e., they are degenerate. The degeneracy is maintained if a spherically symmetrical field of negative charges surrounds the metal atom/ion. However, when this negative field is due to ligands (either anions or negative ends of dipolar molecules like NH₃ and H₂O) in a complex, it becomes asymmetrical and the degeneracy of d-orbitals is lifted. It results in splitting of the d-orbitals energies. The pattern of splitting depend upon the nature of the crystal field. The splitting of d-orbital energies and its effects , form the basis of the crystal field treatment of coordination compounds.

Crystal Field Effects in Octaheral Coordination entities

For convenience, let us assume that the six ligands are positioned symmetrically along the catesian axes, with the metal atom at the origin. As the ligands approach, first there is an increase in energy of d-orbitals relative to that of free ion just as would be the case in spherical field. Next the orbitals lying along the axes( dz2 and dx2 - y2 )get repelled more strongly than d_xy, d_yz and d_xz orbitals, which have lobes directed between the axes. The dz2 and and dx2 - y2 orbitals get raised in energy and d_xy, d_yz and d_xz orbitals are lowered in energy relative to the average energy in the spherical crystal field. Thus, the degenerate set of d-orbitals get split into two sets : the lower energy level set, t_2g and higher energy , e_g set. The energy separation is denoted by D_o (the substript o for octahedral)

Let us consider the significance of D_o by considering first the d1 coordination entity. [Ti(OH₂)₆]3+, formed in aqueous solutions of Ti3+( d1ion). The single electron occupies one of the lower t_2g orbitals. In d2 and d3 coordination entities, the d-electrons occupy the t_2g orbitals singly in accordance with Hund’s rule. For d4 ions,, two possible patterns of electron distribution arises :

(i) the fourth electron may enter an e_g orbital(of higher energy) or

(ii) it may pair an electron in the t_2g level.

The actual configuration adopted is decided by the relative values of D_o and P. P represents the energy required for electron pairing in a single orbital.

If D_o is less than P (D_o < P), we have weak field or high spin situation and the fourth electron enters one of the e_g orbitals giving the configuration t_2g3 e_g1. If the fifth electron is added to a weak field cordination entity, the configuration becomes t_2g3 e_g 2.

When D_o > P , we have the strong field, low spin situation, and pairing will occur in the t_2g level and the e_g level remaining unoccupied in entities of d1 to d6 ions. Calculations show that coordination entities with four to seven d-electrons are more stable for strong field as compared to weak field cases.

Crystal field Effects of Tetrahedral Coordination Entities

In tetrahedral coordination entity, the d-orbital splitting(Fig) is inverted and smaller as compared to octahedral field splitting.

For the same metal, the same ligands and metal-ligand distances, it can be shown that D _t = - 4 / 9 Do. Consequently, the orbital splitting energies are not sufficiently large for forcing pairing and therefore , low spin configurations are rarely observed.

Factors determining the magnitude of the orbital splitting energy (D )

(i) Oxidation state of metal ion : Generally, the higher the ionic charge on the central metal ion, the greater the value of D.

(ii) Nature of the metal ion : For analogous entities within a group, the D values differ. The general trend being 3d < 4d < 5d. Thus while going from Cr to Mn or Co to Rh, the Do value increases by ~50%. As a consequence of this, coordination entities of second and third transition series have greater tendency to be low spin as compared to the first transition series.

(iii) Geometry of the coordination entity : The Dt value is only ~50% as large as that of Do.

(iv) Nature of ligand : The ligands can be arranged in the order of increasing field strength.

