CBSE Class 12 Chemistry –Chapter 9 Coordination Compounds- Study Materials

CBSE Class 12 Chemistry: Chapter 9 – Coordination Compounds. These notes will give you a quick glance of the chapter. These notes are prepared strictly based on the latest CBSE syllabus for CBSE Class 12th Chemistry.

The main topics covered in these quick notes are:

•    Definition of

                o Coordination compounds o Double salt

                o Central atom or ion

•    Ligand and its types

•    Chelate, Chelating ligands and Chelation

•    Coordination number and Oxidation number

•    Nomenclature of coordination compounds

•    Isomerism in coordination compounds and its types

•    Werner’s theory of coordination compounds

•    Valence bond theory (VBT)

                o Principles involved

                o Limitations

•    Crystal field splitting theory (CFT)

                o Crystal field splitting in octahedral complexes

                o Crystal field splitting in tetrahedral complexes

                o Stability of coordination compounds

                o Colour in Coordination Compounds

                o Limitations of crystal field theory

•    Metal carbonyls

                o Structure of some important metal carbonyls

                o Bonding in metal carbonyls

                o Properties of metal carbonyls

•    Applications of coordination compounds

Double salt

When two salts in stoichiometric ratio are crystallised together from their saturated solution they are called double salts. They are stable in solid state but dissociate into constituent  ions when dissolved in water.
For example: FeSO₄. (NH₄) ₂SO₄.6H₂O (Mohr’s salt)

Central atom or ion

The atom or ion to which a fixed number of ions or groups are bound in a definite geometrical arrangement around it, is called the central atom or ion.
For example: In K₄[Fe(CN)₆], Fe₂ is the central metal ion.

Ligands

The atoms or groups which are attached directly to central atoms are called ligands. Ligands are Lewis bases which donates electron pair and forms coordinate bonds with the metal atom.

For example: H₂O, CO, NO₂^‒, etc.

A ligand may be neutral, positively or negatively charged.

Types of ligands

On the basis of charge present, ligands may be categorized into the following three groups:

Anionic ligands: CN^‒ (cyanide), NO₂^‒ (nitrito-N), NO₃^‒ (nitrito), X^‒ (halido), etc.

Cationic ligands: NO₂⁺ (nitrosonium), NO⁺ (nitronium), N₂F₅⁺ (Hydrazenium), etc.

Neutral ligands: NH₃ (ammine), H₂O (aqua), CO (carbonyl), etc.

On the basis of number of coordinating atoms present ligands may be classified as follows:

Monodentate: It is the ligand which has only one coordinating atom.

For example: X^‒ (halido), NH₃ (ammine), H₂O (aqua), CN^‒ (Cynaido)  .

Diddentate: It is the ligand which has two coordinating sites available.

For example: H₂NCH₂CH₂NH₂ (ethane-1, 2-diamine) and C₂O₄^2‒ (oxalate).

Polydentate: Ligand containing several donor sites are called polydentate.

For example: N(CH₂CH₂NH₂)₃ (Nitrilotriethylamine, a tetradentate ligand )

EDTA (Ethylene Diamine tetraetato ion): It is a hexadentate ligand.

Ambidentate ligand: This is the ligand which can ligate through two different atoms present in it.

For examples: NO²⁻ ion can coordinate either through nitrogen or through oxygen to a central metal atom/ion.

Chelate, Chelating ligands and Chelation

Di- and ploydentqte ligands results in the formation of cyclic structure around the central metal atom. Such cyclic metal complex is called chelate and the ligands which gives chelates are called chelating ligands, and the process is known as chelation.

Denticity

Denticity is defined as the number of coordinating atoms present per ligand.

Coordination number

It is defined as the number of coordinate bonds formed by central metal atom, with the ligands. For example: In K₄[Fe(CN)₆], Fe has coordination number 6.

Oxidation number

The oxidation number of the central atom is defined as the charge left on a given atom when all other atom in a complex are removed as ions. The oxidation number is indicated by Roman numerals.

Nomenclature of coordination compounds

Step to name a complex compound are:

Cations are always named before the anions.

The ligands are then listed in alphabetical order.

