9. Coordination compounds class 12-1.pptx

CLASS 12, CHEMISTRY, CHAPTER NO. 9
COORDINATION COMPOUNDS
BY: MEGHA SOLANKI, PGT- CHEMISTRY, KV MANKHURD, MUMBAI

REVISED SYLLABUS FOR 2020-21
• Coordination compounds - Introduction,
• ligands,
• coordination number,
• colour.
• Magnetic properties and shapes.
• IUPAC nomenclature of mononuclear
coordination compounds.
• Bonding - Werner's theory, VBT, and CFT.

COORDINATION COMPOUDS IN OUR DAILY LIFE
• Earlier, these compounds were called as complex compounds.
• Chlorophyll, haemoglobin and
vitamin B12 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. vitamin B12 haemoglobin

COORDINATION COMPOUNDS DISCOVERY
• The sustained and systematic development of modern coordination chemistry
began with the discovery by the French chemist B.M. Tassaert in 1798 that
ammoniacal solutions of cobalt chloride (CoCl3) develops a brownish- red
colour.
• Further experiments were done by other scientists on this solutions and it was
found that it was a mixture of various coloured compounds of Co, Cl and NH3.
• These compounds were separated and it was found that in these compounds,
the no. of Co and Cl atoms remain same but no. of NH3 molecules changes.
And simultaneously the colour of the compounds change.
• When 1 mole of each of these compounds was treated with excess AgNO3
solution, different no. of moles of AgCl were precipitated as shown in the table
below.

The above results can be explained if we consider
the following chemical structure of the above
compounds:

• It is clear from the structural chemical formula of these
compounds that only the ions present outside the square
brackets can react with AgNO3 and they are called primary
valence.
• All the atoms within square brackets behave like a single ion.
• 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, CoCl3.4NH3, but distinct properties. Such
compounds are termed as isomers.

WERNER’S THEORY OF COORDINATION COMPOUNDS
• Alfred Werner (1866-1919), a Swiss chemist was the first to formulate his ideas
about the structures of coordination compounds. Werner in 1898 gave 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 pri. valences are normally ionisable and are satisfied by -ve ions(counter ions)
3. The sec. valences are non ionisable. These are satisfied by neutral molecules or
-ve ions. The sec. valence is equal to the coordination no. and is fixed for a metal.
4. The ions/groups bound by the sec. linkages to the metal have characteristic
spatial arrangements( called polyhedra) corresponding to different coordination
no. The metal along with secondarily linked groups are called coordination
compounds.
5. octahedral, tetrahedral and square planar geometrical shapes are more common
in coordination compounds of transition metals.
Examples: [Co(NH3)6]3+, [CoCl(NH3)5]2+ and [CoCl2(NH3)4]+ are octahedral entities,
while [Ni(CO)4] is tetrahedral and [PtCl4]2– is square planar.

NOW SOLVE THIS
Example 9.1: On the basis of the following observations made with aqueous
solutions, assign secondary valences to metals in the following compounds:
ANSWERS: (i) 4. (ii) 6. (iii) 6. (iv) 6 (v) 4

DIFFERENCE BETWEEN A DOUBLE SALT AND A COMPLEX

DOUBLE SALT VS COORINATION COMPOUNDS

SOME EXAMPLES OF DOUBLE SALT
potash Alum - K2SO4.Al2(SO4)3.24H2O
Ammonium Alum - (NH4)2 Al2(SO4)3.24H2O
Mohr’s Salt - FeSO4(NH4)2 SO4.6H2O
Carnallite - KCl.MgCl2.6H2O,

SOME IMPORTANT TERMS PERTAINING TO 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, [CoCl3(NH3)3] is
a coordination entity in which the cobalt ion is surrounded by 3
ammonia molecules and 3 chloride ions. Other examples are:
[Ni(CO)4], [PtCl2(NH3)2], [Fe(CN)6]4–, [Co(NH3)6]3+.
(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.