COLOUR IN COORDINATION COMPOUNDS

One of the achievement of CFT is its ability to explain the colours of transition metal complexes. Most of the transition metal complexes are coloured in their solid or solution form . The transition metals have the property to absorb certain radiations from the visible region of the spectrum and as a result , the transmitted or reflected light is coloured. The visible light is a mixture of radiations of many wavelengths ranging from blue (about 400 nm) to red (about 700 nm). When some of these wave lengths are removed from a beam of white light by passing the light through a sample, then the transmitted light is no longer white. For example, if red light (long wavelength) is absorbed from white light, then it appears to us to be green (short wavelength). On the other hand if green light (short wave length) is absorbed then it appears to us to be red. Red and green colours are complementary colours. Thus, the colour we see of a complex in solid or in a solution is due to the light which is not absorbed but which is transmitted (complimentary colour)

In the case of transition metal complexes, the energy difference between two sets of d-orbitals is very small. When visible light falls on them, the electron gets raised from lower set of orbitals to higher set of orbitals. For example, in case of octahedral complexes , the electron goes from set of d_xy, d_yz ,d_xz orbitals to set of dz2 and dx2 - y2 orbitals. As a result of absorption of some selected wave length of visible light corresponding to energy difference between these sets of energy levels, the transmitted light gives colour to complexes. For example, the complex [Ti(H₂O)₆]3+

is purple.

In this complex, the metal ion has d1configuration. In ground state this electron occupies one of the lower set of orbitals. When white light passes through the complex, the electron get excited from the lower set of orbitals to one of higher set of orbitals (dz2 or dx2 - y2 orbital) by absorbing light equal to Do.

The energy corresponding to this transition corresponds to green and yellow lights which are absorbed from the white light, while blue and red portions are emitted. The solution of complex [Ti(H₂O)₆]³⁺, therefore looks purple.

By using sphetrocopic data for a number of coordination compounds having same metal ion but different ligands , the crystal field splitting for each ligand has been calculated and the spectrochemical series :

I^- < Br^- < S²^- < SCN^- < Cl^- < F < OH^- < C₂O₄²^- < O²^- < H₂O < NCS^- < Py , NH₃ < CN^- < CO established which helps in understanding the nature of bonding and structures of these compounds.

Since the different sets of orbitals in octahedral complexes vary from one metal ion to another and the nature of the ligands, therefore different complexes absorb different amounts of energies from visible region and exhibit different colours. For example, consider three complexes of Co³⁺ as [Co(H₂O)₆]³⁺, [Co(NH₃)₆]³⁺ and [Co(CN)₆]³^-. According to spectrochemical series, the crystal field splitting energies will be in the order of ligands as :

H₂O < NH₃ < CN^-

Therefore , excitation energy will be smallest (largest wave length absorbed) for [Co(H₂O)₆]³⁺ complex and largest (smallest wavelength absorbed) for [Co(CN)₆]³^- complex. That is why [Co(H₂O)₆]³⁺ absorbs orange colour and appears blue and [Co(CN)₆]³^- absorbs violet colour and appears yellow.

	[Co(H₂O)₆]³⁺	[Co(NH₃)₆]³⁺	[Co(CN)₆]³^-
Do value	Small	Intermediate	Large
Excitation energy (DE)	Small	Intermediate	Large
Absorption wavelength(l)	Large	Intermediate	Small
Colour absorbed	Orange	Blue	Violet
Colour transmitted	Blue	Yellow orange	yellow

It may be noted that only those transition metal complexes are coloured which have incomplete d-subshells. The transition metal ions having completely filled or completely empty d-subshells are colourless. For example , complexes of Cu⁺(d¹⁰), Zn²⁺(d¹⁰), Ag⁺ (d¹⁰), Ti⁴⁺(d⁰) etc are colourless.

Magnetic Properties of Coordination Compounds

Additional information for understanding the nature of coordination entities is provided by magnetic susceptibility measurements. Coordination compounds generally have partially filled d-orbitals and as such they are expected to show characteristic magnetic properties depending upon the oxidation state, electron configuration, coordination number of the central metal and the nature of ligand field. It is experimentally possible to determine the magnetic moments of coordination compounds which can be utilised for understanding the structures of these compounds.

