244Chemistry

In the previous Unit we learnt that the transition metals

form a large number of complex compounds in which

the metal atoms are bound to a number of anions or

neutral molecules by sharing of electrons. In modern

terminology such compounds are called coordination

compounds. The chemistry of coordination compounds

is an important and challenging area of modern

inorganic chemistry. New concepts of chemical bonding

and molecular structure have provided insights into

the functioning of these compounds as vital components

of biological systems. Chlorophyll, haemoglobin and

vitamin B

are coordination compounds of magnesium,

iron and cobalt respectively. Variety of metallurgical

processes, industrial catalysts and analytical reagents

involve the use of coordination compounds.

Coordination compounds also find many applications

in electroplating, textile dyeing and medicinal chemistry.

Coordination

Compounds

After studying this Unit, you will be

able to

• appreciate the postulates of

Werner’s theory of coordination

compounds;

• know the meaning of the terms:

coordination entity, central atom/

ion, ligand, coordination number,

coordination sphere, coordination

polyhedron, oxidation number,

homoleptic and heteroleptic;

• learn the rules of nomenclature

of coordination compounds;

• write the formulas and names

of mononuclear coordination

compounds;

• define different types of isomerism

in coordination compounds;

• understand the nature of bonding

in coordination compounds in

terms of the Valence Bond and

Crystal Field theories;

• appreciate the importance and

applications of coordination

compounds in our day to day life.

Objectives

Coordination Compounds are the backbone of modern inorganic

and bio–inorganic chemistry and chemical industry.

Coordination

Compounds

Alfred Werner (1866-1919), a Swiss chemist was the first to formulate

his ideas about the structures of coordination compounds. He prepared

and characterised a large number of coordination compounds and

studied their physical and chemical behaviour by simple experimental

techniques. Werner proposed the concept of a primary valence and

a secondary valence for a metal ion. Binary compounds such as

CrCl

, CoCl

or PdCl

have primary valence of 3, 2 and 2 respectively.

In a series of compounds of cobalt(III) chloride with ammonia, it was

found that some of the chloride ions could be precipitated as AgCl on

adding excess silver nitrate solution in cold but some remained in

solution.

9.1

Werner’

Theory ofTheory of

Theory of

CoordinationCoordination

Coordination

CompoundsCompounds

Compounds

Unit

2020-21

245 Coordination Compounds

1 mol CoCl

.6NH

(Yellow) gave 3 mol AgCl

1 mol CoCl

.5NH

(Purple) gave 2 mol AgCl

1 mol CoCl

.4NH

(Green) gave 1 mol AgCl

1 mol CoCl

.4NH

(Violet) gave 1 mol AgCl

These observations, together with the results of conductivity

measurements in solution can be explained if (i) six groups in all,

either chloride ions or ammonia molecules or both, remain bonded to

the cobalt ion during the reaction and (ii) the compounds are formulated

as shown in Table 9.1, where the atoms within the square brackets

form a single entity which does not dissociate under the reaction

conditions. Werner proposed the term secondary valence for the

number of groups bound directly to the metal ion; in each of these

examples the secondary valences are six.

Note that the last two compounds in Table 9.1 have identical empirical

formula, CoCl

.4NH

, but distinct properties. Such compounds are

termed as isomers. Werner in 1898, propounded his theory of

coordination compounds. The main postulates are:

1. In coordination compounds metals show two types of linkages

(valences)-primary and secondary.

2. The primary valences are normally ionisable and are satisfied by

negative ions.

3. The secondary valences are non ionisable. These are satisfied by

neutral molecules or negative ions. The secondary valence is equal to

the coordination number and is fixed for a metal.

4. The ions/groups bound by the secondary linkages to the metal have

characteristic spatial arrangements corresponding to different

coordination numbers.

In modern formulations, such spatial arrangements are called

coordination polyhedra. The species within the square bracket are

coordination entities or complexes and the ions outside the square

bracket are called counter ions.

He further postulated that octahedral, tetrahedral and square planar

geometrical shapes are more common in coordination compounds of

transition metals. Thus, [Co(NH

)

]

, [CoCl(NH

)

]

and [CoCl

(NH

)

]

are octahedral entities, while [Ni(CO)

] and [PtCl

]

2–

are tetrahedral and

square planar, respectively.

Colour Formula Solution conductivity

corresponds to

Table 9.1: Formulation of Cobalt(III) Chloride-Ammonia Complexes

Yellow [Co(NH

)

]

3Cl

–

1:3 electrolyte

Purple [CoCl(NH

)

]

2Cl

–

1:2 electrolyte

Green [CoCl

(NH

)

]

–

1:1 electrolyte

Violet [CoCl

(NH

)

]

–

1:1 electrolyte

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246Chemistry

(i) Secondary 4 (ii) Secondary 6

(iii) Secondary 6 (iv) Secondary 6 (v) Secondary 4

On the basis of the following observations made with aqueous solutions,

assign secondary valences to metals in the following compounds:

SolutionSolution

Solution

Difference between a double salt and a complex

Both double salts as well as complexes are formed by the combination

of two or more stable compounds in stoichiometric ratio. However, they

differ in the fact that double salts such as carnallite, KCl.MgCl

.6H

Mohr’s salt, FeSO

.(NH

)

.6H

O, potash alum, KAl(SO

)

.12H

O, etc.

dissociate into simple ions completely when dissolved in water. However,

complex ions such as [Fe(CN)

]

4–

of K

[Fe(CN)

] do not dissociate into

and CN

–

ions.

Formula Moles of AgCl precipitated per mole of

the compounds with excess AgNO

(i) PdCl

.4NH

(ii) NiCl

.6H

O 2

(iii) PtCl

.2HCl 0

(iv) CoCl

.4NH

(v) PtCl

.2NH

Example 9.1

WernerWerner

Werner was born on December 12, 1866, in Mülhouse,

a small community in the French province of Alsace.

His study of chemistry began in Karlsruhe (Germany)

and continued in Zurich (Switzerland), where in his

doctoral thesis in 1890, he explained the difference in

properties of certain nitrogen containing organic

substances on the basis of isomerism. He extended vant

Hoff’s theory of tetrahedral carbon atom and modified

it for nitrogen. Werner showed optical and electrical differences between

complex compounds based on physical measurements. In fact, Werner was

the first to discover optical activity in certain coordination compounds.

He, at the age of 29 years became a full professor at Technische

Hochschule in Zurich in 1895. Alfred Werner was a chemist and educationist.

His accomplishments included the development of the theory of coordination

compounds. This theory, in which Werner proposed revolutionary ideas about

how atoms and molecules are linked together, was formulated in a span of

only three years, from 1890 to 1893. The remainder of his career was spent

gathering the experimental support required to validate his new ideas. Werner

became the first Swiss chemist to win the Nobel Prize in 1913 for his work

on the linkage of atoms and the coordination theory.

