Coordination Compounds Notes CBSE Class 12

Coordination Compounds Terminology

Coordination or Complex Compound

The compound which retain their identity in solid state as well as in dissolved state is called coordination compound. In these compounds. the central metal atom or ion is linked by ions or molecules with coordinate bonds. e.g., Potassium ferrocyanide, K₄[Fe(CN)₆].

Coordination Entity

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.
[CoCl₃(NH₃)₃] and [Fe(CN)₆]^-4 are coordination entities.

Coordination Sphere

The central ion and the ligands attached to it are enclosed in square bracket which is known as coordination sphere. The ionisable group written outside the bracket is known as counter ions.

Central Metal Ions

The central atoms and anions are atoms and ions attached to ligands. In coordination complexes, the central atoms or ions are Lewis acids and can therefore appear as electron-pair acceptors. In [Ni(CO)₄], Ni is central metal atom. It is generally transition element or inner-transition element.

Coordination Polyhedron

The geometric shape created by the attachment of the ligands to the coordination centre is known as the coordination polyhedron. Such kinds of spatial arrangements in coordination compounds include tetrahedral and square planar shapes.
For example, [Co(NH₃)₆]⁺³ is octahedral, [Ni(CO)₄] is tetrahedral and [PtCl₄]^–2 is square planar.

Homoleptic and Heteroleptic Complexes

When the central atom in complex compound is attached to only one type of ligand, the complex compound is known as homoleptic complex.
Example: [Cu(CN)₄]^-3.
When the central atom in complex compound is attached to different types of ligands, the complex compound is known as heteroleptic complex.
Example: [Co(NH₃)₄Cl₂]⁺.

Ligands

It is a donar atom, molecule or ion that donate a pair of electron to the central atom to form complex compound. It may be negative, neutral or even positive. It is of different types such as monodentate, didentate, tridentate and polydentate etc.
Unidentate ligands: Ligands with only one donor atom, e.g. NH₃, Cl^-, F^- etc.
Bidentate ligands: Ligands with two donor atoms, e.g. ethylenediamine, C₂O₄^-2(oxalate ion) etc.
Tridentate ligands: Ligands which have three donor atoms per ligand, e.g. (dien) diethyl triamine.
Hexadentate ligands: Ligands which have six donor atoms per ligand, e.g. EDTA.
Ambidentate ligands: These are the monodentate ligands which can ligate through two different sites, e.g., NO^-2, SCN^-, etc.
Chelating ligands: Di or polydentate ligands cause cyclisation around the metal atom which are known as chelate. Such ligands uses two or more donor atoms to bind a single metal ion and are known as chelating ligands. More number of chelate rings means more stable the complex is.

Denticity

The number of donor groups in a single ligand that bind to a central atom in a coordination complex.

Which of the following is more stable complex and why?
[CO(NH₃)₆]³⁺ and [Co(en)₃]³⁺

Solution:
[Co(en)₃]³⁺ is more stable since ‘en’ is didentate ligand which forms more stable complex than NH₃(unidentate ligand).

Oxidation Number

The charge of the complex if all the ligands are removed along with the electron pairs that are shared with the central atom, is called oxidation number of central atom.
Oxidation number of copper is +1 in [CU(CN)₄^-3 complex and is represented as Cu(I).

Coordination Number

The coordination number in coordination compounds is defined as the number of ligand (donor) atoms/ions surrounding the central metal atom in a complex ion.
For example, the coordination number of cobalt is six in [Co(NH₃)₆]⁺³ complex ion.

Werner's Theory and Limitations

Werner's theory introduced the concept of primary and secondary valencies for metal ions in coordination compounds:

Primary Valency: This corresponds to the oxidation state of the metal ion. It is satisfied by negative ions and is ionizable.
Secondary Valency: This corresponds to the coordination number of the metal ion. It is satisfied by neutral molecules or negative ions (ligands) and is non-ionizable.

Read more in details Werner's Theory and It's Limitations

Which of the following statements is not a postulate of Werner's theory?

