Werner’s Theory & Some Basic Concepts of Coordination Compounds | Chemistry Class 12 - NEET PDF Download

Table of Contents
1. What is Werner's Theory?
2. Werner's Experiments and Historical Notes
3. Postulates of Werner's Theory
4. Isomerism in Coordination Compounds
5. Limitations of Werner's Theory
6. Basic Concepts and Terms Related to Coordination Compounds
7. Representation of Coordination Complexes
8. Oxidation State Examples and Calculations
9. Effective Atomic Number (EAN) - Sidgwick Rule
10. Summary
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What is Werner's Theory?

Werner's Theory, proposed by Swiss chemist Alfred Werner around 1898, established the first clear model for the structure of coordination compounds and laid the basis of modern coordination chemistry. Werner distinguished between two kinds of valences for a metal centre and explained the fixed spatial arrangement of ligands around a metal ion, accounting for many observed isomers and properties of complexes.

The main ideas introduced by Werner that are essential to understanding coordination compounds are listed below.

Dotted line represents the primary valency. Normal line represents the secondary valency.

• Primary and secondary valence: The primary valence corresponds to the oxidation state of the metal and is usually satisfied by simple ions (ionisable). The secondary valence corresponds to the number of ligand attachments to the metal (the coordination number) and is non-ionisable in the complex; it indicates how many ligand donor atoms are directly bonded to the metal.
• Coordination sphere: A complex consists of a central metal atom or ion with ligands directly bonded to it; these together form the coordination sphere, usually written within square brackets.
• Coordination number:  Werner proposed that the coordination number is the total number of ligands directly bonded to the central metal ion within the coordination sphere. The coordination number determines the geometry and stability of the complex.
• Isomerism: Werner's ideas explained the existence of isomers in coordination chemistry - different spatial arrangements of the same set of ligands about the metal can give distinct compounds with different physical and chemical properties.

Werner's experimental and theoretical contributions were recognised with the Nobel Prize in Chemistry in 1913.

Werner's Experiments and Historical Notes

Alfred Werner developed his coordination theory in the early 1890s (major ideas formed during 1890-1893) and then devoted many years to experimental work that supported his model. The experimental evidence included isolation and characterisation of compounds showing distinct isomeric forms and demonstration that certain ligand attachments remained non-ionisable in solution (i.e., part of a coordination sphere rather than simple ionic constituents).

Werner's combined theoretical and experimental approach transformed the understanding of atomic linkage in inorganic chemistry and established coordination chemistry as a systematic discipline.His pioneering work in the field of atomic linkage and coordination theory earned him the prestigious Nobel Prize in Chemistry in 1913, making him the first Swiss chemist to receive this honour.

Postulates of Werner's Theory

The principal postulates of Werner's theory are:

• The central metal atom (or ion) in a coordination compound exhibits two types of valency: primary valence (oxidation state) and secondary valence (coordination number).
• The primary valences are normally ionisable and are satisfied by negative ions; they correspond to the oxidation state of the metal.
• The secondary valences are non-ionisable (in the complex) and are satisfied by ligands (which may be neutral molecules or anions); the number of secondary valences equals the coordination number of the metal.
• The groups or atoms bound by the secondary valences are arranged in space in characteristic geometries determined by the coordination number (for example, octahedral for CN = 6, tetrahedral for CN = 4, square planar for some d8 metal centres, etc.).

Common examples illustrating Werner's assignments of geometry are:

• Octahedral: [Co(NH₃)₆]³⁺, [CoCl(NH₃)₅]²⁺, [CoCl₂(NH₃)₄]⁺.
• Tetrahedral: [Ni(CO)₄] (carbonyls such as Ni(CO)₄ are commonly tetrahedral).
• Square planar: [PtCl₄]^2- (many d8 platinum(II) complexes adopt square-planar geometry).
  Difference between Primary and Secondary Valency in Coordination Compounds:

Isomerism in Coordination Compounds

Werner's theory explained several types of isomerism observed in coordination compounds. Important isomer types are:

• Geometrical isomerism: different spatial arrangements of ligands (e.g., cis and trans isomers in [Co(NH₃)₄Cl₂]⁺).
• Linkage isomerism: when an ambidentate ligand (able to bind through two different donor atoms) bonds through different atoms (for example, NO₂^- can bind through N or O).
• Optical isomerism: non-superposable mirror-image forms (enantiomers) occur when the arrangement of ligands is chiral, for example some octahedral complexes with three bidentate ligands.
• Coordination isomerism: arises in complex salts where the composition of complex cations and anions can interchange (observed in compounds with both cationic and anionic complex ions).

