Coordination Compounds : Daily Practice Problems (DPP) - 2 Class 12 Notes | EduRev

Class 12 : Coordination Compounds : Daily Practice Problems (DPP) - 2 Class 12 Notes | EduRev

 Page 1


1
SCF-07, 1
st
 floor, Sector -15, Panchkula
   9217610408,8699438881,8699438882
     DPP : 02 / Co-ordination compounds
Some important points & facts
1. VALENCE BOND THEORY
It is the simplest of the three theories and was developed mainly by Pauling. It describes the bonding in terms of
hybridized orbitals of the central metal atom or ion. The theory mainly deals with the geometry (i.e., shape) and
magnetic properties of the complexes. The salient features of the theory are
I. The central metal loses a requisite number of electrons to form the ion. The number of electrons lost is the
valency of the resulting cation. In some cases, the metal atom does not lose electrons.
II. The central metal ion or atom (as the case may be) makes available a number of empty s-, p- and d-atomic
orbitals equal to its coordination number. These vacant orbitals hybridize together to form hybrid orbitals
which are same in the number as the atomic orbitals hybridizing together. They are vacant, equivalent in
energy and have definite geometry.
Some of the common hybridized orbitals met in the coordination compounds are listed below
  Coordination        Type of Geometry Examples
      number   hybridization
2 sp Linear [Ag(NH
3
) 2
]
+
, [Ag(CN) 2
]
—
3 sp
2
Trigonal planar [HgI
3
]
—
4 sp
3
Tetrahedral Ni(CO) 4
, [Ni(X) 4
]
2—
, [ZnCl
4
]
2—
, [CuX
4
]
2—
where X = Cl
—
, Br
—
, I
—
dsp
2
Square planar [Ni(CN
4
)]
2—
, [Cu(NH
3
) 4
]
2—
, [Ni(NH
3
) 4
]
2+
5 dsp
3
Trigonal bipyramidal Fe(CO) 5
, [CuCl
5
]
3—
sp
3
d Square pyramidal [SbF
5
]
2—
6 d
2
sp
3
Octahedral [Cr(NH
3
) 6
]
3+
,
  or(Inner orbital) [Fe(CN) 6
]
3—
sp
3
d
2
(Outer orbital) [FeF
6
]
3—
, [Fe(H
2
O) 6
]
2+
, [Ni(NH
3
) 6
]
2+
III. The non-bonding electrons of the metal occupy the inner orbitals. These are grouped in accordance with
Hund’s rule, however, under the influence of some strong ligands, there may be some re-arrangement of
electrons in the atomic orbitals (against Hund’s rule). The d-orbitals participating in this process of
hybridization may be either (n—1)d
2
sp
3
 or nsp
3
d
2
. The complexes thus formed are referred to as inner or
low spin and outer or high spin complexes, respectively.
COORDINATION COMPOUNDS
DAILY  PRACTICE  PROBLEMS - 2
Page 2


1
SCF-07, 1
st
 floor, Sector -15, Panchkula
   9217610408,8699438881,8699438882
     DPP : 02 / Co-ordination compounds
Some important points & facts
1. VALENCE BOND THEORY
It is the simplest of the three theories and was developed mainly by Pauling. It describes the bonding in terms of
hybridized orbitals of the central metal atom or ion. The theory mainly deals with the geometry (i.e., shape) and
magnetic properties of the complexes. The salient features of the theory are
I. The central metal loses a requisite number of electrons to form the ion. The number of electrons lost is the
valency of the resulting cation. In some cases, the metal atom does not lose electrons.
II. The central metal ion or atom (as the case may be) makes available a number of empty s-, p- and d-atomic
orbitals equal to its coordination number. These vacant orbitals hybridize together to form hybrid orbitals
which are same in the number as the atomic orbitals hybridizing together. They are vacant, equivalent in
energy and have definite geometry.
