Introduction to Co-ordination Chemistry: Fundamentals and Applications

Chapter 

1…Co-ordination Chemistry



Contents 

…

1.1 

Introduction-Definition and Formation of Co-ordinate Covalent

        Bond in BF

3 

- NH

3

, [NH

4

]

+

, and H

2

O



1.2

Distinguish between Double Salt and Complex Salt

1.3     Werner’s Theory

1.3.1

 Postulates

1.3.2 The Theory as Applied to Cobalt Ammines viz. CoCl

3

.6H

2

O,

           CoCl

3

.5H

2

O, CoCl

3

.4H

2

O, CoCl

3

.3H

2

O

1.4 

Description of the Terms - Ligands, Co-ordination number,

Co-

         ordination sphere, Effective Atomic Number (EAN)

1.5 

IUPAC Nomenclature of Co-ordination Compounds

1.6

  Isomerism in complexes with C.N. 4 and 6

1.6.1

Geometrical Isomerism

1.6.2 

Optical Isomerism

1.6.3 

Structural Isomerism-Ionisation Isomerism, Hydrate

          Isomerism, Coordination Isomerism, Linkage Isomerism and Co-ordination

          position Isomerism

1.7     Valence Bond Theory of transition metal complex with respect

          to C.N.

=

 4 complexes of Cu, Ni, and C. N.

 =

 6 complexes of Fe, Co

Introduction



The most amazing field of modern inorganic chemistry is



 co-ordination chemistry! It pertains to the study of complexity of

compounds.



A large body of current inorganic research is going on this extremely

attractive field of science which is in the state of rapid



progress. It plays an important role in the chemical industry and in life

processes.



•

The recent study of co-ordination chemistry stems from the work of Alfred

Werner and Sophus Mads Jorgensen.  

Alfred Werner was the first inorganic

chemist who was awarded the Nobel Prize in chemistry was awarded jointly

to Dr. K. Ziegler and Prof. G. Natta for the Ziegler-Natta catalyst, a complex of

aluminium and titanium. The Nobel Prize for chemistry was awarded to

Prof. Wilkinson in 1973 for his work on 



-complexes

. This confirms the

importance of co-ordination chemistry.



•

In this chapter, students will become familiar to the co-ordination

bond and insight of co-ordination chemistry

1.1

Definition and Formation of Co-ordinate Covalent Bond in BF

3 

-

NH

3

 , [NH

4

]* and H

2

O:-

Co-ordinate covalent bond

1. 

Definition

Chemical bond:

The force of attraction between the atoms within a molecule is known as chemical

bond.

Covalent bond:

The force developed by an equal sharing of a pair of electrons between the two atoms

      is known as an electron pair bond or covalent bond.

Co-ordinate bond:

•

According to Lewis, a special kind of electron pair bond formed when both electrons

are contributed by one atom only is known as a co-ordinate covalent band or simply 

a

co-ordinate bond.

•

According to the orbital concept, a co-ordinate bond is formed when an orbital

containing a lone pair of electrons in one atom overlaps with an empty orbital from

the other atom.

Specifications:

(i) 

The co-ordinate bond once formed is indistinguishable from the normal covalent

bond.

(ii)

 The atoms which donate the pair of electron are called the donor atom and the

         other atom which simply accepts the

Formation of Co-ordinate (Covalent) Bond:

The mechanism of formation of a co-ordinate bond may be visualized

as a two step process which is an amalgamation of electrovalent and

covalent bindings.



Suppose A is a donor and B the acceptor.

Step I:

At the initial stage, one electron from the lone pair on A is

transferred to B whereby gains a unit positive charge while B receives a

unit negative charge. In other words, this is a process of transfer of

electrons as observed in ionic bonding.

(with a lone pair of electron)

Step II:

In the second stage, these two electrons one each from A

+

 and B

−

join together and form a pair of electrons which is shared between A

and B just similar to covalent bonding.

A

+

. and . B

+

A  :  B  or  A  B

Co-ordination bond is thus the joint effect of ionic and

covalent bondings.

The co-ordinate bond is therefore termed as co-ionic bond, semi-polar

bond or dative covalent bond.

Conditions for formation of co-ordinate bond:

1.

The atom acting as a donor should have a lone pair of electrons.

2. 

The atom acting as an acceptor should have a vacant orbital to accept the lone pair

of electrons donated by the donor.

3. 

Donor with its lone pair of electrons should overlap appreciably with vacant orbital

of acceptor and should donate and share its electron pair with acceptor.



Examples of Co-ordinate Covalent Bond:

1. Formation of BF

3

:NH

3

 Molecule:

Boron trifluoride reacts with ammonia where a co-ordinate bond is formed

between nitrogen and boron. Here nitrogen from ammonia acts as a 

donor

 atom while

boron from boron trifluoride acts as an acceptor atom. Nitrogen donates and shares

an electron pair with boron and forms a co-ordinate bond.



2. 

Formation of Ammonium Ion, NH

4

+

:



In ammonia molecule, the central atom is nitrogen. It has four    sp

3

orbitals, one of these contains a lone pair of electrons while the three contain

bond pairs as shown in Fig. 1.1. The empty s-orbital from an H

+

 ion overlaps

with the orbital of nitrogen atom containing lone pair of electrons and forms

ammonium ions as shown in Fig. 1.1.(a)

Fig. 1.1(a) Formation of co-ordinate bond in ammonium ion



3.

 Formation of Water molecule [H

2

O]:



In water molecule, the central atom is oxygen. The electronic

configuration of oxygen is, 1s

2

, 2s

2

, 2p

4

.  It has four    sp

3

 orbitals, two

of these contains a lone pair of electrons while the two contain bond

pairs as shown in Fig. 1.1.(b) Here, two loan pairs occupy more volume

in space, since they are under the influence of only one nucleus. Hence,

loan pair–loan pair repulsion is greater than loan pair–bond pair

repulsion which is greater than bond pair–bond pair repulsion.

Consequently, expected bond angle 109.5

0

 is reduced to 104.5

0

. Water

molecule has distorted tetrahedral structure as shown in fig.1.1(b)

Fig. 1.1(b) Structure of water molecule



1.2 Distinguish between Double Salt and Complex Salt





3 Werner’s Theory

(Werner’s concept of Auxiliary Valence):



1.3.1Postulates

(Basic Postulates of Werner’s Theory)



The basic postulates of Werner’s theory may be given as:



1. 

Most of the metallic elements exhibit two types of valencies.



(i) Principal or primary valence and (ii) Secondary valence.



2. 

Every metal tends to satisfy both of its primary and secondary

valencies.



3. 

Every metal has a fixed number of secondary valencies.



4. 

The secondary valence its always directed towards fixed

position in space.



Salient Features or Postulates:



1.

The principal valence also known as 

primary valence or

ionisable valence

 is designated by solid line ( ). In modern terms, it

corresponds to the oxidation state of the metal ion.



The secondary valence also called as 

auxiliary valence, subsidiary

valence

, residual valence or non-ionisable valence, is designated by

dashed line (---). In recent terms, it corresponds to the co-ordination

number of the metal ion, i.e. CN.



2. (i) 

The primary valencies are those which a metal ion exercises towards the

negative groups so that its normal charge is satisfied by the formation of simple

salts.