A critical study of the magnetic data of coordination compounds of metals of the first transition series reveals some complications. For metal ions with upto three electrons in the d-orbitals, like Ti⁴⁺(d⁰) ; Ti³⁺(d¹) ; Ti²⁺(d²) ; V²⁺(d³) and Cr³⁺(d³) ; two vacant d-orbitals are available for octahedral hybridisation with 4s and 4p orbitals. The magnetic behaviour of these free ions and their coordination entities is similar. When more than three 3d electrons are present , the required pair of 3d-orbitals for octahedral hybridisation is not directly available (as a consequence of Hund’s rule). Thus for d⁴ (Cr²⁺, Mn³⁺), d⁵(Mn²⁺, Fe³⁺) , d⁶ (Fe²⁺, Co³⁺) cases , a vacant pair of d-orbital results only by pairing of 3d-electrons which leaves respectively two , one and zero electrons unpaired.

The magnetic data agree with maximum spin pairing in many cases, especially with coordination compounds containing d⁶ ions. However, with species containing d⁴ and d⁵ ions there are complications. [Fe(CN)₆]³^- has magnetic moment of single unpaired electron while [FeF₆]³^- has a paramagnetic moment of five unpaired electrons. [CoF₆]³^- is paramagnetic with four unpaired electrons while [Co(C₂O₄)₃]³^- is diamagnetic. This apparaent anomaly is explained by VB theory in terms of inner orbital and outer orbital cordination entities.

sp³d² hybrids , filled with electron pairs donated by 6 F^- ligands

Paramagnetic , outer orbital entity.

The cystal field model is successful in explaining the formation , structures , optical and magnetic properties of coordination compounds to a large extent. However, from the assumption that the ligands are point charges, it follows that anionic ligands should exert the greatest splitting effect. The anionic ligands actually are found at the low end of spectrochemical series. Also OH^-which lies below H₂O and NH₃ in spectrochemical series, produces a greater splitting. These are the weakness of CFT, which can be explained by Ligand Field theory.

LIGAND FIELD THEORY

According to Ligand field theory :

(i) When ligands approach a metal atom or ion a field is created. The field tends to split the degenerate d-orbitals of metal atom into different energy levels.

(ii) The extent of splitting of d-orbitals of the metal atom depends upon the nature and number of ligands which surround it.

(iii) The magnetic and spectroscopic properties of the complexes depend on the extent of the splitting of d-orbitals of the central metal atom.

STABILITY OF CO-ORDINATION COMPOUNDS

The formation of a co-ordination compound involves reaction between a metal ion and ligands. If the force of attraction of the metal ion with ligand is strong, a stable complex may result. Mostly the complex ions are highly stable. However, their possibility of dissociation in aqueous solutions cannot be ruled out completely, though these may be dissociated to a small extent. Thus, a chemical equilibrium may achieved between undissociated complex and dissociated ions. Let us consider an equilibrium reaction between metal ion Ma+ and n ligands (Lx-) to form the complex.

Here a+, x- and b+ are the charges on the metal atoms, ligand and complex respectively. Thus the equilibrium constant for the above reaction is given by the expression :

The above equilibrium constant is called dissociation constant. Consequently, the smaller is the value of K, greater is the stability of complex and vice versa.

The reciprocal of the dissociation constant is called stability constant. The greater the value of stability constant, the greater will be stability of the complex ion.

Here a+, x- and b+ are the charges on the metal atoms, ligand and complex respectively, the reciprocal of the equilibrium constant K’

The magnitude of stability constant gives an indication of the stability of the complex ion in solution. The values of K’ for few complexes in solution are given below :

System K’

Ag+ + 2 NH₃ ⇌ [Ag(NH₃)₂]+ : 1.6 x 107

Cu2+ + 4 NH₃ ⇌ [Cu(NH₃)₄]2+ : 4.5 x 1011

Cu2+ + 4 CN- ⇌ [Cu(CN)₄]2- : 2.0 x 1027

Co3+ + 6 NH₃ ⇌ [Co(NH₃)₆]3+ : 5.0 x 1033

The values indicate that CN- is a stronger ligand than NH₃. The relative stability of various complexes are important in the analytical applications of co-ordination compounds.

Factors affecting the stability of complexes

The stability of the complex depends on :

(i) The nature of the central ion.

(ii) The nature of ligand.