Types of ligand
There are basically three types of ligands:
1. Monodentate or Unidentate ligands
2. Polydentate ligand or multidentate ligand
3. Ambidentate ligand
4. Chelate ligand
1. Monodentate 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. 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
4. Chelate ligand
Ethylenediaminetetraacetate
ion (EDTA4–) is an important
hexadentate ligand. It can
bind through two nitrogen
and four oxygen atoms to a
central metal ion. EDTA4- EDTA4- in chelation
Ethylenediamine-metal
chelate

(d)Coordination number(CN)
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 or inner
sphere. The ionisable groups are
written outside the bracket and are
called counter ions or outer sphere or
ionization sphere. For example, in the
complex K4[Fe(CN)6], the coordination
sphere is [Fe(CN)6]4– and the counter
ion is K+.

• 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(NH3)6]3+ is octahedral, [Ni(CO)4] is tetrahedral and
[PtCl4]2– is square planar.
(f) Coordination polyhedron
Shapes of different coordination polyhedra

(g) Oxidation number of central atom

(g) Homoleptic and Heteroleptic ligands

LATIN NAMES OF SOME TRANSITION ELEMENT

• Compounds that have the same chemical formula
but different structural arrangements are called
isomers.
• In coordination compounds, due to their
complicated formulae, the variety of bond types
and the number of shapes possible, many
different types of isomerism occur.
ISOMERISM IN COORDINATION COMPOUNDS
***** Isomerism may be skipped as it is not in syllabus for 2020-21 *****

STEREO
ISOMERISM
They have same
chemical formula and
chemical bonds but
they have different
spatial arrangement
STRUCTURAL
ISOMERISM
ISOMERISM
Structural isomers
have the same
chemical formula but
they have different
bonds

STEREOISOMERISM: GEOMETRICAL ISOMERISM
• This type of isomerism arises in heteroleptic complexes due to different
possible geometric arrangements of the ligands.
• It is mostly found in the complex
compounds with coordination
numbers 4 and 6 but not found in
the compounds with CN = 2 and 3.
It is because different geometrical
arrangements of ligands are not
possible around the central metal
in these complexes.
All 3 arrangements are similar to one- another
in CN = 3 with trigonal planar geometry
Both arrangements are similar to each-
other for CN =2, with linear geometry
• Geometrical isomerism is also not
possible in tetrahedral complexes
with CN = 4 because all the four
ligands are adjacent or equidistant to
one another in tetrahedral complex.

GEOMETRICAL ISOMERISM IS CIS-TRANS ISOMERISM

SQUARE PLANAR COMPLEX WITH FORMULA MA2B2

GEOMETRICAL ISOMERISM IN SQUARE PLANAR
COMPLEX WITH FORMULA MA2BC
They also exhibit 1 cis and 1 trans isomer
Cis isomer –When similar ligands(AA) are adjacent.
Trans isomer –When similar ligands are at opposite corners.

GEOMETRICAL ISOMERISM IN SQURE PLANAR
COMPLEX WITH FORMULA MABCD
They show three isomers-two cis and one trans.
Here, we fix the position of one ligand and found its isomers for any another
ligand. For example, in the following figures, a is cis with d two times and
trans with d one time. More than 3 arrangements are not possible
In the following example, Br has a fixed position. With any other ligand, it will
remain cis two times at both the adjacent corners while it remains trans with
every one only once when the ligand takes the opposite corner.

GEOMETRICAL ISOMERISM IN SQUARE PLANAR COMPLEX WITH
UNSYMMETRICAL BIDENTATE LIGAND OF FORMULA M(AB)2
In this, when ligands connect to metal
in same manner, they form cis-isomer
and when connect in opposite
manner they form trans isomer
EXAMPLE:
Square planner complexes with formula MA4 and MA3B have no
geometrical isomers as only one arrangement is possible for them.