For example: 

Primary valencies of Pt, Co, Cu, Ag are 4, 3, 2, 1 respectively whereby

they produce simple salts viz. PtCl

4

, CoCl

3

, CuCl

2

 and AgCl respectively.



(i)The secondary valencies are those which may be satisfied by negative ions,

neutral molecules or both so that complex species is formed.



For example:

 Co

3+

 has secondary valence six. Hence Co

3+

 forms complexes such as

[Co(NH

3

)

6

]

3+

,  [Co(NH

3

)

5

Cl]

2+

,  [Co(NH

3

)

4

Cl

2

]

+

,  [Co(NH

3

)

3

Cl

3

], etc.



Sometimes negative groups show a dual behavior. They may satisfy primary as well

as secondary valencies . It is invariably found that when a negative group satisfies

one secondary valence, it necessarily satisfies one primary valence also.



3.  Every element is characterized by its fixed secondary valence.



For example: 

Fe(II), Co(III) and Pt(IV) have CN = 6, Cu(II) has          CN = 4, Ag(I)

has CN = 2, etc.

4.  The primary  valances  

are non-rigid and non-directional. On the other hand,

secondary  valencies have directional properties. The geometry of the complex is

determined by the 

number and position of the secondary valence in space. 

This

postulate, therefore, decides the stereoisomerism in complexes.



If a metal ion has a secondary valence four, the geometry of the complex can be

either 

square planar or tetrahedral

. On the other hand, if the metal ion has secondary

valence six, the geometry of the complex would be 

octahedral

.

1.3.2

Werner’s Theory in the Light of Modern Electronic Theory of

Valence :



According to Werner, the combining capacity of a metal is considered in the form

of two spheres of attraction or two zones of attraction.

1.

The first or the inner zone surrounding the metal is called as 

co-

ordination sphere

. Neutral molecules or negatively charged ligands are

linked by co-ordinate bonds to the central metal ion in this zone. As the

linking is rigid and tight, this sphere functions as a nonionizable zone.

The secondary valence plays its role in this zone. In terms of

electronic theory of valence, the secondary valence is identified as    co-

ordination number of the metal ion.

2.

 The second or the outer zone, surrounding the co-ordination zone is

called as 

ionization sphere

. In this zone, the metal ion satisfies its primary

valence by linking with negatively charged ions which get ionized on

dissolving the complex in water. Hence the primary valence is identified

as the ionizable valence or electrovalency according to modern theory.

To distinguish between the two spheres of attraction, Werner introduced

the square brackets [   ], to enclose the co-ordination sphere of the

complex compound. Schematically, the two spheres of 

attraction for the

metal ion forming the complex may be as shown in Fig. 1.2.

Fig. 1.2: Two spheres of attraction of a metal

1.3.3

Application of Werner’s Theory to Cobalt Ammine

1.3.3

Application of Werner’s Theory to Cobalt Ammine

(The theory as applied to cobalt ammines viz. CoCl

3

.6H

2

O, CoCl

3

.5H

2

O,

CoCl

3

.4H

2

O, CoCl

3

. 3H

2

O).

Experimental Observations of Werner’s Theory:

In all these co-ordination compounds, according to

Werner:

1. The primary valence (or oxidation state) of cobalt is 3

while secondary valence (or co-ordination number) is 6.

2. When cobalt ammines are treated with hydrochloric

acid and heated at 100°C (373 K), ammonia is not

removed.Because it is bounded by co-ordinate bond.

3. All these compounds are found to differ in their

electrical conductivity.

4. All these complexes differ in their reactivity towards

silver nitrate. With these guidelines, let us study the

individual ammines one by one.



(i) 

CoCl

3

.6NH

3

 or [Co(NH

3

)

6

]Cl

3

:



Thus CoCl

3

.6NH

3

 may be pictured as shown in Fig. 3.3 (A). Here solid one represents primary valence

of cobalt while hollow or broken line represents secondary valence.



1. On treatment with AgNO

3

:



CoCl

3

.6NH

3

 immediately gives precipitate on silver chloride corresponding to all the three chloride ions.

Hence all the three chlorides must be in the ionization sphere and satisfy the primary valence of Co.



CoCl

3

.6NH

3

 + Excess Ag

+



 3AgCl 





2.

On treatment with HCl:



As ammonia is not removed, it must be strongly bound to cobalt. Hence all the six ammonias must be in

the co-ordination sphere of the complex and they must be satisfying secondary valence of cobalt.



3.

Cryoscopic measurement

and molar conductivity data:



CoCl

3

.6NH

3

 indicates the presence of 

four particles

. Corresponds to 

six charges

 for the solution of

CoCl

3

.6NH

3

 in nitrobenzene.



From these observations, formulation of CoCl

3

. 6NH

3

 must be done as [Co(NH

3

)

6

]Cl

3

. The ionization

would be then



               [Co(NH

3

)

6

]Cl

3

                            [Co(NH

3

)

6

]

3+

 + 3Cl−



(ii) 

CoCl

3

.5NH

3

 or [Co(NH

3

)

5

Cl]Cl

2

:



From the above discussion one may predict that CoCl

3

.5NH

3

 should be

formulated as [Co(NH

3

)

5

Cl]Cl

2

 and represented structurally as in Fig. 1.3 (B), in

accordance with Werner’s theory.



The observed properties which support the above formulation are

1.

On treatment with AgNO

3

:

It gives precipitate of AgCl corresponding to two chlorides which must be in the ionization

sphere.

CoCl

3

.5NH

3

 + Excess Ag

+



 2 AgCl 



2. 

On treatment with HCl:

All the ammonias are strongly bound and are not removed on treatment with HCl. They must be

linked to cobalt with strong bonds and must be in co-ordination spheres.

3

.

Cryoscopic measurements and molar conductivity data:

There must be three ions i.e. 

three particles and four charges 

for CoCl

3

.5NH

3

. It is evident from its

ionization.

[Co(NH

3

)

5

Cl]Cl

2

                           [Co(NH

3

)

5

Cl]

2+

 + 2Cl−



(iii)

CoCl

3

.4NH

3

 or [Co(NH

3

)

4

Cl

2

]Cl:



This compound should be formulated as [Co(NH

3

)

4

Cl

2

]Cl and to be represented as shown in

Fig. 1.3 (C), in accordance with Werner’s theory.



 The observed properties which support the above formulation are



1. On treatment with AgNO

3

:



It gives precipitate of AgCl corresponding to one chloride which must be in the

ionization sphere.

                            MCoCl

3

.4NH

3

 + excess Ag

+



        AgCl 





2.

On treatment with HCl:



All the ammonias are strongly bound and are not removed on treatment with HCl.

They must be linked to cobalt with strong bonds and must be in co-ordination spheres.



3. Cryoscopic measurements and molar conductivity data:



Hence all four NH

3 

molecules are strongly bound to cobalt and must be in the co-

ordination sphere. Its ionization leads to 

two ions and two particles

 as shown below:



                      [Co(NH

3

)

4

Cl

2

]Cl                    [Co(NH

3

)

4

Cl

2

]

+

 + Cl−



(iv)

CoCl

3

.3NH

3

 or [Co(NH

3

)

4

Cl

3

]:



This compound should be formulated as [Co(NH

3

)

3

Cl

3

] where all the

three chloride ions are present in the co-ordination sphere. There is no

evolution of NH

3

 on treatment with HCl or no precipitate of AgCl with

AgNO

3

. The complex does not ionize. It is uncharged, neutral molecular

species and exists as a 

single particle

. This can be represented structurally as

shown in Fig. 1.3 (D).



