To find the refractive indices of (a) water (b) oil (transparent) using a plane mirror, an
equiconvex lens (made from a glass of known refractive index) and an adjustable
object needle
3. A)SIMPLE SALT
• A simple salt is formed by the
neutralization of an acid by a base.
• KOH + HCl → KCl + H2O
• Normally, a simple salt ionizes in water
and produces ions in solution.
• The solution of the simple salt exhibits
the properties of its component ions.
5. • ii) Coordination (or complex) compounds
Coordination compound is a compound
formed from a Lewis acid and a Lewis
base. The molecular compounds, do not
dissociate into its constituent ions in
solution are called coordination
compounds.
• Ex : Fe(CN)2 + 4KCN → Fe(CN)2 . 4KCN
(or) K4[Fe(CN)6]
6. Coordination (or complex) compounds
• K4[Fe(CN)6] →4K+ + [Fe(CN) 6]4-
Complex anion
• In K4[Fe (CN)6] the individual
components lose their identity.
• The metal of the complex ion is not
free in solution unlike metal in
double salt in solution.
7. DOUBLE SALTS
They completely ionise
in aqueous solutions
and each ion in the
solution gives the
corresponding
confirmatory test.
Example: Potash Alum
is double sulphate
K2SO4.Al2 (SO4)3.24H2O
on Ionization it gives:K+,
SO4
−2 and Al+3 ions
which response to the
COORDINATION COMPLEX
Co-ordinate complexes
ionise incompletely in the
aqueous solutions. These
give a complex ion which
does not show complete
ionization.
• Example: Potassium
Ferrocyanide. [K4Fe(CN)6
It ionizes to give K+ and
[Fe(CN)6]−4 [ferro cyanide
ions]
8. COORDINATION COMPOUNDS
• The compounds in which the
metal atoms are bound to a
number of anions or neutral
molecules are called as complex
compounds or coordination
compounds.
9. THE IMPORTANT APPLICATIONS OF
COORDINATION COMPOUNDS :
• Due to the formation
• of cyanide complexes
• (dicyanoaurate and
• dicyanoargentate)
noble metals like gold
and silver are extracted
from their ore.
13. • When aqueous
ammonia is mixed
with the copper
sulphate solution, a
deep blue complex
soluble in water is
formed. This
reaction is helpful in
detecting cupric ions
present in the salt.
14. WERNER’S EXPERIMENT
Werner conducted an experiment by
mixing AgNO3(silver nitrate) solution with
CoCl3·6NH3, all three chloride ions got
converted to AgCl (silver chloride).
However, when AgNO3 was mixed with
CoCl3·5NH3, two moles of AgCl were
formed. Further, on mixing
CoCl3·4NH3 with AgNO3, one mole of AgCl
was formed. Based on this observation,
Werner’s theory was postulated.
16. WERNER’S THEORY
• Alfred Werner in 1898 proposed
Werner’s theory explaining the structure
of coordination compounds, based on his
observation.
• POSTULATES OF WERNER’S THEORY:
1.The central metal atom in the
coordination compound exhibits
two types of valency, namely, primary
and secondary linkages or valencies.
2.Primary linkages are ionizable and are
satisfied by the negative ions.
17. POSTULATES OF WERNER’S THEORY:
3.Secondary linkages are non-ionizable.
These are satisfied by negative ions or
neutral molecules. Also, the secondary
valence is fixed for any metal and is
equal to its coordination number.
4.The ions bounded by the secondary
linkages to the metal exhibit
characteristic spatial arrangements
corresponding to different coordination
numbers.
18.
19. LIMITATIONS OF WERNER’S THEORY
• It fails to explain the magnetic, colour and
optical properties shown by coordination
compounds.
• It failed to explain the reason why all
elements don’t form coordination
compounds.
• It failed to explain the directional properties
of bonds in coordination compounds.
• This theory does not explain the stability of
the complex
• This theory could not explain the nature of
complexes
21. COORDINATION ENTITY
A chemical compound in which the
central ion or atom (or the
coordination centre) is bound to a set
number of atoms, molecules, or ions
is called a coordination entity.
