Coordination Compounds Notes

Ruhul Amin Alig

A quick revision of all the important concepts

Werner's Theory of Coordination Compounds

Coordination compounds are compounds linked to a the central metal atom or ion is linked by ions or molecules with coordinate bonds. e.g., Potassium ferrocyanide, $K_{4} [F e (C N)_{6}], [N i (C O)_{4}]$ etc.
- These may be a positively charged, negatively charged or a neutral species.
The chemistry of coordination compounds is an important and challenging area of modern inorganic chemistry.
Chlorophyll, haemoglobin and vitamin $B_{12}$ are coordination compounds of magnesium, iron and cobalt respectively.
In 1898, Alfred Werner, a Swiss chemist was the first to formulate his ideas about the structures of coordination compounds. The main postulates of this theory are:
- Metals possess two types of valencies i.e. primary (ionizable) valency and secondary (non-ionizable) valency.
- Secondary valency of a metal is equal to the number of ligands attached to it i.e. coordination number.
- Primary valencies are satisfied by negative ions, while secondary valencies may be satisfied by neutral, negative or positive ions.
- The ions/groups bound by the secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordination numbers.

For example:In

[C r (H_{2} O)_{6}] C l_{3}

, the primary valency is 3 and the secondary valency is 6.

Counter ions: The species within the square bracket are coordination entities or complexes and the ions outside the square bracket are called counter ions.
The spatial arrangement of the ligands which are directly attached to the central atom or ion, is called coordination polyhedron around the central atom or ion.
Double salt are the addition molecular compounds which are stable in solid state but dissociate into constituent ions in the solution and each ion in the solution gives the corresponding confirmatory test.

For example: Mohr's salt,

F e S O_{4} . (N H_{4})_{2} S O_{4} . 6 H_{2} O

gets dissociated into

F e^{2 +}

N H_{4}^{+}

and

S O_{4}^{2 -}

ions.

Difference between a double salt and a complex ions

Double salts dissociate into simple ions completely when dissolved in water whereas complex ions do not dissociate into ions.
In double salt, the metal atom or ion exhibits the normal valency whereas in complexes, the negative ions or molecules surrounding the central metal atom are beyond its normal valency.

Definitions of Some Important Terms Pertaining to Coordination Compounds Terminology

1. Coordination entity:

A coordination entity comprises of a central metal atom or ion bonded to a fixed number of ions or molecules.
It is written within square brackets.
For example: In $[C o (N H_{3})_{6}] . C l_{3}$ , coordination entity is $[C o (N H_{3})_{6}]$ .

2. Central atom/ion:

The atom or ion to which a fixed number of ions or groups are bound in a definite geometrical arrangement around is called as central atom or ion.
It is also referred to as Lewis acid.
For example: In $[N i C l_{2} (H_{2} O)_{4}]$ , $N i^{2 +}$ is the central atom.

3. Ligands:

The ions or molecules bound to the central atom/ion in the coordination entity are called ligands.
They may be charged or neutral.
Denticity refers to the number of donor groups in a single ligand that bind to a central atom in a coordination complex.
Ligands are of the following types :
- Unidentate- It is a ligand, which has one donor site, i.e., the ligand bound to a metal ion through a single donor site. e.g., $H_{2} O, N H_{3}$ , etc.
- Didentate- When a ligand can bind through two donor atoms. Eg, $H_{2} N C H_{2} C H_{2} N H_{2}$ (ethane-1,2-diamine) or $C_{2} O_{4}^{2 -}$ (oxalate).
- Polydentate- When several donor atoms are present in a single ligand. For eg, Ethylenediamine tetraacetate ion (EDTA $^{4 -}$ ) is a hexadentate ligand. It can bind through two nitrogen and four oxygen atoms to a central metal ion.
- Chelate ligand- When a di- or polydentate ligand uses its two or more donor atoms simultaneously to bind a single metal ion, it is said to be a chelate ligand.
  - More the number of chelate rings, more is the stability of complex.
- Ambidentate ligand- These are the ligands which can ligate through two different sites.
  - Examples of such ligands are the $N O_{2}^{-}$ and $S C N^{-}$ ions.