The observed features of cobalt ammines may be summarized in Table 1.2.

Table 1.2: Properties of Cobalt (III) ammines



1.4

Description of the Terms - Ligands, Co-ordination

Compounds:



i) Ligand :



Definition:

In a complex compound, central metal ion is surrounded by or

directly attached to other atoms, ions or molecules. These surrounding groups are

called as the ligands. (The word stems from Greek, ligare meaning to bind).



For example

, NH

3

 in [Co(NH

3

)

6

]

3+

, is the ligand which is co-ordinated to

Co

3+

 to form the complex. Hence ligand is also called as a co-ordinating group or a

complexing agent.The Werner’s formulation is self-explanatory to account for

ligand

Types:

 The ligands may be atomic, ionic or molecular species.

For example:

1. 

Neutral molecular ligands

2. 

Negatively charged ligands.

3. 

Positively charged ligands.

Table 1.3: Different types of ligands with examples



Some typical illustrations have been given in Table 3.5. All these are called chelate

ligands or chelating agents.

Table 1.5

Ethylenediamine:        H

2

N*–CH

2

–CH

2

–N*H

2

Diethylenetriamine:    H

2

N*–(CH

2

)

2

 – N*H – (CH

2

)

2

 –N*H

2

Triethylenetetramine: H

2

N*–(CH

2

)

2

–N*H–(CH

2

)

2

–N*H–(CH

2

)

2

–N*H

2

where starred (*) atoms are the donor atoms.

5.

Many of the properties of the co-ordination compounds are governed by the charge on

the ligand and the number and nature of the donor groups on the ligand. Based on the

co-ordinating ability, the common ligands may be arranged as follows:

CO, CN

−

 > NO

> en > NH

3

 > NCS

−

 > H

2

O > F

−

 > RCOO

−

 > OH

−

 > Cl

−

> Br

−

 > I

−

           Tsuchida – Fajans Spectrochemical Series.

         >------------ Co-ordinating ability decreases

6. 

The ligands are arranged round the central metal ion inside the first or the inner sphere of

attraction, essentially in a regular geometrical pattern such as linear, triangular, planar, tetrahedral,

square planar, trigonal bipyramidal, square based pyramidal and octahedral. The ligands play the most

important role in deciding the features of the co-ordination compounds.



ii) Co-ordination Number (CN):



The maximum number of groups directly attached to the central metal atom/ion, in an co-

ordination compound, is known as the  co-ordination number of the metal atom/ion.



It is often abbreviated by the symbol (CN). It is the characteristic property of the

metal. It takes the values from 2 to 8 where 4 and 6



co-ordination numbers are very common. It has been noted that light transition metal

prefers to exhibit 

CN = 4 and 6

 while the heavier transition metal prefers higher 

CN, say

8

. Usually a given metal ion has a fixed CN but some metal ions may exhibit more than

one co-ordination number.



The study of CN has revealed that normally the co-ordination number of metal ion is

twice the oxidation state of the metal. All such features of CN can be understood by

studying the Table 3.6.

Table 1.6: Metal Ions and Co-ordination Numbers



iii) Co-ordination Sphere (Non-Ionizable Sphere)



“The central metal ion and the groups (ligands) directly attached to it in

a complex compound, are said to constitute the co-ordination sphere.



This is the first or inner sphere of attraction where ligands are joined to

metal ion by co-ordinate bonds. Hence on dissolving in water ligands are not

ionized. Hence it is called co-ordination sphere or non-ionizable sphere. To

differentiate this inner sphere of attraction from the outer ionization sphere,

Alfred Werner used brackets to enclose co-ordination sphere (See Fig. 1.4).

Fig. 1.4: Co-ordination Sphere (Nonionizable Sphere)

Complexes can retain their molecular identity in solution; it is due to the presence of

co-ordination sphere.In the complex compound [Co(NH

3

)

5

Cl]Cl

2

, the metal ion Co

3+

,

the five ammonias and one chloride attached to it, constitute the           co-ordination

sphere, while the two chloride ions bonded ironically to cobalt ion remain outside

the co-ordination sphere. Werner’s use of brackets was originally as [CoCl(NH

3

)

5

]Cl

2

instead of [Co(NH

3

)

5

Cl]Cl

2

.Parallel explanation may be given for the complexes such

as [Cu(NH

3

)

4

]SO

4

, K

4

[Fe(CN)

6

], [Fe(H

2

O)

6

]Cl

3

, [Pt(NH

3

)

2

Cl

2

], etc.



iv) Effective Atomic Number (EAN)



In 1972, the first attempt was made by Sidgwick to make the Werner’s

concept compatible with the electronic theory of valency. According to

Sidgwick, Werner’s primary valence was regarded as amounting to electron

transfer whereas his secondary valence to  electron pair sharing. In complex

formation, an electron pair is donated by the ligand to the metal ion and

partially shared. This led Sidgwick to introduce the idea of co-ordinate covalent

bonding.



Based on Lewis concept of electron pair bond, Sidgwick tried to clarify

the idea of secondary valence by putting forward the concept of Effective

Atomic Number (EAN). He explained that the metals are engaged in complex

formation where the central metal ion tries to get surrounded by the necessary

number of ligands so that the next inert gas in the periodic table, say 36 (Kr),

54 (Xe), 86 (Rn), etc. Hence EAN may be defined as: The total number of

electrons associated with the central metal ion in a co-ordination compound is

called the effective atomic number of metal.



EAN = [(Atomic Number of metal) − (Number of electrons lost in the

formation of metal ion) + (Number of electrons gained from ligands)]



For example

, in the complex [Co(NH

3

)

6

]

3+



EAN of cobalt = [27 − 3 + 2 



 6]



               = [24 + 12]



               = 36 [Kr]



In some complexes such as metal carbonyls, the ligand act both as electron pair

donor and acceptor

(M             CO   )

 e.g. [Ni(CO)

4

], [Fe(CO)

5

], H

3

N : B(CH

3

)

3

, etc.



1.5

IUPAC Nomenclature of Co-ordination Compound



The Nomenclature Rules

:



1. 

Writing the Formula:



While writing the formula of a complex, the complex entity        (i.e. the co-

ordination sphere) should be enclosed by square brackets. However, the oxoanions and

neutral complexes such as MnO,          Ni(CO)

4

, etc. are often written without brackets.



In the formula, the symbol of metal is placed first, followed by the ligands,

alphabetically in the order: 

negative, neutral and positive

 by using their symbols.



For example



K

4

[Fe(CN)

6

], [Cr(en)

3

]Cl

3

, [Co(NO

2

)

3

(NH

3

)

3

], [Pt(py)

4

] [PtCl

4

]



2. Listing the ions:



In ionic compound, the positive ion is named first followed by the negative ion. The

naming of ions must be in full, without the indication of the number of either type but

separated from one another by a space.