SUMMARY OF GEOMETRICAL ISOMERS IN SQUARE PLANAR COMPLEX

GEOMETRICAL ISOMERISM IN OCTAHEDRAL COMPLEXES
• Octahedral complexes with formula MA4B2,
MA4BC, MA3B3, MA3BCD, MA2B2C2 , MABCDEF,
M(AA)2B2 , M(AA)(BB)C2, M(AA)(BB)(CC)
M(AA)(BB)CD etc. show geometrical isomerism
but MA6 and MA5B do not. Here, AA, BB etc. are
bidentate ligands.
• Here we will study the geometrical isomers of
some of the above octahedral complexes

For MA4B2, the B atoms can be cis or trans to each other.
GEOMETRICAL ISOMERS OF OCTAHEDRAL COMLEXES WITH FORMULA MA4B2

GEOMETRICAL ISOMERS OF OCTAHEDRAL COMLEXES WITH FORMULA MA3B3
• The geometric isomers of these complexes are not called as cis-trans
isomers but are called as Facial (FAC) and Meridional (MER) isomers.
• 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.

GEOMETRICAL ISOMERS OF OCTAHEDRAL COMLEXES WITH FORMULA M(AA)2B2
These complexes also show cis-trans isomerism

SUMMARY OF GEOMETRICAL ISOMERS IN OCTAHEDRAL COMPLEXES

• 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)
SREREOISOMERISM : 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. They don’t have plane of symmetry.
• Optical isomerism is common in
octahedral complexes involving
didentate ligands.

Some examples of optical isomers:
Optical isomers of octahedral compounds with formula M(AA)3

Optical isomers of octahedral complexes with formula M(AA)2B2

ANSWER THE FOLLOWING:
Trans has a plane of symmetry hence it will not give optical isomers.

STRUCTURAL ISOMERISM:- (1) LINKAGE ISOMERISM
• Linkage isomerism arises in a coordination compound containing
ambidentate ligand.
• When different atoms of the same ligand are boded or linked to metal,
then they give rise to linkage isomers as shown below:
• Other ligands capable of switching donor atom are SCN-, CN- etc.

STRUCTURAL ISOMERISM:- (2) COORDINATION ISOMERISM
• 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(NH3)6][Cr(CN)6], in which the NH3 ligands are
bound to Co3+ and the CN– ligands to Cr3+ .In its coordination isomer
[Cr(NH3)6][Co(CN)6], the NH3 ligands are bound to Cr3+ and the CN– ligands
to Co3+ .

• 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 following ionisation isomers
[Co(NH3)5 (SO4)]Br and [Co(NH3)5Br]SO4.
• These isomers give different ions in solutions.
• In the above example, the first one releases Br- ions and can give
precipitation of AgBr with AgNO3.
• The second one releases SO4
2- ions which can give precipitation of
BaSO4 with BaCl2 solution.
STRUCTURAL ISOMERISM:- (3) IONISATION ISOMERISM

STRUCTURAL ISOMERISM:- (4) SOLVATE ISOMERISM
• 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(H2O)6]Cl3 (violet) and its
solvate isomer [Cr(H2O)5Cl]Cl2.H2O (grey-green).
• Another example is given below:

EFFECTIVE ATOMIC NUMBER RULE
• Effective Atomic Number Rule is proposed by Sidgwick. The total number of electrons
possessed by central transition metal ion after the donation of electrons by the ligand is
called as an Effective Atomic Number(EAN)
• A complex is stable if the effective atomic number is equal to the atomic number of nearest
inert gas.
• Example: Calculate the effective atomic number of the following complexes:
(1) K4[Fe(CN)6]
(2) [Co(NH3)]Cl3
1. K4[Fe(CN)6]
Number of electrons in Fe2+ = 24
Number of electrons by Six CN = 2×6 = 12
Total number of electrons possessed by Fe2+ = 24 + 12
Therefore, the effective atomic number = 36.
2. [Co(NH3)]Cl3
Number of electrons in Co+3 = 24
Number of electrons by Six NH3 = 2×6 = 12
Total number of electrons possessed by Co+3 = 24 + 12
Therefore, the effective atomic number = 36.
BONDING IN COORDINATION COMPOUNDS: WHY ARE THEY FORMED