4. Coordination number:

The coordination number (CN) of a metal ion in a complex can be defined as the number of ligand donor atoms to which the metal is directly bonded.
In case of monodentate ligands,Coordination number = number of ligands

In polydentate ligands. Coordination number = number of ligands $\times$ denticity

For example, in the complex ions, $[P t C l_{6}]^{2 -}$ and $[N i (N H_{3})_{4}]^{2 +}$ , the coordination number of Pt and Ni are 6 and 4 respectively.
In the complex ions, $[F e (C_{2} O_{4})_{3}]^{3 -}$ and $[C o (e n)_{3}]^{3 +}$ , the coordination number of both, Fe and Co, is 6 because $C_{2} O_{4}^{2 -}$ and $e n (e t h a n e - 1, 2 - d i a m i n e)$ are didentate ligands.
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.

5. 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.
The ionisable group written outside the bracket is known as counter ions.
For example, in the complex $K_{4} [F e (C N)_{6}]$ , the coordination sphere is $[F e (C N)_{6}]^{4 -}$ . The counter ion is $K^{+}$ .

6. Coordination polyhedron:

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.

7. Oxidation number of central atom:

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.
The oxidation number is represented by a Roman numeral in parenthesis following the name of the coordination entity.
$[C u (C N_{4})]^{3 -}$ , oxidation number of copper is +1, and represented as Cu(I).

8. Homoleptic and heteroleptic complexes:

Complexes in which a metal is bound to only one kind of donor groups in called homoleptic complexes. For example: $[C o (N H_{3})_{6}]^{3 +}$ .
Complexes in which a metal is bound to more than one kind of donor groups are known as heteroleptic complexes. For example: $[C o (N H_{3})_{4} C l_{2}]^{+}$ .

Nomenclature of Coordination Compounds

The formulas and names adopted for coordination entities are based on the recommendations of the International Union of Pure and Applied Chemistry (IUPAC).

Formulas of Coordination Entities

Mononuclear coordination entities contain a single central metal atom.
The formula of a compound is a shorthand tool used to provide basic information about the constitution of the compound in a concise and convenient manner.
Some rules are applied while writing the formulae:
- Name of the compound is written in two parts: (i) name of cation, and (ii) name of anion.
- The ligands are always written before the central metal ion.
- 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.
- 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- etc. When there are many polydentate ligands attached to the central metal ion, the prefixes are of the form bis-, tris-, and so on.
- The names of the anions present in a coordination compound must end with the letter 'o', which generally replaces the letter 'e'. Therefore, the sulfate anion must be written as 'sulfato' and the chloride anion must be written as 'chlorido'.
- The following neutral ligands are assigned specific names in coordination compounds: $N H_{3}$ (ammine), $H_{2} O$ (aqua or aquo), $C O$ (carbonyl), $N O$ (nitrosyl).

- 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.
- While writing the name of the central metal atom in an anionic complex, priority is given to the Latin name of the metal if it exists (with the exception of mercury).
- 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.
- If the coordination compound is accompanied by a counter ion, the cationic entity must be written before the anionic entity.
For example: $[C r (N H_{3})_{3} (H_{2} O)_{3}] C l_{3}$ is named as triamminetriaquachromium(III) chloride
- The complex ion is inside the square bracket, which is a cation. The amine ligands are named before the aqua ligands according to alphabetical order.
- There are three chloride ions in the compound, the charge on the complex ion must be +3 (since the compound is electrically neutral) - $[C r (N H_{3})_{3} (H_{2} O)_{3}]^{3 +}$ .
- From the charge on the complex ion and the charge on the ligands, we can calculate the oxidation number of the metal. Since all the ligands are neutral molecules, therefore, the oxidation number of chromium is the same as the charge of the complex ion, +3.
$[N i C l_{4}]^{2 -}$ is named as tetrachloridonickelate(II) ion.
- The complex ion, an anion, is inside the parentheses. We have to add the suffix $- a t e$ in the name of the metal.

