Coordination Compounds Class 12 Notes – Ligands, Bonding, Isomerism & NCERT Solutions

 

Unit 9 – Coordination Compounds

9.1 Werner’s Theory of Coordination Compounds
9.2 Definitions of Some Important Terms Of Coordination Compounds
9.3 Nomenclature of Coordination Compounds
9.4 Isomerism in Coordination Compounds
9.5 Bonding in Coordination Compounds
9.6 Bonding in Metal Carbonyls
9.7 Importance and Applications of Coordination Compounds

Coordination compounds are a central and challenging area of modern inorganic chemistry, forming the backbone of modern inorganic and bio-inorganic chemistry and the chemical industry. These compounds involve metal atoms bound to various anions or neutral molecules by sharing electrons. They are vital components in biological systems, act as industrial catalysts, and are used in analytical reagents, electroplating, textile dyeing, and medicinal chemistry.

Here are the important pointers regarding coordination compounds, based on the provided sources:

1. Werner’s Theory of Coordination Compounds

Alfred Werner, a Swiss chemist, was the first to formulate ideas about the structures of coordination compounds (between 1890 and 1893). His work earned him the Nobel Prize in 1913. Werner's theory, propounded in 1898, is based on the following postulates:

• Two Types of Linkages (Valences): In coordination compounds, metals exhibit two types of valences: primary and secondary.
  • Primary Valences: These are normally ionisable and are satisfied by negative ions. They correspond to the metal ion's oxidation state [Source A, an external source, not provided. This is general chemistry knowledge]. For example, in binary compounds like CrCl₃, CoCl₂, or PdCl₂, the primary valences are 3, 2, and 2, respectively.
  • Secondary Valences: These are non-ionisable and are satisfied by neutral molecules or negative ions. The secondary valence is equal to the coordination number and is fixed for a given metal. In a series of cobalt(III) chloride-ammonia compounds (e.g., CoCl₃·6NH₃, CoCl₃·5NH₃, CoCl₃·4NH₃), the secondary valences were found to be six.
• Spatial Arrangement: The ions or groups bound by secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordination numbers. In modern terms, these spatial arrangements are called coordination polyhedra. Werner postulated that octahedral, tetrahedral, and square planar geometrical shapes are common in transition metal coordination compounds.
  • Examples of octahedral entities include [Co(NH₃)₆]³⁺, [CoCl(NH₃)₅]²⁺, and [CoCl₂(NH₃)₄]⁺.
  • Examples of tetrahedral and square planar entities are [Ni(CO)₄] and [PtCl₄]²⁻, respectively.

Werner's experiments with cobalt(III) chloride and ammonia showed varying amounts of AgCl precipitated upon adding silver nitrate, indicating that some chloride ions were bonded differently than others (inside vs. outside the coordination sphere). For instance, 1 mol of CoCl₃·6NH₃ (Yellow) gave 3 mol AgCl, while 1 mol of CoCl₃·4NH₃ (Green or Violet) gave 1 mol AgCl. These observations were consistent with specific formulations where atoms within square brackets form a single entity that does not dissociate.