NaCl: Sodium chloride



[Cu(NH

3

)

4

]SO

4

: Tetramminecopper(II) sulphate



3. Naming the Molecular Complexes:



The non-ionic or molecular complexes are given a one-word name.

[Co(NO

2

)

3

(NH

3

)

2

] Triamminetrinitrocobalt (III).



4. Naming the Ligands:



(i)

Order: 

While naming the complex, the ligands are quoted in alphabetical

order, regardless of their number or nature, (hence intermingled), followed by the name of

the metal, without any gap or hyphens.  [CrCl

3

(NH

3

)

3

] Triamminetrichlorochromium (III).



(ii)

Endings:



(1) 

The anionic ligands, whether organic or inorganic,

have names ending in “o”. 

For example

:



(a)  

The ending with

 –

ide changes simply to –“o”.



(b)The ending with 

–ide, -ite or –ate changes to –ido, -ito or.-ato.

(2)

Neutral ligands have no special endings

;

they are named as the molecule. Some exceptions are given below.

Water and ammonia are neutral ligands. In complexes, they are called

aqua (formerly aquo) 

and ammine

 (NH

3

 – term with two m’s; NH

2

 – amine

term with single ‘m’) respectively. The groups NO and CO when act as the

neutral ligands, they are called 

nitrosyl

 and 

carbonyl 

respectively.



(3) Positive ligands end in – 

ium

.



−NH

2

NH

3

+ 

(Hydrazinium), NO

+ 

(Nitrosylium), [Refer: NH   Ammonium, H

3

O

+

 (Hydronium) etc.]



(4) Organic free radicals or organic neutral molecules

are referred by their own names. For such

groups abbreviated short forms are used in complexes.



For example:       −CH

3

: Methyl (Me): - C

2

H

5

: Ethyl (Et)



C

5

H

5

N:  Pyridine (py); Thiourea (tu)



NH

2

.CH

2

.CH

2

.NH

2

: Ethylenediamine (en);



NH

2

-CH

2

-.COO- Glycine ion (gly);



Ethylenediaminetetra acetic acid (EDTA) as (edta).



[F−, Cl−, Br

−

, I−, Halide: (X); Unspecified Ligand: (L), etc.]



(iii) Number of Ligands:



When the number of ligands of one particular type is greater than one, the prefixes di-, tri-, tetra-, penta-, hexa-

are used for the ligands 2, 3, 4, 5, 6, etc. respectively.



If the ligand has a complicated structure or if it is a chelate ligands or if its expression

includes the prefixes di-, tri-, etc. the prefixes bis-(for two), tris-(for three), tetrakis-(for four),

pentakis-(for five) etc. are used and name of the ligand is enclosed in parentheses. 

For example

,

[Ru(NH

3

)

5

(N

2

)]Cl

2

: Pentaammine (dinitrogen) ruthenium (II) chloride. [Pt(py)

4

] [PtCl

4

]: Tetrakis

(pyridine) platinum (II) tetrachloroplatinate (II).



(iv) Point of Attachment: Ambidentate Ligands:



The ligands which have more than one donor atoms are called 

ambidentate ligands

. While naming

such ligands, the point of attachment of ligands is designated, by placing the symbol of attached

element, in 

italics

, after the name of the group and separated by hyphens.



For example:  (i)  M – NCS−

: thiocyanato –N−M SCN−  

: thiocyanato – S−



                  (ii) M – NO

: nitro–N−



                        M − ONO−

: nitro−O−



(NH

4

)

3

[Cr(NCS)

6

]: Ammonium hexathiocyanato –N– chromate (III). or Ammonium

hexaisothiocyanatochromate (III).



(v)

Bridging Ligands:



There are many monodentate ligands such as OH−, NH,

NH−−, NO , Cl−, O, CO, etc. which can                     co-

ordinate simultaneously with more than one metal ion thus

forms a bridge between them. Such ligands are called

bridging ligands

 and the complexes with such ligands are

called Polynuclear (or bridged) complexes



(a)  

A bridging group is indicated by adding the

Greek letter 



(mu) immediately before its name and

separated it by giving hyphens.



(b)  

Two or more bridging groups of the same kind are

indicated by di- 



- , tri- 



, (bis-



), etc.



(c)

The bridging groups are listed with the other

groups in an alphabetical order unless the symmetry of the

molecule allows a simpler name.



(d)  

Where the same ligand is present as a bridging

ligand and as a non-bridging ligand, it is first cited as

bridging ligand.





-amido-octaammine- µ -nitro-dicobalt (III) nitrate,[(NH

3

)

5

Cr.OH.Cr (NH

3

)

5

]Cl

5



5. 

Ending of Complex: Termination of Names:



If complex is anionic

, it has a characteristic ending: -ate. But for

neutral or cationic complexes, the ending occurs with the name of the

central metal, without any modifications.



A distinction is made for the anionic complexes so that they can be

distinguished easily from the corresponding acid (where acids have

characteristic ending – ic as we know).



K

4

[Fe(CN)

6

]

:  Potassium hexacyanoferrate (II)



H

4

[Fe(CN)

6

]

:  Hexacyanoferric (II) acid



 [Fe(H

2

O

6

]

3+

: Hexaaquairon (III) ion



[Ni (dmg)

2

]

: Bis-(dimethyl glyoximato) nickel (II)



For different metals in anionic complexes, the endings may be

considered as follows:



Co  : Cobaltate;

Zn :  Zincate;

Cr : Chromate;



Pt   :   Platinate; 

Re : Rhenate; 

Rh   Rhodate; etc.



For some metals, latin names are used for termination of complexes.



For example:

Ag (Silver – Argentum) 

: Argentate



Cu (Copper-Cuprum)

:  Cuprate



Fe (Iron – Ferrum)

:  Ferrate



Au (Gold – Aurum)

: Aurate



Pb (Lead – Plumbum)

: Plumbate, etc.



6.  

Oxidation State of Metal Ion:



The formal oxidation state (Werner’s principal valence) of the

central metal is designated by a Roman numeral, enclosed in

parentheses, just after the name of the metal, without a space between

the two. An oxidation state of zero is indicated by an Arabic zero (0)

and a negative state by a minus sign, before the Roman numeral.



[Cr(H

2

O)

6

]Cl

3

: Hexaaquachromium (III) chloride



Na

2

[Fe(CN)

5

NO]

: Sodium pentacyanonitrosylferrate (III)



K

4

[Ni(CN)

4

]

:  Potassium tetracyanonickelate (0)



Na [Co(CO)

4

]

:  Sodium tetracarbonylcobaltate (-I)



As a second alternative, it is also permissible to record overall

charge on the complex ion (STOCK number) by Arabic numerals and

sign enclosed in parenthesis, as discussed above.

NH

4

[Cr(NCS)

4

(NH

3

)

2

]: Ammonium diamminetetrakis (isothiocyanato)

chromate (1-) [Co(SCN)(NH

3

)

4

]Cl

2

: Peta (ammine) thiocyanatocobalt

(2+) chloride.



7. 

Designation of Isomers:



(a)  

Geometric isomers are named by using the prefixes cis and trans

but if the structure is complicated, 

number system

 is used to locate the

positions of ligands in the given complex.



(b)

Optical isomers are named by using the prefixes 

d

 and 

l

 to

designate the dextro and laevo rotatory compounds respectively.