A. Metals possess two types of valencies: primary and secondary.
B. Primary valencies are ionizable, while secondary valencies are non-ionizable.
C. Secondary valencies are directed in space, leading to specific geometries.
D. The nature of the bonds formed by secondary valencies is always ionic

Solution:
Option D is correct answer
Werner's theory states that metals have two types of valencies: primary (ionizable) and secondary (non-ionizable), with secondary valencies being directional in space, leading to specific geometries, but not necessarily always ionic in nature.

In the complex [Co(NH₃)₅Cl]Cl₂, the primary and secondary valencies of cobalt are respectively:

A. 3 and 5
B. 5 and 3
C. 3 and 6
D. 6 and 3

Solution:
Option C is correct answer
Cobalt has an oxidation state of +3.
So, primary valency = 3
It is coordinated to five NH₃ molecules and one Cl⁻ ion.
So, secondary valency = 6.

Valence Bond Theory: Assumptions, Merits and Demerits

Valence Bond Theory (VBT) is a fundamental theory developed by Linus Pauling, describes how the central metal atom or ion in a coordination compound forms bonds with surrounding ligands through the hybridization and overlap of atomic orbitals. This theory also helps in predicting the geometry and magnetic properties of complexes. However, it has some limitations given below, which were resolved by Crystal Field Theory.

Limitations of VBT:

It cannot explain the colour of complexes.
It does not provide a quantitative interpretation of magnetic properties.
It cannot explain the relative stabilities of complexes.
It does not predict the exact geometries of some complexes.

Read more in details Valence Bond Theory and Limitations

Crystal Field Theory: Assumptions and Limitations

Crystal field theory was given by Bethe and van Vlack; is a model of the electronic structure of transition-metal complexes that considers how the energies of the d orbitals of a metal ion are affected by the electric field of the ligands. According to this theory, the ligands in a transition-metal complex are treated as point charges. So a ligand anion becomes simply a point of negative charge. A neutral ligand, with its electron pair that it donates to the metal atom, is replaced by a partial negative charge, representing the negative end of the molecular dipole. In an electric field of these negative charges, the five d orbtails of the metal atom no longer have exactly the same energy. The result, as you will see, explains both the paramagnetism and the color observed in certain complexes.

Read more in details Crystal Field Theory: Assumptions and Limitations

IUPAC Nomenclature of Inorganic Compounds

IUPAC Nomenclature of Inorganic Compounds pdf

IUPAC Nomenclature Examples

[CO(NH₃)₆]Cl₃
hexaamminecobalt (III) chloride

[CO(NH₃)₅Cl]Cl₂
pentaamminechloridocobalt (III) chloride

K₃[Fe(CN)₆]
potassium hexacyanoferrate (III)

[K₃[Fe(C₂O₄)₃]
potassium trioxalatoferrate (III)

K₂[PdCl₄]
potassium tetrachloridoplatinum (II)

[Pt(NH₃)₂ClNH₂CH₃]Cl
diamminechlorido (methylamine) platinum(II) chloride

[CO(NH₃)₄(H₂O)₂]Cl₃
Tetraamminediaquacobalt(IlI) chloride

K₂[Ni(CN)₄]
Potassium tetracyanidonickelate(II)

[Cr(en)₃]Cl₃
Tris(ethane-1,2-diamine) chromium(III) chloride

[Pt (NH ₃) Br Cl (N0 ₃)]^–
Amminebromidochloridonitrito-N- platinatc(II) ion

[PtCl₂(en)₂](N0₃)₂
Dichloridobis(ethane-l ,2-diamine) platinum (IV) nitrate

Fe₄[Fe(CN)₆]₃
Iron(III)hexacyanidoferrate(II)

[Zn(OH)₄]^2-
tetrahydroxozincate(II) ion

[Pt(NH₃)₆]⁴⁺
hexaammineplatinum (IV) ion

K₂[PdCl₄]
potassiumtetrachloridopalladate(II)

[Cu(Br)₄]^2-
tetrabromidocuprate (II)

