Chapter 5: Coordination Compounds Long Questions | chemistry Class 12 CBSE

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Chapter 5: Coordination Compounds Long Questions | chemistry Class 12 CBSE | ByteNotes.in

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Question 1

Long Form

Compare and contrast the Valence Bond Theory and Crystal Field Theory in explaining the bonding, magnetic behaviour and colour of coordination compounds.

VBT explains bonding through hybridisation of metal orbitals (d2sp3, sp3d2, dsp2, sp3) to accommodate ligand electron pairs, predicting geometry and qualitatively explaining magnetic behaviour (inner/outer orbital, low/high spin).
CFT is an electrostatic model treating ligands as point charges; it quantitatively explains d orbital splitting (Do in octahedral, Dt in tetrahedral), the spectrochemical series, and the origin of colour through d-d transitions.
VBT fails to explain colour and cannot quantitatively interpret magnetic data or thermodynamic stability, and does not distinguish between strong and weak field ligands.
CFT explains colour (d-d transitions, e.g., [Ti(H2O)6]3+ absorbs at 498 nm and appears violet) and provides quantitative magnetic moment predictions but incorrectly treats anionic ligands as producing the greatest splitting.
CFT also ignores the covalent character of metal-ligand bonding, which is addressed by Ligand Field Theory and Molecular Orbital Theory beyond VBT and CFT.

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Question 2

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Describe the various types of structural isomerism found in coordination compounds with one example each.

Linkage isomerism arises when ambidentate ligands coordinate through different donor atoms; e.g., [Co(NH3)5(NO2)]Cl2 exists as a red form (O-bonded, nitrito) and yellow form (N-bonded, nitro).
Coordination isomerism arises from interchange of ligands between cationic and anionic entities of different metal ions; e.g., [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6].
Ionisation isomerism arises when a counter ion and a coordinated ligand exchange roles; e.g., [Co(NH3)5(SO4)]Br and [Co(NH3)5Br]SO4, which give different precipitates with AgNO3 and BaCl2.
Solvate (hydrate) isomerism arises when a solvent molecule is either within the coordination sphere or as free solvent in the lattice; e.g., [Cr(H2O)6]Cl3 (violet) and [Cr(H2O)5Cl]Cl2.H2O (grey-green).
All structural isomers have the same molecular formula but different bonds; structural isomers differ from stereoisomers (geometric and optical) which have the same bonds but different spatial arrangements.

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Question 3

Long Form

Explain the crystal field splitting in octahedral and tetrahedral coordination entities. How does the magnitude of Do determine the actual configuration of d orbitals?

In an octahedral field, dx2-y2 and dz2 (eg) orbitals pointing along axes experience greater repulsion from ligands and are raised by 3/5 Do; dxy, dyz, dxz (t2g) orbitals are lowered by 2/5 Do. The splitting energy is Do.
In a tetrahedral field, splitting is inverted (t2 higher, e lower) and Dt = 4/9 Do for the same metal-ligand system; low spin tetrahedral complexes are rare because the small splitting cannot overcome pairing energy.
For d4-d7 configurations, if Do > P (pairing energy), electrons pair in t2g giving low spin (strong field) configuration; if Do < P, electrons enter eg giving high spin (weak field) configuration.
Example: for d4, strong field gives t2g4eg0 (1 unpaired) while weak field gives t2g3eg1 (4 unpaired); similarly, d5 Mn2+ gives 1 unpaired (strong field CN-) or 5 unpaired (weak field H2O).
The spectrochemical series (I- to CO) represents ligands from weakest to strongest field; the actual d orbital configuration and magnetic moment of a complex depend on which ligands are present relative to the pairing energy P.

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Question 4

Long Form

Analyse the Werner theory of coordination compounds using experimental data from cobalt(III) chloride-ammonia complexes, explaining the concepts of primary and secondary valence.

Werner proposed primary valence (ionisable, satisfied by negative ions) and secondary valence (non-ionisable, equal to coordination number, satisfied by neutral or anionic ligands) for metal ions in coordination compounds.
In CoCl3.6NH3 (yellow), 3 mol AgCl precipitate per mole with AgNO3, indicating 3 free Cl- ions; the formula is [Co(NH3)6]3+3Cl-, acting as a 1:3 electrolyte with coordination number 6.
In CoCl3.5NH3 (purple), 2 mol AgCl precipitate, giving [CoCl(NH3)5]2+2Cl- (1:2 electrolyte); in CoCl3.4NH3 (green or violet), 1 mol AgCl precipitates, giving [CoCl2(NH3)4]+Cl- (1:1 electrolyte).
The atoms within square brackets form the coordination sphere and do not dissociate under normal conditions; Werner proposed these are held by secondary valence.
The last two (green and violet) have identical empirical formula CoCl3.4NH3 but distinct properties, demonstrating isomerism; this validated Werner's structural interpretation and his concept of coordination polyhedra.