Some of the common hybridized orbitals met in the coordination compounds are listed below
  Coordination        Type of Geometry Examples
      number   hybridization
2 sp Linear [Ag(NH
3
) 2
]
+
, [Ag(CN) 2
]
—
3 sp
2
Trigonal planar [HgI
3
]
—
4 sp
3
Tetrahedral Ni(CO) 4
, [Ni(X) 4
]
2—
, [ZnCl
4
]
2—
, [CuX
4
]
2—
where X = Cl
—
, Br
—
, I
—
dsp
2
Square planar [Ni(CN
4
)]
2—
, [Cu(NH
3
) 4
]
2—
, [Ni(NH
3
) 4
]
2+
5 dsp
3
Trigonal bipyramidal Fe(CO) 5
, [CuCl
5
]
3—
sp
3
d Square pyramidal [SbF
5
]
2—
6 d
2
sp
3
Octahedral [Cr(NH
3
) 6
]
3+
,
  or(Inner orbital) [Fe(CN) 6
]
3—
sp
3
d
2
(Outer orbital) [FeF
6
]
3—
, [Fe(H
2
O) 6
]
2+
, [Ni(NH
3
) 6
]
2+
III. The non-bonding electrons of the metal occupy the inner orbitals. These are grouped in accordance with
Hund’s rule, however, under the influence of some strong ligands, there may be some re-arrangement of
electrons in the atomic orbitals (against Hund’s rule). The d-orbitals participating in this process of
hybridization may be either (n—1)d
2
sp
3
 or nsp
3
d
2
. The complexes thus formed are referred to as inner or
low spin and outer or high spin complexes, respectively.
COORDINATION COMPOUNDS
DAILY  PRACTICE  PROBLEMS - 2
2
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IV . The ligands have at least one s-orbital containing a lone pair of electrons. Vacant hybrid orbitals of the
metal atom or ion overlap with the s-orbitals containing lone pair or electrons of the ligands to form
M ? ligand s-bond. This bond is called coordinate bond (a special type of covalent bond) and possesses
a considerable amount of polarity.
V . It is possible to predict the magnetic properties of the complex if the geometry of the complex ion (or
vice-versa) is known. If the complex contains unpaired electrons, it is paramagnetic in nature whereas if it
does not contain unpaired electrons, i.e., all are paired, the complex if diamagnetic in nature. Complexes
having unpaired electron are coloured.
Geometry (shape) nd magnetic nature of some of the complexes
(Application of valence bond theory) Atom/ion/ Configuration Oxidation Type of       Geometry         No. of Magnetic
complex state of hybridiz-        shape unpaired  nature
      (I)       (II)  metal -ation            (V) electrons     (VII)    (III)   (IV)      (VI) Ni
2+
 (d
8
) 3d 4s 4p
    +2       2            Paramagnetic
[ N i C l
4
]
2—
sp
3
    +2   sp
3
     Tetrahedral       2            Paramagnetic
[Ni(CN) 4
]
2—
dsp
2
Rearrangement
    +2 dsp
2
     Square planar       0            Diamagnetic
Ni     0       2            Paramagnetic
Ni(CO) 4
sp
3
Rearrangement
     0 sp
3
     Tetrahedral       0            Diamagnetic
[Ni(NH
3
) 6
]
2+
3d
4s
4p
sp
3
d
2
4d
    +2            sp
3
d
2
     Octahedral       2            Paramagnetic
(Outer) Mn
2+
(d
5
) 3d 4s 4p
    +2       5            Paramagnetic
[Mn(CN) 6
]
4—
d
2
sp
3
Rearrangement
+2 d
2
sp
3
     Octahedral       1            Paramagnetic
(Inner) [MnCl
4
]
2—
sp
3
    +2 sp
3
     Tetrahedral       5            Paramagnetic
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SCF-07, 1
st
 floor, Sector -15, Panchkula
   9217610408,8699438881,8699438882
     DPP : 02 / Co-ordination compounds
Some important points & facts
1. VALENCE BOND THEORY
It is the simplest of the three theories and was developed mainly by Pauling. It describes the bonding in terms of
hybridized orbitals of the central metal atom or ion. The theory mainly deals with the geometry (i.e., shape) and
magnetic properties of the complexes. The salient features of the theory are
I. The central metal loses a requisite number of electrons to form the ion. The number of electrons lost is the
valency of the resulting cation. In some cases, the metal atom does not lose electrons.
II. The central metal ion or atom (as the case may be) makes available a number of empty s-, p- and d-atomic
orbitals equal to its coordination number. These vacant orbitals hybridize together to form hybrid orbitals
which are same in the number as the atomic orbitals hybridizing together. They are vacant, equivalent in
energy and have definite geometry.