(1866-1919)(1866-1919)

(1866-1919)

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247 Coordination Compounds

(a) Coordination entity

A coordination entity constitutes a central metal atom or ion bonded

to a fixed number of ions or molecules. For example, [CoCl

(NH

)

]

is a coordination entity in which the cobalt ion is surrounded by

three ammonia molecules and three chloride ions. Other examples

are [Ni(CO)

], [PtCl

(NH

)

], [Fe(CN)

]

4–

, [Co(NH

)

]

(b) Central atom/ion

In a coordination entity, the atom/ion to which a fixed number

of ions/groups are bound in a definite geometrical arrangement

around it, is called the central atom or ion. For example, the

central atom/ion in the coordination entities: [NiCl

[CoCl(NH

)

]

and [Fe(CN)

]

3–

are Ni

, Co

and Fe

, respectively.

These central atoms/ions are also referred to as Lewis acids.

The ions or molecules bound to the central atom/ion in the

coordination entity are called ligands. These may be simple ions

such as Cl

–

, small molecules such as H

O or NH

, larger molecules

such as H

NCH

or N(CH

)

or even macromolecules,

such as proteins.

When a ligand is bound to a metal ion through a single donor

atom, as with Cl

–

, H

O or NH

, the ligand is said to be unidentate.

When a ligand can bind through two donor atoms as in

NCH

(ethane-1,2-diamine) or C

2–

(oxalate), the

ligand is said to be didentate and when several donor atoms are

present in a single ligand as in N(CH

)

, the ligand is said

to be polydentate. Ethylenediaminetetraacetate ion (EDTA

4–

) is

an important hexadentate ligand. It can bind through two

nitrogen and four oxygen atoms to a central metal ion.

When a di- or polydentate ligand uses its two or more donor

atoms simultaneously to bind a single metal ion, it is said to be a

chelate ligand. The number of such ligating groups is called the

denticity of the ligand. Such complexes, called chelate complexes

tend to be more stable than similar complexes containing unidentate

ligands. Ligand which has two different donor atoms and either of

the two ligetes in the complex is called ambidentate

ligand. Examples of such ligands are the NO

–

and

SCN

–

ions. NO

–

ion can coordinate either through

nitrogen or through oxygen to a central metal

atom/ion.

Similarly, SCN

–

ion can coordinate through the

sulphur or nitrogen atom.

(d) Coordination number

The coordination number (CN) of a metal ion in a complex can be

defined as the number of ligand donor atoms to which the metal is

directly bonded. For example, in the complex ions, [PtCl

]

2–

and

[Ni(NH

)

]

, the coordination number of Pt and Ni are 6 and 4

respectively. Similarly, in the complex ions, [Fe(C

)

]

3–

and

[Co(en)

]

, the coordination number of both, Fe and Co, is 6 because

2–

and en (ethane-1,2-diamine) are didentate ligands.

9.29.2

9.2

Definitions ofDefinitions of

Definitions of

SomeSome

Some

ImportantImportant

Important

TermsTerms

Terms

Pertaining toPertaining to

Pertaining to

CoordinationCoordination

Coordination

CompoundsCompounds

Compounds

2020-21

248Chemistry

It is important to note here that coordination number of the central

atom/ion is determined only by the number of sigma bonds formed by

the ligand with the central atom/ion. Pi bonds, if formed between the

ligand and the central atom/ion, are not counted for this purpose.

(e) Coordination sphere

The central atom/ion and the ligands attached to it are enclosed in

square bracket and is collectively termed as the coordination

sphere. The ionisable groups are written outside the bracket and

are called counter ions. For example, in the complex K

[Fe(CN)

the coordination sphere is [Fe(CN)

]

4–

and the counter ion is K

(f) Coordination polyhedron

The spatial arrangement of the ligand atoms which are directly

attached to the central atom/ion defines a coordination

polyhedron about the central atom. The most common

coordination polyhedra are octahedral, square planar and

tetrahedral. For example, [Co(NH

)

]

is octahedral, [Ni(CO)

] is

tetrahedral and [PtCl

]

2–

is square planar. Fig. 9.1 shows the

shapes of different coordination polyhedra.

9.39.3

9.3

NomenclatureNomenclature

Nomenclature

CoordinationCoordination

Coordination

Compounds

(g) Oxidation number of central atom

The oxidation number of the central atom in a complex is defined

as the charge it would carry if all the ligands are removed along

with the electron pairs that are shared with the central atom. The

oxidation number is represented by a Roman numeral in parenthesis

following the name of the coordination entity. For example, oxidation

number of copper in [Cu(CN)

]

3–

is +1 and it is written as Cu(I).

(h) Homoleptic and heteroleptic complexes

Complexes in which a metal is bound to only one kind of donor

groups, e.g., [Co(NH

)

]

, are known as homoleptic. Complexes in

which a metal is bound to more than one kind of donor groups,

e.g., [Co(NH

)

]

, are known as heteroleptic.

Nomenclature is important in Coordination Chemistry because of the

need to have an unambiguous method of describing formulas and

writing systematic names, particularly when dealing with isomers. The

formulas and names adopted for coordination entities are based on the

recommendations of the International Union of Pure and Applied

Chemistry (IUPAC).

Fig. 9.1: Shapes of different coordination polyhedra. M represents

the central atom/ion and L, a unidentate ligand.

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249 Coordination Compounds

The formula of a compound is a shorthand tool used to provide basic

information about the constitution of the compound in a concise and

convenient manner. Mononuclear coordination entities contain a single

central metal atom. The following rules are applied while writing the formulas:

(i) The central atom is listed first.

(ii) The ligands are then listed in alphabetical order. The placement of

a ligand in the list does not depend on its charge.

(iii) Polydentate ligands are also listed alphabetically. In case of

abbreviated ligand, the first letter of the abbreviation is used to

determine the position of the ligand in the alphabetical order.

(iv) The formula for the entire coordination entity, whether charged or

not, is enclosed in square brackets. When ligands are polyatomic,

their formulas are enclosed in parentheses. Ligand abbreviations

are also enclosed in parentheses.

(v) There should be no space between the ligands and the metal

within a coordination sphere.

(vi) When the formula of a charged coordination entity is to be written

without that of the counter ion, the charge is indicated outside the

square brackets as a right superscript with the number before the

sign. For example, [Co(CN)

]

3–

, [Cr(H

]

, etc.

(vii) The charge of the cation(s) is balanced by the charge of the anion(s).

The names of coordination compounds are derived by following the

principles of additive nomenclature. Thus, the groups that surround the

central atom must be identified in the name. They are listed as prefixes

to the name of the central atom along with any appropriate multipliers.

The following rules are used when naming coordination compounds:

(i) The cation is named first in both positively and negatively charged

coordination entities.

(ii) The ligands are named in an alphabetical order before the name of the

central atom/ion. (This procedure is reversed from writing formula).

(iii) Names of the anionic ligands end in –o, those of neutral and cationic

ligands are the same except aqua for H

O, ammine for NH

, carbonyl

for CO and nitrosyl for NO. While writing the formula of coordination

entity, these are enclosed in brackets ( ).