2. Definitions of Some Important Terms Pertaining to Coordination Compounds

• Coordination Entity: A central metal atom or ion bonded to a fixed number of ions or molecules. It is enclosed in square brackets, e.g., [CoCl₃(NH₃)₃], [Ni(CO)₄], [Fe(CN)₆]⁴⁻. The species within the square brackets are also called complexes.
• Central Atom/Ion: The atom or ion in a coordination entity to which a fixed number of ions/groups are bound in a definite geometrical arrangement. These central atoms/ions are also referred to as Lewis acids. Examples include Ni²⁺ in [NiCl₂(H₂O)₄] and Fe³⁺ in [Fe(CN)₆]³⁻.
• Ligands: The ions or molecules bound to the central atom/ion in the coordination entity. Ligands can be simple ions (e.g., Cl⁻), small molecules (e.g., H₂O, NH₃), larger molecules (e.g., H₂NCH₂CH₂NH₂), or even macromolecules like proteins.
  • Unidentate Ligands: Bind through a single donor atom (e.g., Cl⁻, H₂O, NH₃).
  • Didentate Ligands: Bind through two donor atoms (e.g., H₂NCH₂CH₂NH₂ (ethane-1,2-diamine), C₂O₄²⁻ (oxalate)).
  • Polydentate Ligands: Have several donor atoms in a single ligand (e.g., N(CH₂CH₂NH₂)₃). Ethylenediaminetetraacetate ion (EDTA⁴⁻) is an important hexadentate ligand, binding through two nitrogen and four oxygen atoms.
  • Chelate Ligands: Di- or polydentate ligands that use two or more donor atoms simultaneously to bind a single metal ion. Complexes formed by chelate ligands, called chelate complexes, tend to be more stable than similar complexes with unidentate ligands. The number of ligating groups is called the denticity of the ligand.
  • Ambidentate Ligands: Ligands with two different donor atoms where either can ligate in the complex (e.g., NO₂⁻ which can coordinate through nitrogen or oxygen; SCN⁻ which can coordinate through sulfur or nitrogen).
• Coordination Number (CN): The number of ligand donor atoms directly bonded to the central metal ion in a complex. This is determined only by the number of sigma bonds formed. Pi bonds are not counted.
  • Examples: Pt has a CN of 6 in [PtCl₆]²⁻, Ni has a CN of 4 in [Ni(NH₃)₄]²⁺.
  • For didentate ligands like C₂O₄²⁻ and ethane-1,2-diamine (en), each contributes two donor atoms. So, in [Fe(C₂O₄)₃]³⁻ and [Co(en)₃]³⁺, the CN of Fe and Co is 6.
• Coordination Sphere: The central atom/ion and its directly attached ligands, enclosed in square brackets.
• Counter Ions: Ionisable groups written outside the square bracket of the coordination sphere. For example, in K₄[Fe(CN)₆], [Fe(CN)₆]⁴⁻ is the coordination sphere and K⁺ is the counter ion.
• Coordination Polyhedron: The spatial arrangement of ligand atoms directly attached to the central atom/ion. Common shapes include octahedral, square planar, and tetrahedral.
  • [Co(NH₃)₆]³⁺ is octahedral, [Ni(CO)₄] is tetrahedral, and [PtCl₄]²⁻ is square planar.
• Oxidation Number of Central Atom: The charge the central atom would carry if all ligands were removed along with their shared electron pairs. It is indicated by a Roman numeral in parenthesis after the coordination entity's name (e.g., Cu(I) for copper in [Cu(CN)₄]³⁻).
• Homoleptic Complexes: Complexes where a metal is bound to only one kind of donor group (e.g., [Co(NH₃)₆]³⁺).
• Heteroleptic Complexes: Complexes where a metal is bound to more than one kind of donor group (e.g., [Co(NH₃)₄Cl₂]⁺).

3. Nomenclature of Coordination Compounds (IUPAC Recommendations)

Nomenclature is crucial for unambiguously describing formulas and systematic names, especially for isomers.

• Formulas of Mononuclear Coordination Entities:
  • The central atom is listed first.
  • Ligands are listed in alphabetical order, regardless of their charge.
  • Polydentate ligands and ligand abbreviations are also listed alphabetically, using the first letter of their abbreviation.
  • The entire coordination entity is enclosed in square brackets. Polyatomic ligands and abbreviations are enclosed in parentheses within the brackets.
  • No space exists between ligands and the metal within the coordination sphere.
  • For charged entities without a counter ion, the charge is indicated as a right superscript outside the square brackets (e.g., [Co(CN)₆]³⁻).
  • The charge of cations is balanced by the charge of anions.
• Naming of Mononuclear Coordination Compounds:
  • The cation is named first, followed by the anion (if present).
  • Ligands are named in alphabetical order before the central atom/ion. (This is the reverse of formula writing).
  • Anionic ligands end in "-o" (e.g., chlorido, cyanido, sulfato). Neutral and cationic ligands retain their common names, with exceptions like aqua for H₂O, ammine for NH₃, carbonyl for CO, and nitrosyl for NO.
  • Prefixes like mono, di, tri are used for the number of individual ligands. If the ligand's name already includes a numerical prefix (e.g., ethane-1,2-diamine), then bis, tris, tetrakis are used, and the ligand name is placed in parentheses.
    • Example: [NiCl₂(PPh₃)₂] is named dichloridobis(triphenylphosphine)nickel(II).
  • The oxidation state of the metal is indicated by a Roman numeral in parenthesis following the name of the coordination entity.
  • If the complex ion is a cation, the metal is named the same as the element (e.g., cobalt for Co, platinum for Pt).
  • If the complex ion is an anion, the metal name ends with the suffix "-ate" (e.g., cobaltate for Co, ferrate for Fe). Latin names are used for some metals in complex anions (e.g., ferrate for Fe).
  • A neutral complex molecule is named similarly to a complex cation.
  • Examples:
    • [Cr(NH₃)₃(H₂O)₃]Cl₃ is triamminetriaquachromium(III) chloride.
    • [Co(H₂NCH₂CH₂NH₂)₃]₂(SO₄)₃ is tris(ethane-1,2–diamine)cobalt(III) sulphate.
    • [Ag(NH₃)₂][Ag(CN)₂] is diamminesilver(I)dicyanidoargentate(I).