Some examples of such coordination
entities include [CoCl3 (NH3) 3] and
[Fe(CN) 6]4-.
23. CENTRAL ATOMS AND CENTRAL IONS
• The atoms and ions to which a set
number of atoms, molecules, or ions are
bound are referred to as the central
atoms and the central ions.
• In coordination compounds, the central
atoms or ions are typically Lewis
Acids and can, therefore, act as electron-
pair acceptors.
25. LIGANDS
• The atoms, molecules, or ions that are
bound to the coordination centre or the
central atom/ion are referred to
as ligands.
• These ligands can either be a simple ion
or molecule (such as Cl- or NH3) or in the
form of relatively large molecules, such as
ethane-1,2-diamine (NH2-CH2-CH2-NH2).
26.
27. CLASSIFICATION OF LIGANDS
• Based on the nature of the charge on the
ligand and the central atom, ligands are
classified as follows:
• Anionic ligands: CN–, Br–, Cl–
• Cationic ligands: NO+ (Nitrosonium ion)
• Neutral ligands: CO, H2O, NH3
28.
29. DENTICITY
DENTICITY: It is the number of donor
groups in a single ligand that bind to a
central atom in a coordination complex.
• Based on the denticity, ligands are
classified as follows:
• UNIDENTATE LIGANDS
• BIDENTATE LIGANDS
• POLYDENTATE LIGANDS
• AMBIDENTATE LIGAND
• CHELATE LIGANDS
30. MONO/UNIDENTATE LIGANDS
• The ligands which only have one
atom that can bind to the
coordination centre are called
unidentate ligands. Ammonia
(NH3 ) is an example of a unidentate
ligand. Some common unidentate
are Cl–, H2O etc.
31.
32.
33. BIDENTATE LIGANDS
Ligands which have the ability to bind to
the central atom via two separate donor
atoms, such as ethane-1,2-diamine and
Oxalate ion are called bidentate as it can
bond through two atoms to the central
atom in a coordination compound and
Ethane-1, 2-diamine
36. POLYDENTATE LIGANDS
• Some ligands have many donor atoms that
can bind to the coordination centre. These
ligands are often referred to as polydentate
ligands.
• A great example of a polydentate ligand is
the EDTA4–ion (ethylene diamine
tetraacetate ion), which can bind to the
coordination centre via its four oxygen
atoms and two nitrogen atoms.
38. CHELATE LIGANDS
• When a polydentate ligand attaches
itself to the same central metal atom
through two or more donor atoms, it
is known as a chelate ligand. The
number of atoms that ligate to the
metal ion are termed as the denticity
of such ligands.
39.
40.
41. • Di or polydentate ligands cause cyclisation
around the metal atom which is known as
chelation . Such ligands use two or more
donor atoms to bind a single metal ion and
are known as chelating ligands.
• More the number of chelate rings, more is
the stability of complex.
• The stabilisation of coordination
compounds due to chelation is known as
chelate effect.
42. AMBIDENTATE LIGAND
• Some ligands have the ability to bind to
the central atom via the atoms of two
different elements.
• For example, the SCN– ion can bind to a
ligand via the nitrogen atom or via the
sulphur atom. Such ligands are known as
ambidentate ligands.
45. COORDINATION NUMBER
The coordination numberof the central
atom in the coordination compound
refers to the total number of bonds
through which the ligands are bound to
the coordination centre.
For example, in the coordination
complex given by [Ni(NH3)4] 2+,
the coordination number of nickel is 4.
46. CALCULATION OF COORDINATION NUMBER
IN CASE OF MONODENTATE LIGANDS,
• Coordination number = number of ligands
IN POLYDENTATE LIGANDS.
• Coordination number = number of ligands * denticity
47. COORDINATION SPHERE
• The non-ionizable part of a complex compound
which consists of central transition metal ion
surrounded by neighbouring atoms or groups
enclosed in square bracket.
• The coordination centre, the ligands attached to
the coordination centre, and the net charge of
the chemical compound as a whole, form
the coordination sphere when written together.
