1

9. Coordination Compounds
Double Salt :
When to salts in equimolar ratio are crystallised together from their
saturated solution they are called double salts.
These exist as such in the crystalline state. When dissolved in water (in
aqueous solution), it loses its identity and dissociates to give individual ions.
Examples: K2SO4.Al2(SO4)3.24H2O (Potash alum)
Fe SO4.(NH4)2 SO4 .6H2O (Mohr’s Salt)

Co-ordination compound:
Co-ordination compounds retain their identities in solid state as well as in
aqueous solution. In aqueous solution, these compounds do not lose their identity
and give at least one complex ion.

Sl.No.
Distinction between double salts and co-ordination compounds
Double Salts Co-ordination Compounds
1 Double salts are normally formed In co-ordination compounds the
from two salts mixed in equimolar simple salts may or may not be mixed
proportions from their solution. in equimolar proportions.
2 Double salts exist only in solid states Co-ordination compounds do not lose
and in aqueous solution, they their identity in the solid state as
dissociates into ions. well in aqueous solutions ie., they do
not dissociated into ions.
3 The constituting salts are of ionic In the complex part, species present
nature. are linked with metal atom or ion by
co-ordinate bonds.
4 The properties of the double salts The properties of complex part
are of those of the constituting ions enclosed in square bracket are quite
present. different from the constituents
present.
5 In double salts, the metal ions show In co-ordination compounds, metal
their normal valencies. ions show primary and secondary
valencies.
6 Example: Example:
K2SO4.Al2(SO4)3.24H2O (Potash alum) K4[Fe(CN)6] ,
Fe SO4.(NH4)2 SO4 .6H2O (Mohr’s Salt) [Co(NH3)6]Cl3.
Complex ion:
A complex ion is an electrically charged (cationic or anionic) species which
consists of a central metal atom or ion surrounded by a group of ions or neutral
molecules e.g., [Co(NH3)6]3+. [Fe(CN)6]3-, [Cu(NH3)4]2+.
 2

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.
Precipitation Studies:
Corresponding to 3Cl-
Corresponding to 2Cl-
Corresponding to 1Cl-
Corresponding to 1Cl-
Conductance measurements:
Formulation of Cobalt(III) Chloride-Ammonia Complexes

Note that, the last two compounds in the above Table have identical empirical
formula, CoCl3.4NH3, but distinct properties(called Isomers).
Based on these observations, the compounds are formulated as shown in the
above Table, where the atoms within the square brackets form a single entity
which does not dissociate under the reaction conditions.

Werner 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.
4. The secondary valence is equal to the coordination number and is fixed for a
metal.
5. Secondary valence is responsible their structure (or) spatial arrangements.

He further postulated that

octahedral, tetrahedral and square planar geometrical shapes are more common
in coordination compounds of transition metals.
 3

Example:
[Co(NH3)6]3+, [CoCl(NH3)5]2+ and [CoCl2(NH3)4]+ are octahedral

[Ni(CO)4] and [PtCl4]2– are tetrahedral and square planar, respectively.

Ligands

K4 [ Fe (CN)6 ]
Co-ordination
Number
Ionisation Sphere Central
metal ion

Co-ordination Sphere (or)

Non- Ionisation Sphere

Some important terms involved in Co-ordination Chemistry

a) Complex ion (or) Coordination entity:

A complex ion is an electrically charged (cationic or anionic) species which
consists of a central metal atom or ion surrounded by a group of ions or neutral
molecules e.g., [Co(NH3)6]3+. [Fe(CN)6]3-, [Cu(NH3)4]2+.

Cationic Complex:
A complex carrying a net positive charge is known as cationic complex.
Example: [Co(NH3)6]3+.

