COORDINATION COMPOUNDS

By

D.GAYATHRI DEVI,PGT-CHEMISTRY,

JAWAHARNAVODAYA VIDYALAYA,N.R.PALLI.

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CLASSIFICATION OF SALTS

• A salt is formed by the neutralization of an acid by a base.

There are different types of salts. They are:-

a) Simple salt

b) Molecular (or) addition compounds

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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.

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MOLECULAR (OR) ADDITION COMPOUNDS �

• i)Double salts or lattice compounds
• ii) Coordination (or complex) compounds

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• i)Double salts or lattice compounds: These are molecular compounds which are formed by the evaporation of solution containing two (or) more salts in Stoichiometric proportions. The molecular compounds which dissociate in solution into its constituent ions are known as double salts. Double salts retain their properties only in solid state.
• Example : K₂SO₄ . Al₂(SO₄)₃ . 24H₂O - Potash alum
• FeSO₄ . (NH₄)₂ SO₄. 6H₂O - Mohr’s salt

K₂SO₄ . Al₂(SO₄)₃ . 24H₂O → 2K⁺ + 2Al³⁺ + 4SO₄^2- +24H₂O

• The double salts give the test of all their constituent
• ions in solution.

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• 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.
• Example : Fe(CN)₂ + 4KCN → Fe(CN)₂ . 4KCN (or) K₄[Fe(CN)₆] Ferrous cyanide

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.

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DOUBLE SALTS

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. 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 corresponding tests.

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COORDINATION COMPLEX

Co-ordinate complexes are incompletely ionizable in the aqueous solutions. These give a complexion 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]

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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.

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.

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���������������The hemoglobin is a coordination compound of iron.�

• In the polymerization of ethene, The Ziegler Natta catalyst (combination of triethyl aluminum and titanium tetrachloride) is used.

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• A complex metal catalyst is used in the hydrogenation of alkenes.

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• 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.�

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.

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� WERNER’S OBSERVATIONS�Many coordination compounds are brightly colored, but again, same metal, same ligands, different colors.�

WERNER’S THEORY

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• 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.

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��POSTULATES OF WERNER’S THEORY:�

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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.

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Difference between Primary and Secondary Valency in Coordination Compounds�

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

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IMPORTANT TERMS

INVOLVING

COORDINATION COMPOUNDS

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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-.

COORDINATION ENTITY�

CENTRAL ATOMS AND CENTRAL IONS

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• 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.

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CENTRAL ATOMS AND CENTRAL IONS

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).

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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

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� � 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�

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MONO/UNIDENTATE LIGANDS�

• The ligands which only have one atom that can bind to the coordination centre are called unidentate ligands. Ammonia (NH₃ ) is a great example of a unidentate ligand. Some common unidentate are Cl^–, H₂O etc.

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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

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ETHANE-1, 2-DIAMINE

OXALATE ION

OTHER BIDENTATE LIGANDS

POLYDENTATE LIGANDS

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• 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 EDTA^4–ion (ethylene diamine tetraacetate ion), which can bind to the coordination centre via its four oxygen atoms and two nitrogen atoms.

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EDTA^4– or ethylene diamine tetraacetate ion

CHELATE LIGANDS

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• 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.

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• 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.

AMBIDENTATE LIGAND

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• 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.�

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AMBIDENTATE LIGAND

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.

CALCULATION OF COORDINATION NUMBER

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• IN CASE OF MONODENTATE LIGANDS,
• Coordination number = number of ligands
• IN POLYDENTATE LIGANDS.
• Coordination number = number of ligands * denticity

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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

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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.

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COORDINATION POLYHEDRON�

tetrahedral square planar Octahedral

OXIDATION NUMBER

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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.

HOMOLEPTIC AND HETEROLEPTIC COMPLEX�

Homoleptic and Heteroleptic Complex

When the coordination centre is bound to only one type of electron pair donating ligand group, the coordination complex is called a homoleptic complex, for example: [Cu(CN)₄]^3-.

When the central atom is bound to many different types of ligands, the coordination compound in question is called a heteroleptic complex, an example for which is [Co(NH₃)₄Cl₂]⁺.

Properties of Coordination Compounds

The general properties of coordination compounds are discussed in this subsection.

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.

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.