TABLE OF EAN OF SOME COMPLEXES

Werner’s Theory was the first to describe the bonding in
coordination compounds. But his theory was not satisfactory.
Drawbacks of Werner’s theory: This theory failed to 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?
BONDING IN COORDINATION COMPOUNDS

Many approaches have been put forth to explain the nature of
bonding in coordination compounds viz.
(i) Valence Bond Theory (VBT),
(ii)Crystal Field Theory (CFT),
(iii)Ligand Field Theory (LFT)
(iv)Molecular Orbital Theory (MOT).
Here, we shall focus mainly on VBT and CFT
OTHER THEORIES TO EXPLAIN BONDING IN COORDINATION COMPOUNDS

VALANCE BOND THEORY
This theory was given by Linus Pauling. 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 as shown in the following table:
• These hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron
pairs for bonding. Such a bond is called coordinate covalent bond.

• These complexes can be of two types, depending upon the nature of ligand:
(a) Low spin complexes:
 These are formed when ligand is strong field ligand like NH3, CN- etc.
 These ligands push the unpaired electrons of (n-1)d orbitals inside to be paired up. This
process is called spin pairing. Thus metal gets one or two (n-1)d orbitals vacant which
take part in hybridization.
 Due to pairing of electrons, such compounds are generally diamagnetic.
 They are also called as inner sphere or inner orbital or spin paired complexes.
 Examples are: dsp2 hybridization and d2sp3 hybridization.

(b) High spin complexes:
 These are formed when ligand is weak field ligand like F-, Cl- etc.
 These ligands can not push the unpaired electrons of (n-1)d orbitals to be paired up.
That’s why metal uses one or two nd vacant orbitals for hybridization.
 As pairing of (n-1)d electron could not be done, unpaired electrons are present and
hence these complexes are paramagnetic.
 They are also called as outer sphere or outer orbitals or spin free complexes.
 Examples are: sp3d and sp3d2 hybridization.
• Spin pairing is not done in metal ions with d1, d2 and d3 configuration in penultimate shell
as already vacant orbitals are available.
• The total number of metal orbitals undergone hybridization is equal to its coordination no.(CN).

SOME MORE EXAMPLES OF COMPLEXES EXPLAINED THROUGH VBT

TABLE OF SOME COMPLEXES WITH THEIR DETAILS

• The series of increasing order of field strength
of ligands is known as Spectrochemical series.
• It is given as below:
I– < Br– < SCN– < Cl– < S-2 < F– < OH– < C2O4
-2 <
H2O < NCS− < (EDTA)-4 < NH3 < en < CN- < CO
SPECTRO CHEMICAL SERIES

1. The complex in which central transition metal ion has unpaired electrons
is Paramagnetic.
2. The complex in which central transition metal ion has no unpaired electrons
is diamagnetic.
3. The magnetic moment of a complex is calculated by the spin only formula
M = √[n(n+2)] BM
BM = Bohr Magneton
The magnetic moment of complex compounds depends upon:
• Type of hybridization.
• The oxidation state of central transition metal ion.
• The number of unpaired electrons.
MAGNETIC PROPERTIES OF COORDINATION COMPOUNDS

MAGNETIC PROPERTIES OF COORDINATION COMPOUNDS:-
EXPERIMENTAL RESULTS AND THEIR EXPLANATION BY VBT
• The magnetic moment of coordination compounds can be measured by the magnetic
susceptibility experiments. The results are discussed below:
RESULT / OBSERVATIONS EXPLANATION BY VBT
For metal ions with up to three electrons in the d
orbitals, like Ti3+ (d1); V3+ (d2); Cr3+ (d3); the
magnetic behaviour of these free ions and their
coordination entities is similar.
two vacant d orbitals are available for octahedral
hybridisation with 4s and 4p orbitals. So no pairing
is required.
[Mn(CN)6]3– has magnetic moment of two unpaired
electrons while [MnCl6]3– has a paramagnetic
moment of four unpaired electrons. [Fe(CN)6]3– has
magnetic moment of a single unpaired electron
while [FeF6]3– has a paramagnetic moment of five
unpaired electrons. [CoF6]3– is paramagnetic with
four unpaired electrons while [Co(C2O4)3]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)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
sp3d2 hybridisation and are paramagnetic
corresponding to four, five and four unpaired
electrons.