Anionic Complex:
If the complex has a net negative charge on the square bracket, it is called anionic
complex.
Example: [Fe(CN)6]3-

Neutral Complex :
Neutral complexes have not net charge.
For example: 1) [Ni(CO)4] 2) [Pt(NH3)2Cl2]
 4

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

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

(c) Ligands :
The ions or molecules bound to the central atom/ion in the coordination entity are
called ligands.
Ligands may be simple ions such as Cl–,
small molecules such as H2O or NH3,
larger molecules such as H2NCH2CH2NH2 or N(CH2CH2NH2)3 or
even macromolecules, such as proteins.
Types of Ligands
Unidentate:
When a ligand is bound to a metal ion through a single donor atom, as with Cl–, H2O
or NH3, the ligand is said to be unidentate.
Didentate :
When a ligand can bind through two donor atoms as in H2NCH2CH2NH2 (ethane-1,2-
diamine) or C2O42– (oxalate), the ligand is said to be didentate and
Polydentate :
When several donor atoms are present in a single ligand as in N(CH2CH2NH2)3, the
ligand is said to be polydentate.
Example: Ethylenediaminetetraacetate ion (EDTA4–) is an important hexadentate
ligand. It can bind through two nitrogen and four oxygen atoms to a central metal
ion. (totally 6 Places)

Chelate ligand :
When a di- or polydentate ligand uses its two or more donor atoms to bind a single
metal ion, it is said to be a chelate ligand.

Denticity of the ligand :

The number of such ligating groups is called the denticity of the ligand.
 5

Such complexes, called chelate complexes tend to be more stable than similar
complexes containing unidentate ligands.
Ambidentate ligand :
Ligand which can ligate through two different atoms is called ambidentate ligand.
Examples :

NO2– 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 (CN):

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, [PtCl6]2– and [Ni(NH3)4]2+, the coordination
number of Pt and Ni are 6 and 4 respectively.

Similarly,
in the complex ions, [Fe(C2O4)3]3– and [Co(en)3]3+, the coordination number of both,
Fe and Co, is 6 because C2O42– and en (ethane-1,2-diamine) are didentate ligands.

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 K4[Fe(CN)6], the coordination sphere is [Fe(CN)6]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.
 6

For example, [Co(NH3)6]3+ is octahedral, [Ni(CO)4] is tetrahedral and [PtCl4]2– is

square planar.
Shapes of different coordination polyhedra.

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

(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)4]3– is +1 and it is written as Cu(I).

(h) Homoleptic and heteroleptic complexes :

i)Homoleptic complexes
Complexes in which a metal is bound to only one kind of donor groups,
e.g., [Co(NH3)6]3+, are known as homoleptic.
ii)Heteroleptic complexes
Complexes in which a metal is bound to more than one kind of donor groups, e.g.,
[Co(NH3)4Cl2]+, are known as heteroleptic.

Nomenclature of Coordination Compounds :

The formulas and names adopted for coordination entities are based on the
recommendations of the International Union of Pure and Applied Chemistry
(IUPAC).
 7

Formulas of Mononuclear Coordination Entities : [Cu(CN)4]3–

(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.
[Cr(NH3)3(H2O)3]Cl3

(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.

[Co(H2NCH2CH2NH2)3]2(SO4)3
(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)6]3–, [Cr(H2O)6]3+, etc.
(vii) The charge of the cation(s) is balanced by the charge of the anion(s).
K4[Fe(CN)6], [Cr(NH3)3(H2O)3]Cl3

Naming of Mononuclear Coordination Compounds :

[Cr(NH3)3(H2O)3]Cl3 is named as:
triamminetriaquachromium(III) chloride
(Note: The Name of the complex compound should not start with a Capital letter)

(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 H2O, ammine for NH3, carbonyl for CO and nitrosyl
for NO. These are placed within enclosing marks ( ).
(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, [NiCl2(PPh3)2] is named as
 8

dichlorobis(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 [Co(SCN)4]2-is called cobaltate.
For some metals, the Latin names are used in the complex anions, e.g., ferrate
for Fe.
(vii) The neutral complex molecule is named similar to that of the complex cation.