[CO(NH₃)₆]₂ (SO₄)₃
hexaaminecobalt(III) sulphate

K₂[Ni(CN)₄]
potassiumtetracyanonicklate (II)

K₃[Cr(OX)₃]
potassiumtrioxalatochromate(III)

[CO(NH₃)₅ONO]²⁺
pentaamminenitrito-O-cobalt(III)

[Pt(NH₃)₂Cl₂]
diamminedichloridoplatinum(II)

[CO(NH₃)₅NO₂]²⁺
pentaamminenitrito-N-cobalt (III)

[Pt(NH₃)₂CI (NH₂CH₃)]Cl
Diammine chlorido (methylamine) platinum (II) chloride

[Ti(H₂O)₆]³⁺
Hexaaqua titanium(III)ion

[Ni(CO)₄]
Tetra carbonyl nickel (0)

|Mn(H₂O)₆]²⁺
Hexaaquamanganese (II) ion

Fe₃(Fe(CN)₆)₂
Iron(III) hexacyanoferrate(II)

The hypothetical complex chlorodiaquatriammine cobalt (III) chloride can be represented as

A. [CoCl(NH₃)₃(H₂O)₂ ]Cl₂
B. [Co(NH₃)₃(H₂O)Cl₃]
C. [Co(NH₂)₃(H₂O)₂Cl]
D. [Co(NH₃)₃(H₂O)₃]Cl₃

Solution:
Chlorodiaquatriamminecobalt(Ill) chloride can be represented as
[CoCl(NH₃)₃(H₂O)₂ ]Cl₂

Effective Atomic Number (EAN) Rule

EAN rule is a theory given by Sidgwick which gives an idea about the stability of the coordination compound that forms. EAN represents the total number of electrons surrounding the nucleus of a metal atom in a complex compound or coordination sphere. The compound which follows EAN rule is compreratively more stable. This rule is also called 18 –electron rule.

Generally, the EAN of the central metal is numerically equal to the atomic number of the noble gas element found in the same period in which the central metal atom is located.

Effective Atomic Number Formula

EAN = (Z – X) + (C.N × 2)

EAN = (Z – X) + (L × D × 2)

Z is the atomic number of the central metal ion.
X is the oxidation number of the central metal ion.
L is the total number of ligands bound to the central metal atom.
D is the denticity of the ligand.

Example-
EAN of Co in [Co(NH₃)₆]Cl₃ complex is 36.
Co is in +3 oxidation state. So it has 24 electrons.
6 NH₃ donates 12 electrons (2 electrons by each NH₃)
So, total electrons around the nucleus of a Co atom in [Co(NH₃)₆]Cl₃ complex is-
EAN = 24 + 12 = 36.

EAN Calculator

EAN: 0

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Complexes which do not obey effective atomic number rule

There are many exceptions to the effective atomic number rule. Many stable complexes are known in which the effective atomic number rule is not obeyed.

EAN for [Ni(CN)₄]^-2 complex is 34
EAN for [Ni(NH₃)₆]⁺² complex is 38
EAN for [Cr(NH₃)₆]⁺³ complex is 33
EAN for [CoCl₄]^-2 complex is 33
EAN for [Cu(NH₃)₄]⁺² complex is 35
EAN for [Fe(CN)₆]^-3 complex is 35

Significance

The EAN rule is used to predict the stability of coordination compounds, the oxidizing and reducing character of carbonyl compounds etc.