While the VB theory, to a larger extent, explains the formation, structures and
magnetic behaviour of coordination compounds, it has 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.
DRAWBACKS OF VBT

The crystal Field
Theory was proposed
by Hans Bethe and
Van Vleck. This
theory gives much
more satisfactory
explanation for the
bonding and the
properties of complex
compounds than VBT.
CRYSTAL FIELD THEORY

1. The interaction between metal ion and the ligands are purely electrostatic. It
means metal-ligand bond is 100% ionic in nature.
2. Negative ligands are treated as point charges. While the neutral ligands are
treated as dipoles. Thus, the bonding in the complex may be an ion-ion
interaction or an ion-dipole interaction.
ASSUMPTIONS OF CRYSTAL FIELD THEORY

3. The five d-orbitals in an isolated gaseous metal atom or ion are degenerate.
Means they all have the same energy.

4. When ligands approach to metal ion to form a complex, the electrons in the
d- orbitals of the metal ion will be repelled by the negative charge or the lone
pair electrons of the ligands due to the repulsion between the like charges.
ligands
approach
to metal

5. As a result the energy of the d- orbitals increases and the d-orbitals no
longer remain degenerate. It results in the splitting of d- orbitals.
energy of the d-orbitals
increases
Ligands approaches,
repulsion occurs.

6. the pattern of splitting of the d-orbitals depends on the number of ligands
and their arrangement around the central metal atom or ion.

APPLICATION OF CFT IN OCTAHEDRAL COMPLEXES
1. In octahedral complexes, the metal ion is in the centre of the octahedral and
the six ligands lie at the six corners of the octahedral along the three axes x, y
and z.

2. The approach of the ligands towards the central metal ion is considered a
two step process:
(a) In the first step the ligands approach the metal spherically. i.e. at the same
distance from all the five d-orbitals. At this stage, the energy of all the d-
orbitals is raised by the same amount i.e. the five d- orbitals still remain
degenerate. At this stage whatever is the value of energy of d-orbitals is
considered zero and it is called Barry centre.

(b) However, as the ligands come closer the dx2-y2, and dz2 orbitals experience more
repulsion as they lie on axes and ligands are also approaching in the direction of
axes and hence their energy is raised up from Barry centre. The set of these two
orbitals with higher energy is called as eg orbitals. But rest of the three d- orbitals,
dxy, dyz and dzx , which lie in between the axes experience less repulsion and hence
their energy is lowered as compared to Barry centre. These three orbitals are
known as T2g orbitals. This is called Crystal Field Splitting.
(the names of these
orbitals come from
Mulliken symbols. T
stand for triple
degenerate orbital and e
stand for double
degenerate orbital.
Subscript g (origin:
German word gerade)
means – symmetric with
respect to inversion
center. Subscript 2 – anti-
symmetric with respect
to perpendicular C2 axis.)

3. The extent of splitting in energies is
denoted by symbol ∆o where O
represents Octahedral geometry.
4. Due to this splitting, the energy of
two eg orbitals will increase by 3/5th
of ∆o while that of three t2g orbitals
will decrease by 2/5th of ∆o.
5. The magnitude of crystal field
splitting depends upon the field
strength of the ligand and the charge
on the metal ion.
6. Strong field ligands cause large
splitting whereas weak field ligands
cause small splitting.