[Cr(NH3)3(H2O)3]Cl3 is named as:

triamminetriaquachromium(III) chloride

[Co(H2NCH2CH2NH2)3]2(SO4)3 is named as:

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

[Ag(NH3)2][Ag(CN)2] is named as diamminesilver(I) dicyanoargentate(I)

Note:
1.The name of the complex should not start with a capital letter.
2.The full name of the complex ion (or) entity should be written as one word without any gap.
3.In the ionic complexes, there should be a gap between the name of the counter ion and of the
complex entity.
4.The full name of non-ionic complexes should be written as one word without any gap.
9
 10

Isomerism in Coordination Compounds

Isomerism in Coordination Compounds

(a) Stereoisomerism (b) Structural isomerism

(i) Geometrical isomerism (i) Linkage isomerism

(ii) Optical isomerism (ii) Coordination isomerism
(iii) Ionisation isomerism
(iv) Solvate isomerism

Isomers are two or more compounds that have the same chemical formula
but a different Structural formula.
They differ in one or more physical or chemical properties.
(a) Stereoisomerism :
Stereoisomers have the same chemical formula and chemical bonds but they have
different spatial arrangement.
(i) Geometrical isomerism
This type of isomerism arises in heteroleptic square planar complexes like [MX2L2]
Geometrical isomers (cis and trans) of Pt(NH3)2Cl2)

cis trans
In a square planar complex of formula [MX2L2] (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.

“Geometrical isomerism is not possible for a tetrahedral geometry”.

 11

Geometrical isomerism is possible in octahedral complexes of formula [MX2L4] in

which the two ligands X may be oriented cis or trans to each other.

Geometrical isomers (cis and trans) of [Co(NH3)4Cl2]+

This type of isomerism also arises when didentate ligands L – L [e.g., NH2 CH2 CH2 NH2 (en)]
are present in complexes of formula [MX2(L – L)2].
Geometrical isomers (cis and trans) of [CoCl2(en)2]

cis trans
Another type of geometrical isomerism occurs in octahedral coordination entities
of the type [Ma3b3] like [Co(NH3)3(NO2)3].
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.
The facial (fac) and meridional (mer) isomers of [Co(NH3
)3(NO2 )3]

fac mer
 12

(ii) Optical isomerism

Optical isomers are mirror images that cannot be superimposed on one another.
These are called as enantiomers.
The molecules or ions that cannot be superimposed are called chiral.

3+
Optical isomers (d and l) of [Co(en)3 ]

Optical isomers (d and l) of cis-[PtCl2(en)2]2+

The two forms are called dextro (d) and laevo (l) depending upon the direction
they rotate the plane of polarised light. (‘d’ rotates to the right, ‘l’ to the
left)

Optical isomerism is common in octahedral complexes involving didentate ligands.

In a coordination entity of the type [PtCl2(en)2]2+, only the cis-isomer shows optical
activity. (the trans-isomer will not show optical activity)
 13

(b) Structural isomerism :

Structural isomers have the same chemical formula and have different chemical
bonds . i.e., have different structures.
(i) Linkage isomerism
Linkage isomerism arises in a coordination compound containing
ambidentate ligand.

Example:
In the complex [Co(NH3)5(NO2)]Cl2, which is obtained in two forms i.e., Red form
and Yellow form.
1. In Red form, nitrite ligand is bound through oxygen (–ONO) and
2. In the yellow form, the nitrite ligand is bound through nitrogen (–NO2).

Thiocyanate ligand, NCS–, which may bind through the nitrogen to give M–NCS or
through sulphur to give M–SCN.

(ii) Coordination isomerism

This isomerism arises from the interchange of ligands between cationic and anionic
entities of different metal ions present in a complex.

Example:
[Co(NH3)6][Cr(CN)6], in which the NH3 ligands are bound to Co3+ and the CN–
ligands to Cr3+.

[Co(CN)6][Cr(NH3)6], in which the NH3 ligands are bound to Cr3+ and the CN–
ligands to Co3+.