 

Q. Which of the following complexes do not follow EAN rule

A. [Fe(CN)₆]^2-
B. [Fe(CN)₆]^3-
C. Ni(CO)₄
D. [Co(NH₃)₆]³⁺

Solution:
Option B is correct answer
EAN = (26 - 3) + (2 x 6) = 35

Q. What is the EAN for [Mo(CO)₆] complex

A. 34
B. 36
C. 54     ✔
D. none of these

Solution:
Option C is correct answer
EAN = 42 + 12 = 54

Inner Orbital Complexes

If the complex is formed by the use of inner d-orbitals for hybridisation (written as dⁿsp³), it is called inner orbital complex. In the formation of inner orbital complex, the electrons of the metal are forced to pair up and hence the complex will be either diamagnetic or will have lesser number of unpaired electrons. Such a complex is also called low spin complex.
For example, [Fe(CN)₆]^-3 and [Co(NH₃)₆]⁺³ are inner orbital complexes.
[Co(NH₃)₆]⁺³

Outer Orbital Complexes

If the complex is formed by the use of outer d-orbitals for hybridisation (written as sp³dⁿ), it is called an outer orbital complex. The outer orbital complex will have larger number of unpaired electrons since the configuration of the metal ion remains undisturbed. Such a complex is also called high spin complex.
For example, [Fe(H₂O)₆]⁺³, [CoF₆]^{- 3} and [Co(NH₃)₆]⁺² are outer orbital complexes.
[Co(NH₃)₆]⁺²

Which one of the following is an outer orbital complex and exhibits paramagnetic behaviour?

A. [Ni(NH₃)₆]²⁺
B. [Zn(NH₃)₆]²⁺
C. [Cr(NH₃)₆]³⁺
D. [Co(NH₃)₆]³⁺

Solution:
[Ni(NH₃)₆]²⁺ is right answer.
Ni is in +2 state so, it is a 3d⁸ syste.
Ammonia is a strong field ligand and can cause pairing of electrons. However, after pairation, only one 3d orbital left vacant for ligand's electron. Since the number of ligands are six; so, we consider 4s, 4p and 4d orbitals for hybridization and it will be sp³d².
It has two unpaired electrons in 3d orbital so it is paramagnetic and its magnetic moment is 2.828BM.

High Spin and Low Spin Complexes

High spin and low spin complexes can be determine by using crystal field theory and ligand field theory. Generally, octahedral complexes and tetrahedral complexes are high spin, while square planar complexes are low spin.

High Spin Complexes: High spin complexes are formed when the crystal field splitting energy (Δ) is smaller than the pairing energy of electrons, due to this, pairation does not occur resulting in a higher number of unpaired electrons and paramagnetic character.

Weak field ligands (e.g., Cl⁻, F⁻) creates high spin complexes.
Example: [FeCl₆]^3⁻

Low Spin Complexes: Low spin complexes are formed when the crystal field splitting energy (Δ) is greater than the pairing energy of electrons, due to this, pairation occur resulting in a lower number of unpaired electrons and diamagnetic character.

Strong field ligands (e.g., CN⁻, CO) creates low spin complexes.
Example: [Fe(CN)₆]^4⁻

Remember:
☛ If Δ > P, the complex will be low-spin
☛ If P > Δ, the complex will be high-spin

Colour of Coordination Compounds

The colour of complex compounds is due to absorption of light in visible region of sepctrum and radiation of complementary colour. The energy is absorbed by electrons present in d-orbitals and they get excited to higher energy d-orbitals from lower energy d-orbitals. They radiate energy, when they come back to lower energy d-orbitals. The color of coordinate compounds is also due to charge transfer transitions. It involves the transfer of an electron from ligand to metal (LMCT) or from metal to ligand (MLCT). These transitions are generally very intense colors.

The colour in the coordination compounds can be readily explained in terms of the crystal field theory. Let us consider the complex [Ti(H₂O)₆]³⁺, which is violet in colour. This is an octahedral complex where the single electron (Ti³⁺ is a 3d¹ system) in the metal d orbital is in the t_2g level in the ground state of the complex. The next higher state available for the electron is the empty e_g level. If light corresponding to the energy of blue-green region is absorbed by the complex, it would excite the electron from t_2g level to the e_g level (t_2g¹ e_g⁰ → t_2g⁰ e_g¹). Consequently, the complex appears violet in colour.

[Ti(H₂O)₆]³⁺

The color of complex compounds due to charge transfer transitions is intense because they are both spin-allowed and Laporte-allowed. This means that there are no restrictions on these transitions based on the spin or orbital symmetry of the electrons involved. This allows for a high probability of the transition occurring, leading to strong light absorption and intense colors.