When the coordination centre is bound to only one type of electron pair donating ligand group, the coordination complex is called a homoleptic complex, for example: [Cu(CN)₄]^3-.

When the central atom is bound to many different types of ligands, the coordination compound in question is called a heteroleptic complex, an example for which is [Co(NH₃)₄Cl₂]⁺.

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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.

• 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.
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DOUBLE SALTS

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. 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 corresponding tests.

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COORDINATION COMPLEX

Co-ordinate complexes are incompletely ionizable in the aqueous solutions. These give a complexion 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]

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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(CH)6]
• 3.Neutral Complexes: In this co-ordination sphere is neither cation or anion. Example: [Ni(CO)4]

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TYPES OF COORDINATION COMPLEXES�based on whether complex ion is a cation/anion

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• 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:

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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

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IUPAC

NOMENCLATURE

OF

COORDINATION COMPOUNDS

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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.

Rules For Naming Coordination Compound�

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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.

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).

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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).

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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.

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��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 ���

IUPAC NAMES OF SOME COORDINATION COMPOUNDS�

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• [Co (NH₃)₄(H₂O) ₂] Cl₃ = Tetraamminediaquacobalt(III) chloride
• [Cr(en)₃] Cl₃ = Tris(ethane-1,2-diamine)chromium(III) chloride
• [Pt(NH₃)BrCl(NO₂)]^-= Amminebromidochloridonitrito-N-platinate(II)
• [PtCl₂(en)₂](NO₃)₂ = Dichloridobis(ethane-1,2-diamine)platinum(IV) nitrate
• (NH₄)₃[Cr(SCN)₆] = Ammonium hexathiocyanato-S-chromate(III)
• Na₂[Cr(CH₃COO)₄(en)] = Sodium ethylenediaminetetraacetatochromate(II)
• [Co(NH₃)₅(CO₃)]Cl = Pentaamminecarbonatocobalt(III) chloride
•  [Pt(py)₄][PtCl₄] = Tetrapyridineplatinum(II)tetrachloridoplatinate(II)
• (NH₄)₃[Cr(SCN)₆] = Ammonium hexathiocyanato-S-chromate(III)
• Na₂[Cr(CH₃COO)₄(en)] = Sodium ethylenediaminetetraacetatochromate(II)
• [Co(NH₃)₅(CO₃)]Cl = Pentaamminecarbonatocobalt(III) chloride
•  [Pt(py)₄][PtCl₄] = Tetrapyridineplatinum(II)tetrachloridoplatinate(II)

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[Co(NH₃)₄Cl₂]₃[Cr(CN)₆] = Tetraamminedichloridocobalt(III)hexacyanochromate(III)

 Na₂[Fe(CN)₅NO] = Sodium pentacyanonitrosoniumferrate(II)

K₃[Co(CN)₅NO] = Potassium pentacyanonitrosylcobaltate(II)

 Na₂[CrF₄O] = Sodium tetrafluoridooxochromate(IV)

 [Cr(H₂O)₄Cl₂]NO₃ = Tetraaquadichloridochromium(III) nitrate

(NH₄)₃[Cr(SCN)₆] = Ammonium hexathiocyanato-S-chromate(III)

Na₂[Cr(CH₃COO)₄(en)] =

Sodium ethylenediaminetetraacetatochromate(II)

[Co(NH₃)₅(CO₃)]Cl = Pentaamminecarbonatocobalt(III) chloride

 [Pt(py)₄][PtCl₄] = Tetrapyridineplatinum(II)tetrachloridoplatinate(II)

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��� 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.

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• 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.

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• 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).

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ISOMERISM

IN

COORDINATION COMPOUNDS�

[Co(H2O)4Cl2]+

ISOMERISM IN COORDINATION COMPOUNDS�

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• 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.

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STRUCTURAL ISOMERISM

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• 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

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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)

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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].

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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.

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IONISATION ISOMERISM�

�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

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

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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.

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.

GEOMETRIC OR CIS-TRANS ISOMERS

Example of MA2B2 � complex�

• ML4 tetrahedral complexes do not show cis-trans isomerism since ligands are in different directions.
• MABCD has 3 geometrical isomers. 2-cis and 1-trans.
• MA2B2 complex shows cis and trans isomers.