Some of the common hybridized orbitals met in the coordination compounds are listed below
  Coordination        Type of Geometry Examples
      number   hybridization
2 sp Linear [Ag(NH
3
) 2
]
+
, [Ag(CN) 2
]
—
3 sp
2
Trigonal planar [HgI
3
]
—
4 sp
3
Tetrahedral Ni(CO) 4
, [Ni(X) 4
]
2—
, [ZnCl
4
]
2—
, [CuX
4
]
2—
where X = Cl
—
, Br
—
, I
—
dsp
2
Square planar [Ni(CN
4
)]
2—
, [Cu(NH
3
) 4
]
2—
, [Ni(NH
3
) 4
]
2+
5 dsp
3
Trigonal bipyramidal Fe(CO) 5
, [CuCl
5
]
3—
sp
3
d Square pyramidal [SbF
5
]
2—
6 d
2
sp
3
Octahedral [Cr(NH
3
) 6
]
3+
,
  or(Inner orbital) [Fe(CN) 6
]
3—
sp
3
d
2
(Outer orbital) [FeF
6
]
3—
, [Fe(H
2
O) 6
]
2+
, [Ni(NH
3
) 6
]
2+
III. The non-bonding electrons of the metal occupy the inner orbitals. These are grouped in accordance with
Hund’s rule, however, under the influence of some strong ligands, there may be some re-arrangement of
electrons in the atomic orbitals (against Hund’s rule). The d-orbitals participating in this process of
hybridization may be either (n—1)d
2
sp
3
 or nsp
3
d
2
. The complexes thus formed are referred to as inner or
low spin and outer or high spin complexes, respectively.
COORDINATION COMPOUNDS
DAILY  PRACTICE  PROBLEMS - 2
2
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     DPP : 02 / Co-ordination compounds
IV . The ligands have at least one s-orbital containing a lone pair of electrons. Vacant hybrid orbitals of the
metal atom or ion overlap with the s-orbitals containing lone pair or electrons of the ligands to form
M ? ligand s-bond. This bond is called coordinate bond (a special type of covalent bond) and possesses
a considerable amount of polarity.
V . It is possible to predict the magnetic properties of the complex if the geometry of the complex ion (or
vice-versa) is known. If the complex contains unpaired electrons, it is paramagnetic in nature whereas if it
does not contain unpaired electrons, i.e., all are paired, the complex if diamagnetic in nature. Complexes
having unpaired electron are coloured.
Geometry (shape) nd magnetic nature of some of the complexes
(Application of valence bond theory) Atom/ion/ Configuration Oxidation Type of       Geometry         No. of Magnetic
complex state of hybridiz-        shape unpaired  nature
      (I)       (II)  metal -ation            (V) electrons     (VII)    (III)   (IV)      (VI) Ni
2+
 (d
8
) 3d 4s 4p
    +2       2            Paramagnetic
[ N i C l
4
]
2—
sp
3
    +2   sp
3
     Tetrahedral       2            Paramagnetic
[Ni(CN) 4
]
2—
dsp
2
Rearrangement
    +2 dsp
2
     Square planar       0            Diamagnetic
Ni     0       2            Paramagnetic
Ni(CO) 4
sp
3
Rearrangement
     0 sp
3
     Tetrahedral       0            Diamagnetic
[Ni(NH
3
) 6
]
2+
3d
4s
4p
sp
3
d
2
4d
    +2            sp
3
d
2
     Octahedral       2            Paramagnetic
(Outer) Mn
2+
(d
5
) 3d 4s 4p
    +2       5            Paramagnetic
[Mn(CN) 6
]
4—
d
2
sp
3
Rearrangement
+2 d
2
sp
3
     Octahedral       1            Paramagnetic
(Inner) [MnCl
4
]
2—
sp
3
    +2 sp
3
     Tetrahedral       5            Paramagnetic
3
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Cu
2+
(d
9
)    +2       1            Paramagnetic
[CuCl
4
]
2—
sp
3
    +2 sp
3
    Tetrahedral       1            Paramagnetic
[Cu(NH
3
) 4
]
2+
dsp
2
    +2 dsp
2
    Square planar       1            Paramagnetic
One electron is shifted from
3d- to 4p- orbital
Cr
3+
(d
3
) 3d 4s 4p
    +3       3            Paramagnetic
[Cr(NH
3
) 6
]
3+
d
2
sp
3
    +3 d
2
sp
3
     Octahedral       3            Paramagnetic
2. CRYSTAL FIELD THEORY
Crystal field theory was proposed by H.Bethe (1929) and Van Vleck (1932) and was originally applied to ionic
crystals to explain their optical properties and is, therefore, called crystal field theory. However, this theory was
applied to the  study of coordination compounds in 1950. The valence bond theory represents the ligand metal
bond as covalent, with an electron pair, shared between the metal and the ligand donor atom. The valence bond
theory is frequently used but it is not adequate to explain properties of complexes such as colour and magnetism.
I. According to the crystal field theory, the bonding in complex ions is purely electrostatic. This theory regards
the ligand atoms of ionic ligands such as F
—
, Cl
—
 or CN
—
 as negative point charges (also called charges) and if the ligand molecules are neutral these are regarded as point dipoles or simply dipoles, the negative
end pointing towards central metal ion.