The first three electrons will occupy low
energy t2g orbitals according to Aufbau
principle and Hund’s rule. But the fourth
electron have two paths to follow:
(a) If the value of ∆o is lesser than pairing
energy(in case of splitting by weak
field ligands), then the fourth electron
will not be paired but will occupy one
of the eg orbitals. Thus we get a high
spin complex with weak field ligands.
(b) If the value of ∆o is greater than
pairing energy(in case of splitting by
strong field ligands), then the fourth
electron will be paired up in t2g and
will not move in the higher energy eg
orbitals. Thus we get a low spin
complex with strong field ligands.
Assigning electrons in the d-orbitals of the metal ion
in the octahedral coordination entities

It occurs in following steps :
1. Metal remains at the centre of tetrahedron.
2. Ligands approaches spherically towards
the metal and the energy of all the five d-
orbitals is raised equally due to repulsion
between ligand electrons and d- orbital
electrons of metal.
3. Crystal field splitting- In tetrahedral
symmetry, the ligands approach to central
metal not along the x, y, z axes but they lie
in between them. Therefore the d- orbitals
which lie in between the axes like dxy, dyz
and dzx experience more repulsion than dx2-
y2 and dz2. Thus dxy, dyz and dzx make higher
energy level t2g whereas the remaining two
make lower enery level called eg.
APPLICATION OF CFT IN TETRAHEDRAL COMPLEXES

6. Due to smaller value
of ∆t, pairing energy
is always greater
than ∆t. Therefore the
third electron choses
to go in higher t2g
orbitals instead of
being paired. Hence
low spin tetrahedral
complexes are rarely
observed.
4. Thus tetrahedral splitting is reversed of octahedral complexes.
5. The crystal field splitting energy(CFSE) in tetrahedral complexes is denoted
by ∆t and it is smaller than CFSE in octahedral complexes. ( ∆t = 4/9 ∆o).

• for example [Ti(H2O)6 ]+3 absorbs
wavelengths mostly in blue green
portion of visible light (𝝀= 498 nm).
The transmitted light does not have
blue green light. Therefore, the
remaining wavelengths together
generate a pinkish purple colour.
COLOURS IN COORDINATION COMPOUNDS
• Transition metal atoms and their ions with one or more unpaired electrons
and their complexes exhibit colours in both solid state and their solutions.
• How a substance gets its colour: When white light passes through a
substance, it absorbs some of the wavelengths. The light transmitted by the
substances is deprived of the wavelengths that have been absorbed. If
absorption occurs in the visible region of the spectrum, then the transmitted
light has a colour complementary to the colour absorbed. Complementary
colour is a colour generated from the wavelengths left after the absorption.

• The origin of colours of the coordination complexes can be very well explained by
Crystal Field Theory.
COLOURS IN COORDINATION COMPOUNDS: EXPLAINED BY CFT
Eabs = Eeg - Et2g
• When light is incident on the metal ion in
octahedral complexes the electrons in the
lower t2g orbitals absorb energy in the form
of some wavelengths and get promoted to
the higher eg sets of orbitals. The
complementary colour of the absorbed
colour becomes the colour of the complex.
Thus the colour of the coordination
compounds arises due to the transition of
electrons between the split (n-1)d orbitals.
Thus absorbed energy of light =
• We know that in complexes, the five degenerate (n-1)d- orbitals split into two
separate energy levels namely t2g and eg.

• 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(H2O)6]Cl3 on heating
renders it colourless. Similarly, anhydrous CuSO4 is white, but
CuSO4.5H2O is blue in colour.
Purple coloured Blue coloured
Colourless Colourless

• [Ni(H2O)6]2+ complex is formed when NiCl2 is dissolved in water. It is green in colour.
• If ethane-1,2-diamine(en) is progressively added in the molar ratios en:Ni = 1:1, 2:1,
3:1, the following series of complexes and their associated colour changes occur:
THE INFLUENCE OF THE LIGAND ON THE COLOUR OF A COMPLEX
No en added

COLOURS IN GEM STONES
(1) Ruby is aluminium oxide (Al2O3)
containing about 0.5-1% Cr3+ ions
(d3), which are randomly distributed
in positions normally occupied by
Al3+ .We may view these Cr3+ ions as
octahedral chromium(III) complexes
incorporated into the alumina
lattice; d–d transitions at these
centres give rise to the colour.
(2) In emerald Cr3+ ions occupy octahedral
sites in the mineral beryl (Be3Al2Si6O18).
The absorption bands here are yellow-
red and blue, causing emerald to
transmit light in the green region.