1. Nature of central ion : The term ‘nature’ means the charge density on the central ion i.e., greater the charge density or larger the (charge / radius) ratio, the more is the stability of a complex. For example, out of Fe2+ and Fe3+, those of Fe3+ are more stable. For example,

Fe3+ + 6 CN- ⇌ [Fe(CN)₆ ]3- : K’ = 1.2 x 1031

Fe2+ + 6 CN- ⇌ [Fe(CN)₆]4- : K’ = 1.8 x 106

For the ions which carry the same charge, the one with smaller size gives more stable complex. For example, among Cu2+, Ni2+, Fe2+ complexes as the size of copper is the smallest and thus it gives the most stable complexes.

2. Nature of Ligand : More basic is the ligand, greater is the ease with it can donate its electrons and therefore more is the stability of the complex. For example, complexes having F- ions are more stable than those involving Cl- ions or Br- ions.

Preparation of co-ordination compounds

The co-ordination compounds are generally prepared by the following methods.

1. By substitution reactions : A large number of complexes are prepared by this method. In general , this method of synthesis involves the replacement of water molecules or ligands around the metal ion by some other ligands. For example, an aqueous solution of CuSO₄ on heating with excess of ammonia gives [Cu(NH₃)₄]2+ species having deep blue colour.

[Cu(H₂O)4]2+ + 4 NH₃ ® [Cu(NH₃)₄]2+ + 4 H₂O

deep blue

Some other examples are :

[Ni(H₂O) 6]2+ + 6 NH₃ ® [Ni(NH₃)6]2+ + 6 H₂O

[Ni(H₂O)6]2+ + 3 en ® [Ni(en)₃]2+ + 6 H₂O

[Co(NO₂)₆]3- + 2 en ® [Co(en)₂ (NO₂)₂]+ + 4 NO₂-

(ii) By direct combination : Metal amines can be obtained by direct addition of salt to liquid ammonia. A few examples are :

NiCl₂ + 6 NH₃ ® Ni(NH₃)₆]Cl2

AgCl + 2 NH₃ ® [Ag(NH₃)₂]Cl

CdCl₂ + 4 NH₃ ® [Cd(NH₃)₄]Cl₂

CoCl₃ + 6 NH₃ ® [Co(NH₃)₆]Cl₃

Apart from amines, a number of other compounds can be obtained by similar addition reactions. For example,

PtCl2 + 2 en ® [Pt(en)₂]Cl₂

PtCl₄ + 2 KCl ® K₂ [PtCl₆]

Ni + 4 CO ® Ni(CO)₄

(iii) By redox reaction : The complex [Co(NH₃)₆]Cl₃ can be prepared by the oxidation of an aqueous solution of cobalt chloride made alkaline with ammonia in the presence of NH₄Cl and H₂O₂.

2 CoCl₂ + 2 NH₄Cl + 10 NH₃ + H₂O₂ ® 2 [Co(NH₃)₆]Cl₃+2 H₂O

Applications of co-ordination compounds

1. By complex formation, water insoluble species are brought into solution. For example, red bauxite, Al2O₃ is separated from Fe₂O₃ by heating with conc. NaOH. Al₂O₃ dissolves due to the formation of [Al(OH)₄]- complex ion.

Al₂O₃ (s) + 3 H₂O(ℓ) + 2 OH-(aq) ® 2 [Al(OH)₄]- (aq)

2. In photography, developed film is fixed by washing it with a solution of sodium thiosulphate.

AgBr(s)+2Na₂S₂O₃(aq) ® Na₃[Ag(S₂O₃)₂](aq) + NaBr(aq)

3. The hardness of water is estimated by simple titration against EDTA solution. EDTA forms stable complexes with metal ions present in the hard water. Since stability constants of calcium and magnesium complexes of EDTA are different, even the selective estimation of these ions is possible.

4. Many complexes are used as electrolytes for electroplating. These complexes deliver the metal ions in a controlled manner. For example, for silver plating the complex K[Ag(CN)₂] is used.