4. Isomerism in Coordination Compounds

Isomers are compounds with the same chemical formula but different arrangements of atoms, leading to different physical or chemical properties. Two main types exist:

• Stereoisomerism: Isomers with the same chemical formula and bonds but different spatial arrangements.
  • Geometrical Isomerism: Arises in heteroleptic complexes due to different possible geometric arrangements of ligands.
    • Common in coordination numbers 4 (square planar) and 6 (octahedral).
    • In square planar [MX₂L₂] type complexes, cis (ligands X adjacent) and trans (ligands X opposite) isomers exist. Not possible for tetrahedral geometry.
    • In octahedral [MX₂L₄] type complexes, cis and trans isomers exist.
    • Octahedral [Ma₃b₃] complexes (e.g., [Co(NH₃)₃(NO₂)₃]) can show facial (fac) isomers (three donor atoms of the same ligand occupy corners of an octahedral face) or meridional (mer) isomers (positions are around the octahedron's meridian).
  • Optical Isomerism (Chirality): Occurs when a complex is non-superimposable on its mirror image.
• Structural Isomerism: Isomers that have different chemical bonds.
  • Linkage Isomerism: Arises with ambidentate ligands, which can coordinate through different donor atoms (e.g., [Co(NH₃)₅(NO₂)]²⁺ with nitro-N or nitrito-O).
  • Coordination Isomerism: Occurs in compounds where both cation and anion are complex entities, and the ligands are interchanged between the two metal centers (e.g., [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆]).
  • Ionisation Isomerism: Arises when the counter ion of a complex salt is a potential ligand and can displace a ligand from within the coordination sphere, making the displaced ligand the new counter ion (e.g., [Co(NH₃)₅(SO₄)]Br and [Co(NH₃)₅Br]SO₄). These isomers yield different ions in solution and react differently with reagents.
  • Solvate Isomerism: A specific type of ionisation isomerism where solvent molecules are involved as ligands or as lattice solvent.

5. Bonding in Coordination Compounds

Werner's theory described bonding features but couldn't explain why only certain elements form coordination compounds, the directional properties of their bonds, or their characteristic magnetic and optical properties. To address these, other theories were developed:

• Valence Bond Theory (VBT):

  • Concept: States that the metal atom or ion, under the influence of ligands, uses its (n-1)d, ns, np or ns, np, nd orbitals for hybridisation to form a set of equivalent orbitals with a definite geometry (e.g., octahedral, tetrahedral, square planar). These hybrid orbitals overlap with ligand orbitals that donate electron pairs for bonding.
  • Geometry Prediction:
    • Coordination Number 4: Can be sp³ (tetrahedral) or dsp² (square planar).
      • [NiCl₄]²⁻ (Ni is d⁸, +2 oxidation state) is tetrahedral and paramagnetic (two unpaired electrons) because Cl⁻ is a weak ligand and doesn't force electron pairing.
      • [Ni(CN)₄]²⁻ (Ni is d⁸, +2 oxidation state) is square planar and diamagnetic (no unpaired electrons) because CN⁻ is a strong ligand, forcing d-electron pairing to make a d-orbital available for dsp² hybridization.
      • [Ni(CO)₄] (Ni is d¹⁰, 0 oxidation state) is tetrahedral and diamagnetic (no unpaired electrons).
    • Coordination Number 6: Forms sp³d² (outer orbital/high spin/spin-free complex) or d²sp³ (inner orbital/low spin/spin-paired complex), both having octahedral geometry.
      • [Co(NH₃)₆]³⁺ (Co is d⁶, +3 oxidation state) is diamagnetic (no unpaired electrons) and an inner orbital complex (d²sp³ hybridization) because NH₃ is a strong ligand, forcing 3d electrons to pair up.
      • [CoF₆]³⁻ (Co is d⁶, +3 oxidation state) is paramagnetic (four unpaired electrons) and an outer orbital complex (sp³d² hybridization) because F⁻ is a weak ligand and does not force electron pairing.
      • [Fe(CN)₆]³⁻ (Fe is d⁵, +3 oxidation state) has a magnetic moment of one unpaired electron (low spin).
      • [FeF₆]³⁻ (Fe is d⁵, +3 oxidation state) has a paramagnetic moment of five unpaired electrons (high spin).
  • Limitations of VBT:
    • Involves a number of assumptions.
    • Does not provide a quantitative interpretation of magnetic data.
    • Does not explain the color exhibited by coordination compounds.
    • Does not give a quantitative interpretation of thermodynamic or kinetic stabilities.
    • Does not make exact predictions regarding tetrahedral vs. square planar structures for 4-coordinate complexes.
    • Does not distinguish between weak and strong ligands.