Geometrical isomers



Cis-diamminedichloroplatinum (II).



trans-diamminedichloroplatinum (II)



Cis-tetraamminedibromorhodium (III) ion.



d-K

3

[Ir(C

2

O

4

)

3

] potassium d-trioxalatoiridate (III)



8.

Solvent of Crystallization:



If in a complex, any lattice components such as water or solvent of

crystallization is present, they follow the name, and are preceded by the

number of such groups in Arabic number.



For example:              CuSO

4

.5H

2

O

: Copper (II) sulphate-5-water



                 AlCl

3

.(EtOH)

4

: Aluminium (III) chloride-4-ethanol



      [Cu(NH

3

)

4

]SO

4

.H

2

O 

: Tetra (ammine) copper (II) sulphate-1 water



IUPAC Nomenclature at a Glance



A brief and precise summary of the rules, useful for systematic naming

of co-ordination compounds has been given in the Chemical Society’s Hand-

book for Authors (special publication No. 11). An abstract of it may be given for

ready reference as follows:



The cationic part is named first followed by the anionic one.



The ligands are named prior to the metal ion.



Negative ligands end in-‘o’-, neutral ligands have no special endings, while

positive ligands end in –

ium.



The ligands of several types are listed alphabetically, irrespective of their type.



The prefixes 

di-, tri-, tetra-, penta-, and hexa-

 are used to indicate the number

of simple monodentate ligands.  If the ligand is having complicated structure,

the prefixes, 

bis, tris, tetrakis

, etc. are used.



The oxidation state of metal is specified by a Roman numeral enclosed in

brackets immediately following its name.



Some appropriate illustrations of chemical nomenclature

have been given below to digest the rules laid by IUPAC

Nomenclature committee. This may be regarded as name

to formula or formula to name exercise.



[Co(NH

3

)

6

]Cl

2

Hexamminecobalt (II) chloride or



Hexammine cobalt (II) chloride



[Co(CN)

 6

]

3-

            Hexacyanocobaltate (III) ion



[CrCl

3

(NH

3

)

3

]

Triamminetrichlorochromium (III)



K

4

[Ni(CN)

4

]

Potassium tetracyanonickelate (0)



[Ag(NH

3

)

2

]

2

 [PtI

4

]Diamminesilver (I) tetraiodoplatinate(II)



[Cr(NH

3

)

5

H

2

O][Co(CN)

6

]

Petaammineaquachromium (III)



   hexacyanocobaltate (III)



[Pt(NH

3

)

4

(en)]Cl

4

Tetra (ammine) (ethylenediamine)



                                 platinum (IV) chloride.



K

3

[Al(C

2

O

4

)

3

]

Potassium trioxalatoaluminate (III)



[Zn(OH)

4

]

2−

Tetrahydroxozincate (II) ion



[Al(OH) (H

2

O)

5

]

2+

Pentaaquahydroxoaluminium (III) ion



µ-amido-µ-peroxo-bisTriamminechlorocobalt (III)

chloride.



[CoCl.CN.NO

2

 (NH

3

)

3

]

Triamminechlorocyanonitrocobalt (III)



Li[AlH

4

]

Lithium tetrahydridoaluminate (III)



(Lithium aluminium hydride)



K

2

[O

5

Cl

5

N]Potassium pentachloronitridoosmate (VI)



K

2

[Cr(CN)

2

O

2

(O

2

)NH

3

]

Potassium

amminedicyanodioxoperoxo-chromate (VI)



[Cr(C

6

H

6

)

2

]

Bis-(benzene) chromium (0)



[Be

4

O(CH

3

COO)

6

]Hexa-µ-acetato-(O, O') - µ-oxo-

tetra- beryllium (II)



AlK(SO

4

)

2

.12H

2

O

Aluminium potassium sulphate-12-

water



1.6 ISOMERISM IN COMPLEXES WITH C.N. = 4 AND 6 :-



Isomerism :



The compound having the same molecular formula but different

structural arrangements (of their atoms) are called isomers (Greek, iso 



equal; mer 



 mass). The phenomenon that give rise to structural

differences and hence differences in properties of the compound is called

as isomerism.



Stereoisomerism:



When two co-ordination compounds contain the same ligands attached to the

same metal ion having different arrangement for their ligands in space, are

said to be stereo-isomers and the phenomenon is called stereoisomerism

 (or

space isomerism). It is of two types geometrical (or cis-trans) isomerism and

optical (or mirror image) isomerism

.



1.6.1 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 form and trans form,

respectively. This type of isomerism is, therefore, also referred to as 

cis-trans

isomerism

.





I. Geometrical isomerism in complexes of

coordination number (C.N. = 4)



Complexes with co-ordination number four may be either tetrahedral or

square planar. For tetrahedral complexes, geometrical isomerism is not possible

as all the four ligands are equidistant from each other (i.e. at the angle 109.5



).

Further,



(i)Square planar complexes of type [Ma

4

], [Ma

3

b], [Mab

3

] do not show geometric

isomers as the ligands have equivalent disposition around the metal ion.



(ii)However, square planar complexes of the type [Ma

2

b

2

], [Mabc

2

], [Mabcd] etc.

do show cis-trans isomerism.



Here M stands for central metal while a, b, c and d stands for monodentate

ligands. Some important types may be discussed as follows:



(a)

[Ma

2

b

2

] complexes :



When a and b are arranged cis and trans to each other, then two isomers

are obtained. Best known example of this type is [Pt(NH

3

)

2

Cl

2

] and is

represented in Fig. 1.5. The cis and trans forms of the complex type Ma

2

b

2

.



Fig. 1.5 : The cis and trans forms of the complex [PtCl

2

(NH

3

)

2

]



(b)

[Ma

2

bc] complex :



Many compounds of platinum (II) of this type are known. Here ‘a’ is a

neutral ligand such as NH

3

, H

2

O, pyridine, etc. ‘b’ and ‘c’ are anionic



ligands such Cl



, Br



, SCN



, NO

2

 etc. One of the example of [PtBrI(NH

2

)

2

] is

shown in Fig. 1.6

Fig. 1.6 : 

t

he cis and trans forms of the complex of type Ma

2

bc

(c)[Mabcd] complexes :

A few compounds of platinum (II) of this type are known. We may select

one ligand from either of a, b, c and d and the remaining three may be placed trans to

it one by one, and we can get three isomers. For example, [PtBrClNH

3

py] can be

shown as follows. Let us select NH

3

 and place other trans to it.

Fig. 1.7 : cis and trans isomers of [PtBrClNH

3

py]



(d) [M(AB)

2

] complexes :



Here AB is an unsymmetrical bidentate chelate ligand in which A and B

are the two different donor atoms. Square planar complexes of [M(AB)

2

] type

also exhibit cis and trans isomers.



For example [Pt(gly)

2

], where gly is glycinato ligand NH

2

CH

2

COO



. The

two isomers may be shown as follows :

                                     Fig. 1.8 : cis and trans isomers of [Pt(gly)

2

]

The chemistry of Pt(II) complexes has been studied extensively by Russian

chemists.Complexes of Pt(II), Pd(II), Ni(II), and Cu(II) are square planar.

II. 

Geometrical isomerism in complexes of coordination number 6

(C.N. = 6).