5. Silver and gold are extracted from their respective ores by treatment with sodium cyanide solution.

Ag+(aq)+2NaCN(aq) ® Na[Ag(CN)₂](aq) + Na+(aq)

Au+(aq) + 2 NaCN(aq) ® Na[Au(CN)₂](aq) + Na+(aq)

6. Co-ordination compounds also find application in the qualitative as well as quantitative estimation of metal ions. For example, Ni2+ ions are estimated as their red complex with dimethyl glyoxime(DMG). Ca2+ and Mg2+ ions are estimated as their complexes with EDTA.

7. The complex of calcium with EDTA is used to treat lead poisoning. Inside the body calcium in the complex is replaced by lead. The more stable Pb-EDTA complex is eliminated in urine.

8. The platinum complex cis-[Pt(NH₃)₂Cl₂ ] is known as cisplatin is used as anti -tumour agent in the treatment of cancer

9. Many natural compounds, exist as co-ordination complexes. For example, haemoglobin(a complex of Fe2+), chlorophyll (a complex of Mg2+) and vitamin B₁₂ ( a complex of Co2+).

10. Certain complex compounds act as catalysts for different reactions. For example, pentacarbonylcobalt(II) acts as a catalyst in the hydrogenation of alkene.

ORGANOMETALLIC COMPOUNDS

Organometallic compounds are defined as those compounds in which the carbon atoms of organic groups are directly bonded to metal atoms. The compounds of elements such as boron, phosphorus, silicon, germanium and antimony with organic groups are also included in organometallics. By convention metal cyanides possessing an M-C bond are not included in organometallics. Many organic compounds are important reagents which are used for the synthesis of organic compounds.

Classification of organometallic compounds

Two broad divisions of organometallics are :

(i) Main group organometallics

(ii) d- and f-block organometallics.

Main Group Organometallics

E.C Frankland in 1848 synthesised dimethyl zinc (CH₃)₂Zn . He has also prepared Zn(C₂H₅)₂, Hg(CH₃)₂, Sn(CH₃)₄ and B(CH₃)₃. He has also introduced the term ’organometallic’ in the language of chemistry. Organometallic compounds of Li,Mg,B,Al and Si are compounds of industrial importance.

The s- and p-block organometallic are named according to the substituent names used in organic chemistry. For example, methyl lithium for CH₃Li and trimethyl boron for B(CH₃)₃ which is also called trimethylborane, taking it to be a derivative of its hydrogen counterpart. Thus Si(CH₃)₄ and As(CH₃)₃ are tetramethylsilane and trimethylarsane respectively.

The oxidation number of the metallic element in an organometallic compound is based on the organic moiety being considered to be anionic. For example Zn(CH₃)₂ , CH₃ group is taken to be negatively charged (1-) . Thus, oxidation number of zinc is 2+. The bond in alkyls of s-block elements is highly polar (Md+- Cd-). In organometallic compounds of groups 14, 15 and 16, the M -C bonds are relatively low polarity. Methyl compounds of Li, Na, Be, Mg and Al are associated through alkyl bridges and multicentre two-electron bonds. The structures of some representative main group organometallics are shown in Fig.

Structures of some representative group organometallic compounds

Organometallic compounds of electropositive elements are strong reducing agents. They are pyrophoric and ignite spontaneously in air.

d- and f-block Organometallic compounds

The first organometallic compound of a d-block element (platinum) was prepared by W.C. Zeise in 1827. The compound , trichloroentheneplatinate(II), [PtCl₃(C₂H₄)]-, has the structure shown below.

In this structure, the three chlorine and the middle point of ethylene double bond form a square plane. The platinum atom is present in the centre of the square and C=C double bond of ethylene molecule is perpendicular to the plane containing Pt and Cl atoms.

Tetracarbonyl nickel [NiCO)₄] was synthesised by Mond and Quinke in 1899. The synthesis of highly stable bis(cyclopentadienyl iron) , Ferrocene in 1951 was a land mark in the advancement of modern organometallic chemistry. For explaining the stability, structure and bonding of ferrocene, the new concept of p-bonding of carbocyclic rings to metal atoms was invoked. This led to the preparation of a large number of organometallic compounds containing benzene and other carbocyclic ligands. Collectively these compounds are called metallocenes.