• Crystal Field Theory (CFT):

  • Concept: An electrostatic model that considers the metal-ligand bond as purely ionic, arising from electrostatic interactions. Ligands are treated as point charges (anions) or point dipoles (neutral molecules).
  • d-orbital Degeneracy Splitting: In an isolated gaseous metal atom/ion, the five d orbitals are degenerate (have the same energy). When ligands surround the metal ion in a complex, the negative field becomes asymmetrical, lifting the d-orbital degeneracy and causing splitting of the d orbitals. The pattern of splitting depends on the crystal field geometry.
  • Octahedral Splitting: In an octahedral field, d-orbitals pointing towards ligands (dₓ²₋ᵧ² and dz²) experience more repulsion and are raised in energy (forming the e_g_ set), while orbitals directed between the axes (dxy, dyz, dxz) are lowered in energy (forming the t₂_g_ set). The energy separation between these sets is called crystal field splitting energy, denoted as Δ_o_ (delta-octahedral). The e_g_ orbitals increase by (3/5)Δ_o_, and the t₂_g_ orbitals decrease by (2/5)Δ_o_ relative to the barycenter (average energy in a spherical field).
  • Spectrochemical Series: Ligands can be arranged in a series by their increasing field strength, which determines the magnitude of Δ_o_. This series is experimentally determined from complex light absorption: I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < EDTA⁴⁻ < NH₃ < en < CN⁻ < CO
  • High Spin vs. Low Spin Complexes (d⁴ to d⁷ ions): For d⁴ ions (and generally d⁴ to d⁷), two electron distribution patterns depend on the relative magnitudes of Δ_o_ and pairing energy (P, energy required for electron pairing):
    • Weak Field Ligands (High Spin): If Δ_o_ < P, the fourth electron enters the e_g_ orbital (configuration t₂_g_³ e_g_¹), forming a high spin complex.
    • Strong Field Ligands (Low Spin): If Δ_o_ > P, the fourth electron pairs up in a t₂_g_ orbital (configuration t₂_g_⁴ e_g_⁰), forming a low spin complex.
  • Tetrahedral Splitting: The d-orbital splitting in tetrahedral coordination entities is inverted compared to octahedral splitting, and the energy splitting (Δ_t_) is smaller (Δ_t_ = (4/9)Δ_o_). Because Δ_t_ is generally not large enough to force pairing, low spin configurations are rarely observed in tetrahedral complexes. The "g" subscript is not used for tetrahedral complexes as they lack a center of symmetry.
  • Color in Coordination Compounds: CFT explains color by d-d transitions. When a complex absorbs light of a specific wavelength, an electron is excited from a lower energy d-orbital set to a higher energy d-orbital set (e.g., t₂_g_ to e_g_ in octahedral). The color observed is the complementary color of the light absorbed.
    • Example: [Ti(H₂O)₆]³⁺ (a 3d¹ system) is violet because it absorbs light in the blue-green region to excite its electron from t₂_g_ to e_g_.
    • In the absence of ligands, no crystal field splitting occurs, and the substance is colorless (e.g., anhydrous CuSO₄ is white, but CuSO₄·5H₂O is blue). Ligand changes also affect color (e.g., green [Ni(H₂O)₆]²⁺ becomes violet [Ni(en)₃]²⁺ with increasing ethane-1,2-diamine).
  • Limitations of CFT:
    • Considers metal-ligand bonds as purely ionic, which is often an oversimplification.
    • Does not account for bonding in metal carbonyls or other complexes with strong covalent character [Source A, an external source, not provided].

6. Bonding in Metal Carbonyls

Metal carbonyls are compounds containing only carbonyl (CO) ligands, typically formed by transition metals.