Isomerism in Coordination Compounds

Isomers are two or more compounds that have the same chemical formula but differ in the arrangement of atoms. They differ in one or more physical or chemical properties.
Coordination compounds exhibit two principal types of isomerism:

Stereo Isomerism
Structural Isomerism

Stereo Isomerism

Stereoisomers have the same chemical formula and chemical bonds but they have different spatial arrangement.
They are classified into two types:
1. Geometric Isomerism: Geometric isomerism arises in heteroleptic complexes due to different possible geometric arrangements of the ligands. This type of isomerism is mainly found in coordination compounds with coordination numbers 4 and 6. Some examples are:
  1. In a square planar complex, (i.e. coordination compounds with coordination number 4) which has $[M A_{2} B_{2}]$ type formula (A and B are unidentate ligands), the two ligands may be present adjacent to each other in a cis isomer, or opposite to each other to form a trans isomer. For example: $P t [(N H_{3})_{2} C l_{2}]$

ii. Square planar complex with MABCD type formula (where A, B, C, D are unidentate ligands) show three isomers-two cis and one trans.

iii. Octahedral complexes, with formula

[M A_{2} B_{4}]

, can show cis-trans geometry.

This type of isomerism also arises when didentate ligands L - L are present in complexes of formula

[M A_{2} (L - L)_{2}]

. For example:

[C o C l_{2} (e n)_{2}]

iv. Another type of geometrical isomerism occurs in octahedral coordination entities of the type

[M A_{3} B_{3}]

. It shows two types of isomers:

1.Facial (fac) isomer: If three donor atoms of the same ligands occupy adjacent positions at the corners of an octahedral face, it is a facial isomer.

2.Meridional (mer) isomer: When the positions are around the meridian of the octahedron, we get the meridional (mer) isomer.

b. Optical Isomerism: The complexes which are non-superimposable on their mirror images are optically active. Enantiomers exist when the molecules of the substances are mirror images but are not superimposable upon one another.

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

Levorotatory (l) - the compound which rotates plane polarised light to left hand side.
Dextrorotatory (d) - the compound which rotates plane polarised light to right hand side.

Optical isomerism is common in octahedral complexes involving didentate ligands.
In a coordination entity of the type $[P t C l_{2} (e n)_{2}]^{2 +}$ , only the cis-isomer shows optical activity.

Structural Isomerism

The isomers having same molecular formula but different structural arrangement of atoms or groups of atoms around the central metal ion are called structural isomers & phenomenon is called structural isomerism. It can be classified as:
1. Linkage Isomerism: Linkage isomerism arises in a coordination compound containing ambidentate ligand.
  - The best known cases involve the ligands $S C N^{-} / N C S^{-}$ and $N O_{2}^{-} / O N O^{-}$ .

- Jorgensen discovered such behaviour in the complex $[C o (N H_{3})_{5} (N O_{2})] C l_{2}$ , which is obtained as the red form, in which the nitrite ligand is bound through oxygen $(- O N O)$ , and as the yellow form, in which the nitrite ligand is bound through nitrogen $(- N O_{2}$ ).
- Example: $[C o (N H_{3})_{5} (O N O)] C l$ , the nitrito isomer is connected through O, $[C o (N H_{3})_{5} (N O_{2})] C l$ , the nitro isomer is connected through N

2. Coordination Isomerism: Isomerism arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex.

- Example: One isomer $[C o (N H_{3})_{6}] [C r (C_{2} O_{4})_{3}]$ and another isomer $[C o (C_{2} O_{4})_{3}] [C r (N H_{3})_{6}]$ .

3. Ionisation Isomerism: This kind of isomerism arises when the counter ion is a potential ligand(exchange of anions between the coordination sphere and ionization sphere.)