The complexes having coordination number 6 have octahedral geometry. The six

positions in this geometry are considered by numbering them from 1 to 6 as

shown in Fig. 1.9.

Cis positions : 

1-2,2-3,3-4,4-5,1-3,1-5,2-6,3-6,4-6,5-6 and 2-5.

Trans positions : 

1-6, 2-4 and 3-5.



II. 

Geometrical isomerism in complexes of coordination

number 6 (C.N. = 6).



The complexes having coordination number 6 have octahedral

geometry. The six positions in this geometry are considered by

numbering them from 1 to 6 as shown in Fig. 1.9.



Cis positions 

: 

1-2,2-3,3-4,4-5,1-3,1-5,2-6,3-6,4-6,5-6 and 2-5.



Trans positions : 

1-6, 2-4 and 3-5.

Fig. 1.9 : Regular octahedron with its six positions numbered from 1 to 6

The octahedral complexes also exhibit geometrical isomerism. The complexes of the type

[Ma

6

]

 and 

[Ma

5

b] or [M(A-A)

3

] 

do not show geometrical isomerism. The different types of

octahedral complexes exhibiting geometrical isomerism are discussed below.

(a)

[Ma

4

b

2

] complexes :

The two similar ligands b may occupy 1 and 2 positions and 1 and 6 positions to give cis

and trans isomers respectively. The most familiar examples of octahedral geometrical

isomers of this type are violet (cis) and green (trans) isomers of tetraamminedichlorocobalt

(III) and chromium (III) cations as shown in Fig. 1.10.

                         Fig. 1.10 : Cis-trans isomers of [CoCl

2

(NH

3

)

4

]

+

(b)[Ma

3

b

3

] complexes :

In such complexes, cis form arises when the three similar groups are placed on

the corners of one of the triangular faces of octahedron and the remaining three

identical groups on the opposite triangular face of the octahedron.In trans-isomer, the

sets of identical groups lie on the edges opposite to each other as shown in Fig. 1.11.

Fig. 1.11 : Cis-trans isomers of [CoCl

3

(NH

3

)

3

]

Other examples are

[CrF

3

(H

2

O)

3

], [RhCl

3

(py)

3

(NH

3

)

3

]

+

, [Pt(NO

2

)

3

(NH

3

)

3

]

+

,

[IrCl

3

(H

2

O)

3

], [PtI

3

(NH

3

)

3

]

+

 ,[RuCl

3

(H

2

O)

3

],



(c)

 [Ma

2

b

2

c

2

] complexes :



Theoretically, five geometrical isomers are

expected for this type of complex. Out of these, only

three have been prepared. Five possible isomers have

been shown for [PtCl

2

(NH

3

)

2

(py)

2

]

2+

 in Fig. 1.12.

Fig. 1.12 : [PtCl

2

(NH

3

)

2

(py)

2

]

2+



(d)

[Mabcdef] complexes :



Octahedral complexes with six different monodentate ligands give 15

geometrical isomers theoretically. Only one example of this type of complex is

known.



It is of Pt(IV) having the formula: [PtBrClINO

2

NH

3

py]. Only three isomers of

these complexes have been isolated.



(e)[M(A-A)

2

a

2

] complexes :



Here (A-A) is a symmetrical bidentate ligand with identical donor groups A

and A. The complex of this type can show cis-trans isomers whereas the two

monodentate ligands (a) will be cis to each other in the cis form, while trans to each

other in the trans form as shown in Fig. 1.13



Some examples of this type are :



[Co(NH

3

)

2

(en)

2

]

3+

, [Co(NO

2

)

2

(en)

2

]

+

, [CrCl

2

(en)

2

]

+

,



[CrCl

2

(en)

2

]

+

, [IrCl

2

(C2O

4

)

2

]1

3



, [RhCl

2

(en)

2

]

+

, etc.

Fig. 1.13 : cis and trans forms complex [Co(en)

2

Cl

2

]

+

 ion



(f)

[M(A-A)

2

ab] complexes :



We may have complexes of the type [M(A-A)

2

ab]. For example,

[CoClNH

3

(en)

2

]

2+

, [CrClNH

3

(en)

2

]

2+

, [RhClNH

3

(en)

2

]

2+

, etc. All such complexes

show cis and trans isomerism.



(g)

[M(A-A)

a

2b

2

] complexes :



Further, we can have geometrical isomerism for the complexes of the

type [M(A-A)a

2

b

2

]. For example, [CoCl

2

(NH

3

)

2

en]

+

 ion having both

NH

3

 - NH

3

  and Cl



 - Cl



 ions in cis give cis form while NH

3

 and Cl



 in trans

positions give rise to trans isomer.



(h)

[M(A-B)

3

] complexes :



Here (A-B) is an unsymmetrical bidentate chelate ligand having A and B

as two different co-ordinating groups. Octahedral complexes of the type [M(A-

B)

3

] exist in cis and trans isomers. For example, triglycinatochromium (III).



1.6.2 Optical Isomerism



There are certain substances which can rotate the plane of polarized

light. These are called 

optically active substances

. The optically active

isomers of a compound which rotate the plane of polarized light equally but in

opposite directions are called 

enantiomers

. The isomer which rotates the

plane of polarized light to the right is called 

dextro rotatory

 designated as d

and the one which rotates the plane of polarized light to the left is called 

laevo

rotatory

 designated as 

l

.



(i)

Optical isomerism in complexes with

coordination number 4 :



Tetrahedral complexes with coordination number 4 should exhibit

optical isomerism because there is no plane of symmetry in their molecules.

However, this has not been found to be the case. The first tetrahedral metal

complex in which the central ion was surrounded by four different ligands was

prepared in 1963 but it has not been possible to resolve it into d and l forms.

However, tetrahedral complexes of Pt(II) Be(II), B(III) and Zn(II), containing

two symmetrical), bidentate ligands, can be resolved into optically active

forms. For example, the complex bis (benzoylacetonato) beryllium (II) has

been shown to exist in two optical forms, as shown in Fig. 1.14.

                            Fig. 1.14 : d- and 

l

- isomers of [PtBrClINO

2

NH

3

Py]

It may be noted that 4 different groups around the central atom are not

essentially required for optical activity. All that is required is that the molecule

should be asymmetric so that it can exist in two forms which are mirror images of

each other, as shown above. Evidently, the two structures do not superimpose on

each other.



(ii) Optical isomerism in complexes with coordination number

(CN = 6) :



The most well known cases of optical isomerism occur among

octahedral complexes. The common examples are complexes which

contain bidentate ligands. These complexes fall under the following

categories:



(A)

[M(L

2

)

3

] type complexes:



Where, M = metal ion, L = bidentate ligand. The complexes such as

[Co(en)

3

]

3+

,  [Co(en)

3

]Cl

3

,  [Cr(ox)

3

]

3



, etc., exist as optical isomers

because they form non-superimposable mirror images, as illustrated in

Fig. 1.16

.