Charge transfer transitions are involved the movement of an electron from metal to ligand or ligand to metal by following the selection rules.

When a transition is spin-allowed as well as Laporte-allowed, a large number of photons can be absorbed during the transition, leading to a strong absorption of light.

The deep purple color of KMnO₄ is due to the charge transfer from the ligand (O₂^-) to the metal (Mn⁺⁷) center. This is ligand to metal charge transfer (LMCT).

The pale yellow color of [Fe(CN)₆]^4- is due to transfer of an electron from the iron to the π* orbitals of the cyanide ligands. This is metal to ligand charge transfer (MLCT).

Factors Affecting the Colour of Complex Compounds

Several factors influence the color of complex compounds, some of which are discussed below-

Nature of Ligands: Strong field ligands, such as CN^⁻ and CO, cause large splitting of d-orbitals, which can result in absorption of higher energy (shorter wavelength) light and thus different colors.

Nature of Metal Ion: Different metal ions have different d-orbital structures and energies, affecting the absorption and thus the observed color.

Oxidation State of Metal Ion: Higher the oxidation state of the metal ion, greater the splitting of d-orbitals, leading to different absorptions and different colors.

Geometry of Complex: The shape of the coordination compound (e.g., octahedral, tetrahedral) affects the energy levels of d-orbitals, which affects the color.

Magnetic Properties of Coordination Compounds

Magnetic properties of coordination compounds depend on the electronic structure and the presence of unpaired electrons in the central metal atom. Coordination compounds may be paramagnetic or diamagnetic depending on the number of unpaired electrons.

Paramagnetic Compounds: Compounds that have unpaired electrons in the d-orbitals of the central metal atom exhibit paramagnetism and are called paramagnetic compounds. The presence of unpaired electrons generates a magnetic moment, which causes the compound to be attracted to a magnetic field. The magnetic moment can be measured by the magnetic susceptibility experiments. The degree of paramagnetism is directly related to the number of unpaired electrons. Larger the number of unpaired electrons, greater the paramagnetic behaviour.
[Fe(CO)₅] is paramagnetic because the iron atom has four unpaired electrons in its d-orbitals.

Diamagnetic Compounds: Compounds that contain no unpaired electrons in the d-orbitals of the central metal atom exhibit diamagnetism and are known as diamagnetic compounds. Since there are no unpaired electrons, these compounds do not possess a net magnetic moment and are weakly repelled by a magnetic field.
[Ni(CO)₄] is diamagnetic due to the absence of unpaired electrons in the d-orbitals of Ni atom.

Paramgnetic and diamagnetic can be expained by Valence bond theory and calculated by the formula:

μ = n(n + 2)^1/2 BM or μ = 4s(s + 1)^1/2 BM

where, n is the number of unpaired electrons, s is Spin magnetic moment and BM is Bohr Magneton, unit of magnetic moment.

[Ti(H₂O)₆]³⁺. Ti is in 3+ oxidation state, so it has one unpaired electron and it is paramagnetic and its magnetic moment is 1.732BM

Magnetic Moment Calculator

Magnetic Moment: 0 BM

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Factors Affecting the Magnetic Properties

Oxidation State of the Metal: Higher oxidation state of the metal have fewer unpaired electrons and might lead to a more diamagnetic or weakly paramagnetic behavior.

Ligands: The nature of the ligands can affects the splitting of the metal's d-orbitals.
Strong field ligands (like CN^⁻, CO) can cause pairing of electrons, leading to a diamagnetic or low magnetic moment complex.
Weak field ligands (like Cl^⁻, H₂O) cause less orbital splitting, which may result in more unpaired electrons and a higher magnetic moment.

Why do compounds having similar geometry have a different magnetic moment?

Solution:
Compounds with similar geometry can have different magnetic moments because the number of unpaired electrons in the central metal ion can vary depending on the strength of ligands attached. A strong field ligand will cause pairing of electrons while a weak field ligand will not cause pairing. Pairing or not pairing will change the number of unpaired electrons, which affects the magnetic moment.