+
d
- d +
d
+
d
- d
+
d
+
d
N O
H
(Ammonia molecule)(Water molecule) H
H
;
H
H
II. The complex is regarded as a combination of a central metal ion surrounded by ligands which act as point
dipoles. The arrangement of ligands around the central netal ion or atom is such that the repulsion
between these negative points or dipoles is minimum.
III. Interactions between positively charged nucleus of the central metal or atom and the negatively charged
ligands are of two types
(a) The attractive forces arise due to the positive metal ion and the negatively charged ligands or the
negative end of a polar neutral molecules. For example, in the case of complex ion, [Fe(CN) 6
]
3—
, the
interactions are between Fe
3+
 ion and negatively charged CN
—
 ions whereas in the complex;
[Cr(NH
3
) 6
]
3+
, the interactions are between Cr
3+
 ion and negatively charged ends of ammonia
molecules. These attractive forces bind the ligands to the metal ion. The bonds between metal and
the surrounding ligands are purely ionic. This theory does not consider any orbital overlapping.
Page 4


1
SCF-07, 1
st
 floor, Sector -15, Panchkula
   9217610408,8699438881,8699438882
     DPP : 02 / Co-ordination compounds
Some important points & facts
1. VALENCE BOND THEORY
It is the simplest of the three theories and was developed mainly by Pauling. It describes the bonding in terms of
hybridized orbitals of the central metal atom or ion. The theory mainly deals with the geometry (i.e., shape) and
magnetic properties of the complexes. The salient features of the theory are
I. The central metal loses a requisite number of electrons to form the ion. The number of electrons lost is the
valency of the resulting cation. In some cases, the metal atom does not lose electrons.
II. The central metal ion or atom (as the case may be) makes available a number of empty s-, p- and d-atomic
orbitals equal to its coordination number. These vacant orbitals hybridize together to form hybrid orbitals
which are same in the number as the atomic orbitals hybridizing together. They are vacant, equivalent in
energy and have definite geometry.
Some of the common hybridized orbitals met in the coordination compounds are listed below
  Coordination        Type of Geometry Examples
      number   hybridization
2 sp Linear [Ag(NH
3
) 2
]
+
, [Ag(CN) 2
]
—
3 sp
2
Trigonal planar [HgI
3
]
—
4 sp
3
Tetrahedral Ni(CO) 4
, [Ni(X) 4
]
2—
, [ZnCl
4
]
2—
, [CuX
4
]
2—
where X = Cl
—
, Br
—
, I
—
dsp
2
Square planar [Ni(CN
4
)]
2—
, [Cu(NH
3
) 4
]
2—
, [Ni(NH
3
) 4
]
2+
5 dsp
3
Trigonal bipyramidal Fe(CO) 5
, [CuCl
5
]
3—
sp
3
d Square pyramidal [SbF
5
]
2—
6 d
2
sp
3
Octahedral [Cr(NH
3
) 6
]
3+
,
  or(Inner orbital) [Fe(CN) 6
]
3—
sp
3
d
2
(Outer orbital) [FeF
6
]
3—
, [Fe(H
2
O) 6
]
2+
, [Ni(NH
3
) 6
]
2+
III. The non-bonding electrons of the metal occupy the inner orbitals. These are grouped in accordance with
Hund’s rule, however, under the influence of some strong ligands, there may be some re-arrangement of
electrons in the atomic orbitals (against Hund’s rule). The d-orbitals participating in this process of
hybridization may be either (n—1)d
2
sp
3
 or nsp
3
d
2
. The complexes thus formed are referred to as inner or
low spin and outer or high spin complexes, respectively.
COORDINATION COMPOUNDS
DAILY  PRACTICE  PROBLEMS - 2
2
SCF-07, 1
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 floor, Sector -15, Panchkula
   9217610408,8699438881,8699438882
     DPP : 02 / Co-ordination compounds
IV . The ligands have at least one s-orbital containing a lone pair of electrons. Vacant hybrid orbitals of the
metal atom or ion overlap with the s-orbitals containing lone pair or electrons of the ligands to form
M ? ligand s-bond. This bond is called coordinate bond (a special type of covalent bond) and possesses
a considerable amount of polarity.
V . It is possible to predict the magnetic properties of the complex if the geometry of the complex ion (or
vice-versa) is known. If the complex contains unpaired electrons, it is paramagnetic in nature whereas if it
does not contain unpaired electrons, i.e., all are paired, the complex if diamagnetic in nature. Complexes
having unpaired electron are coloured.