Limitations of Werner's Theory

Werner's theory successfully explained many structural and isomeric properties of coordination compounds but had limitations:

• Limited treatment of bonding: It did not describe the electronic nature of metal-ligand bonding (no orbital explanation), and therefore could not distinguish the covalent versus ionic character of coordinate bonds.
• Neglect of electronic structure: The theory did not explicitly use the concepts of metal and ligand orbital interactions, crystal field splitting, or ligand field theory, which are necessary to explain electronic spectra, magnetic properties and many thermodynamic trends.
• Magnetic properties: Werner's model could not explain magnetic behaviour arising from unpaired d-electrons; later theories (crystal field and ligand field theories) account for such observations.
• Exceptions and extensions: Some coordination geometries and stabilities require consideration of electronic configuration, steric effects and ligand field effects; these were outside Werner's original scope.
• Simplistic treatment of isomerism: Werner's theory recognized geometric and linkage isomerism but did not offer detailed explanations or predictions of these isomeric forms. It lacked a comprehensive framework for understanding the factors governing isomerism in coordination compounds.

Despite these limitations, Werner's theory provided the correct structural framework and remains a fundamental step in inorganic chemistry.

Basic Concepts and Terms Related to Coordination Compounds

Coordination compounds (also called complex compounds) contain a central metal atom or ion bonded to a set of ligands (atoms, ions or molecules that donate electron pairs). The following terms and concepts are essential.

Coordination Entity

A coordination entity is the chemical species made up of the central metal atom or ion together with its directly bonded ligands. Examples: [CoCl₃(NH₃)₃] and [Fe(CN)₆]^4-.

Central Atom / Central Ion

The central atom or ion is the metal centre to which ligands are attached. In coordination chemistry the central species typically acts as a Lewis acid (electron-pair acceptor).

Ligands

The atoms, molecules, or ions that are bound to the coordination centre or the central atom/ion are referred to as ligands. Ligands are atoms, ions or molecules that donate one or more lone pairs to the metal to form coordinate (dative) bonds. Ligands may be neutral (e.g., NH₃, H₂O, CO) or anionic (e.g., Cl^-, CN^-). Ligands are Lewis bases.

Coordination Number (CN)

The coordination number of a metal centre is the number of donor atoms from ligands that are directly bonded to the metal; it determines the coordination geometry.
For example, [Co(NH₃)₆]³⁺ has coordination number 6 (octahedral).

Coordination Sphere

The coordination sphere consists of the central metal and the ligands bound to it, written inside square brackets.

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: [Fe(CN)₆]^4- in K₄[Fe(CN)₆]. Species outside the brackets (counter-ions) are part of the ionisation sphere.

Coordination Polyhedron

The coordination polyhedron is the geometric shape formed by joining the positions of ligand donor atoms around the central metal. Common polyhedra are octahedral (CN = 6), tetrahedral (CN = 4) and square planar (often CN = 4 for d8 metals).

Examples: [Co(NH₃)₆]³⁺ (octahedral), [Ni(CO)₄] (tetrahedral), [PtCl₄]^2- (square planar).

Oxidation Number

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.

Example: The oxidation number of the platinum atom in the complex [PtCl₆]^2- is +4.

Homoleptic and Heteroleptic Complexes

• Homoleptic complex: a complex in which the metal is bound to only one kind of ligand, e.g., [Cu(CN)₄]^3-.
• Heteroleptic complex: a complex in which the metal is bound to different kinds of ligands, e.g., [Co(NH₃)₄Cl₂]⁺.

Double Salts and Complexes

Double salts remain fully dissociated into simple ions in solution and retain individual ionic behaviour (examples: carnallite KCl·MgCl₂·6H₂O, potassium alum KAl(SO₄)₂·12H₂O). 
Complexes (or coordination compounds) lose their crystalline identity on dissolution and give complex ions rather than simple free metal ions and ligands (example: K₄[Fe(CN)₆] gives [Fe(CN)₆]^4- not free Fe²⁺ and CN^-).

When crystals of carnallite are dissolved in water, the solution shows properties of K⁺, Mg²⁺  and Cl^- ions. In a similar way, a solution of potassium alum shows the properties of K⁺, Al³⁺  and SO₄^2- ions. These are both examples of double salts which exist only in the crystalline state. When the other two examples of coordination compounds are dissolved they do not form simple ions, Cu²⁺, Fe²⁺  and CN^-, but instead, their complex ions are formed.