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Question 5

Long Form

Discuss the importance and applications of coordination compounds in biological systems, medicinal chemistry, analytical chemistry and metallurgy.

In biological systems: chlorophyll (Mg complex) enables photosynthesis; haemoglobin (Fe complex) carries oxygen in blood; vitamin B12 or cyanocobalamine (Co complex) prevents pernicious anaemia; metalloenzymes like carboxypeptidase A and carbonic anhydrase rely on coordinated metal ions.
In medicinal chemistry: chelate therapy using D-penicillamine removes excess copper and iron; EDTA treats lead poisoning; desferrioxime B chelates excess iron; cis-platin and related platinum coordination compounds effectively inhibit tumour growth.
In analytical chemistry: EDTA, DMG (dimethylglyoxime), alpha-nitroso-beta-naphthol and cupron are chelating reagents used for detection and quantitative estimation of metal ions; water hardness is estimated by titration with Na2EDTA exploiting stable Ca2+ and Mg2+ EDTA complexes.
In metallurgy: gold is extracted from ores using cyanide to form [Au(CN)2]-, from which gold is recovered by adding zinc; impure nickel is purified by converting it to volatile [Ni(CO)4], which is then decomposed to give pure nickel.
In electroplating and photography: smooth silver and gold electroplating uses [Ag(CN)2]- and [Au(CN)2]-; in black-and-white photography, hypo (sodium thiosulphate) fixes the developed film by dissolving undecomposed AgBr to form [Ag(S2O3)2]3-.

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Question 6

Long Form

Discuss the nature of bonding and geometry of [Co(NH3)6]3+, [CoF6]3-, [NiCl4]2- and [Ni(CN)4]2- using Valence Bond Theory.

[Co(NH3)6]3+: Co3+ has 3d6 configuration; NH3 (strong ligand) causes electron pairing in 3d, leaving two 3d orbitals empty for d2sp3 hybridisation; octahedral geometry; inner orbital/low spin; diamagnetic (0 unpaired electrons).
[CoF6]3-: Co3+ has 3d6; F- (weak ligand) cannot cause pairing; sp3d2 hybridisation uses outer 4d orbitals; octahedral geometry; outer orbital/high spin; paramagnetic with 4 unpaired electrons.
[NiCl4]2-: Ni2+ has 3d8; Cl- is a weak ligand; sp3 hybridisation; tetrahedral geometry; paramagnetic with 2 unpaired electrons (high spin complex).
[Ni(CN)4]2-: Ni2+ has 3d8; CN- is a strong ligand that causes the 2 unpaired 3d electrons to pair, making one 3d orbital empty; dsp2 hybridisation; square planar geometry; diamagnetic (0 unpaired electrons).
The key difference between [NiCl4]2- and [Ni(CN)4]2- is due to the field strength: Cl- cannot force electron pairing (tetrahedral, sp3) while CN- forces pairing (square planar, dsp2); this shows how ligand strength determines geometry and magnetic behaviour.

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Question 7

Long Form

Describe the rules of IUPAC nomenclature for mononuclear coordination compounds and apply them to name: [Cr(NH3)3(H2O)3]Cl3, K3[Cr(C2O4)3] and [Ag(NH3)2][Ag(CN)2].

IUPAC rules: (i) cation named before anion; (ii) ligands listed alphabetically before central metal; (iii) anionic ligands end in -o, neutral ligands use special names (aqua, ammine, carbonyl, nitrosyl); (iv) oxidation state of metal in Roman numerals in parentheses; (v) complex anion metal name ends in -ate.
Prefixes mono, di, tri for simple ligands; bis, tris, tetrakis for ligands with numerical prefix in their name; charge of cation(s) balanced by anion(s) to determine oxidation state.
[Cr(NH3)3(H2O)3]Cl3 = triamminetriaquachromium(III) chloride; all ligands neutral, complex cation with charge +3 requires 3 Cl- counter ions, so Cr is in +3 state; ammine listed before aqua alphabetically.
K3[Cr(C2O4)3] = potassium trioxalatochromate(III); complex anion contains Cr3+ and three oxalate didentate ligands; metal anion name uses -ate suffix (chromate); three K+ balance the 3- charge.
[Ag(NH3)2][Ag(CN)2] = diamminesilver(I) dicyanidoargentate(I); cation Ag+ with two ammine ligands named as silver (element name); anion Ag+ with two cyanido ligands named as argentate (Latin: argent for silver); both Ag are in +1 state.

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Question 8

Long Form

Explain geometric isomerism in square planar and octahedral coordination compounds with appropriate examples and diagrams description.