Geometry (shape) nd magnetic nature of some of the complexes
(Application of valence bond theory) Atom/ion/ Configuration Oxidation Type of       Geometry         No. of Magnetic
complex state of hybridiz-        shape unpaired  nature
      (I)       (II)  metal -ation            (V) electrons     (VII)    (III)   (IV)      (VI) Ni
2+
 (d
8
) 3d 4s 4p
    +2       2            Paramagnetic
[ N i C l
4
]
2—
sp
3
    +2   sp
3
     Tetrahedral       2            Paramagnetic
[Ni(CN) 4
]
2—
dsp
2
Rearrangement
    +2 dsp
2
     Square planar       0            Diamagnetic
Ni     0       2            Paramagnetic
Ni(CO) 4
sp
3
Rearrangement
     0 sp
3
     Tetrahedral       0            Diamagnetic
[Ni(NH
3
) 6
]
2+
3d
4s
4p
sp
3
d
2
4d
    +2            sp
3
d
2
     Octahedral       2            Paramagnetic
(Outer) Mn
2+
(d
5
) 3d 4s 4p
    +2       5            Paramagnetic
[Mn(CN) 6
]
4—
d
2
sp
3
Rearrangement
+2 d
2
sp
3
     Octahedral       1            Paramagnetic
(Inner) [MnCl
4
]
2—
sp
3
    +2 sp
3
     Tetrahedral       5            Paramagnetic
3
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Cu
2+
(d
9
)    +2       1            Paramagnetic
[CuCl
4
]
2—
sp
3
    +2 sp
3
    Tetrahedral       1            Paramagnetic
[Cu(NH
3
) 4
]
2+
dsp
2
    +2 dsp
2
    Square planar       1            Paramagnetic
One electron is shifted from
3d- to 4p- orbital
Cr
3+
(d
3
) 3d 4s 4p
    +3       3            Paramagnetic
[Cr(NH
3
) 6
]
3+
d
2
sp
3
    +3 d
2
sp
3
     Octahedral       3            Paramagnetic
2. CRYSTAL FIELD THEORY
Crystal field theory was proposed by H.Bethe (1929) and Van Vleck (1932) and was originally applied to ionic
crystals to explain their optical properties and is, therefore, called crystal field theory. However, this theory was
applied to the  study of coordination compounds in 1950. The valence bond theory represents the ligand metal
bond as covalent, with an electron pair, shared between the metal and the ligand donor atom. The valence bond
theory is frequently used but it is not adequate to explain properties of complexes such as colour and magnetism.
I. According to the crystal field theory, the bonding in complex ions is purely electrostatic. This theory regards
the ligand atoms of ionic ligands such as F
—
, Cl
—
 or CN
—
 as negative point charges (also called charges) and if the ligand molecules are neutral these are regarded as point dipoles or simply dipoles, the negative
end pointing towards central metal ion.
+
d
- d +
d
+
d
- d
+
d
+
d
N O
H
(Ammonia molecule)(Water molecule) H
H
;
H
H
II. The complex is regarded as a combination of a central metal ion surrounded by ligands which act as point
dipoles. The arrangement of ligands around the central netal ion or atom is such that the repulsion
between these negative points or dipoles is minimum.
III. Interactions between positively charged nucleus of the central metal or atom and the negatively charged
ligands are of two types
(a) The attractive forces arise due to the positive metal ion and the negatively charged ligands or the
negative end of a polar neutral molecules. For example, in the case of complex ion, [Fe(CN) 6
]
3—
, the
interactions are between Fe
3+
 ion and negatively charged CN
—
 ions whereas in the complex;
[Cr(NH
3
) 6
]
3+
, the interactions are between Cr
3+
 ion and negatively charged ends of ammonia
molecules. These attractive forces bind the ligands to the metal ion. The bonds between metal and
the surrounding ligands are purely ionic. This theory does not consider any orbital overlapping.
4
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(b) The repulsive forces arise between the lone pairs on the ligands and electrons in the d-orbitals of
the metal or atom. The crystal field theory mainly focuses on these repulsive forces. These forces
are responsible for causing a considerable effect on the relative energies of the d-orbitals of the
central metal ion or atom.
IV . In a free transition metal or ion, there are five d-orbitals which are designated as d
xy
, d
yz
, d
zx
, d
x
2
—y
2 and d
x
2.
The five d-orbitals are divided into two sets depending on the nature of their orientation in space.
(a) The three d-orbitals (d
xy
, d
yz
, d
zx
) which orient in the regions between the coordinate axes are
designated as t
2g
–orbitals (pronounced as “t-two-g”). t
2g
–orbitals are three-fold degenerate. These
are non-axial orbitals.