• This coordination sphere is usually accompanied
by a counter ion (the ionizable groups that attach
to charged coordination complexes).
• Example: [Co(NH3)6]Cl3
48.
49. COORDINATION POLYHEDRON
• The geometric shape formed by the
attachment of the ligands to the
coordination centre is called
the coordination polyhedron.
• Examples of such spatial
arrangements in coordination
compounds include tetrahedral and
square planar.
51. OXIDATION NUMBER
The oxidation number of the central
atom can be calculated by finding the
charge associated with it when all the
electron pairs that are donated by the
ligands are removed from it.
For example, the oxidation number of
the platinum atom in the complex
[PtCl6]2- is +4.
52.
53. TYPES OF COORDINATION COMPLEXES
based on whether complex ion is a
cation/anion
• 1.Cationic complexes: In this co-ordination
sphere is a cation. Example: [Co(NH3) 6]Cl3
• 2.Anionic complexes: In this co-ordination
sphere is Anion. Example: K4[Fe(CN) 6]
• 3.Neutral Complexes: In this co-ordination
sphere is neither cation or anion. Example:
[Ni(CO) 4]
54. TYPES OF COORDINATION COMPLEXES
based on the no.of central atoms/ions
present
• Mononuclear complexes: In this co-
ordination sphere has single transition
metal ion. Example: K4[Fe(CN) 6]
• Polynuclear complexes:
• More than one transition
• metal ion is present.
• Example:
55. TYPES OF COORDINATION COMPLEXES
based on the types of ligands present
• Homoleptic complex: The complex
consist of a similar type of ligands.
Example: K4[Fe(CN)6]
• Heteroleptic complexes: These
consists of different types of ligands.
Example: [Co(NH3)5Cl]SO4
56. PROPERTIES OF COORDINATION
COMPOUNDS
The coordination compounds formed by the
transition elements are coloured due to the
presence of unpaired electrons that absorb
light in their electronic transitions. For
example, the complexes containing Iron(II)
can exhibit green and pale green colours,
but the coordination compounds containing
iron(III) have a brown or yellowish-brown
colour.
57. • When the coordination centre is a metal,
the corresponding coordination
complexes have a magnetic nature due to
the presence of unpaired electrons.
• Coordination compounds exhibit a variety
of chemical reactivity. They can be a part
of inner-sphere electron transfer
reactions as well as outer-sphere electron
transfers.
• Complex compounds with certain ligands
have the ability to aid in the
transformation of molecules in a catalytic
or a stoichiometric manner.
•
59. Rules For Naming Coordination Compound
1.The ligands are always written before the
central metal ion in the naming of
complex coordination complexes.
2.When the coordination centre is bound
to more than one ligand, the names of
the ligands are written in an alphabetical
order which is not affected by the
numerical prefixes that must be applied
to the ligands.
60. Rules For Naming Coordination
Compound
3.When there are many monodentate
ligands present in the coordination
compound, the prefixes that give
insight into the number of ligands are
of the type: di-, tri-, tetra-, and so on.
4.When there are many polydentate
ligands attached to the central metal
ion, the prefixes are of the form bis-,
tris-, etc.
61. Rules For Naming Coordination
Compound
5.The names of the anions present in a
coordination compound must end with the
letter ‘o’, which generally replaces the letter
‘e’. Therefore, the sulphate anion must be
written as ‘sulfato’ and the chloride anion
must be written as ‘chlorido’.
6. The following neutral ligands are assigned
specific names in coordination compounds:
NH3 (ammine), H2O (aqua or aquo), CO
(carbonyl), NO (nitrosyl).
62. Rules For Naming Coordination Compound
7.After the ligands are named, the name of
the central metal atom is written. If the
complex has an anionic charge associated
with it, the suffix ‘-ate’ is applied.
8.When writing the name of the central
metallic atom in an anionic complex,
priority is given to the Latin name of the
metal if it exists (with the exception of
mercury).