Geometric isomerism arises in heteroleptic complexes when different spatial arrangements of ligands are possible without breaking bonds; it occurs in square planar [MX2L2] and octahedral [MX2L4] types.
In square planar [MX2L2] (e.g., [Pt(NH3)2Cl2]): cis isomer has the two Cl ligands adjacent (90 degrees) and trans isomer has them opposite (180 degrees); these are structurally and chemically distinct.
For octahedral [MX2L4] (e.g., [Co(NH3)4Cl2]+): cis has two Cl adjacent; trans has two Cl diametrically opposite; similarly, [MX2(L-L)2] type with didentate ligands (e.g., [CoCl2(en)2]+) also shows cis/trans isomerism.
Octahedral [Ma3b3] type (e.g., [Co(NH3)3(NO2)3]) shows fac (facial) and mer (meridional) isomers: fac has three a ligands on one triangular face; mer has three a ligands in a plane including the central atom.
Tetrahedral complexes do not show geometric isomerism because all four positions are equivalent relative to each other; isomerism is possible only when the relative positions of ligands can differ.

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Question 9

Long Form

Analyse the bonding in metal carbonyls discussing the structure, hybridisation and the synergic bonding concept with examples.

Metal carbonyls are homoleptic compounds containing only carbonyl (CO) ligands, formed by most transition metals; they have simple, well-defined structures.
Structures: Ni(CO)4 is tetrahedral (sp3, Ni in zero oxidation state); Fe(CO)5 is trigonal bipyramidal; Cr(CO)6 is octahedral; Mn2(CO)10 has two square pyramidal Mn(CO)5 units joined by a Mn-Mn bond; Co2(CO)8 has a Co-Co bond bridged by two CO groups.
The M-C sigma bond is formed by donation of a lone pair from the carbon of CO into a vacant orbital of the metal; this is ligand-to-metal donation.
The M-C pi bond is formed by back-donation of a pair of electrons from a filled d orbital of the metal into the vacant antibonding pi* orbital of CO; this is metal-to-ligand donation.
The sigma donation increases electron density on the metal, which is relieved by pi back-donation; this mutual reinforcement is called synergic bonding and strengthens the overall M-C bond, stabilising the carbonyl complex.

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Question 10

Long Form

Explain the types and properties of ligands in coordination chemistry, covering unidentate, didentate, polydentate, chelate and ambidentate ligands with examples.

Unidentate ligands bond through a single donor atom: e.g., Cl-, H2O, NH3; didentate ligands use two donor atoms: e.g., ethane-1,2-diamine (en) through two N, and oxalate (C2O42-) through two O.
Polydentate ligands have several donor atoms: e.g., N(CH2CH2NH2)3 (tren); EDTA4- is a hexadentate ligand binding through two N and four O atoms to a single metal ion simultaneously.
Chelate ligands are di- or polydentate ligands that use two or more donor atoms simultaneously to bind a single metal ion, forming ring structures; chelate complexes are more stable than corresponding unidentate complexes.
Ambidentate ligands possess two different donor atoms, either of which can coordinate in a given complex: NO2- bonds through N (nitro) or O (nitrito); SCN- bonds through S (thiocyanato-S) or N (thiocyanato-N).
The denticity is the number of donor groups of a chelate ligand that are simultaneously bonded to a metal; high-denticity chelating ligands like EDTA form the most stable complexes, which is exploited in analytical chemistry and chelate therapy.

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Chapter 5: Coordination Compounds | Long Questions

Compare and contrast the Valence Bond Theory and Crystal Field Theory in explaining the bonding, magnetic behaviour and colour of coordination compounds.

Describe the various types of structural isomerism found in coordination compounds with one example each.

Explain the crystal field splitting in octahedral and tetrahedral coordination entities. How does the magnitude of Do determine the actual configuration of d orbitals?

Analyse the Werner theory of coordination compounds using experimental data from cobalt(III) chloride-ammonia complexes, explaining the concepts of primary and secondary valence.

Discuss the importance and applications of coordination compounds in biological systems, medicinal chemistry, analytical chemistry and metallurgy.

Discuss the nature of bonding and geometry of [Co(NH3)6]3+, [CoF6]3-, [NiCl4]2- and [Ni(CN)4]2- using Valence Bond Theory.

Describe the rules of IUPAC nomenclature for mononuclear coordination compounds and apply them to name: [Cr(NH3)3(H2O)3]Cl3, K3[Cr(C2O4)3] and [Ag(NH3)2][Ag(CN)2].

Explain geometric isomerism in square planar and octahedral coordination compounds with appropriate examples and diagrams description.

Analyse the bonding in metal carbonyls discussing the structure, hybridisation and the synergic bonding concept with examples.

Explain the types and properties of ligands in coordination chemistry, covering unidentate, didentate, polydentate, chelate and ambidentate ligands with examples.

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