(b) The other two orbitals (d
x
2
—y
2), d
z
2 which orient along the axes are called e
g
, (e
g
 orbitals are two fold
degenerate). These are also called axial-orbitals.
[The names t
2g
 and e
g
 are derived from spectroscopic terms.]
In a free transition metal ion or atom, all the five d-orbitals same energy, i.e., they are degenerate. However, when
the ligands approach the central metal or atom are repelled by lone pairs of the ligands. As a result of these
interactions, the degeneracy of d-orbitals of the metal ion is lost depending on the orientation of ligands in space.
The d-orbitals split into two sets of orbitals having different energies. This is called crystals field splitting. It is the
basis of crystal field theory. The extent of splitting depends on the number of ligands and their position around the
metal atom or ion. The splitting is different in different structures with different coordination numbers.
Splitting of d-orbitals in octahedral complexes.
0
?
e
g
t
2g
Average energy
(t
2g
+ e
g
) of orbitals in
spherical crystal field
e
g
t
2g
(t
2g
+ e
g
) orbitals in
free ion
(a)(b)(c) 10 Dq
Energy
No splitting state
Crystal field splitting in tetrahedral complexes.
t
?
Average energy
(t
2g
+ e
g
) of orbitals in
spherical crystal field
e
g
t
2g
(t
2g
+ e
g
) orbitals in
free ion
10 Dq
Energy
4 Dq
6 Dq
Page 5


1
SCF-07, 1
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   9217610408,8699438881,8699438882
     DPP : 02 / Co-ordination compounds
Some important points & facts
1. VALENCE BOND THEORY
It is the simplest of the three theories and was developed mainly by Pauling. It describes the bonding in terms of
hybridized orbitals of the central metal atom or ion. The theory mainly deals with the geometry (i.e., shape) and
magnetic properties of the complexes. The salient features of the theory are
I. The central metal loses a requisite number of electrons to form the ion. The number of electrons lost is the
valency of the resulting cation. In some cases, the metal atom does not lose electrons.
II. The central metal ion or atom (as the case may be) makes available a number of empty s-, p- and d-atomic
orbitals equal to its coordination number. These vacant orbitals hybridize together to form hybrid orbitals
which are same in the number as the atomic orbitals hybridizing together. They are vacant, equivalent in
energy and have definite geometry.
Some of the common hybridized orbitals met in the coordination compounds are listed below
  Coordination        Type of Geometry Examples
      number   hybridization
2 sp Linear [Ag(NH
3
) 2
]
+
, [Ag(CN) 2
]
—
3 sp
2
Trigonal planar [HgI
3
]
—
4 sp
3
Tetrahedral Ni(CO) 4
, [Ni(X) 4
]
2—
, [ZnCl
4
]
2—
, [CuX
4
]
2—
where X = Cl
—
, Br
—
, I
—
dsp
2
Square planar [Ni(CN
4
)]
2—
, [Cu(NH
3
) 4
]
2—
, [Ni(NH
3
) 4
]
2+
5 dsp
3
Trigonal bipyramidal Fe(CO) 5
, [CuCl
5
]
3—
sp
3
d Square pyramidal [SbF
5
]
2—
6 d
2
sp
3
Octahedral [Cr(NH
3
) 6
]
3+
,
  or(Inner orbital) [Fe(CN) 6
]
3—
sp
3
d
2
(Outer orbital) [FeF
6
]
3—
, [Fe(H
2
O) 6
]
2+
, [Ni(NH
3
) 6
]
2+
III. The non-bonding electrons of the metal occupy the inner orbitals. These are grouped in accordance with
Hund’s rule, however, under the influence of some strong ligands, there may be some re-arrangement of
electrons in the atomic orbitals (against Hund’s rule). The d-orbitals participating in this process of
hybridization may be either (n—1)d
2
sp
3
 or nsp
3
d
2
. The complexes thus formed are referred to as inner or
low spin and outer or high spin complexes, respectively.
COORDINATION COMPOUNDS
DAILY  PRACTICE  PROBLEMS - 2
2
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IV . The ligands have at least one s-orbital containing a lone pair of electrons. Vacant hybrid orbitals of the
metal atom or ion overlap with the s-orbitals containing lone pair or electrons of the ligands to form
M ? ligand s-bond. This bond is called coordinate bond (a special type of covalent bond) and possesses
a considerable amount of polarity.