(iv) Prefixes mono, di, tri, etc., are used to indicate the number of the

individual ligands in the coordination entity. When the names of

the ligands include a numerical prefix, then the terms, bis, tris,

tetrakis are used, the ligand to which they refer being placed in

parentheses. For example, [NiCl

(PPh

)

] is named as

dichloridobis(triphenylphosphine)nickel(II).

(v) Oxidation state of the metal in cation, anion or neutral coordination

entity is indicated by Roman numeral in parenthesis.

(vi) If the complex ion is a cation, the metal is named same as the

element. For example, Co in a complex cation is called cobalt and

Pt is called platinum. If the complex ion is an anion, the name of

the metal ends with the suffix – ate. For example, Co in a complex

anion,

( )

SCN

−

 

 

is called cobaltate. For some metals, the Latin

names are used in the complex anions, e.g., ferrate for Fe.

9.3.2 Naming of

Mononuclear

Coordination

Compounds

Note: The 2004 IUPAC

draft recommends that

ligands will be sorted

alphabetically,

irrespective of charge.

Note: The 2004 IUPAC

draft recommends that

anionic ligands will end

with–ido so that chloro

would become chlorido,

etc.

9.3.1 Formulas of

Mononuclear

Coordination

Entities

2020-21

250Chemistry

(vii) The neutral complex molecule is named similar to that of the

complex cation.

The following examples illustrate the nomenclature for coordination

compounds.

1. [Cr(NH

)

]Cl

is named as:

triamminetriaquachromium(III) chloride

Explanation: The complex ion is inside the square bracket, which is

a cation. The amine ligands are named before the aqua ligands

according to alphabetical order. Since there are three chloride ions in

the compound, the charge on the complex ion must be +3 (since the

compound is electrically neutral). From the charge on the complex

ion and the charge on the ligands, we can calculate the oxidation

number of the metal. In this example, all the ligands are neutral

molecules. Therefore, the oxidation number of chromium must be

the same as the charge of the complex ion, +3.

2. [Co(H

NCH

)

]

(SO

)

is named as:

tris(ethane-1,2–diamine)cobalt(III) sulphate

Explanation: The sulphate is the counter anion in this molecule.

Since it takes 3 sulphates to bond with two complex cations, the

charge on each complex cation must be +3. Further, ethane-1,2–

diamine is a neutral molecule, so the oxidation number of cobalt

in the complex ion must be +3. Remember that you never have to

indicate the number of cations and anions in the name of an

ionic compound.

3. [Ag(NH

)

][Ag(CN)

] is named as:

diamminesilver(I) dicyanidoargentate(I)

Write the formulas for the following coordination compounds:

(a) Tetraammineaquachloridocobalt(III) chloride

(b) Potassium tetrahydroxidozincate(II)

(d) Dichloridobis(ethane-1,2-diamine)cobalt(III)

(e) Tetracarbonylnickel(0)

(a) [Co(NH

)

O)Cl]Cl

(b) K

[Zn(OH)

] (c) K

[Al(C

)

]

(d) [CoCl

(en)

]

(e) [Ni(CO)

]

Write the IUPAC names of the following coordination compounds:

(a) [Pt(NH

)

Cl(NO

)] (b) K

[Cr(C

)

] (c) [CoCl

(en)

]Cl

(d) [Co(NH

)

(CO

)]Cl (e) Hg[Co(SCN)

]

(a) Diamminechloridonitrito-N-platinum(II)

(b) Potassium trioxalatochromate(III)

(d) Pentaamminecarbonatocobalt(III) chloride

(e) Mercury (I) tetrathiocyanato-S-cobaltate(III)

Example 9.2Example 9.2

Example 9.2

SolutionSolution

Solution

Example 9.3Example 9.3

Example 9.3

SolutionSolution

Solution

Notice how the name of

the metal differs in

cation and anion even

though they contain the

same metal ions.

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251 Coordination Compounds

9.4.1 Geometric Isomerism

Isomers are two or more compounds that have the same chemical

formula but a different arrangement of atoms. Because of the different

arrangement of atoms, they differ in one or more physical or chemical

properties. Two principal types of isomerism are known among

coordination compounds. Each of which can be further subdivided.

(a) Stereoisomerism

(i) Geometrical isomerism (ii) Optical isomerism

(b) Structural isomerism

(i) Linkage isomerism (ii) Coordination isomerism

(iii) Ionisation isomerism (iv) Solvate isomerism

Stereoisomers have the same chemical formula and chemical

bonds but they have different spatial arrangement. Structural isomers

have different bonds. A detailed account of these isomers are

given below.

This type of isomerism arises in heteroleptic

complexes due to different possible geometric

arrangements of the ligands. Important examples of

this behaviour are found with coordination numbers

4 and 6. In a square planar complex of formula

[MX

] (X and L are unidentate), the two ligands X

may be arranged adjacent to each other in a cis

isomer, or opposite to each other in a trans isomer

as depicted in Fig. 9.2.

Other square planar complex of the type

MABXL (where A, B, X, L are unidentates)

shows three isomers-two cis and one trans.

You may attempt to draw these structures.

Such isomerism is not possible for a tetrahedral

geometry but similar behaviour is possible in

octahedral complexes of formula [MX

] in

which the two ligands X may be oriented cis

or trans to each other (Fig. 9.3).

9.49.4

9.4

Isomerism inIsomerism in

Isomerism in

CoordinationCoordination

Coordination

CompoundsCompounds

Compounds

Fig. 9.2: Geometrical isomers (cis and

trans) of Pt [NH

)

]

N H

cis trans

Fig. 9.3: Geometrical isomers (cis and trans)

of [Co(NH

)

]

Intext QuestionsIntext Questions

Intext Questions

9.1 Write the formulas for the following coordination compounds:

(i) Tetraamminediaquacobalt(III) chloride

(ii) Potassium tetracyanidonickelate(II)

(iii) Tris(ethane–1,2–diamine) chromium(III) chloride

(iv) Amminebromidochloridonitrito-N-platinate(II)

(v) Dichloridobis(ethane–1,2–diamine)platinum(IV) nitrate

(vi) Iron(III) hexacyanidoferrate(II)

9.2 Write the IUPAC names of the following coordination compounds:

(i) [Co(NH

)

]Cl

(ii) [Co(NH

)

Cl]Cl

(iii) K

[Fe(CN)

]

(iv) K

[Fe(C

)

] (v) K

[PdCl

] (vi) [Pt(NH

)

Cl(NH

)]Cl

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252Chemistry

This type of isomerism also

arises when didentate ligands

L–L [e.g., NH

(en)]

are present in complexes of formula

[MX

(L–L)

] (Fig. 9.4).

Another type of geometrical

isomerism occurs in octahedral

coordination entities of the type

[Ma

] like [Co(NH

)

(NO

)

]. If

three donor atoms of the same

ligands occupy adjacent positions

at the corners of an octahedral

face, we have the facial (fac)

isomer. When the positions are

around the meridian of the

octahedron, we get the meridional

(mer) isomer (Fig. 9.5).