In case of polydentate ligands, ligands are named alphabetically using a prefix di, tri, tetra, penta etc, to indicate the number of ligands of that type present.

Then the name the metal atom is written followed by its oxidation state in Roman numerals.

Finallythe anion is named.

For example: [Cu(NH₃)₄]SO₄  Tetra ammine copper (II) sulphate

If the ligands itself include di, tri etc. then we use bis – (for two), tris (for three) as prefix. For example:

[PtCl₂(en)₂]²⁺  Dichloridobis (1, 2 – ethanediamine) platinum (IV) ion.

[Co(en)₃]³⁺ Tris(ethanediamine)cobalt (III) ion.

If the complex ion is an anion, the name of the metal ends with the suffix – ate.

For example: [Cr(C₂O₄)₃]^3‒ Trioxalatochromate (III) ion.

Isomerism in coordination compounds

Isomers are two or more compounds that have the same chemical formula but a different arrangement of atoms.

Types of Isomerism

There are two types of isomerism:

(a) Stereoisomerism

(i) Geometrical isomerism (ii) Optical isomerism

(b) Structural isomerism

(i) Linkage isomerism       (ii) Coordination isomerism

(iii) Ionisation isomerism  (iv) Solvate isomerism Structural isomerism

Stereoisomerism

It arises due to different arrangement of atoms or groups in space. Two different types of stereoisomerism are described as follows:

(i) Geometrical isomerisms:

It arises in heteroleptic complexes due to different possible geometrical arrangements of ligands.  In square planar complex of formula [MX₂L₂] (where 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.

For example: Geometrical isomers of Pt(NH₃)₂Cl₂) are shown below:

Similarly in octahedral complexes of formula [MX₂L₄], two ligands X may be oriented cis or trans to each other.

For example: Geometrical isomers of [Co(NH₃)₄Cl₂] are shown below:

(ii) Optical isomerism:

It arises due to the presence of non-super imposable mirror images. The non-superimposable mirror images are called enantiomers.

The enantiomers reacts differently with the plane polarised light.

The enantiomer which rotate the plane polarised light in a clockwise direction is called dextrorotatry (d) or (+) and the enantiomer which rotate the plane polarised light in anticlockwise direction is called lavorotatory (l) or (‒).

For example: Optical isomers of [Co(en)₃ ] ³⁺ are shown below:

Structural isomerism

It arises due to the difference in structures of coordination compounds. Four different types of structural isomerism are described as follows:

Ionisation isomerism: These are the isomers having the same molecular formula but gives different ions in solution.

For example: [Cr(NH₃)Br]SO₄ and [Cr(NH₃)(SO₄)]Br are ionization isomers.

Linkage isomerism: Isomers having the same molecular formula but different linking atoms are named as linkage isomers. This arises due to the presence of ambident ligands.

For example:[CO(NH₃)₅(NO₂)]²⁺ and [CO(NH₃)₅ (ONO)]²⁺ are linkage isomers.

Coordination isomerism: This type of isomerism is possible when cation and anion both are complex and isomerism arises due to complete exchange of coordination sphere.

For example: [Pt (NH₃)₄] [Ni (CN)₄] and [Ni (NH₃)₄] [Pt (CN)₄] are coordination isomers.

Hydration isomerism: Isomers having the same molecular formula but different water molecules of hydration are known as hydration isomers of each other.

For example: [Cr (H₂O)₅Cl]Cl₂.H₂O and [Cr(H₂O)₄Cl₂]Cl.2H₂O are hydration isomers.

Werner’s theory of coordination compounds

According to Werner’s theory of coordination compounds, there are two types of valencies in coordination compounds:

  o   Primary valencies: These are ionizable valencies, satisfied by anions and determines the charge on the complex ions.

  o   Secondary valencies: These are non-ionisable valencies, satisfied by ligands and determines the coordination number of the metal atom.

Valence Bond Theory

Valence bond theory states that 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. These hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron pairs for bonding.

Following table shows the combination of different number of orbitals to give different types of hybridization:

Thus the basic principles involved in the valence bond theory are:

•    Hybridisation of orbitals

•    Bonding between the ligands and the metal atoms/ion

•    Relation between the observed magnetic moment and the bond type

Limitation of the VB theory

•    It does not tell anything about the spectral properties of the complexes.