63. Rules For Naming Coordination Compound
9.The oxidation state of the central metal
atom/ion must be specified with the help
of roman numerals that are enclosed in a
set of parentheses.
10.If the coordination compound is
accompanied by a counter ion, the
cationic entity must be written before the
anionic entity.
64. Examples of Naming Coordination
Compounds
K4[Fe(CN)6]:Potassium hexacyanidoferrate (II)
[Ni(CN)4]−2:Tetra cyanidonickelate (II) ion.
[Zn(OH)4]−2:Tetra hydroxidozincate(II) ion.
[Ni(CO)4]: Tetra carbonyl Nickel (O).
[Co(NH3)4(H2O)2]Cl3:
Tetraamminediaquacobalt(IlI) chloride
[Cr(en)3]Cl3: Tris(ethane-1,2-diamine)
chromium(III) chloride
70. FORMULAS OF MONONUCLEAR
COORDINATION ENTITIES:
• The following rules are applied while
writing the formulas:
• Central atom is listed first.
• Ligands are then listed in alphabetical
order. The placement of a ligand in the
list does not depend on its charge.
71. • 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.
• 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.
72. • There should be no space between the
ligands and the metal within a coordination
sphere.
• 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.
• The charge of the cation(s) is balanced by
the charge of the anion(s).
73. Write the formulas for the following
coordination compounds:
(i) Tetraamminediaquacobalt(III)chloride-
[Co(NH3)4(H2O)2]Cl3
(ii) Potassium tetracyanidonickelate(II) ANS: K2[Ni(CN)4]
(iii)Tris(ethane–1,2–diamine) chromium(III) chloride
[Cr(en) 3]Cl3
(iv) Amminebromidochloridonitrito-N-platinate(II)
[Pt(NH3)(Br)(Cl)(NO2)] -
(v) Dichloridobis(ethane–1,2–diamine)platinum(IV) nitrate
[PtCl2(en)2](NO3)2
(vi) Iron(III) hexacyanidoferrate(II) ANS: Fe4[Fe(CN)6]3
75. ISOMERISM IN COORDINATION COMPOUNDS
• Two or more compounds that have the
same chemical formula but a different
arrangement of atoms are known as
isomers. Due to this difference in the
arrangement of atoms, coordination
compounds pre-dominantly exhibit two
types of isomerism namely, stereo-
isomerism and structural isomerism.
76.
77. STRUCTURAL ISOMERISM
• Structural isomerism is exhibited by the
coordination compounds having the
same chemical formula but a different
arrangement of atoms. These are further
divided into four types:
• 1. Linkage Isomerism
• 2. Coordination Isomerism
• 3. Ionisation Isomerism
• 4. Solvate Isomerism
78. LINKAGE ISOMERISM
• Linkage isomerism is exhibited by
coordination compounds having
Ambidentate ligands, which may bind to
the central metal atom through different
atoms of the ligand like SCN& NCS, NO2
&ONO,etc.
• For example:[Co(NH3)5NO2]Cl2(RED)and
[Co(NH3)5ONO] Cl2(YELLOW)
79.
80. COORDINATION ISOMERISM
• In coordination isomerism, the
interchange of ligands between
cationic and anionic entities of
different metal ions present in
coordination compounds takes place.
• For example: [Co(NH3)6][Cr(CN)6] and
[Cr(NH3)6][Co(CN)6].
81.
82. IONISATION ISOMERISM
• Ionisation isomerism arises when
the counter ion in a complex salt
which is a potential ligand
replaces the ligand.
• For example: [Co(NH3)5(SO4)]Br
and [Co(NH3)5Br]SO4.
84. SOLVATE ISOMERISM
Solvate isomers are a special case of
ionisation isomerism in which compounds
differ depending on the number of the
solvent molecules directly bonded to the
metal ion. If water molecules are the
solvent molecules present, it is called
HYDRATE ISOMERISM.