Fig. 9.4: Geometrical isomers (cis and trans)

of [CoCl

(en)

]

Why is geometrical isomerism not possible in tetrahedral complexes

having two different types of unidentate ligands coordinated with

the central metal ion ?

Tetrahedral complexes do not show geometrical isomerism because

the relative positions of the unidentate ligands attached to the central

metal atom are the same with respect to each other.

SolutionSolution

Solution

Optical isomers are mirror images that

cannot be superimposed on one

another. These are called as

enantiomers. The molecules or ions

that cannot be superimposed are

called chiral. The two forms are called

dextro (d) and laevo (l) depending

upon the direction they rotate the

plane of polarised light in a

polarimeter (d rotates to the right, l to

the left). Optical isomerism is common

in octahedral complexes involving

didentate ligands (Fig. 9.6).

In a coordination

entity of the type

[PtCl

(en)

]

, only the

cis-isomer shows optical

activity (Fig. 9.7).

9.4.2 Optical Isomerism

Fig.9.6: Optical isomers (d and l) of [Co(en)

]

Fig.9.7

Optical isomers (d

and l) of cis-

[PtCl

(en)

]

Fig. 9.5

The facial (fac) and

meridional (mer)

isomers of

[Co(NH

)

(NO

)

]

Example 9.4Example 9.4

Example 9.4

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253 Coordination Compounds

Linkage isomerism arises in a coordination compound containing

ambidentate ligand. A simple example is provided by complexes

containing the thiocyanate ligand, NCS

–

, which may bind through the

nitrogen to give M–NCS or through sulphur to give M–SCN. Jørgensen

discovered such behaviour in the complex [Co(NH

)

(NO

)]Cl

, which is

obtained as the red form, in which the nitrite ligand

is bound through

oxygen (–ONO), and as the yellow form, in which the nitrite ligand is

bound through nitrogen (–NO

This type of isomerism arises from the interchange of ligands between

cationic and anionic entities of different metal ions present in a complex.

An example is provided by [Co(NH

)

][Cr(CN)

], in which the NH

ligands

are bound to Co

and the CN

–

ligands to Cr

. In its coordination

isomer [Cr(NH

)

][Co(CN)

], the NH

ligands are bound to Cr

and the

–

ligands to Co

This form of isomerism arises when the counter ion in a complex salt

is itself a potential ligand and can displace a ligand which can then

become the counter ion. An example is provided by the ionisation

isomers [Co(NH

)

(SO

)]Br and [Co(NH

)

Br]SO

9.4.3 Linkage

Isomerism

9.4.4 Coordination

Isomerism

9.4.5 Ionisation

Isomerism

Out of the following two coordination entities which is chiral

(optically active)?

(a) cis-[CrCl

(ox)

]

3–

(b) trans-[CrCl

(ox)

]

3–

The two entities are represented as

Draw structures of geometrical isomers of [Fe(NH

)

(CN)

]

–

SolutionSolution

Solution

Out of the two, (a) cis - [CrCl

(ox)

]

is chiral (optically active).

Example 9.5Example 9.5

Example 9.5

SolutionSolution

Solution

Example 9.6Example 9.6

Example 9.6

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254Chemistry

This form of isomerism is known as ‘hydrate isomerism’ in case where

water is involved as a solvent. This is similar to ionisation isomerism.

Solvate isomers differ by whether or not a solvent molecule is directly

bonded to the metal ion or merely present as free solvent molecules

in the crystal lattice. An example is provided by the aqua

complex [Cr(H

]Cl

(violet) and its solvate isomer [Cr(H

Cl]Cl

(grey-green).

9.4.6 Solvate

Isomerism

Werner was the first to describe the bonding features in coordination

compounds. But his theory could not answer basic questions like:

(i) Why only certain elements possess the remarkable property of

forming coordination compounds?

(ii) Why the bonds in coordination compounds have directional

properties?

(iii) Why coordination compounds have characteristic magnetic and

optical properties?

Many approaches have been put forth to explain the nature of

bonding in coordination compounds viz. Valence Bond Theory (VBT),

Crystal Field Theory (CFT), Ligand Field Theory (LFT) and Molecular

Orbital Theory (MOT). We shall focus our attention on elementary

treatment of the application of VBT and CFT to coordination compounds.

According to this theory, the metal atom or ion under the influence of

ligands can use its (n-1)d, ns, np or ns, np, nd orbitals for hybridisation

to yield a set of equivalent orbitals of definite geometry such as octahedral,

tetrahedral, square planar and so on (Table 9.2). These hybridised orbitals

are allowed to overlap with ligand orbitals that can donate electron pairs

for bonding. This is illustrated by the following examples.

9.59.5

9.5

Bonding inBonding in

Bonding in

CoordinationCoordination

Coordination

CompoundsCompounds

Compounds

9.5.1 Valence

Bond

Theory

Table 9.2: Number of Orbitals and Types of Hybridisations

4 sp

Tetrahedral

4 dsp

Square planar

5 sp

d Trigonal bipyramidal

6 sp

Octahedral

6 d

Octahedral

Coordination

number

Type of

hybridisation

Distribution of hybrid

orbitals in space

Intext QuestionsIntext Questions

Intext Questions

9.3 Indicate the types of isomerism exhibited by the following complexes and

draw the structures for these isomers:

(i) K[Cr(H

)

(ii) [Co(en)

]Cl

(iii) [Co(NH

)

(NO

)](NO

)

(iv) [Pt(NH

)(H

O)Cl

]

9.4 Give evidence that [Co(NH

)

Cl]SO

and [Co(NH

)

(SO

)]Cl are ionisation

isomers.

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255 Coordination Compounds

It is usually possible to predict the geometry of a complex from

the knowledge of its

magnetic behaviour on

the basis of the valence

bond theory.

In the diamagnetic

octahedral complex,

[Co(NH

)

]

, the cobalt ion

is in +3 oxidation state

and has the electronic

configuration 3d

. The

hybridisation scheme is as

shown in diagram.

Orbitals of Co ion

sp d

3 2

hybridised

orbitals of Co

[CoF ]

(outer orbital or

high spin complex)

3–

Six pairs of electrons

from six F ions

–

4s 4p

sp d

3 3

hybrid

Six pairs of electrons, one from each NH

molecule, occupy the six

hybrid orbitals. Thus, the complex has octahedral geometry and is

diamagnetic because of the absence of unpaired electron. In the formation

of this complex, since the inner d orbital (3d) is used in hybridisation,

the complex, [Co(NH

)

]

is called an inner orbital or low spin or spin

paired complex. The paramagnetic octahedral complex, [CoF

]

3–

uses

outer orbital (4d ) in hybridisation (sp

). It is thus called outer orbital

or high spin or spin free complex. Thus:

In tetrahedral complexes

one s and three p orbitals

are hybridised to form four

equivalent orbitals oriented

tetrahedrally. This is ill-

ustrated below for [NiCl

]

Here nickel is in +2

oxidation state and the ion

has the electronic

configuration 3d

. The

hybridisation scheme is as

shown in diagram.