Representation of Coordination Complexes

General notation used for coordination complexes is:

• M = central metal atom/ion (usually a d-block element)
• L = ligand
• x = number of ligands of type L
• 
  = overall charge on the coordination entity

The part outside the square brackets is the ionisation sphere (counter ions), the part inside brackets is the coordination sphere.

Central metal atoms/ions must have vacant orbitals to accept electron pairs from ligands which is why transition metals form many coordination compounds. Ligands donate lone pairs and are Lewis bases.

Ligands

Types of Ligands by Denticity

Ligands are classified by the number of donor atoms they use to bind to the metal:

• Monodentate (unidentate): one donor atom per ligand (e.g., F^-, Cl^-, Br^-, H₂O, NH₃, CN^-, NO₂^-, OH^-, CO).Unidentate Ligands
• Bidentate: two donor atoms per ligand (examples shown below).
• Tridentate: three donor atoms per ligand.diethylenetriamine (dien)
• Tetradentate Ligands These ligands possess four donor atoms. Examples are

  

• 
  

• Pentadentate Ligands They have five donor atoms. For example, ethylenediamine triacetate ion.

  

• Hexadentate Ligands They have six donor atoms. The most important example is the ethylenediaminetetraacetate ion.

  

• Ambidentate: ligands having more than one potential donor atom but coordinating through only one of them in a given complex (e.g., NO₂^- can bind through N or O).

• Polydentate (chelating): ligands that attach through several donor atoms and form chelate rings, increasing complex stability.

Coordination Sphere Reinforced

The central metal and its directly bound ligands constitute the coordination sphere and behave as a single entity in many reactions and in solution. Examples: [Co(NH₃)₆]³⁺.

Coordination Number - Further Examples

The coordination number equals the number of donor atoms bonded to the metal. Many common coordination numbers and associated geometries are:

• 4 - tetrahedral or square planar
• 6 - octahedral
• 2 - linear (rare for transition metals)
• 5 - square pyramidal or trigonal bipyramidal (less common)

Oxidation State Examples and Calculations

The oxidation state (O.S.) of the central metal is obtained by assigning charges to ligands and using the overall charge of the coordination entity. Worked examples:

(i) Potassium ferrocyanide, K₄[Fe(CN)₆]

The complex ion is [Fe(CN)₆]^4-. Each CN^- is univalent; total ligand charge = -6. Let x be oxidation state of Fe:

x + (-6) = -4

Therefore x = +2. Iron is Fe(II) in this complex.

(ii) [Cr(C₂O₄)₃]^3-

Oxalate (C₂O₄) is a bidentate ligand with charge -2. Three oxalates give total -6. Let x be oxidation state of Cr:

x + (-6) = -3

Therefore x = +3. Chromium is Cr(III).

(iii) Nickel carbonyl, Ni(CO)₄

CO is a neutral ligand. Total complex charge is zero, so nickel must be Ni(0): oxidation state of Ni = 0.

Effective Atomic Number (EAN) - Sidgwick Rule

Effective Atomic Number (EAN) is the total number of electrons on the central metal after accepting electron pairs from ligands through coordinate bonds. Sidgwick proposed that a complex tends to be stable if the metal's EAN equals the atomic number of the nearest noble gas.

EAN is calculated as:

EAN = Z - (oxidation state) + 2 × (coordination number)

Example: For [Co(NH₃)₆]³⁺

Atomic number of Co, Z = 27

Oxidation state = +3

Coordination number = 6

EAN = 27 - 3 + 2×6 = 27 - 3 + 12 = 36

36 corresponds to krypton (Kr), a noble gas; thus the complex is considered stable by the EAN criterion. Note that the EAN rule is useful for many metal-carbonyls and other complexes but is not universally obeyed.

Summary

Werner's theory introduced the key structural ideas - primary and secondary valences, coordination sphere, coordination number and characteristic coordination geometries - that explain the formation and isomerism of coordination compounds. Later theories (crystal/ligand field theory, molecular orbital approaches) extended Werner's framework by explaining electronic structure, magnetic behaviour and spectral properties. Understanding the basic terminology (ligand types, denticity, coordination number, oxidation state, coordination polyhedron and EAN) is essential for further study of transition-metal chemistry.