ADVANTAGES OF CFT
The crystal field model is successful in explaining the
following properties of coordination compounds:
(1) Formation of coordination compounds
(2) Structures of coordination compounds
(3) colour of coordination compounds
(4) magnetic properties of coordination compounds
ADVANTAGES OF CFT

DRAWBACKS OF CFT
There are mainly two drawbacks of CFT:
1. A negatively charged ligand should repel the d-electrons of
metal more strongly as compared to neutral ligands. We know
that greater the repulsion, larger will be the splitting of d- orbitals
and such a ligand will be strong field ligand. But actually
negatively charged ligands like F-, Cl- etc. are weak field ligands
whereas neutral ligands like NH3, CO etc. are strong field ligands.
This could not be explained by CFT.
2. In CFT, the metal-ligand bond is considered purely ionic but
actually they exhibit covalent characters also. This also could not
be explained by CFT.

BONDING IN METAL CARBONYLS
The followings are some example of complexes with CO group as homoleptic
ligand:

In a metal carbonyl, the metal-carbon bond possesses both σ and π character. The bond
between the carbonyl molecule and the metal is further strengthened by the synergic effect (π
back bonding) produced by the metal-ligand bond. These two types of bonding that exist in
metal carbonyls are explained below:
2. When this happens there is very good overlap between the metal’s other d orbitals (dxy, dxz,
dyz) with an empty pi* (antibonding) orbital on the carbonyl group. The overlap allows the
metal to donate some of it’s excess electron density into the empty pi* (antibonding)
carbonyl orbital. This forms a pi bond. The carbonyl is acting like a Lewis acid.
This synergistic effect strengthens the M-CO interaction. It also weakens the CO triple bond (as
electron density is being inserted into an antibonding orbital on CO).
Structure of Metal Carbonyls:
1. A carbonyl group donates its lone pair of
electrons to an empty metal d orbital. This is a
sigma bond and involves the lone pair in the
valence sigma orbital of the carbonyl and the
dx2-y2 or dz2 (or sometimes p) metal orbitals.
The carbonyl is acting like a Lewis base.

STABILITY OF COORDINATION COMPOUNDS
The stability of a complex in solution refers to the degree of association between the
two species involved in the state of equilibrium. The magnitude (stability or
formation) of the equilibrium constant for the association, quantitatively expresses
the stability. For a complex MLn , we can write:
The instability constant or the dissociation constant of coordination compounds
is defined as the reciprocal of the formation constant.

Free metal ions rarely exist in the solution so that M will usually be surrounded by
solvent molecules which will compete with the ligand molecules, L, and be
successively replaced by them. Thus we can write the formation of complex MLn in
n steps and can write stability constant for each step as shown below:
STEPWISE STABILITY CONSTANT

Since, water is in access, we can ignore the concentration of water and can write these
reactions in terms of only ligand and metal as shown below:

Taking log of both the side we get the following expression:
Example: We can calculate the overall stability constant of [Cu(NH3)4]2+ as follows:

APPLICATIONS OF COORDINATION COMPOUNDS
For the year 2020-21, the applications in analytical Chemistry, metallurgy and biological
systems are removed. So we will study only the applications in industry and in medicine.

APPLICATIONS OF COORDINATION COMPOUNDS IN INDUSTRY
•Coordination compounds are used as
catalysts for many industrial processes.
Examples include rhodium complex,
[(Ph3P)3RhCl], 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)2]– and [Au(CN)2]– 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(S2O3)2]3– .

APPLICATIONS OF COORDINATION COMPOUNDS IN MEDICINE
•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.

9. Coordination compounds class 12-1.pptx

9. Coordination compounds class 12-1.pptx

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