V . It is possible to predict the magnetic properties of the complex if the geometry of the complex ion (or
vice-versa) is known. If the complex contains unpaired electrons, it is paramagnetic in nature whereas if it
does not contain unpaired electrons, i.e., all are paired, the complex if diamagnetic in nature. Complexes
having unpaired electron are coloured.
Geometry (shape) nd magnetic nature of some of the complexes
(Application of valence bond theory) Atom/ion/ Configuration Oxidation Type of       Geometry         No. of Magnetic
complex state of hybridiz-        shape unpaired  nature
      (I)       (II)  metal -ation            (V) electrons     (VII)    (III)   (IV)      (VI) Ni
2+
 (d
8
) 3d 4s 4p
    +2       2            Paramagnetic
[ N i C l
4
]
2—
sp
3
    +2   sp
3
     Tetrahedral       2            Paramagnetic
[Ni(CN) 4
]
2—
dsp
2
Rearrangement
    +2 dsp
2
     Square planar       0            Diamagnetic
Ni     0       2            Paramagnetic
Ni(CO) 4
sp
3
Rearrangement
     0 sp
3
     Tetrahedral       0            Diamagnetic
[Ni(NH
3
) 6
]
2+
3d
4s
4p
sp
3
d
2
4d
    +2            sp
3
d
2
     Octahedral       2            Paramagnetic
(Outer) Mn
2+
(d
5
) 3d 4s 4p
    +2       5            Paramagnetic
[Mn(CN) 6
]
4—
d
2
sp
3
Rearrangement
+2 d
2
sp
3
     Octahedral       1            Paramagnetic
(Inner) [MnCl
4
]
2—
sp
3
    +2 sp
3
     Tetrahedral       5            Paramagnetic
3
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Cu
2+
(d
9
)    +2       1            Paramagnetic
[CuCl
4
]
2—
sp
3
    +2 sp
3
    Tetrahedral       1            Paramagnetic
[Cu(NH
3
) 4
]
2+
dsp
2
    +2 dsp
2
    Square planar       1            Paramagnetic
One electron is shifted from
3d- to 4p- orbital
Cr
3+
(d
3
) 3d 4s 4p
    +3       3            Paramagnetic
[Cr(NH
3
) 6
]
3+
d
2
sp
3
    +3 d
2
sp
3
     Octahedral       3            Paramagnetic
2. CRYSTAL FIELD THEORY
Crystal field theory was proposed by H.Bethe (1929) and Van Vleck (1932) and was originally applied to ionic
crystals to explain their optical properties and is, therefore, called crystal field theory. However, this theory was
applied to the  study of coordination compounds in 1950. The valence bond theory represents the ligand metal
bond as covalent, with an electron pair, shared between the metal and the ligand donor atom. The valence bond
theory is frequently used but it is not adequate to explain properties of complexes such as colour and magnetism.
I. According to the crystal field theory, the bonding in complex ions is purely electrostatic. This theory regards
the ligand atoms of ionic ligands such as F
—
, Cl
—
 or CN
—
 as negative point charges (also called charges) and if the ligand molecules are neutral these are regarded as point dipoles or simply dipoles, the negative
end pointing towards central metal ion.
+
d
- d +
d
+
d
- d
+
d
+
d
N O
H
(Ammonia molecule)(Water molecule) H
H
;
H
H
II. The complex is regarded as a combination of a central metal ion surrounded by ligands which act as point
dipoles. The arrangement of ligands around the central netal ion or atom is such that the repulsion
between these negative points or dipoles is minimum.
III. Interactions between positively charged nucleus of the central metal or atom and the negatively charged
ligands are of two types
(a) The attractive forces arise due to the positive metal ion and the negatively charged ligands or the
negative end of a polar neutral molecules. For example, in the case of complex ion, [Fe(CN) 6
]
3—
, the
interactions are between Fe
3+
 ion and negatively charged CN
—
 ions whereas in the complex;
[Cr(NH
3
) 6
]
3+
, the interactions are between Cr
3+
 ion and negatively charged ends of ammonia
molecules. These attractive forces bind the ligands to the metal ion. The bonds between metal and
the surrounding ligands are purely ionic. This theory does not consider any orbital overlapping.
4
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     DPP : 02 / Co-ordination compounds
(b) The repulsive forces arise between the lone pairs on the ligands and electrons in the d-orbitals of
the metal or atom. The crystal field theory mainly focuses on these repulsive forces. These forces
are responsible for causing a considerable effect on the relative energies of the d-orbitals of the
central metal ion or atom.
IV . In a free transition metal or ion, there are five d-orbitals which are designated as d
xy
, d
yz
, d
zx
, d
x
2
—y
2 and d
x
2.