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FACIAL AND MERIDIONAL ISOMERISM� ( fac- and mer-isomers) � by Ma₃b₃ Type of Complexes

• 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. 

[Co(NH₃)₃ (NO₂)₃] [CoCl₃(CN)₃]

Optical isomers

• The optical isomers or enantiomers, are mirror images of each other and two enantiomers cannot be superimposed on each other

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Chemistry of

Coordination

Compounds

OPTICAL ISOMERISM

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-)

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Laevo(l-) dextro(d-)

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.

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.

The following table gives the coordination number, orbital hybridisation and geometry �

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
• μ_s = spin-only magnetic moment , n=number of unpaired electrons

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[Co(NH₃)₆] ³⁺

Hybridisation: d²sp ³,Shape: octahedral, Diamagnetic,Low spin compex, Innerorbital complex

Hybridisation: sp ³d²,Shape: octahedral, Paramagnetic.High spin/Outer compex orbital complex

Hybridisation: sp ³,Shape: tetrahedral, Paramagnetic,Low spin compex, Innerorbital complex

[Ni(Cl₄] ^2-

Hybridisation: dsp ²,Shape: Square planar, Diamagnetic,Low spin compex, Innerorbital complex

[Ni(CN₄] ^2-

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Hybridisation: d²sp ³,Shape: octahedral, Paramagnetic

Low spin compex, Innerorbital complex

Hybridisation: d²sp ³,Shape: octahedral, diamagnetic

Low spin compex/ Innerorbital complex

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 strong ligands.

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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 ligands

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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 _o for octahedral and _t for tetrahedral
• The magnitude of _o and _t depends upon the field strength of ligand around the metal

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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.

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• 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

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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

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As a result, a set of d-orbitals split into two sets: e_g orbitals of higher energy including d(x² –y²) and d(z²) and t_2g orbitals of lower energy including d(xy), d(yx) and d(xz)

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The crystal field splitting is measured in terms of energy difference between t_2g and e_g orbital and is denoted by a symbol _o . It is generally measured in terms of D_q. It is called as crystal field splitting energy or crystal field stabilization energy E_g orbitals are 6Dq above the average energy level and t_2g orbitals are 4Dq below the average energy level �

The energy of e_g set of orbitals > energy of t_2g set of orbitals.

Ligands for which energy separation, Δ_o < 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, Δ_o > P, form low spin complex.

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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.

• For the same metal, the same ligands and metal-ligand distances, it can
• be shown that Dt = (4/9) D0.
• Consequently, the orbital splitting energies are not sufficiently large for forcing pairing and therefore, low spin configurations are rarely

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• The energy of t_2g set of orbitals > Energy of e_g 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.

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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.

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THE FACTORS AFFECTING THE COLOUR OF COMPLEXES

• Number of unpaired electrons in transition metal 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)4en]⁺² 
• Green Pale blue

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• 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.

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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).

• 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

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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

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OVERALL STABILITY CONSTANT(β)

• M + 4L 🡪 ML₄
• The overall formation constant or stability constant, β = K₁ × K₂ × K₃ × K₄
• INSTABILITY CONSTANT
• Instability constant= 1/β �

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.

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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.

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• Some complex compounds containing cyanide as a ligand are used in the process of electroplating. These compounds are also very useful in photography.
• 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.

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APPLICATIONS IN BIOLOGY�

• Haemoglobin consists of Haeme complex-ion which has tetrapyrrole Porphyrin ring structure with central Fe²⁺ ion.
• Vitamin B12 consists of tetrapyrrole porphyrin ring complex with central Co⁺³ ion and its coordination number is 6.

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APPLICATIONS IN LABORATORY�

• Ni⁺² is estimated using a complexing agent Dimethylglyoxime (DMG). The hardness of water is estimated using complexes of Ca⁺⁺, Mg⁺⁺ with ED↑A
• In Medicine: Cisplatin is used in the treatment of cancer.
• In Photography: Developing of the film involves complex formation.
• In Metallurgy: In the extraction of gold, silver by Mac Arthur Forest Process involves a complex of cyanide ions.

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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 – (ox.state of the metal) + 2(coordination number of the metal)].
• for example.�[Co(NH3)6]3+ →EAN = [27 – 3 + 2(6)] = 36

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