• Structure Examples:
  • Tetracarbonylnickel(0) [Ni(CO)₄] is tetrahedral.
  • Pentacarbonyliron(0) [Fe(CO)₅] is trigonal bipyramidal.
  • Hexacarbonyl chromium(0) [Cr(CO)₆] is octahedral.
  • Decacarbonyldimanganese(0) [Mn₂(CO)₁₀] has two square pyramidal Mn(CO)₅ units joined by a Mn–Mn bond.
  • Octacarbonyldicobalt(0) [Co₂(CO)₈] has a Co–Co bond bridged by two CO groups.
• Synergic Bonding: The metal-carbon bond in metal carbonyls has both sigma (σ) and pi (π) character.
  • M–C σ bond: Formed by the donation of a lone pair of electrons from the carbonyl carbon into a vacant orbital of the metal.
  • M–C π bond: Formed by the donation of a pair of electrons from a filled d orbital of the metal into the vacant antibonding π* orbital of carbon monoxide.
  • This metal-to-ligand bonding creates a synergic effect that strengthens the bond between CO and the metal.

7. Importance and Applications of Coordination Compounds

Coordination compounds are highly important in various fields:

• Analytical Chemistry: Used in qualitative and quantitative chemical analysis due to their characteristic color reactions with metal ions, especially with chelating ligands. Reagents like EDTA, DMG (dimethylglyoxime) are examples. Hardness of water is estimated using Na₂EDTA.
• Metallurgy: Utilized in important metal extraction processes, such as those for silver and gold, through complex formation. Impure nickel is purified via [Ni(CO)₄] formation and decomposition.
• Biological Systems: Play critical roles in living organisms.
  • Chlorophyll (for photosynthesis) is a coordination compound of magnesium.
  • Haemoglobin (oxygen carrier in blood) is a coordination compound of iron.
  • Vitamin B₁₂ (cyanocobalamine) is a coordination compound of cobalt.
  • Enzymes like carboxypeptidase A and carbonic anhydrase also involve coordinated metal ions as biological catalysts.
• Industry: Used as catalysts in various industrial processes. For example, Wilkinson catalyst [(Ph₃P)₃RhCl] is used for alkene hydrogenation.
• Medicine (Chelate Therapy): Growing interest in using coordination compounds for medicinal purposes.
  • Chelate therapy removes toxic metals from biological systems (e.g., D-penicillamine and desferrioxime B for excess copper and iron, EDTA for lead poisoning).
  • Some platinum coordination compounds, such as cis-platin, effectively inhibit tumor growth.
• Electroplating: Metals like silver and gold can be electroplated more smoothly and evenly from solutions of their complexes, such as [Ag(CN)₂]⁻ and [Au(CN)₂]⁻, rather than from simple metal ion solutions.
• Photography: In black and white photography, undecomposed AgBr is dissolved by hypo solution (sodium thiosulfate) to form a complex ion, [Ag(S₂O₃)₂]³⁻.

Understanding coordination compounds is like learning to decipher a complex code. Each component—the central metal, the surrounding ligands, and their specific arrangement—contributes to the compound's unique "message" in terms of its properties, reactivity, and role in various fields, from biology to industrial processes.

___________________________________________________________________________________

Here is a comprehensive collection of NCERT solutions for Class 12 Chemistry Chapter 9 (Coordination Compounds), along with past CBSE board exam questions from the chapter organized by marks (1M, 2M, 3M, etc.) with detailed answers. The exercises and exam questions include all sections from the NCERT textbook and commonly appeared CBSE questions.

NCERT Solutions for Class 12 Chemistry Chapter 9 – Coordination Compounds

Key Resources

• Exercises from NCERT Class 12 Chemistry Chapter 9 Coordination Compounds

• Past CBSE questions collected from previous board exams

• Detailed step-by-step solutions for each question

NCERT Exercise Questions with Solutions

Q1. Write formulas for the following coordination compounds: (1M each)
(i) Tetraamminediaquacobalt(III) chloride
(ii) Potassium tetracyanidonickelate(II)
(iii) Tris(ethane-1,2-diamine) chromium(III) chloride
(iv) Amminebromidochloridonitrito-N-platinate(II)
(v) Dichloridobis(ethane-1,2-diamine)platinum(IV) nitrate
(vi) Iron(III) hexacyanoferrate(II)

Answer:
(i) [Co(NH3)4(H2O)2]Cl3
(ii) K2[Ni(CN)4]
(iii) [Cr(en)3]Cl3
(iv) [Pt(NH3)BrCl(NO2)]
(v) PtCl2(en)2
(vi) Fe3[Fe(CN)6]2

Q2. Explain Werner’s theory of coordination compounds. (2M)
Answer:
Werner's theory explains the composition and structure of coordination compounds. It distinguishes between primary valence (oxidation state) and secondary valence (coordination number). According to Werner, coordination compounds contain central metal atoms or ions surrounded by definite numbers of ions or molecules called ligands, which are directly attached to the metal ion.