Each Cl

–

ion donates a pair of electrons. The compound is

paramagnetic since it contains two unpaired electrons. Similarly,

[Ni(CO)

] has tetrahedral geometry but is diamagnetic since nickel is in

zero oxidation state and contains no unpaired electron.

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256Chemistry

Orbitals of Ni ion

dsp hybridised

orbitals of Ni

[Ni(CN) ]

(low spin complex)

2–

4s 4p

Four pairs of electrons

from 4 CN groups

–

dsp

hydrid3d

9.5.2 Magnetic

Properties

Coordination

Compounds

In the square planar complexes, the hybridisation involved is dsp

An example is [Ni(CN)

]

2–

. Here nickel is in +2 oxidation state and has

the electronic configuration 3d

. The hybridisation scheme is as shown

in diagram:

Each of the hybridised orbitals receives a pair of electrons from a

cyanide ion. The compound is diamagnetic as evident from the absence

of unpaired electron.

It is important to note that the hybrid orbitals do not actually exist.

In fact, hybridisation is a mathematical manipulation of wave equation

for the atomic orbitals involved.

The magnetic moment of coordination compounds can be measured

by the magnetic susceptibility experiments. The results can be used to

obtain information about the number of unpaired electrons (page 228)

and hence structures adopted by metal complexes.

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 Ti

); V

); Cr

); 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 d

(Cr

, Mn

), d

(Mn

, Fe

), d

(Fe

, Co

)

cases, a vacant pair of d orbitals results only by pairing of 3d electrons

which leaves two, one and zero unpaired electrons, respectively.

The magnetic data agree with maximum spin pairing in many cases,

especially with coordination compounds containing d

ions. However,

with species containing d

and d

ions there are complications. [Mn(CN)

]

3–

has magnetic moment of two unpaired electrons while [MnCl

]

3–

has a

paramagnetic moment of four unpaired electrons. [Fe(CN)

]

3–

has magnetic

moment of a single unpaired electron while [FeF

]

3–

has a paramagnetic

moment of five unpaired electrons. [CoF

]

3–

is paramagnetic with four

unpaired electrons while [Co(C

)

]

3–

is diamagnetic. This apparent

anomaly is explained by valence bond theory in terms of formation of

inner orbital and outer orbital coordination entities. [Mn(CN)

]

3–

, [Fe(CN)

]

3–

and [Co(C

)

]

3–

are inner orbital complexes involving d

hybridisation,

the former two complexes are paramagnetic and the latter diamagnetic.

On the other hand, [MnCl

]

3–

, [FeF

]

3–

and [CoF

]

3–

are outer orbital

complexes involving sp

hybridisation and are paramagnetic

corresponding to four, five and four unpaired electrons.

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257 Coordination Compounds

The spin only magnetic moment of [MnBr

]

2–

is 5.9 BM. Predict the

geometry of the complex ion ?

Since the coordination number of Mn

ion in the complex ion is 4, it

will be either tetrahedral (sp

hybridisation) or square planar (dsp

hybridisation). But the fact that the magnetic moment of the complex

ion is 5.9 BM, it should be tetrahedral in shape rather than square

planar because of the presence of five unpaired electrons in the d orbitals.

Example 9.7Example 9.7

Example 9.7

SolutionSolution

Solution

9.5.3 Limitations

of Valence

Bond

Theory

9.5.4 Crystal Field

Theory

While the VB theory, to a larger extent, explains the formation, structures

and magnetic behaviour of coordination compounds, it suffers from

the following shortcomings:

(i) It involves a number of assumptions.

(ii) It does not give quantitative interpretation of magnetic data.

(iii) It does not explain the colour exhibited by coordination compounds.

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

The crystal field theory (CFT) is an electrostatic model which considers

the metal-ligand bond to be ionic arising purely from electrostatic

interactions between the metal ion and the ligand. Ligands are treated

as point charges in case of anions or point dipoles in case of neutral

molecules. The five d orbitals in an isolated gaseous metal atom/ion

have same energy, i.e., they are degenerate. This 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 the negative ends of dipolar molecules

like NH

and H

O) in a complex, it becomes asymmetrical and the

degeneracy of the d orbitals is lifted. It results in splitting of the d

orbitals. The pattern of splitting depends upon the nature of the crystal

field. Let us explain this splitting in different crystal fields.

(a) Crystal field splitting in octahedral coordination entities

In an octahedral coordination entity with six ligands surrounding

the metal atom/ion, there will be repulsion between the electrons in

metal d orbitals and the electrons (or negative charges) of the ligands.

Such a repulsion is more when the metal d orbital is directed towards

the ligand than when it is away from the ligand. Thus, the

2 2

x y

−

and

orbitals which point towards the axes along the direction of

the ligand will experience more repulsion and will be raised in

energy; and the d

, d

and d

orbitals which are directed between

the axes will be lowered in energy relative to the average energy in

the spherical crystal field. Thus, the degeneracy of the d orbitals

has been removed due to ligand electron-metal electron repulsions

in the octahedral complex to yield three orbitals of lower energy, t

set and two orbitals of higher energy, e

set. This splitting of the

2020-21

258Chemistry

degenerate levels due to the

presence of ligands in a

definite geometry is termed as

crystal field splitting and the

energy separation is denoted

by ∆

(the subscript o is for

octahedral) (Fig.9.8). Thus, the

energy of the two e

orbitals

will increase by (3/5) ∆

and

that of the three t

will

decrease by (2/5)∆

The crystal field splitting,

∆

, depends upon the field

produced by the ligand and

charge on the metal ion. Some

ligands are able to produce

strong fields in which case, the

splitting will be large whereas

others produce weak fields

and consequently result in

small splitting of d orbitals.

In general, ligands can be arranged in a series in the order of increasing

field strength as given below:

–

< Br

–

< SCN

–

< Cl

–

< S

2–

< F

–

< OH

–

< C

2–

< H

O < NCS

–

< edta

4–

< NH

< en < CN

–

< CO

Such a series is termed as spectrochemical series. It is an

experimentally determined series based on the absorption of light

by complexes with different ligands. Let us assign electrons in the d

orbitals of metal ion in octahedral coordination entities. Obviously,

the single d electron occupies one of the lower energy t

orbitals. In

and d

coordination entities, the d electrons occupy the t

orbitals

singly in accordance with the Hund’s rule. For d

ions, two possible

patterns of electron distribution arise: (i) the fourth electron could

either enter the t

level and pair with an existing electron, or (ii) it

could avoid paying the price of the pairing energy by occupying the

level. Which of these possibilities occurs, depends on the relative

magnitude of the crystal field splitting, ∆

and the pairing energy, P

(P represents the energy required for electron pairing in a single

orbital). The two options are:

(i) If ∆

< P, the fourth electron enters one of the e

orbitals giving the

configuration

3 1

2g g

t e

. Ligands for which ∆

< P are known as weak

field ligands and form high spin complexes.

(ii) If ∆

> P, it becomes more energetically favourable for the fourth

electron to occupy a t

orbital with configuration t

. Ligands

which produce this effect are known as strong field ligands and

form low spin complexes.