•    It does not give quantitative interpretation of magnetic data.

•    It does not distinguish between strong and weak ligands.

•    It does not explain the colour exhibited by coordination compounds.

•    It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.

Magnetic properties of coordination compounds

A coordination compound is paramagnetic in nature if it has unpaired electrons and diamagnetic if all the electrons in the coordination compound are paired.

Crystal Field Theory

In crystal field theory (CFT), ligands are considered as point charges and the interaction between the ligands and the metal ion is purely electrostatic in nature. The five d-orbitals in an isolated gaseous metal atom/ion have same energy, i.e., they are degenerate. The degeneracy is lost in the presence of the ligand field.

The five d-orbitals are classified as:

Three d-orbitals, d_xy, d_yz and d_zx that are oriented in between the coordinate axes and are called t_2g –orbitals.

Two d-orbitals, dx² ‒ y² and dz² that are oriented along the x – y axes and are called e_g – orbitals.

Factors affecting the splitting of d-orbitals

•    Nature of the ligand

•    Nature of the metal ions

•    Geometry of complex whether it is octahedral or tetrahedral

•    Oxidation state of the metal ion

Crystal field splitting in octahedral complexes

Here 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.

Crystal field splitting in tetrahedral complexes

Here 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.

Stability of coordination compounds

The stability of the coordination compound depends on

Nature of the ligand

Chelating ligands form strong and more stable complexes than the monodentate ligands. The π–  bond ligands forms more stable complexes than the σ– bonded complex.

Nature of the metal atom/ion

Small, highly charged metal ions form more stable complexes than large size, lowly charged metal ion.

Colour in Coordination Compounds

The crystal field theory attributes the colour of the coordination compounds to d-d transition of the electron, i.e., the transiton of electron from t_2g level to the higher e_g level which accompanies the absorption of light in visible spectrum.
In the absence of ligands, crystal field splitting does not occur and hence the substance is colourless.

Limitations of crystal field theory

It does not take into account the partly covalent character of bonding between the ligand and the central atom.

It is also unable to explain the relative strengths of ligands e.g., it does not explain why H₂O is stronger ligand than OH⁻.

Metal carbonyls

The homoleptic complexes in which carbon monoxide (CO) acts as the ligand are called metal carbonyls.
For example: Ni(CO)₄

Structure of some important metal carbonyls are:

Bonding in metal carbonyls

•   The metal-carbon bond in metal carbonyls possess both s and p character. CO as a ligand binds itself to metal atoms through the carbon atom to form the metal-carbon (M-C) bond. It is a weak donor.

•   The M–C σ bond is formed by the donation of lone pair of electrons on the carbonyl carbon into a vacant orbital of the metal.

•   The M–C π bond is formed by the donation of a pair of electrons from a filled d orbital of metal into the vacant antibonding π* orbital of carbon monoxide. This characteristic property of back bonding which stabilises the metalligand interaction is termed as synergic effect.

Properties of metal carbonyls

•   They are generally solids at room temperature and pressure except Ni(CO)₄ and Fe (CO)₅.

•   Mononuclear carbonyls are volatile and toxic.

•   Mononuclear carbonyls are either colourless or light coloured

Applications of coordination compounds

•   EDTA is used in the estimation of Ca²⁺ and Mg²⁺ in hardwater.  The Ca²⁺ and Mg²⁺ ions form stable complexes with EDTA.

•   Metals can be purified  by the formation and subsequent decomposition of their coordination compounds. For example, impure nickel is converted to [Ni(CO)₄], which is decomposed to yield pure nickel.

•    In analytical chemistry [Ni(DMG)₂]²⁺ complex is used in the detection of Ni in chocolates.

•    In medicine, cisplatin, a cis isomer of [Pt(Cl)₂(NH₃)₂] is used in the treatment of cancer.

•     Solutions of the complexes like [Ag(CN)₂]^– and [Au(CN)₂]^–  can be used for the smooth and even electroplating of metals by gold or silver.

•     Chlorophyll, a pigment responsible for photosynthesis, is a coordination compound of magnesium. Also haemoglobin, the red pigment of blood which acts as oxygen carrier is a coordination compound of iron.