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

Solution:
Since the coordination number of Mn²⁺ ion in the complex ion is 4, it will be either tetrahedral (sp³ hybridisation) or square planar (dsp² hybridisation). But the fact that the magnetic moment of the complex ion is 5.9 BM, it should be tetrahedral in shape rather than square planar because of the presence of five unpaired electrons in the d orbitals.

Spectrochemical Series

Spectrochemical Series Ligands can be arranged in increasing order of their strength (ability to cause crystal field splitting) and this series is called spectrochemical series. It is as follows:
I^– < Br^– < S^2– < SCN^– < Cl^– < F^– < OH^– < Ox^2– < O^2– < H₂O < NCS^– < py < NH₃ < en < NO₂^– < CN^– < CO

Bonding in Metal Carbonyls

Metal carbonyls are homoleptic complexes in which carbon monoxide (CO) acts as the ligand. Bonds in Metal carbonyl are a type of chemical bond in which metal atoms are bonded to carbon monoxide (CO) ligands. These bonds exhibit unique properties due to the synergic interaction between the metal and the CO ligand. Ludwig Mond created the first metal carbonyl compound, Ni(CO)₄, in 1884.

The carbon monoxide ligand donates electron density from its carbon atom's lone pair into an empty orbital of the metal atom. This forms a σ-bond between the carbon atom of CO and the metal atom.

The metal atom donates electron density from its filled d-orbitals into the empty π* (anti-bonding) orbitals of the CO ligand. This is known as π-backbonding or π-acceptor interaction.

The combination of these two interactions—σ-donation from the ligand to the metal and π-back bonding from the metal to the ligand forms a synergistic bond that significantly increases the overall stability of the complex.

Let us explain with example of [Fe(CO)₅]
In iron pentacarbonyl, the iron atom forms five synergic bonds with five CO ligands.
The structure is trigonal bipyramidal, where three CO ligands are equatorial and two are axial.

What is the cause of back bonding in metal carbonyl?

Solution:
Back bonding in metal carbonyls is caused by the interaction between the filled d orbitals of the transition metal and the empty π* antibonding orbitals of the CO ligand. This results in the formation of a π bond, which strengthens the metal-carbon bond and weakens the carbon-oxygen bond.

Isomerism

Isomers are two or more compounds that have the same chemical formula but a different arrangement of atoms. Because of the different arrangement of atoms, they differ in one or more physical or chemical properties. Two principal types of isomerism are known among coordination compounds. Each of which can be further subdivided.

1. Stereoisomerism
A. Geometrical isomerism
B. Optical isomerism

2. Structural isomerism
A, Linkage isomerism
B. Coordination isomerism
C. Ionisation isomerism
D. Solvate isomerism

Stereoisomers have the same chemical formula and chemical bonds but they have different spatial arrangement. Structural isomers have different bonds.

Geometrical isomerism

This type of isomerism arises in heteroleptic complexes due to different possible geometric arrangements of the ligands. Important examples of this behaviour are found with coordination numbers 4 and 6. In a square planar complex of formula [MX₂L₂] (X and L are unidentate), the two ligands X may be arranged adjacent to each other in a cis isomer, or opposite to each other in a trans isomer as shown in figure.

Other square planar complex of the type MABXL (where A, B, X, L are unidentates) shows three isomers-two cis and one trans. Such isomerism is not possible for a tetrahedral geometry but similar behaviour is possible in octahedral complexes of formula [MX₂L₄] in which the two ligands X may be oriented cis or trans to each other.

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

Another type of geometrical isomerism occurs in octahedral coordination entities of the type [Ma₃b₃] like [Co(NH₃)₃ (NO₂)₃].

If three donor atoms of the same ligands occupy adjacent positions at the corners of an octahedral face, we have the facial (fac) isomer. When the positions are around the meridian of the octahedron, we get the meridional (mer) isomer

Why tetrahedral complexes don't exhibit geometrical isomerism?