Fig. 1.16 : [Co(en)

3

]

3+



(B)

[M(L

2

)

2

L

2

'] or [M(L

2

)

2

bc] type complexes :



It has already been shown (Fig. 30 that these complexes form

geometrical isomers (cis and trans forms). It is interesting to note that

the trans form does now show optical isomerism, i.e., it cannot be

resolved into optical isomers. The reason is that the molecule in this

case is symmetric. On the other hand, the cis isomer is asymmetric and

cm into d and/isomers. For example, the complex ion 

cis-dichloro bis

(ethylenediamine) chromium (III) : 

cis-

  [CrCl

2

(en)

2

]

+

 exists in the

form of three

                                             Fig. 1.17 : cis- [CrCl

2

(en)

2

]

+

isomers (trans optically inactive; cis (d) and cis (

l

). The resolution of the cis

isomer of the complex into d and 

l

 forms is shown in Fig. 1.17.



(C)

cis-cis [ML

2

L

2

'L

2

"] type of complexes :



The complexes containing only one symmetric bidentate ligand

also show optical isomerism. For example, the complex ion

[CoCl

2

(NH

3

)

2

(en)]

+

 exists in d and l forms, as shown in Fig. 1.18.

Fig. 1.18 : Structure of cis-cis-[CoCl

2

(NH

3

)

2

(en)]

+

(D)

Complexes containing polydentate ligand like EDTA

(haxadentate) ligands :

The complexes containing haxadentate ligands such as

ethylenediaminetetraacetato cobaltate (III) also show optical

activity. For example, the [Co(edta)]



 exists in two forms, as

shown in Fig. 1.19.

Fig. 1.19 :

Ethylenediaminetetraa

cetato cobaltate (III)



(E)

Optical isomerism is also possible among

polynuclear complexes containing :



The complexes of the type [Mabcdef] are also

expected to show optical isomerism. A complex of this

type can theoretically have 15 geometrical forms, each

of which should be optically active.

Fig. 1.20 : m-amido-tetrakis-(ethylenediamine)-m-nitro-dicobalt (III) ion



1.6.3 Structural Isomerism-Ionisation Isomerism, Hydrate Isomerism,

Coordination Isomerism, Linkage Isomerism and Co-ordination

position Isomerism



Isomerism in Coordination Compounds :



It may be recalled that the compounds having the same

molecular formula but, different structures and hence different

physical and chemical properties are called isomers. The phenomenon

of the existence of such compounds is known as isomerism.

Coordination compounds give rise to a wide variety of isomers due to

different types of linkages and different structural arrangements of the

constituent atoms.



Isomerism in coordination compounds may be divided into two main

types :



(A) 

Structural Isomerism.



(B) 

Stereo Isomerism



These two types are further subdivided as shown below :

Isomerism in Coordination

Compounds



(

A)Structural Isomerism 

:

The different types of structural isomerism are discussed below :

1. Ionisation Isomerism :

In this type of isomerism, the difference arises from the

positions of groups within or outside the coordination sphere. Therefore, these

isomers give different ions in solution, hence the name ionisation isomerism. For

example, there are two distinct compounds of the formula Co(NH

3

)

5

BrSO

4

. One of

them is red-violet and gives a precipitate with BaCl2 indicating that sulphate ion is

outside the coordination sphere. The second one is red and does not give

precipitate with BaCl

2

 but gives a precipitate of AgBr with silver nitrate indicating

that now bromide ion is outside the coordination sphere.

2.

Hydrate Isomerism :

This type of isomerism, which is similar to ionisation isomerism discussed

above, arises from replacement of a coordinated group by water of hydration. As an

example, consider the following three isomers having the molecular formula CrCl

3



6H

2

O :

[Cr(H

2

O)

6

]Cl

3

, [CrCl(H

2

O)

5

]Cl

2



 H

2

O and [CrCl

2

(H

2

O)

4

]Cl 



 2H

2

O

They differ largely from one another in their physical and chemical properties, as

illustrated below.

Another example is furnished by the following two isomers :

[CoCl(en)

2

(H

2

O)]Cl

2

  and  [CoCl

2

(en)

2

]Cl 



 H

2

O



3.  Coordination Isomerism :



This type of isomerism is observed in the case of compounds comprising of

both cationic and anionic complexes. The examples are :



(i)[Co(NH

3

)

6

] [Cr(CN)

6

]

and

[Cr(NH

3

)

6

] [Co(CN)

6

]



(ii)[Cu(NH

3

)

4

] [PtCl

4

]

and

[Pt(NH

3

)

4

] [CuCl

4

]



This type of isomerism is also shown by compounds in which the metal ion

is the same in both cationic and anionic complexes. For example,



[Cr(NH

3

)

6

] [Cr(CN)

6

]

and

[Cr(CN)

2

(NH

3

)

4

] [Cr(CN)

4

(NH

3

)

2

]



[Pt(NH

3

)

4

] [PtCl

4

]

and

[PtCl(NH

3

)

3

] [PtCl

3

(NH

3

)]



4. 

Linkage Isomerism :



In some ligands, there are two atoms which can donate their lone pairs. For

example, in NO ion, the nitrogen atom as well as the oxygen atom can donate their

lone pairs. This gives rise to 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. For example, two

different pentaamminecobalt (III) chlorides, each containing the NO

2

 group in the

complex ion, have been prepared. These are:



[Co(NO

2

)(NH

3

)

5

Cl

2

 Pentaamminenitro-N cobalt (III) chloride



[Co(ONO)(NH

3

)

5

]Cl

2

 Pentaamminenitrito-O cobalt (III) chloride



Other ligands which give rise to linkage isomerism are CN (cyano) and NC

(isocyano) and SCN (thiocyanato) and NCS (isothiocyanato).



5. Coordination Position Isomerism :



This type of isomerism is exhibited by bridged complexes and is from different

placement of ligands, as shown below :



Fig. 1.21 : Ammonia and Chloro ligands



It will be seen that in (a) and (b), ammonia and



chloro ligands are differently placed relative to the atoms.



1.7

Valence Bond Theory of Transition Metal Complex with respect

to C.N. = 4, C.N. = 6



1.7.1 General



The valence bond theory was developed and then applied to

metal complexes mainly by Prof. Linus Pauling, of the California

Institute of Technology. In 1931, it was the most widely used theory

though now rarely used in co-ordination chemistry. Today VBT is

replaced by the more modern theories like CFT and LFT. But at the

same time it is not untrue that much of the recent research in          co-

ordination chemistry has been inspired from VBT. It is still useful to

explain various experimental observations qualitatively. It deals with

the electronic structure of the central metal atom in its ground state

and is concerned mainly with the study of:



The kind of bonding.



Stereochemistry, and



The gross magnetic properties of the metal complexes.



1.6.2 Salient Features (Pauling’s) of VBT



The salient features (or highlights) of VBT as assumed by            Prof. Linus

Pauling may be outlined as follows:



 1.  According to VBT, the central metal atom or ion makes available an appropriate

number of empty orbitals, in accordance with its co-ordination number.



 2.   The metal atom orbitals undergo hybridization to form new orbitals of

equivalent energy and suitable orientation.



 3.The d-orbitals involved in bonding for complex formation may be inner d-

orbitals i.e. (n-1) d or outer d-orbitals i.e. nd.



[The resulting complexes are therefore referred to inner orbital complexes (i.e. 

low

spin or spin paired complexes

) or/and outer orbital complexes (i.e. 

high spin or

spin-free

 complexes) respectively.]