(iii) Ionisation isomerism:

This 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.
Example:
Ionisation isomers [Co(NH3)5SO4]Br and [Co(NH3)5Br]SO4.
(iv) Solvate isomerism :
This is similar to ionisation isomerism.
Solvate isomers differ by whether or not a solvent molecule is directly bonded to
the metal ion or present as free solvent molecules in the crystal lattice.
This form of isomerism is known as ‘hydrate isomerism’ in case where water is
involved as a solvent.
Example:
[Cr(H2O)6]Cl3 (violet) and its solvate isomer [Cr(H2O)5Cl]Cl2.H2O (grey-green).
 14

Bonding in Coordination Compounds :

Application of VBT and CFT to coordination compounds.

(Molecular Orbital Theory (MOT)-out of syllabus).

Valence Bond Theory:

According to this theory, the metal atom or ion under the influence of ligands
undergoes hydbridisation using (n-1)d, ns, np or ns, np, nd orbitals (i.e., orbitals of
nearly equal energies) to yield a set of equivalent orbitals of definite geometry.
e.g., octahedral, tetrahedral, square planar and so on.
These hybridised orbitals overlap with ligand orbitals.
Number of Orbitals and Types of Hybridisations

Magnetic Properties of Coordination Compounds :

The magnetic moment of coordination compounds can be measured
by the magnetic susceptibility experiments.
The results can be used to obtain information about the structures adopted by
metal complexes.

No. of No. of unpaired No. of vacant ‘d’

Ion electrons in electrons orbital available
‘d’ orbital for bonding
3+ 1
Ti d 1 4
3+ 2
V d 2 3
3+ 3
Cr d 3 2
2+ 3+ 4
Cr , Mn d 4 1
Mn2+, Fe3+ d5 5 0
2+ 3+ 6
Fe , Co d 4 0
For metal ions with upto three electrons in the ‘d’ orbitals, 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 per Hund’s rule). 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 d6 ions.
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However, with species containing d4 and d5 ions there are complications.

Mn3+ with d4 configuration
[Mn(CN)6]3– has magnetic moment of two unpaired electrons while [MnCl6]3– has a
paramagnetic moment of four unpaired electrons. (Both are having Paramaagnetic
moment only)
Fe3+ with d5configuration
[Fe(CN)6]3– has magnetic moment of a single unpaired electron while [FeF6]3– has a
paramagnetic moment of five unpaired electrons. (Both are having Paramaagnetic
moment only)
Co3+ with d6configuration
[CoF6]3– is paramagnetic with four unpaired electrons while [Co(C2O4)3]3– is
diamagnetic i.e., no unpaired electrons.

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

orbital and outer orbital coordination entities.

[Mn(CN)6]3–, [Fe(CN)6]3–and [Co(C2O4)3]3– are inner orbital complexes involving d2sp3

hybridisation.

The former two complexes are paramagnetic and the latter diamagnetic. On the
other hand, [MnCl6]3–, [FeF6]3– and [CoF6]3– are outer orbital complexes involving
sp3d2hybridisation and are paramagnetic corresponding to four, five and four
unpaired electrons.

The spin only magnetic moment of [MnBr4]2– is 5.9 BM. Predict the
geometry of the complex ion ?

No. of No. of unpaired No. of vacant ‘d’

Ion electrons in electrons orbital available
‘d’ orbital for bonding
Mn2+ d5 5 0

Since the coordination number of Mn2+ ion in the complex ion is 4,

it will be either tetrahedral (sp3 hybridisation) or square planar (dsp2

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.
 16

Limitations of Valence Bond Theory :

(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.

Crystal Field Theory :

The valence bond theory regards the metal at atom/ion and ligand as
purely covalent. However,
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.
In case the ligand are anionic the interactions are ion-ion. ( F-, CN- etc.,)
If the ligands are neutral then the interactions are ion-dipolar nature. (NH3, H2O, etc.,)

We know that in a free metal atom or ion, all the five ‘d’ orbitals have the same
energy i.e., these are degenerate orbitals.