The so called sandwich structure of ferrocene was confirmed with IR spectrum and X-ray diffraction analysis. The structure of Ferrocene is regarded as Sandwich structure in which the iron atom is sandwiched between two C₅H₅ organic rings. The planes of the rings are parallel so that all the carbon atoms are at the same distance from the iron atom. Dibenzene chromium also has sandwich structure. Its structure is shown below.

Walter Kaminsky and Hans Brintzinger showed that metallocenes are highly specific for homogeneous catalysis processes like carbonylation, hydrogenation and polymerization.

The first f-block organometallic compound [ThH(OR)(C₅Me₅)₂] was prepared in 1970s. Pentamethylcyclopentadienyl ligand , C₅H₅Me₅, forms stable f-block compounds.

METAL CARBONYLS

A widely studied and important class of organometallic compounds is that of metal carbonyls. The homoleptic carbonyls (compounds containing carbonyl ligands only) are formed by most of the transition metals(d-metals). The metals constituting the central part of d-block form stable , neutral binary carbonyls like : [V(CO)₆], [Cr(CO)₆], [Mo(CO)₆], [W(CO)₆], [Mn₂(CO)₁₀], [Fe(CO)₅], [Fe₂(CO)₉], [Co₂(CO)₈], [Co₄(CO)₁₂], [Ni(CO)₄], etc. Outside the central part of d-block , the metal carbonyls are unusally unstable.

Metal Carbonyls – Structure and Bonding

Homoleptic binary carbonyls have simple , well defined structures. Tetracarbonylnickel(0) is trigonal bipyramidal while hexacarbonylchromium(0) is octahedral. Decacarbonyldimanganese(0) is made up of two square pyramidal Mn(CO)₅ units joined by a Mn – Mn bond. Octacarbonyldicobalt(0) has a Co–Co bond bridgeb by two CO groups.

Structures of some representative homoleptic metal carbonyls.

Metal carbonyls are mostly solids at room temperature and pressure. Exceptions being iron and nickel carbonyls which are liquids. The mononuclear carbonyls are volatile and toxic. With exception of F₂(CO)₉, ennacarbonyldiiron(0), carbonyls are soluble in hydrocarbon solvents. Mononuclear carbonyls are either colouless or light coloured. For example, Fe(CO)₅ is light straw coloured liquid. Polynuclear carbonyls are deeply coloured. Fe₃(CO)₁₂, dodecacarbonyltriiron(0) , for example, is deep grass green solid. The reactivity of metal carbonyls is due to:

(a) the metal centre of the carbonyl, and (b) the CO ligands.

Metal carbonyls find use as industrial catalysts and as precursors in organic synthesis.

BONDING IN ORGANOMETALLIC COMPOUNDS

Bonding in Metal carbonyls

The bonding in metal carbonyls involves the following steps.

(i) There is a donation of lone pair of electrons of carbon(of CO) into suitable empty orbital of the metal atom. There is a dative overlap and forms a sigma M¬ C bond.

(ii) There is a p-overlap involving donation of electrons from filled metal d-orbitals into vacant antibonding p-molecular orbitals of CO. This results into the formation of M® C , p-bond. This is also called back bonding or back donation.

The bonding in metal carbonyl is shown in Fig.

Bonding in metal carbonyls

Representation of bonding in alkene complexes

Bonding in Alkene complexes

The bonding in alkene involves the following steps :

(i) The p-electron density of alkene is donated into vacant orbital of metal atom forming alkene®M bond.

(ii) The back-bonding resulting from the donation of filled d-orbitals on metal into the vacant anti-bonding orbitals of alkene. As a result of back-bonding, the double bond character of alkene is retained, though the bond becomes weak. The bonding is shown in Fig.

Synthesis of Organometallic Compounds

(i) Grignard reagents are prepared by slow addition of organic halide to stirred suspension magnesium turnings in appropriate solvent in absence of air and moisture.

The reactivity of alkyl halides is :

RI > RBr > RCl

and alkyl is more reactive than aryl.

(ii) Tetra alkyl organometallics are prepared by treating metal halides with organometallic reagents.