 4. Each ligand becomes ready with its lone pair of electrons,     so that a sigma type

of covalent bond will be formed.



5. The empty hybrid orbital of metal ion overlaps linearly with the completely full

orbital of the ligand to form co-ordinate covalent bond.



6. The electrons of metal ion usually remain non-bonding and are accommodated

in the inner d-orbitals (i.e. the 



-orbitals, viz. d

xy

, d

yz

, d

zx

).



7. In addition to the σ-bond, a 



-bond may be formed. Lateral overlap of

completely full orbitals of metal with completely empty orbitals of ligands or vice

versa cause 



-bonding in complexes.



 8. A substance without unpaired electrons is said to be diamagnetic. On the other

hand, the substance containing one or more unpaired electrons is said to be

paramagnetic.



 9. 

Geometry (shape) of the complex

: Due to hybridization, the orbitals attain

definite orientation in space. This in turn imparts definite geometry to the

resulting complex.



e.g. sp

3

 (tetrahedral), dsp

2

 (square planar), dsp

3

 or sp

3

d (trigonal

bipyramidal or square pyramidal), d

2

sp

3

 or sp

3

d

2

 (octahedral),

etc.



10. 

The strength of co-ordinate bond depends on the

extent of overlap of orbitals of metal ion and complexing agent.



In short, formation of a complex is a reaction between a

Lewis acid (metal) and a Lewis base (ligand) to form a co-

ordinate covalent (dative) bond between the tow (M

←

L).



1.6.3

Valence Bond Theory of Transition Metal Complex with

respect to C.N. = 4



There are two possible configurations for metal complexes

with co-ordination number four. These are tetrahedral and

square planar. Tetrahedral structure arises from sp

3

 hybridization

while square planar configuration results from dsp

2

hybridization. Let us consider nickel (II) and copper (II)

complexes. The pictorial representations are depicted in Fig. 1.5

where electron distributions in metal, metal ion and in the

resulting complex species are shown step-by-step.



Ni (II) Complexes:



(i)Tetrahedral complex

 [NiCl

4

]

2−



Ni

+2

 Ion-Excited state configuration on approach of Cl−

ligands, leading to

sp

3

 hybridization.



The electronic configuration of Ni (Z = 28)   is,  (3d

8

4s

2

4p

0

)



Ni Ground state



Ni

+2

 Ion configuration after gaining four pairs (xx) of

electrons  from 4Cl−

ions, forming 

tetrahedral complex

[NiCl

4

]

2−



(2) Cu (II) Complexes:



(i)Tetrahedral 

 Complexes: [CuCl

4

]

2−

 :



Cu (Z = 29) Atom-Ground state configuration (3d

10

4s

1

4p

0

)



Cu

+2

 Ion-Ground state configuration (3d

9

4s

0

4p

0

) after loss of two

electrons, one from 4s and another from 3d.



ii) S

quare planar

 complex.  [Cu(NH

3

)

4

]

2− 

 :

up, it is a 

low spin or spin-paired octahedral complex

 and it is diamagnetic complex.



(D)

Choose the correct alternative and rewrite the sentence:



1. 

Co-ordinate bond is 





(a)  

semi-polar bond        

(b) 

double bond



(c) 

van der Waal’s bond 

(d) 

weak bond



2. 

Double salt retains its identity only in the form of 





(a)  

crystal lattice

(b) 

aqueous solution



(c) 

complex ion 

(d) 

coloured compound



3.  

Concept of auxiliary valence was put forth by 





(a)  

Alfred Werner

(b) 

S.M. Jorgensen



(c)  

Linus Pauling 

(d) 

Sidgwick



4.  

CoCl

3

.4NH

3

 may be formulated as 





(a) 

[Co(NH

3

)

4

Cl

2

]Cl

(b) 

[Co(NH

3

)

3

Cl

3

](NH

3

)



(c) 

[Co(NH

3

)

4

Cl

3

]

(d) 

[(NH

3

)

3

CoCl(NH)

3

]



5.  

Effective atomic number of Fe in K

4

 [Fe(CN)

6

] is 





(a) 

36

(b) 

26



(c) 

24

(d) 

28



6. 

Secondary valence of Co in [Co(NH

3

)

4

Cl

2

]

+

 is 





(a) 

6

(b) 

4



(c) 

2

(d) 

0



7.  

Every transition element is characterized by two types of valencies 





(a) 

primary and secondary   

(b) 

principle and primary



(c) 

primary and ionizable

(d) 

secondary and subsidiary



8. 

Valence bond theory was applied to complex by 





(a) 

Linus Pauling

(b) 

Alfred Werner



(c) 

Sidgwick

(d) 

S.M. Jorgensen



9.  

Valence bond theory is simple, conceptual and 





(a) 

qualitative

(b) 

semi-qualitative



(c) 

quantitative

(d) 

logical



10. 

Valence bond theory is based on 





(a) 

concept of hybridization

(b) 

secondary valence



(c) 

Ligand orbitals

(d) 

crystal field splitting



11. 

Secondary valence of Cr in [Cr(NH

3

)

4

Cl

2

]

+

 is 





(a) 

6

(b) 

4



(c) 

2

(d) 

0



12.

Effective atomic number of Fe in [Fe(CN)

6

]

−4

 is 





(a) 

36

(b) 

26



(c) 

24

(d) 

28



13 As the electrons in [Co(NH

3

)

6

]

3+

 are paired up, it is a 

low spin or



 spin-paired octahedral complex

 and it is



(a) 

diamagnetic complex

.(b) paramagnetic complex



(c)

antiferomagnetic complex.(d) feromagnetic complex.





14. 

Co

3+

Ion-Excited state configuration on approach of F

−

 ligands, leading to  octahedral



(a) 

sp

3

 d

2 

hybridization

. (b) d

2

sp

3

 hybridization.



(c) sp

3

 dhybridization. (a) dsp

3

 hybridization.



15. 

Co

3+

Ion-Excited state configuration on approach of NH

3

 ligands, leading to  octahedral ………….





(a) 

d

2 

sp

3

 hybridization

. (b) sp

3

 d

2 

 hybridization.



(c) sp

3

d dhybridization. (a) dsp

3

 hybridization.



16. As the electrons in [Fe(CN)

6

]

3−

 are paired up, it is a ………………..

complex

 and is essentially a slightly

paramagnetic due to a single unpaired electron.



(

a)

low spin or spin-paired

(a)

high- spin or spin-paired



(a)

 low spin or spin-free     

(a)

 high spin or spin-free



17. 

Cu

+2

 Ion-Excited state configuration on approach of 4NH

3

ligands,



causing to ……………..



(a)dsp

3

 hybridization 

(a) sp

2

d hybridization



(a)sp

3

 hybridization (a) sp

3

d  hybridization



18. 

Cu

+2

 Ion-Excited state configuration on approach of 4Cl

-  

ligands,



causing to ……………..



(a)sp

3

 hybridization 

(a) sp

2

d hybridization



(a)sp

3

 hybridization (a) sp

3

d  hybridization



19. 

Ni

+2

 Ion-Excited state configuration on approach of 4NH

3

ligands,



causing to ……………..



(a)dsp

3

 hybridization 

(a) sp

2

d hybridization



(a)sp

3

 hybridization (a) sp

3

d  hybridization