For example:CrCl3.6H2O
85. EXAMPLE FOR SOLVATE AND HYDRATE
ISOMERISM
[Cr(H2O)4Cl2]Cl.2H2O - Bright green
Tetraaquadichlorochromium(III) chloride
dihydrate
[Cr(H2O)5Cl]Cl2.H2O - grey-green
Pentaaquachlorochromium(III) chloride
monohydrate
[Cr(H2O)6]Cl3 - Violet
Hexaaquachromium(III) chloride
86. STEREOISOMERISM
Coordination compounds which have
the same chemical formula
and chemical bonds but have different
spatial arrangement are known as
stereoisomers. These are further
divided into optical isomerism and
geometrical isomerism.
87. GEOMETRIC OR CIS-TRANS ISOMERS
• Geometrical isomerism is observed in
heteroleptic complexes (complexes with
more than one type of ligands) due to
different possible geometric
arrangements of the ligands.
• This behaviour is mainly observed in
coordination compounds having
coordination numbers equal to 4 and 6.
91. Example of MA 2B2
complex
• ML4 and
tetrahedral
complexes do
not show cis-
trans isomerism.
• MABCD has 3
geometrical
isomers. 2-cis
and 1-trans.
• MA 2B2 complex
shows cis and
trans isomers.
92. If two ligands in an octahedral complex are
different from the other four, giving an
Ma4b2 complex, two isomers are possible.
The two b ligands can be cis or trans. Cis-
and trans-[Co(NH3)4Cl2] Cl
CIS- AND TRANS- ISOMERISM
IN
OCTAHEDRAL COMPLEXES
93.
94.
95. The octahedral complexes of the type
[M(AA) 3]n+ havingsymmetrical
bidentate ligands do not show
geometrical isomerism.
However, complexes, [M(AA)2B2]n+
and [M(AA)B2C2]n+ give two geometric
isomers each.
97. FACIAL AND MERIDIONAL ISOMERISM
Replacing another A ligand
by B gives MA3B3 complex
for which there are two
possibleisomers.When three
identical ligands occupy one
face, the isomer is called
facial, or fac. If the three
ligands and the metal ion
are in one plane, it is
meridional/ mer-isomer.
100. Chemistry of
Coordination
Optical isomers
• The optical isomers or enantiomers, are
mirror images of each other and two
enantiomers cannot be superimposed on
each other
Compounds
102. Enantiomers
A molecule or ion that exists as a pair of
enantiomers is said to be chiral.Each
form is called –Laevo(l-) and dextro(d-)
Laevo(l-) dextro(d-)
103.
104.
105. VALENCE BOND THEORY (VB THEORY)
It primarily the work of Linus Pauling
The postulates of valence bond theory:
The central metal atom/ion makes available
a number of vacant orbitals equal to its
coordination number. These vacant orbitals
form covalent bonds with the ligand orbitals.
A covalent bond is formed by the overlap of
a vacant metal orbital and filled ligand
orbitals. This complete overlap leads to the
formation of a metal ligand,σ (sigma) bond.
106. VALENCE BOND THEORY (Continued)
A strong covalent bond is formed only when
the orbitals overlap to the maximum extent.
This maximum overlapping is possible only
when the metal vacant orbitals undergo a
process called ‘hybridisation’. A hybridised
orbital has a better directional characteristics
than an unhybridized one.
107. The following table gives the coordination
number, orbital hybridisation and geometry
Coordination
number
Types of
hybridization
Geometry
2 Sp Linear
4 sp3 Tetrahedral
4 dsp2 square planar
6 d2sp3 Octahedral
6 sp3d2 Octahedral
108.
109. MAGNETIC MOMENT
A species having at least one unpaired
electron, is said to be paramagnetic.
• It is attracted by an external field. The
paramagnetic moment is given by the
following spin-only formula.
• BM
• μs = spin-only magnetic moment , n=number
of unpaired electrons
116. LIMITATIONS OF VALENCE BOND THEORY:
• It involves a number of assumptions.
• It does not give quantitative interpretation of
magnetic data.
• It does not explain the colour exhibited by
coordination compounds.
• It does not give a quantitative interpretation
of the thermodynamic or kinetic stabilities of
coordinationcompounds.