The five d-orbitals are divided into two sets depending on the nature of their orientation in space.
(a) The three d-orbitals (d
xy
, d
yz
, d
zx
) which orient in the regions between the coordinate axes are
designated as t
2g
–orbitals (pronounced as “t-two-g”). t
2g
–orbitals are three-fold degenerate. These
are non-axial orbitals.
(b) The other two orbitals (d
x
2
—y
2), d
z
2 which orient along the axes are called e
g
, (e
g
 orbitals are two fold
degenerate). These are also called axial-orbitals.
[The names t
2g
 and e
g
 are derived from spectroscopic terms.]
In a free transition metal ion or atom, all the five d-orbitals same energy, i.e., they are degenerate. However, when
the ligands approach the central metal or atom are repelled by lone pairs of the ligands. As a result of these
interactions, the degeneracy of d-orbitals of the metal ion is lost depending on the orientation of ligands in space.
The d-orbitals split into two sets of orbitals having different energies. This is called crystals field splitting. It is the
basis of crystal field theory. The extent of splitting depends on the number of ligands and their position around the
metal atom or ion. The splitting is different in different structures with different coordination numbers.
Splitting of d-orbitals in octahedral complexes.
0
?
e
g
t
2g
Average energy
(t
2g
+ e
g
) of orbitals in
spherical crystal field
e
g
t
2g
(t
2g
+ e
g
) orbitals in
free ion
(a)(b)(c) 10 Dq
Energy
No splitting state
Crystal field splitting in tetrahedral complexes.
t
?
Average energy
(t
2g
+ e
g
) of orbitals in
spherical crystal field
e
g
t
2g
(t
2g
+ e
g
) orbitals in
free ion
10 Dq
Energy
4 Dq
6 Dq
5
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Crystal field splitting in square planar complexes.
2 2
x y
d
- 2
z
d
Octahedral
d
xy
d
xz d
yz
e
g
t
2g
Average energy
of orbitals in
spherical crystal field
d-orbitals in
free ion
Spectrochemical Series. When the ligands are arranged in order of the magnitude of crystal field splitting, the
arrangement, thus, obtained is called spectrochemical series.
I
—
 < Br
—
 < Cl
—
 <
3
NO
- < F
—
 < OH
—
 < OX
2—
 < H
2
O < py = NH
3
< en < dipy < o-phen <
2
NO
- < CN
—
 < CO
Weak field ligands Increasing crystal field Strong field ligands
Objective Questions
1. The unpaired electrons in Ni(CO) 4
 are
(a) zero(b) one(c) three(d) four
2. Which of the following compounds is not coloured ?
(a) Na
2
[CuCl
4
](b) Na
2
[CdCl
4
](c) K
4
[Co(CN) 6
](d) K
3
[Fe(CN) 6
]
3. Which of the following has square planar structure ?
(a) Na
2
[CuCl
4
](b) [NiCl
4
]
2–
(c) [Ni(CN) 4
]
2–
(d) all of these
4. The atomic numbers of chromium and iron are 24 and 26 respectively. Which one of the following complexes
exhibits paramagnetic character due to electronic spin?
(a) [Fe(CO) 5
](b) [Cr(NH
3
) 6
]
3+
(c) [Fe(CN) 6
]
4–
(d) [Cr(CO) 6
]
5. Which one of the following complexes is diamagnetic in nature?
(a) [Cr(NH
3
) 6
]
3+
(b) [FeF
6
]
3–
(c) [Co(NH
3
) 6
]
3+
(d) [Fe(H
2
O) 6
]
2+
6. Amongst Ni(CO) 4
,[Ni(CN) 4
]
–2
 and [NiCl
4
]
–2
(a) Ni(CO) 4
 and [Ni(Cl) 4
]
–2
are diamagnetic and  [Ni(CN) 4
]
–2
 is paramagnetic
(b) [Ni(CN) 4
]
–2
 and [Ni(Cl) 4
]
–2
are diamagnetic and Ni(CO) 4
 is paramagnetic
(c) Ni(CO) 4
 and [Ni(CN) 4
]
–2
are diamagnetic and [Ni(Cl) 4
]
–2
 is paramagnetic
(d) Ni(CO) 4
 is diamagnetic  and [NiCl
4
]
–2
 &[Ni(CN) 4
]
–2
 is paramagnetic
7. Which of the following system has maximum number of unpaired electrons ?
(a) d
4
(octahedral, low spin)(b) d
6
(tetrahedral, high spin)(c) d
6
(octahedral, low spin)(d) d
9
(octahedral, high spin)
Read More
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