Solution:
Tetrahedral complexes do not exhibit geometrical isomerism because in a tetrahedral arrangement, all four ligands are positioned symmetrically around the central metal ion, that means their relative positions are always identical which prevent the formation of distinct geometric isomers, no matter how you rotate the molecule, the ligands will always appear in the same relative positions to each other with same bond angle.

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

Optical isomerism is common in octahedral complexes involving didentate ligands. In a coordination entity of the type [PtCl₂(en)₂]⁺², only the cis-isomer shows optical activity

Linkage Isomerism

Linkage isomerism arises in a coordination compound containing ambidentate ligand. A simple example is provided by complexes containing the thiocyanate ligand, NCS^–, which may bind through the nitrogen to give M–NCS or through sulphur to give M–SCN. Jørgensen discovered such behaviour in the complex [Co(NH₃)₅ (NO₂)]Cl₂, which is obtained as the red form, in which the nitrite ligand is bound through oxygen (–ONO), and as the yellow form, in which the nitrite ligand is bound through nitrogen (–NO₂).

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 Co⁺³ and the CN^– ligands to Cr⁺³. In its coordination isomer [Cr(NH₃)₆][Co(CN)₆], the NH₃ ligands are bound to Cr⁺³ and the CN^– ligands to Co⁺³.

Ionization Isomerism

This form of isomerism arises when the counter ion in a complex salt is itself a potential ligand and can displace a ligand which can then become the counter ion. An example is provided by the ionisation isomers [Co(NH₃)₅(SO₄)]Br and [Co(NH₃)₅Br]SO₄.

Solvation 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).

Can in a complex compound water molecules behave in two ways? Explain in brief.

Solution:
Yes, in a complex compound, water molecules can behave in two ways: as a neutral ligand donating a lone pair of electrons to a central metal ion, forming a coordinate bond (example: [Cr(H₂O)₆]³⁺), and as a solvent molecule interacting with other ions or molecules in the solution through hydrogen bonding. Water acting as both a coordinating molecule and a surrounding solvent molecule depending on the situation within the complex compound.

Applications of Coordination Compounds

Coordination compounds are found in living systems and have many uses in the home, in industry and in medicines. A few examples are given below:

Extraction of metals

Cyanide ions are used for the for the extraction of gold and silver. The crushed ore is heated with an aq. cyanide solution in the presence of air to dissolve the gold by forming the soluble complex ion [Au(CN)₂]^–.

4Au(s) + 8CN^–(aq) + O₂(g) + 2H₂O(l) → 4[Au(CN)₂]^–(aq) + 4 OH^–(aq)
Zn(s) + 2[Au(CN)₂]^–(aq) → [Zn(CN)₄]^–2 (aq) + 2Au(s)

Complex formation is also useful for the purification of metals. Nickel is purified by converting the metal to the gaseous compound Ni(CO)₄ and then decomposing the latter to pure nickel.

Medicines

EDTA is a chelating agent which is used in the treatment of lead poisoning. Cis platin cis [Pt(NH₃)₂Cl₂] is used in the treatment of cancer. Sodium nitroprusside, Na2[Fe(CN)₅NO] is used to lower blood pressure during surgery.

Qualitative Analyses

Complex formation is useful for qualitative analyses.
1. Separation of Ag⁺ from Pb⁺² & Hg⁺²
Ag⁺ + 2NH₃(aq.) → [Ag(NH₃)₂]⁺(Soluble)
2. Separation of IIA and IIB groups: The cations of IIB group form soluble complex with yellow ammonium sulphide.
3. Cu⁺² ion forms complex on addition of ammonia [Cu(NH₃)₄]⁺²
4. Fe⁺² forms a blue complex with K₃Fe(CN)₆, i.e. K Fe^II[Fe^III(CN)₆].
5. Cobalt(II) gives color with HCl due to the formation of complex [CoCl₄]^–2.
6. Nickel forms a red complex [Ni(DMG)₂] with dimethylglyoxime (H₂DMG).

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