Calculations show that d

to d

coordination entities are more

stable for strong field as compared to weak field cases.

Fig.9.8: d orbital splitting in an octahedral crystal field

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259 Coordination Compounds

Fig.9.9: d orbital splitting in a tetrahedral crystal

field.

In the previous Unit, we learnt that one of the most distinctive

properties of transition metal complexes is their wide range of colours.

This means that some of the visible spectrum is being removed from

white light as it passes through the sample, so the light that emerges

is no longer white. The colour of the complex is complementary to

that which is absorbed. The complementary colour is the colour

generated from the wavelength left over; if green light is absorbed by

the complex, it appears red. Table 9.3 gives the relationship of the

different wavelength absorbed and the colour observed.

9.5.5 Colour in

Coordination

Compounds

Coordinaton

entity

Wavelength of light

absorbed (nm)

Colour of light

absorbed

Colour of coordination

entity

Table 9.3: Relationship between the Wavelength of Light absorbed and the

Colour observed in some Coordination Entities

[CoCl(NH

)

]

535 Yellow Violet

[Co(NH

)

O)]

500 Blue Green Red

[Co(NH

)

]

475 Blue Yellow Orange

[Co(CN)

]

3–

310 Ultraviolet Pale Yellow

[Cu(H

]

600 Red Blue

[Ti(H

]

498 Blue Green Violet

The colour in the coordination compounds can be readily explained

in terms of the crystal field theory. Consider, for example, the complex

[Ti(H

]

, which is violet in colour. This is an octahedral complex

where the single electron (Ti

is a 3d

system) in the metal d orbital is

in the t

level in the ground state of the complex. The next higher state

available for the electron is the empty e

level. If light corresponding to

the energy of blue-green region is absorbed by the complex, it would

excite the electron from t

level to the e

level (t

→ t

Consequently, the complex appears violet in colour (Fig. 9.10). The

crystal field theory attributes the colour of the coordination compounds

to d-d transition of the electron.

(b) Crystal field splitting in tetrahedral coordination entities

In tetrahedral coordination entity formation,

the d orbital splitting (Fig. 9.9) is inverted

and is smaller as compared to the octahedral

field splitting. For the same metal, the same

ligands and metal-ligand distances, it can

be shown that ∆

= (4/9) ∆

. Consequently,

the orbital splitting energies are not

sufficiently large for forcing pairing and,

therefore, low spin configurations are rarely

observed. The ‘g’ subscript is used for the

octahedral and square planar complexes

which have centre of symmetry. Since

tetrahedral complexes lack symmetry, ‘g’

subscript is not used with energy levels.

Not in visible

region

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260Chemistry

It is important to note that

in the absence of ligand,

crystal field splitting does

not occur and hence the

substance is colourless. For

example, removal of water

from [Ti(H

]Cl

on heating

renders it colourless.

Similarly, anhydrous CuSO

is white, but CuSO

.5H

O is

blue in colour. The influence

of the ligand on the colour

of a complex may be illustrated by considering the [Ni(H

]

complex,

which forms when nickel(II) chloride is dissolved in water. If the

didentate ligand, ethane-1,2-diamine(en) is progressively added in the

molar ratios en:Ni, 1:1, 2:1, 3:1, the following series of reactions and

their associated colour changes occur:

[Ni(H

]

(aq) + en (aq) = [Ni(H

(en)]

(aq) + 2H

green pale blue

[Ni(H

(en)]

(aq) + en (aq) = [Ni(H

(en)

]

(aq) + 2H

blue/purple

[Ni(H

(en)

]

(aq) + en (aq) = [Ni(en)

]

(aq) + 2H

violet

This sequence is shown in Fig. 9.11.

Fig.9.11

Aqueous solutions of

complexes of

nickel(II) with an

increasing number of

ethane-1,

2-diamine ligands.

[Ni(H O) ] (aq)

2 6

[Ni(H O) ] (aq)

2 4

en [Ni(H O) ] (aq)

2 4

[Ni(en) ] (aq)

Colour of Some Gem Stones

The colours produced by electronic transitions within the d orbitals of a

transition metal ion occur frequently in everyday life. Ruby [Fig.9.12(a)] is

aluminium oxide (Al

) containing about 0.5-1% Cr

ions (d

), which are

randomly distributed in positions normally occupied by Al

. We may view

these chromium(III) species as octahedral chromium(III) complexes incorporated

into the alumina lattice; d–d transitions at these centres give rise to the colour.

Fig.9.10: Transition of an electron in [Ti(H

]

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261 Coordination Compounds

The crystal field model is successful in explaining the formation,

structures, colour and magnetic properties of coordination compounds

to a large extent. However, from the assumptions 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 the spectrochemical series. Further, it does not take into account

the covalent character of bonding between the ligand and the central

atom. These are some of the weaknesses of CFT, which are explained

by ligand field theory (LFT) and molecular orbital theory which are

beyond the scope of the present study.

9.5.6 Limitations

of Crystal

Field

Theory

The homoleptic carbonyls (compounds containing carbonyl ligands

only) are formed by most of the transition metals. These carbonyls

have simple, well defined structures. Tetracarbonylnickel(0) is

tetrahedral, pentacarbonyliron(0) is trigonalbipyramidal while

hexacarbonyl chromium(0) is octahedral.

Decacarbonyldimanganese(0) is made up of two square pyramidal

Mn(CO)

units joined by a Mn – Mn bond. Octacarbonyldicobalt(0)

has a Co – Co bond bridged by two CO groups (Fig.9.13).

9.6 Bonding in9.6 Bonding in

9.6 Bonding in

Metal

Carbonyls

In emerald [Fig.9.12(b)], Cr

ions occupy octahedral sites

in the mineral beryl

(Be

). The absorption

bands seen in the ruby shift

to longer wavelength, namely

yellow-red and blue, causing

emerald to transmit light in

the green region.

Fig.9.12: (a) Ruby: this gemstone was found in

marble from Mogok, Myanmar; (b) Emerald:

this gemstone was found in Muzo,

Columbia.

(a) (b)

Intext QuestionsIntext Questions

Intext Questions

9.5 Explain on the basis of valence bond theory that [Ni(CN)

]

2–

ion with square

planar structure is diamagnetic and the [NiCl

]

2–

ion with tetrahedral

geometry is paramagnetic.

9.6 [NiCl

]

2–

is paramagnetic while [Ni(CO)

] is diamagnetic though both are

tetrahedral. Why?

9.7 [Fe(H

]

is strongly paramagnetic whereas [Fe(CN)

]

3–

is weakly

paramagnetic. Explain.

9.8 Explain [Co(NH

)

]

is an inner orbital complex whereas [Ni(NH

)

]

is an

outer orbital complex.

9.9 Predict the number of unpaired electrons in the square planar [Pt(CN)

]

2–

ion.

9.10 The hexaquo manganese(II) ion contains five unpaired electrons, while the

hexacyanoion contains only one unpaired electron. Explain using Crystal

Field Theory.