Preparation of Metal carbonyls and p-complexes

(i) Metal carbonyls are prepared by the direct interaction of finely divided metal and carbonmonoxide.

323K

Ni + 4 CO ® Ni(CO)₄

(ii) Zeise’s salt is prepared by the replacement of less stronly bound ligand such as halide ion by alkene by directly heating ethylene and [PtCl₄]2-.

H₂C=CH₂ + K₂ [PtCl₄] ® K[(C₂H₄)PtCl₃] + KCl

Ziese’s salt

(iii) Ferrocene is prepared by the following reaction.

2 C₅H₅MgBr + FeCl₂ ® [(C₅H₅)₂Fe] + 2 MgBrCl

(iv) Dibenzene chromium is obtained by heating benzene with chromium vapour.

2 C₆H₆ + Cr(vapour) ® (C₆H₆)₂Cr

Applications of Organometallic compounds

1. The soluble organometallic complexes of transition metals act as homogeneous catalysts. Some examples are :

(a) Selective hydrogenation of certain double bonds using Wilkinson’s catalyst, (Ph₃P)₃ RhCl

(b) (Et₃P)₂ NiCl₂ acts as catalyst for isomerisation of ethylene and other alkenes.

2. Silicone rubbers because of their high thermal stability, resistance to oxidation and chemical attack are used in modern surgery for the purpose of production of artifical body parts.

3. Tetraethyl lead is used as an antiknock compound in gasoline.

4. A number of organometallics also find application in agriculture. For example, ethyl mercury chloride C₂H₅HgCl is used as a fungicide for the protection of young plants and seeds against fungal infection.

5. Organometallic compounds of magnesium (RMgX) , cadmium (R₂Cd) and lithium (RLi) are extensively used in organic synthesis.

6. Aryl arsenic compounds are used as chemotherapeutic agents.

7. Organometallic compounds can also act as hetergeneous catalysts. Zeigler-Natta catalyst (a solution of TiCl₄ containing triethyl aluminium) for polymerisation of ethylene and other alkenes.

8. Purification of Nickel, free from other metals, particularly cobalt has been obtained by the reaction with carbon monoxide at 323 K to Ni(CO)₄. At this temperature no other metal forms a carbonyl . Ni(CO)₄ is then decomposed at 473 K to get very pure form of nickel.

QUESTIONS

1. Explain the terms : (i) Co-ordination compound (ii) Ligand (iii) Co-ordination number and (iv) Complexation

2. Name (i) one neutral monodentate (ii) one neutral bidentate (ii) one anionic monodentate and (iv) one cataionic monodentate ligand.

3. What is a chelate ? Give an example

4. A cyclic complex is more stable than an open one. Substantiate your answer with an example.

5. How is a double salt distinguished from a complex compound ? Explain taking a suitable example.

6. Give

7. the rules of nomenclature of complex compounds.

8. How do you arrive at the formula of a given named complex ?

9. Describe briefly any two methods which are employed in the preparation of co-ordination compounds.

10. Explain with suitable example geometrical isomerism exhibited by co-ordinated compounds.

11. Explain with an example ionisation isomerism exhibited by co-ordination compounds.

12. Give a chemical test to distinguish between [Co(NH₃)₅Br]SO₄ and [Co(NH₃)₅ SO₄]Br. Name the type of isomerism exhibited by these compounds.

13. Explain how [Pb(NH₃)₂Cl₂] and [Pb(NH₃)₆]Cl₄ will differ in their electrolytic conductances ?

14. Write notes on (a) Linkage isomerism exhibited by complexes (b) Optical isomerism exhibited by co-ordination compounds.

15. Write all the isomers of : (a) [Co(NH₃)₄Cl₂]+ (b) [Cr(Ox)₃]3-

16. What is meant by co-ordinate isomerism ?

17. Write all isomers of : (i) [Co(NH₃)₆][Cr(C₂O4)₃] (ii) [Co(NH₃)₆SCN]Cl₂ (iii) Pt(SCN)(NH₃)₃](SCN)

18. Give the postulates of Werner’s theory of co-ordination compounds.

19. Square planar complexes with co-ordination number of four exhibit geometrical isomerism; whereas tetrahedral complexes do not. Why ?