• It does not make exact predictions regarding
the tetrahedral and square planar structures
of 4-coordinate complexes.
• It does not distinguish between weak and
117. CRYSTAL FIELD THEORY (CFT)
Main postulates of crystal field theory are
In a coordination compound there are
electrostatic interaction between metal
atom/ion and ligands. Ligand assumed to be
a point charge
In an isolated metal atom or ion all five d-
orbitals have equal energy i.e. they are
degenerate
When metal atom/ion gets surrounded by
ligands, there occur interaction between d-
electron cloud of metal atom/ion and ligand
118. CRYSTAL FIELD THEORY (Continued)
• If the field due to ligand around metal atom is
spherically symmetrical, d-orbitals of metal remains
degenerated
• If field due to ligand surrounding metal is
unsymmetrical ( as in octahedral and tetrahedral
complexes) the degenaracy of d-orbitals is splitted
into two sets of orbitals
• Orbitals lying in the direction of ligands (point
charges) are raised to higher energy state than those
orbitals lying between the ligands ( point charges)
• The energy difference between two sets of orbitals is
denoted by Δo for octahedral and Δt for tetrahedral
• The magnitude of Δo and Δt depends upon the field
strength of ligand around the metal
119. SPECTROCHEMICAL SERIES.
The arrangement of ligands in order of their
increasing CFSE values is known as
spectrochemical series. The ligands with
small CFSE values are called weak field
ligands, whereas those with large value of
CFSE are called strong field ligands.
120. • The spectrochemical series is an
experimentally determined series. It is
difficult to explain the order as it
incorporates both the effect of σ and π
bonding.
• A pattern of increasing σ donation is as
follows-
Halides donors < O donors < N donors <
C donors
121. CRYSTAL FIELD SPLITTING IN
OCTAHEDRAL COMPLEXES
• ligands approaching the x, y, and
• z axis. The two d orbitals namely
• d(x2 –y2) and d(z2) will suffer
• More electrostatic repulsion
and hence their energy will be
• greater than other three orbitals
• d(xy), d(yx) and d(xz) which will
• have their lobes lying between the axis
122. As a result, a set of d-orbitals split into
two sets: eg orbitals of higher energy
including d(x2 –y2) and d(z2) and t2g
orbitals of lower energy including d(xy),
d(yx) and d(xz)
123. The crystal field splitting is measured in
terms of energy difference between t2g
and eg orbital and is denoted by a symbol
Δo . It is generally measured in terms of
Dq. It is called as crystal field splitting
energy or crystal field stabilization
energy Eg orbitals are 6Dq above the
average energy level and t2g orbitals are
4Dq below the average energy level
124. The energy of eg set of orbitals > energy of
t2g set of orbitals.
Ligands for which energy separation, Δo < P
(the pairing energy, i.e., energy required for
electron pairing in a single orbital) form a
high spin complex.
Ligands for which energy separation, Δo > P,
form low spin complex.
125. CRYSTAL FIELD SPLITTING IN
TETRAHEDRAL COMPLEXES
(b) Crystal field splitting in
tetrahedral coordination
entities
In tetrahedral coordination
entity formation,the d
orbital splitting is inverted
and is smaller as
compared to the
octahedral field splitting.
126. For the same metal, the
same ligands and metal-
ligand distances, it can be
shown that Dt=(4/9)Do
Consequently, the orbital
splitting energies are
not sufficiently large
for forcing pairing and
therefore, low spin
configurations are rarely
127. The energy of t2g set of orbitals >
Energy of eg set of orbitals.
In such complexes d-orbital splitting is
inverted and is smaller as compared to
the octahedral field splitting.
No pairing of electrons is possible due
to the lowest splitting energies which
leads to high spin complexes.
130. COLOUR IN COORDINATION COMPOUNDS
• Complexes in which central transition metal ion
contains unpaired electrons shows colour. It is ‘d – d’
transition.