- Example: $[C o (N H_{3})_{5} (S O_{4})] B r$ and $[C o (N H_{3})_{5} B r] S O_{4}$ .

4. Solvate Isomerism: This form of isomerism is known as 'hydrate isomerism' in case where water is involved as a solvent. Hydrate isomers are the type of isomers which have similar composition but differ in the presence of number of water molecules as ligands.

- Example: $[C r C l_{2} (H_{2} O)_{4}] C l . 2 H_{2} O$ bright-green, $[C r C l (H_{2} O)_{5}] C l_{2} . H_{2} O$ grey-green, $[C r (H_{2} O)_{6}] C l_{3}$ violet

Bonding in Coordination Compounds

Bond formation in coordination compounds can be explained by Valence Bond Theory (VBT), Crystal Field Theory (CFT), Ligand Field Theory (LFT) and Molecular Orbital Theory (MOT).

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.
The hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron pairs for bonding.
The valence bond theory is used to predict the geometry of a complex by its magnetic behaviour.

Inner Orbital Complex: When the inner d orbitals take part in hybridization, it is called an inner orbital complex. It is also known as low spin or spin paired complex.

In the diamagnetic octahedral complex, $[C o (N H_{3})_{6}]^{3 +}$ , the cobalt ion is in $+ 3$ oxidation state and has the electronic configuration $3 d^{6}$ . The hybridisation is shown below:
Six pairs of electrons, one from each $N H_{3}$ molecule, occupy the six hybrid orbitals. Thus, the complex has octahedral geometry and is diamagnetic because of the absence of unpaired electron.

Outer orbital complex: When the outer d orbitals take part in hybridization, it is called an outer orbital complex.

The octahedral complex, $[C o F_{6}]^{3 -}$ uses outer orbital (4d ) in hybridisation ( $s p^{3} d^{2}$ ). Complex has unpaired electrons, therefore, it will be paramagnetic in nature.

In tetrahedral complexes, one s and three p orbitals are hybridised to form four equivalent orbitals oriented tetrahedrally.

In $[N i C l_{4}]^{2 -}$ , nickel is in +2 oxidation state and the ion has the electronic configuration $3 d^{8}$ .
The compound is paramagnetic since it contains two unpaired electrons.

In square planar complexes, the hybridisation involved is

d s p^{2}

. In

[N i (C N)_{4}]^{2 -}

, nickel is in +2 oxidation state and has the electronic configuration

3 d^{8}

The compound is diamagnetic due to the absence of unpaired electron.

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.

Limitations of Valence Bond Theory

It does not explain the colour exhibited by coordination compounds.
It involves a number of assumptions.
It does not give quantitative interpretation of magnetic data.
It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.
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.

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, $d_{x y}$ , $d_{y z}$ and $d_{z x}$ that are oriented in between the coordinate axes and are called $t_{2 g}$ -orbitals.
- Two d-orbitals, $d_{x^{2} - y^{2}}$ and $d_{z^{2}}$ that are oriented along the x - y axes and are called $e_{g}$ - orbitals.
The pattern of splitting depends upon the nature of the crystal field.

(a) Crystal field splitting in octahedral coordination entities

In an octahedral coordination entity with six ligands surrounding the metal atom/ion, there will be repulsion between the electrons in metal d orbitals and the electrons (or negative charges) of the ligands.
Thus, $d_{x^{2} - y^{2}}$ and $d_{z^{2}}$ orbitals which point towards the axes along the direction of the ligand will experience more repulsion and will be raised in energy; and the dxy, dyz and dxz orbitals which are directed between the axes will be lowered in energy relative to the average energy in the spherical crystal field.
$e_{g}$ orbital are of higher energy than $t_{2 g}$ orbital.
The splitting of degenerate levels due to the presence of ligands in a definite geometry is termed as crystal field splitting and the energy separation is denoted by $Δ_{o}$ (the subscript o is for octahedral).