*When ligands approach the metal/ion, repulsion will arise between the
electrons present in the ‘d’ orbitals and the lone pairs on the ligands. This will
raise the energy of the ‘d’ orbitals. If all ligands are equidistant from the ‘d’
orbitals involved, the increase in the energy of these orbitals will be still the same
 17

i.e., they will remain degenerate. But we know that ‘d’ orbitals differ in their
orientations and are therefore likely to experience different repulsive interactions
from the ligands.

The orbitals that are located in the direction

of the approaching ligands will be raised to higher
energy level as compared to the orbitals which lie
away from them. This means that the ‘d’ orbitals will
be no longer degenerate or have the same enengy
and will split up into different sets on the basis of
energy. This splitting of the degenerate ‘d’ orbitals
under the influence of approaching ligands into
different sets is called ‘Crystal field splitting’ and it
plays an important role in explaining the geometry
and characteristics of co-ordinate complexes.

Splitting of ‘d’ orbitals in Octahedral Complexes:

 18

Splitting of ‘d’ orbitals in tetrahedral Complexes:

Spectrochemical series :
For a particular metal ion, the ligands can be arranged in a series in the order
of increasing field strength (ΔO values). This series is termed as spectrochemical
series.
It is an experimentally determined series based on the absorption of light
by complexes with different ligands.

When we assign electrons in the d orbitals of metal ion in octahedral coordination

entities. Obviously, the single d electron occupies one of the lower energy t2g
orbitals.
In d2 and d3 coordination entities, the d electrons occupy the t2g orbitals
singly in accordance with the Hund’s rule. For d4 ions, two possible patterns of
electron distribution arise:
(i) the fourth electron could either enter the t2g level and pair with an existing electron
(or)
(ii) it could avoid paying the price of the pairing energy by occupying the eg level.

The two options are:

(i) If Δo < P, the fourth electron enters one of the eg orbitals giving the
configuration (t2g)3 (eg)1 .
Ligands for which Δo < P are known as weak field ligands and form high spin complexes.
(ii) If Δo > P, it becomes more energetically favourable for the fourth electron to
occupy a t2g orbital with configuration (t2g)4 (eg)0.
Ligands which Δo > P, are known as strong field ligands and form low spin complexes.
 19

Comparison of Δt < ΔO for the same metal, the same ligands :

For the same metal, the same ligands

and metal-ligand distances

Δt < ΔO
The energy difference is so small that is
not sufficient to force the pairing of
electrons. Hence, tetrahedral
complexes have high spin configuration
ie., pairing is not takes place.

9Δt = 4ΔO

4
Δt = ΔO
9

Colour in Coordination Compounds:

The colour in the coordination compounds can be readily explained in terms of
the crystal field theory.
The crystal field theory attributes the colour of the coordination compounds to
‘d-d’ transition of the electron.
Consider, for example, the complex [Ti(H2O)6]3+, which is violet in colour.

(t2g)1(eg)0 -----------------------------(t2g)0(eg)1

It is important to note that in the absence of ligand, crystal field splitting does
not occur and hence the substance is colourless.
 20

For example,
1) Removal of water from [Ti(H2O)6]Cl3 on heating renders it colourless.

2) CuSO4.5H2O is blue in colour but anhydrous CuSO4 is white in colour.

Limitations of Crystal Field Theory :

Crystal Field Theory has been assumed that the ligands are point charges,
this means that the anionic ligands should cause large splitting effect. But these
are mostly present in the beginning of the spectrochemical series. Similarly, OH-
ion is placed before H2O and NH3.

This limitations can be explained with the help of Ligand Field Theory and
molecular orbital theory.
. ************************
 21

Influence of the ligand on the colour of a complex :

The influence of the ligand on the colour of a complex may be illustrated by

considering the [Ni(H2O)6]2+ 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 occurs,
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Prepared by

Shri. R.Sekar, M.Sc, M.Phil., M.Ed.,

PGT Chemistry,
Kendriya Vidyalaya No.2,
KALPAKKAM-603 102