Q3. Define the following terms with examples: (3M total or 1M each)

• Central atom/ion

• Ligand

• Coordination number

• Coordination sphere

• Oxidation number

Answer:

• Central atom/ion: The metal ion or atom to which ligands are attached; e.g., Co in [Co(NH3)6]3+.

• Ligand: A molecule or ion that donates a pair of electrons to the central metal atom; e.g., NH3, CN-.

• Coordination number: Number of ligand donor atoms attached to the central metal ion; e.g., 6 in [Fe(CN)6]4-.

• Coordination sphere: The central metal atom/ion and its attached ligands as a unit; e.g., [Cr(en)3]3+.

• Oxidation number: Charge assigned to the central atom assuming ionic ligand bonding; e.g., +3 for Cr in [Cr(en)3]3+.

Q4. Write IUPAC names for the following: (1M each)
(i) [Co(NH3)5Cl]Cl2
(ii) K4[Fe(CN)6]
(iii) Ni(H2O)62
(iv) [Cr(H2O)4Cl2]Cl

Answer:
(i) Pentaamminechloridocobalt(III) chloride
(ii) Potassium hexacyanoferrate(II)
(iii) Hexaaquanickel(II) nitrate
(iv) Tetraaquadichloridochromium(III) chloride

Q5. Differentiate between: (2M)
(i) Coordination number and oxidation state
(ii) Ligand and coordination entity

Answer:
(i) Coordination number is the number of donor atoms attached to central metal, oxidation state is the charge on the metal ion.
(ii) Ligand is the donor molecule or ion, coordination entity includes central metal plus its bound ligands.

Past CBSE Board Questions for Coordination Compounds

1 Mark Questions (1M)

• Define coordination number with example.

• Write the formula of potassium hexacyanoferrate(II).

• Name the ligands in [Co(NH3)6]Cl3.

• What is the oxidation state of Cr in [Cr(H2O)4Cl2]Cl?

2 Mark Questions (2M)

• State Werner’s theory of coordination compounds.

• Write IUPAC names for [Ni(CO)4] and [Fe(CN)6]4-.

• Differentiate between homoleptic and heteroleptic complexes with examples.

• Explain the term ‘coordination sphere’.

3 Mark Questions (3M)

• Explain the valence bond theory for coordination compounds.

• Describe the types of isomerism in coordination compounds with examples.

• Write notes on bonding in coordination compounds.

• What are ligands? Differentiate between monodentate, bidentate and ambidentate ligands with examples.

5 Mark Questions (5M)

• Discuss crystal field theory and explain the splitting of d-orbitals in octahedral complexes.

• Explain the importance and applications of coordination compounds in industry and medicine.

• Describe the stereoisomerism in coordination compounds with suitable examples.

• Write a detailed note on the nomenclature of coordination compounds.

Additional Important CBSE Questions from Recent Years (Detailed Answers)

1. Explain the bonding in coordination compounds with respect to Werner’s theory and Valence Bond Theory. (5M)
   Answer:
   Werner’s theory establishes the basics of coordination compounds distinguishing primary and secondary valencies. Valence Bond Theory extends this by explaining hybridization of the central metal’s orbitals to accommodate ligands, describing shapes and magnetic properties. For example, in [Fe(CN)6]4-, d2sp3 hybridization leads to octahedral geometry.

2. Write the IUPAC name for [Co(en)2Cl2]Cl and mention its possible isomers. (3M)
   Answer:
   Name: Dichlorido-bis(ethane-1,2-diamine)cobalt(III) chloride. It shows geometrical isomerism (cis and trans).

3. What are ambidentate ligands? Give examples and explain how they differ from bidentate ligands. (2M)
   Answer:
   Ambidentate ligands can coordinate through two different atoms but only one donor atom at a time, example: NO2- (can bind via N or O). Bidentate ligands coordinate simultaneously through two donor atoms like ethylenediamine (en).

Hashtags:

#CoordinationCompounds #Class12Chemistry #NCERTSolutions #CBSE2025 #Ligands #WernerTheory #CrystalFieldTheory #JEEPreparation #NEETChemistry #ChemistryNotes