An experimentally determined series based on the absorption of light by complexes with different ligands is termed as spectrochemical series. The arrangement of ligands in a series in the order of increasing field strength as given below:

C O > C N^{-} > N O_{2}^{-} > e n > N H_{3} > H_{2} O > O H^{-} > F^{-} > C l^{-} > B r^{-} > I^{-}

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 and are known as weak field ligands.
Ligands for which energy separation or $Δ_{o} > P$ , form low spin complex and are known as strong field ligands.

(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 $Δ_{t} = \frac{4}{9} Δ_{0}$ .
Here, $e_{g}$ orbitals are of lower energy than $t_{2 g}$ orbitals.
No pairing of electrons is possible due to the low splitting energies leading to high spin complexes.

Colour in Coordination Compounds

The colour in the coordination compounds can be readily explained in terms of the crystal field theory.
The theory attributes the colour of the coordination compounds to d-d transition of the electron, i.e., the transition of electron from $t_{2 g}$ level to the higher $e_{g}$ level which accompanies the absorption of light in visible spectrum.
In the absence of ligand, crystal field splitting does not occur and hence the substance is colourless. Example: $[T i (H_{2} O)_{6}] C l$ .

Transition of an electron in $[T i (H_{2} O)_{6}]^{3 +}$

$[T i (H_{2} O) 6]^{3 +}$ is an octahedral complex which is violet in colour. The single electron in the metal $(3 d^{1})$ d orbital is in the $t_{2 g}$ 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 $t_{2 g}$ level to the $e_{g}$ level.

Let us consider $[N i (H_{2} O)_{6}]^{2 +}$ complex, which forms when nickel(II) chloride is dissolved in water. If the didentate ligand, ethane-1,2-diamine(en) is progressively added in the molar ratios en:Ni, 1:1, 2:1, 3:1, change in color is observed.

Limitations of Crystal Field Theory

From the assumption that ligands are point charges, it follows that anionic ligands should exert greatest splitting effect. But anionic ligands actually are found at the low end of the spectrochemical series.
It does not take into account for the covalent character of bonding between the ligand and the central atom.

Bonding in Metal Carbonyls

The homoleptic complexes in which carbonyl group (C=O) acts as the ligand are called metal carbonyls. For example: $N i (C O)_{4}$ . These carbonyls have simple, well defined structures.
Tetracarbonylnickel(0) is tetrahedral, pentacarbonyl iron(0) is trigonal bipyramidal while hexacarbonyl chromium(0) is octahedral.
Decacarbonyldimanganese(0) is made up of two square pyramidal $M n (C O)_{5}$ units joined by a Mn - Mn bond.
Octacarbonyldicobalt(0) has a Co - Co bond bridged by two CO groups.

The metal-carbon bond in metal carbonyls possess both $σ$ and $π$ character. 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 metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal.

Importance and Applications of Coordination Compounds

Coordination compounds play a vital role in analytical chemistry, metallurgy, biological systems, industry and medicine. These are described below:

Coordination compounds have many applications in qualitative as well as quantitative chemical analysis. The familiar colour reactions given by metal ions with a number of ligands (especially chelating ligands), as a result of formation of coordination entities, form the basis for their detection and estimation by classical and instrumental methods of analysis.

The hardness of water is estimated by titration with the sodium salt of EDTA. During titration, the calcium and magnesium ions in hard water form the stable complexes, Calcium EDTA and Magnesium EDTA. These ions can be selectively estimated due to the difference in the stability constants of calcium and magnesium EDTA complexes.

Metals can be purified by the formation and subsequent decomposition of their coordination compounds
It is used as catalysts for many industrial processes. Example: rhodium complex, $[(P h_{3} P)_{3} R h C l]$ , a Wilkinson catalyst, is used for the hydrogenation of alkenes.

Solutions of the complexes like $[A g (C N)_{2}]^{-}$ and $[A u (C N)_{2}]^{-}$ can be used for the smooth and even electroplating of metals by gold or silver.

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Coordination Compounds Notes