• In coordination complexes energy difference (∆)
between two d-sets of d-orbitals is small. Radiations of
appropriate frequency absorbed from visible region can
cause excitation of d-electron from lower energy orbital
to higher energy orbital. Remaining light is transmitted
and compound appears coloured
• This frequency generally lies in the visible region. The
colour observed corresponds to the complementary
colour of the light absorbed. The frequency of the light
absorbed is determined by the nature of the ligand.
131.
132. THE FACTORS AFFECTING THE COLOUR OF COMPLEXES
Number of unpaired electrons in transition meta
ion
Nature of ligands
The oxidation state of central metal ion
The wavelength of light absorbed and emitted
The proportion of ligands in the coordination
sphere
Ex: [Ni(H2O)6] +2 + en(aq)→[Ni(H2O) 4(en)]+2
Green Pale blue
133. • It is important to note that in the
absence of ligand, crystal field
splitting does not occur and hence
the substance is colourless. For e.g.
removal of water from [Ti (H2O) 6] Cl3
on heating renders it colourless.
Similarly, anhydrous copper sulphate
is white, but copper sulphate
pentahydrate is blue in colour.
134. BONDING IN METAL COMPLEXES [METAL
CARBONYLS]
• Complexes in which carbon monoxide
acts as ligands are metal carbonyls
• Example: [Ni(CO)4] Tetracarbonyl Nickel
(0) and [Fe(CO)5] Penta Carbonyl Iron (0)
In these complexes, complexes, a′σ‘ bond
is formed by the overlapping of vacant ‘d’
orbital of metal ion and filled orbital of C-
atom (carbon).
135. A π bond is formed
by the lateral
overlapping of
filled inner orbitals
of metal ion and
vacant of the
carbon atom. Thus
synergic bonding
exist in metal
carbonyls
136.
137. STABILITY OF COMPLEXES
A complex is formed
in several steps.
Each process step is
reversible and the
equilibrium constant
is known as
stepwise formation
constant. Let us
consider the
formation of
complex ML4
139. The factors on which stability of the
complex depends :
(i) Charge on the central metal atom As
the magnitude of charge on metal atom
increases, stability of the complex
increases.
(ii) Nature of metal ion The stability order
is 3d < 4d < 5d series.
(iii) Basic nature of ligands Strong field
ligands form stable complex.
140. APPLICATIONS OF COORDINATION COMPOUNDS
• The colour of the
coordination
• compounds containing
transition metals causes
them to be extensively
used in industries for the
colouration of materials.
They find applications in
the dye and pigment
industries.
142. IN METALLURGY
• In the extraction
of gold, silver by
Mac Arthur
Forest Process
involves a
complex of
cyanide ions.
143. Coordination complexes are very useful
in the extraction of many metals from
their ores. For example, nickel and cobalt
can be extracted from their ores via
hydro- metallurgical processes involving
ions of coordination compounds.
• For example-
144. APPLICATIONS IN BIOLOGY
• Haemoglobin
consists of
Haeme complex-
ion which has
tetrapyrrole
Porphyrin ring
structure with
central Fe2+ ion.
145. • Chlorophyll is a
coordination
compound of
magnesium which is
present in the plants
and plays an
imporatant role in the
preparation of food by
photosynthesis .
146. Vitamin B-12 consists
of tetrapyrrole
porphyrin ring complex
with central Co+3 ion
and its coordination
number is 6.
151. BEYOND THE TEXT BOOK
• EFFECTIVE ATOMIC NUMBER (EAN)
• The sum of the number of electrons, donated by all ligands and those
present on the central metal ion or atom in complex is called as effective
atomic number (EAN).
• • Generally EAN of central metal ion will be equal to the number of
electrons in the nearest noble gas.
• • If the EAN of the central metal is equal to the number of electrons in the
nearest noble gas then the complex possess greater stability.
• EAN = [(atomic number of central metal) – (the oxidation state of the
metal) + (the number of electrons gained by the metal from the ligands
through co-ordination)]
• EAN= [Z metal – (oxidation state of the metal) + 2(coordination number of
the metal)].
• for example.
[Co(NH3)6]+3 →EAN = [27 – 3 + 2(6)] = 36