20. Write a note on Pauling’s valence bond scheme on bonding in co-ordination complex.

21. Mention the geometrical shapes attained by the following types of hybrid orbitals (a) sp3 (b) dsp2 ( c) d2sp3 . Give an example each for each of the above.

22. Explain the geometry of Ni(CO)₄ by valence bond theory. Why is this ion not paramagnetic ?

23. Explain the magnetic property and geometry of [NiCl₄]2- ion on the basis of valence bond theory .

24. Using valence bond theory, explain the diamagnetic nature and square planar structure of [Ni(CN)₄]2- ion..

25. Explain the geometry of [Cr(NH₃)₆]3+ ion by valence bond theory of complexes. Why is the ion paramagnetic ?

26. Using valence bond theory, deduce the structure of [Ni(NH₃)₆]2+ . Is it paramagnetic or diamagnetic ?

27. Using valence bond theory of complexes , explain the geometry and diamagnetic nature of ion [Co(NH₃)₆]3+ ion.

28. Using valence bond theory, explain why [Pt(NH₃)₂Cl₂] molecule is square planar ?

29. [NiCl₄]2- is paramagnetic ; while [Ni(CO)₄] is diamagnetic though both are tetrahedral. Why ?

30. Explain the structure of [CoF₆]3- ion on the basis of valence bond theory.

31. Using valence bond theory, predict the magnetic nature and character of [Co(CN)₆]3- ion.

32. On the basis of valence bond theory, explain why is [Fe(CN)₆]3- weakly paramagnetic ; while [Fe(CN)₆]4- is diamagnetic ?

33. [Ni(CO)₄] possesses tetrahedral geometry ; while Pt(NH₃)₄Cl₂ is square planar. Why ?

34. [Ti(H₂O)₆]2+ is coloured ; while [Sc(H₂O)₆]3+ is colourless. Account for it.

35. What do you understand by the term stability constant of a complex ?

36. The value of K for [Cu(NH₃)₄]2+ is 4.5 x 1011 and for [Cu(CN)₄]2- is 2.0 x 1027, suggest (a) Which complex species will furnish less Cu2+ ions in solution and (b) Which out of NH3 and CN- is stronger ligand ?

37. What are the factors which affect the stability of complexes ?

38. Co(II) is stable in aqueous solution, but in the presence of strong ligands and air , it gets oxidised to Co(III). Account for it.

39. Mention the application of co-ordination compounds in the following areas, giving an example for each. : (a) analytical chemistry (b) extraction of metals.

40. What is the importance of co-ordination compounds in industry and chemotherapy ?

41. What is meant by hexadentate ligand ? Give its example. How is such a ligand useful for measuring hardness of water ?

42. Define organometallic compounds.

43. Write the structures of each of the following : (a) Trimethyl aluminium (b) Fe(CO)₅ ( c) [Co(NH₃)₆]2- (d) [NiCl₄]2- (e) ferrocene.

44. How are organometallic compounds classsified ? Write a chemical equation for the preparation of Zeise’s salt.

45. Describe the bonding of an alkene to a transition metal.

46. Briefly describe the nature of bonding in metal carbonyls.

47. Describe one method of preparation for each of the following :

(a) Tertiary butyl tin (b) Tetraethyl lead

(c ) n-Butyl lithium (d) Ferrocene.

48. Mention the application of organometallic compounds in the following areas :

(a) Homogeneous catalysis

(b) Heterogeneous catalysis.

49. Describe one method for the preparation of p-complexes.

50. What are s-bonded organometallic compounds ?

UNIT 11 CO-ORDINATION COMPOUNDS AND ORGANOMETALLICS

Some common unidentate ligands

Neutal molecule

Structure Mode of ionisation

VALENCE BOND THEORY

ORGANOMETALLIC COMPOUNDS

METAL CARBONYLS

Bonding in Alkene complexes

Synthesis of Organometallic Compounds

QUESTIONS

PLUS TWO

QUESTIONS