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262Chemistry

9.79.7

9.7

ImportanceImportance

Importance

andand

and

ApplicationsApplications

Applications

ofof

CoordinationCoordination

Coordination

CompoundsCompounds

Compounds

Ni(CO)

Tetrahedral

Fe(CO)

Trigonal bipyramidal

Cr(CO) Octahedral

[Mn (CO) ]

2 10

[Co (CO) ]

2 8

Fig. 9.13Fig. 9.13

Fig. 9.13

Structures of some

representative

homoleptic metal

carbonyls.

The metal-carbon bond in metal carbonyls

possess both σ and π character. The M–C σ bond

is formed by the donation of lone pair of electrons

on the carbonyl carbon into a vacant orbital of

the metal. The M–C π bond is formed by the

donation of a pair of electrons from a filled d orbital

of metal into the vacant antibonding π*

orbital of

carbon monoxide. The metal to ligand bonding

creates a synergic effect which strengthens the

bond between CO and the metal (Fig.9.14).

Fig. 9.14:Fig. 9.14:

Fig. 9.14:

Example of synergic bonding

interactions in a carbonyl

complex.

The coordination compounds are of great importance. These compounds

are widely present in the mineral, plant and animal worlds and are

known to play many important functions in the area of analytical

chemistry, metallurgy, biological systems, industry and medicine. These

are described below:

• Coordination compounds find use in many qualitative and

quantitative chemical analysis. The familiar colour reactions given

by metal ions with a number of ligands (especially chelating ligands),

as a result of formation of coordination entities, form the basis for

their detection and estimation by classical and instrumental methods

of analysis. Examples of such reagents include EDTA, DMG

(dimethylglyoxime), α–nitroso–β–naphthol, cupron, etc.

• Hardness of water is estimated by simple titration with Na

EDTA.

The Ca

and Mg

ions form stable complexes with EDTA. The

selective estimation of these ions can be done due to difference in

the stability constants of calcium and magnesium complexes.

• Some important extraction processes of metals, like those of silver and

gold, make use of complex formation. Gold, for example, combines with

cyanide in the presence of oxygen and water to form the coordination

entity [Au(CN)

]

–

in aqueous solution. Gold can be separated in metallic

form from this solution by the addition of zinc (Unit 6).

• Similarly, purification of metals can be achieved through formation

and subsequent decomposition of their coordination compounds.

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263 Coordination Compounds

SummarySummary

Summary

The chemistry of coordination compounds is an important and challenging

area of modern inorganic chemistry. During the last fifty years, advances in this

area, have provided development of new concepts and models of bonding and

molecular structure, novel breakthroughs in chemical industry and vital

insights into the functioning of critical components of biological systems.

The first systematic attempt at explaining the formation, reactions, structure

and bonding of a coordination compound was made by A. Werner. His theory

postulated the use of two types of linkages (primary and secondary) by a

metal atom/ion in a coordination compound. In the modern language of chemistry

these linkages are recognised as the ionisable (ionic) and non-ionisable (covalent)

bonds, respectively. Using the property of isomerism, Werner predicted the

geometrical shapes of a large number of coordination entities.

The Valence Bond Theory (VBT) explains with reasonable success, the

formation, magnetic behaviour and geometrical shapes of coordination compounds.

It, however, fails to provide a quantitative interpretation of magnetic behaviour

and has nothing to say about the optical properties of these compounds.

The Crystal Field Theory (CFT) to coordination compounds is based on

the effect of different crystal fields (provided by the ligands taken as point charges),

For example, impure nickel is converted to [Ni(CO)

], which is

decomposed to yield pure nickel.

• Coordination compounds are of great importance in biological

systems. The pigment responsible for photosynthesis, chlorophyll,

is a coordination compound of magnesium. Haemoglobin, the red

pigment of blood which acts as oxygen carrier is a coordination

compound of iron. Vitamin B

, cyanocobalamine, the anti–

pernicious anaemia factor, is a coordination compound of cobalt.

Among the other compounds of biological importance with

coordinated metal ions are the enzymes like, carboxypeptidase A

and carbonic anhydrase (catalysts of biological systems).

• Coordination compounds are used as catalysts for many industrial

processes. Examples include rhodium complex, [(Ph

RhCl], a

Wilkinson catalyst, is used for the hydrogenation of alkenes.

• Articles can be electroplated with silver and gold much more

smoothly and evenly from solutions of the complexes, [Ag(CN)

]

–

and [Au(CN)

]

–

than from a solution of simple metal ions.

• In black and white photography, the developed film is fixed by

washing with hypo solution which dissolves the undecomposed

AgBr to form a complex ion, [Ag(S

)

]

3–

• There is growing interest in the use of chelate therapy in medicinal

chemistry. An example is the treatment of problems caused by the

presence of metals in toxic proportions in plant/animal systems.

Thus, excess of copper and iron are removed by the chelating ligands

D–penicillamine and desferrioxime B via the formation of coordination

compounds. EDTA is used in the treatment of lead poisoning. Some

coordination compounds of platinum effectively inhibit the growth

of tumours. Examples are: cis–platin and related compounds.

2020-21

264Chemistry

on the degeneracy of d orbital energies of the central metal atom/ion. The

splitting of the d orbitals provides different electronic arrangements in strong

and weak crystal fields. The treatment provides for quantitative estimations of

orbital separation energies, magnetic moments and spectral and stability

parameters. However, the assumption that ligands consititute point charges

creates many theoretical difficulties.

The metal–carbon bond in metal carbonyls possesses both σ and π character.

The ligand to metal is σ bond and metal to ligand is π bond. This unique synergic

bonding provides stability to metal carbonyls.

Coordination compounds are of great importance. These compounds provide

critical insights into the functioning and structures of vital components of

biological systems. Coordination compounds also find extensive applications in

metallurgical processes, analytical and medicinal chemistry.

Exercises

9.1 Explain the bonding in coordination compounds in terms of Werner’s postulates.

9.2 FeSO

solution mixed with (NH

)

solution in 1:1 molar ratio gives the

test of Fe

ion but CuSO

solution mixed with aqueous ammonia in 1:4

molar ratio does not give the test of Cu

ion. Explain why?

9.3 Explain with two examples each of the following: coordination entity, ligand,

coordination number, coordination polyhedron, homoleptic and heteroleptic.

9.4 What is meant by unidentate, didentate and ambidentate ligands? Give two

examples for each.

9.5 Specify the oxidation numbers of the metals in the following coordination entities:

(i) [Co(H

O)(CN)(en)

]

(iii) [PtCl

]

2–

(v) [Cr(NH

)

]

(ii) [CoBr

(en)

]

(iv) K

[Fe(CN)

]

9.6 Using IUPAC norms write the formulas for the following:

(i) Tetrahydroxidozincate(II) (vi) Hexaamminecobalt(III) sulphate

(ii) Potassium tetrachloridopalladate(II) (vii) Potassium tri(oxalato)chromate(III)

(iii) Diamminedichloridoplatinum(II) (viii) Hexaammineplatinum(IV)

(iv) Potassium tetracyanidonickelate(II) (ix) Tetrabromidocuprate(II)

(v) Pentaamminenitrito-O-cobalt(III) (x) Pentaamminenitrito-N-cobalt(III)

9.7 Using IUPAC norms write the systematic names of the following:

(i) [Co(NH

)

]Cl

(iv) [Co(NH

)

Cl(NO

)]Cl (vii) [Ni(NH

)

]Cl

(ii) [Pt(NH

)

Cl(NH

)]Cl (v) [Mn(H

]

(viii) [Co(en)

]

(iii) [Ti(H

]

(vi) [NiCl

]

2–

(ix) [Ni(CO)

]

9.8 List various types of isomerism possible for coordination compounds, giving

an example of each.

9.9 How many geometrical isomers are possible in the following coordination entities?

(i) [Cr(C

)

]

3–

(ii) [Co(NH

)

]

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265 Coordination Compounds

9.10 Draw the structures of optical isomers of:

(i) [Cr(C

)

]

3–

(ii) [PtCl

(en)

]

(iii) [Cr(NH

)

(en)]

9.11 Draw all the isomers (geometrical and optical) of:

(i) [CoCl

(en)

]

(ii) [Co(NH

)Cl(en)

]

(iii) [Co(NH

)

(en)]

9.12Write all the geometrical isomers of [Pt(NH

)(Br)(Cl)(py)] and how many of these

will exhibit optical isomers?

9.13 Aqueous copper sulphate solution (blue in colour) gives:

(i) a green precipitate with aqueous potassium fluoride and

(ii) a bright green solution with aqueous potassium chloride. Explain these

experimental results.

9.14 What is the coordination entity formed when excess of aqueous KCN is

added to an aqueous solution of copper sulphate? Why is it that no precipitate

of copper sulphide is obtained when H

S(g) is passed through this solution?

9.15 Discuss the nature of bonding in the following coordination entities on the

basis of valence bond theory:

(i) [Fe(CN)

]

4–

(ii) [FeF

]

3–

(iii) [Co(C

)

]

3–

(iv) [CoF

]

3–

9.16 Draw figure to show the splitting of d orbitals in an octahedral crystal field.

9.17 What is spectrochemical series? Explain the difference between a weak

field ligand and a strong field ligand.

9.18 What is crystal field splitting energy? How does the magnitude of ∆

decide

the actual configuration of d orbitals in a coordination entity?

9.19 [Cr(NH

)

]

is paramagnetic while [Ni(CN)

]

2–

is diamagnetic. Explain why?

9.20 A solution of [Ni(H

]

is green but a solution of [Ni(CN)

]

2–

is colourless.

Explain.

9.21 [Fe(CN)

]

4–

and [Fe(H

]

are of different colours in dilute solutions. Why?

9.22 Discuss the nature of bonding in metal carbonyls.

9.23 Give the oxidation state, d orbital occupation and coordination number of

the central metal ion in the following complexes:

(i) K

[Co(C

)

] (iii) (NH

)

[CoF

]

(ii) cis-[CrCl

(en)

]Cl (iv) [Mn(H

]SO

9.24 Write down the IUPAC name for each of the following complexes and indicate

the oxidation state, electronic configuration and coordination number. Also

give stereochemistry and magnetic moment of the complex:

(i) K[Cr(H

)

].3H

O (iii) [CrCl

(py)

] (v) K

[Mn(CN)

]

(ii) [Co(NH

)

]Cl

(iv) Cs[FeCl

]

9.25 Explain the violet colour of the complex [Ti(H

]

on the basis of crystal

field theory.

9.26 What is meant by the chelate effect? Give an example.

9.27 Discuss briefly giving an example in each case the role of coordination

compounds in:

(i) biological systems (iii) analytical chemistry

(ii) medicinal chemistry and (iv) extraction/metallurgy of metals.

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266Chemistry

Answers to Some Intext Questions

9.1 (i) [Co(NH

)

]Cl

(iv) [Pt(NH

)BrCl(NO

)]

–

(ii) K

[Ni(CN)

] (v) [PtCl

(en)

](NO

)

(iii) [Cr(en)

]Cl

(vi) Fe

[Fe(CN)

]

9.2 (i) Hexaamminecobalt(III) chloride

(ii) Pentaamminechloridocobalt(III) chloride

(iii) Potassium hexacyanidoferrate(III)

(iv) Potassium trioxalatoferrate(III)

(v) Potassium tetrachloridopalladate(II)

(vi) Diamminechlorido(methanamine)platinum(II) chloride

9.3 (i) Both geometrical (cis-, trans-) and optical isomers for cis can exist.

(ii) Two optical isomers can exist.

(iii) There are 10 possible isomers. (Hint: There are geometrical, ionisation

and linkage isomers possible).

(iv) Geometrical (cis-, trans-) isomers can exist.

9.4 The ionisation isomers dissolve in water to yield different ions and thus

react differently to various reagents:

[Co(NH

)

Br]SO

+ Ba

→ BaSO

(s)

[Co(NH

)

]Br + Ba

→ No reaction

[Co(NH

)

Br]SO

+ Ag

→ No reaction

[Co(NH

)

]Br + Ag

→ AgBr (s)

9.6 In Ni(CO)

, Ni is in zero oxidation state whereas in NiCl

2–

, it is in +2 oxidation

state. In the presence of CO ligand, the unpaired d electrons of Ni pair up

but Cl

–

being a weak ligand is unable to pair up the unpaired electrons.

9.7 In presence of CN

–

, (a strong ligand) the 3d electrons pair up leaving only

one unpaired electron. The hybridisation is d

forming inner orbital

complex. In the presence of H

O, (a weak ligand), 3d electrons do not pair

up. The hybridisation is sp

forming an outer orbital complex containing

five unpaired electrons, it is strongly paramagnetic.

9.8 In the presence of NH

, the 3d electrons pair up leaving two d orbitals

empty to be involved in d

hybridisation forming inner orbital complex

in case of [Co(NH

)

]

9.28 How many ions are produced from the complex Co(NH

)

in solution?

(i) 6 (ii) 4 (iii) 3 (iv) 2

9.29 Amongst the following ions which one has the highest magnetic moment value?

(i) [Cr(H

]

(ii) [Fe(H

]

(iii) [Zn(H

]

9.31 Amongst the following, the most stable complex is

(i) [Fe(H

]

(ii) [Fe(NH

)

]

(iii) [Fe(C

)

]

3–

(iv) [FeCl

]

3–

9.32 What will be the correct order for the wavelengths of absorption in the visible

region for the following:

[Ni(NO

)

]

4–

, [Ni(NH

)

]

, [Ni(H

]

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267 Coordination Compounds

In Ni(NH

)

, Ni is in +2 oxidation state and has d

configuration,

the hybridisation involved is sp

forming outer orbital complex.

9.9 For square planar shape, the hybridisation is dsp

. Hence the

unpaired electrons in 5d orbital pair up to make one d orbital

empty for dsp

hybridisation. Thus there is no unpaired electron.

2020-21