From the red of blood to the green of chlorophyll — the chemistry of transition metal complexes. With interactive nomenclature helper, CFSE calculator, orbital splitting diagrams, and 20+ real-life examples.
Before diving in, here's the entire chapter on one page. From Werner's theory → Nomenclature → Isomerism → Bonding theories (VBT, CFT) → Applications.
Coordination compounds = central metal ion + ligands (Lewis bases donating electron pairs). Werner's theory distinguishes primary (ionizable) from secondary (non-ionizable) valency. Naming follows IUPAC rules. Isomerism is extensive. VBT and CFT explain bonding — CFT is better (explains colour, magnetism, and CFSE). Applications range from medicine (cisplatin) to industry (extraction, catalysis).
Alfred Werner (1893) proposed his coordination theory to explain the properties of metal-ammonia complexes like CoCl₃·6NH₃, CoCl₃·5NH₃, etc.
1. Two types of valency:
2. Every metal has a fixed secondary valency (coordination number): Common values are 2, 4, 6.
3. Secondary valencies are directional: They determine the spatial arrangement (geometry) of ligands — octahedral (CN = 6), tetrahedral or square planar (CN = 4), linear (CN = 2).
| Compound | Formula | Cl⁻ ionized | AgCl formed | Conductivity |
|---|---|---|---|---|
| Hexaamminecobalt(III) chloride | [Co(NH₃)₆]Cl₃ | 3 | 3 | High (4 ions) |
| Pentaamminechloridocobalt(III) chloride | [Co(NH₃)₅Cl]Cl₂ | 2 | 2 | Medium (3 ions) |
| Tetraamminedichloridocobalt(III) chloride | [Co(NH₃)₄Cl₂]Cl | 1 | 1 | Low (2 ions) |
| Triamminetrichloridocobalt(III) | [Co(NH₃)₃Cl₃] | 0 | 0 | None (non-electrolyte) |
Observation: As NH₃ replaces Cl⁻, the number of ionizable Cl⁻ decreases. This supports Werner's theory — some Cl⁻ are in the coordination sphere (non-ionizable) while others are outside (ionizable).
Werner's theory: primary valency (ionizable, = oxidation state) + secondary valency (non-ionizable, = coordination number). Coordination sphere fixes ligand geometry. Cl⁻ inside sphere → non-ionizable; outside → ionizable. Explains [Co(NH₃)₆]Cl₃ → 3 AgCl, [Co(NH₃)₅Cl]Cl₂ → 2 AgCl.
Coordination entity (complex): A central metal atom/ion bonded to a fixed number of other atoms, ions, or groups.
Central atom/ion: The metal (usually transition metal) that accepts electron pairs. Acts as a Lewis acid.
Ligand: An ion/molecule that donates electron pair(s) to central metal. Acts as Lewis base. Examples: H₂O, NH₃, Cl⁻, CN⁻, CO, en, EDTA.
Coordination number (CN): Number of ligand donor atoms bonded to central metal.
Coordination sphere: The central metal + all ligands directly attached, enclosed in [ ].
Coordination polyhedron: The spatial arrangement of ligand donor atoms around the metal (octahedral, tetrahedral, square planar, linear).
Denticity: Number of donor atoms per ligand.
Chelate: A complex with a polydentate ligand that forms a ring with the metal (like a crab's claw — chela).
Chelate effect: Enhanced stability of chelate complexes compared to similar monodentate complexes.
Ambidentate ligand: Can bind through two different donor atoms but only one at a time. Examples: NO₂⁻ (bind via N or O), SCN⁻ (bind via S or N).
Homo-leptic complex: All ligands identical (e.g., [Co(NH₃)₆]³⁺).
Hetero-leptic complex: Different ligands (e.g., [Co(NH₃)₄Cl₂]⁺).
| Ligand | Type | Charge | Denticity |
|---|---|---|---|
| H₂O | Aqua | Neutral | Monodentate |
| NH₃ | Ammine | Neutral | Monodentate |
| Cl⁻ | Chlorido | −1 | Monodentate |
| CN⁻ | Cyanido | −1 | Monodentate |
| CO | Carbonyl | Neutral | Monodentate |
| OH⁻ | Hydroxido | −1 | Monodentate |
| NO₂⁻ | Nitrito-N/Nitrito-O | −1 | Monodentate (ambidentate) |
| SCN⁻ | Thiocyanato-S / Thiocyanato-N | −1 | Monodentate (ambidentate) |
| en (ethylenediamine) | Ethane-1,2-diamine | Neutral | Bidentate |
| C₂O₄²⁻ | Oxalato | −2 | Bidentate |
| EDTA | Ethylenediaminetetraacetato | −4 | Hexadentate |
Central metal = Lewis acid. Ligand = Lewis base. CN = number of donor atoms. Denticity = number of donor atoms per ligand. Chelate = polydentate ring. Homo-leptic = all same ligands. Ambidentate = can bind via two different atoms.
Rule 1 — Order of ions: Cation first, anion second. Just like salts.
Example: [Co(NH₃)₆]Cl₃ → hexaamminecobalt(III) chloride
Rule 2 — Order of ligands: Ligands named in alphabetical order (ignoring prefixes like di-, tri-).
Examples: [CoCl₂(NH₃)₄]⁺ → tetraamminedichloridocobalt(III) (ammine before chlorido)
Rule 3 — Prefixes: Di-, tri-, tetra-, penta-, hexa- for multiple identical ligands. If ligand name already has a prefix, use bis-, tris-, tetrakis-, etc.
Examples: [Ni(CN)₄]²⁻ → tetracyanonickelate(II)
Example with bis: [Co(en)₃]³⁺ → tris(ethane-1,2-diamine)cobalt(III)
Rule 4 — Complex as cation: Metal name as-is + oxidation state in Roman numerals in parentheses.
Example: [Cu(NH₃)₄]²⁺ → tetraamminecopper(II) ion
Rule 5 — Complex as anion: Metal name + -ate suffix + oxidation state in Roman numerals.
Examples: [Fe(CN)₆]⁴⁻ → hexacyanoferrate(II) ion
Special cases: Cu → cuprate, Ag → argentate, Au → aurate, Fe → ferrate, Pb → plumbate, Sn → stannate, Sb → antimonate
Rule 6 — Roman numerals: Oxidation state of metal, calculated from charges.
Rule 7 — No spaces: Name is one continuous word for the complex part.
Rule 8 — Bridge ligands: If a ligand bridges two metal centers, use μ- (mu) prefix.
| Ligand | Name in Complex | Ligand | Name in Complex |
|---|---|---|---|
| H₂O | aqua | CO | carbonyl |
| NH₃ | ammine | NO | nitrosyl |
| Cl⁻ | chlorido | OH⁻ | hydroxido |
| F⁻ | fluorido | O²⁻ | oxido |
| Br⁻ | bromido | SO₄²⁻ | sulfato |
| I⁻ | iodido | NO₃⁻ | nitrato |
| CN⁻ | cyanido | NO₂⁻ | nitrito-N / nitrito-O |
| SCN⁻ | thiocyanato-S / thiocyanato-N | en | ethane-1,2-diamine |
| C₂O₄²⁻ | oxalato | EDTA | ethylenediaminetetraacetato |
Choose a well-known complex to see its IUPAC name, structure, and oxidation state.
1. [Cr(NH₃)₆]Cl₃ — hexaamminechromium(III) chloride
2. [Co(en)₃]₂(SO₄)₃ — tris(ethane-1,2-diamine)cobalt(III) sulfate
3. [Pt(NH₃)₂Cl₂] — diamminedichloridoplatinum(II)
4. K₃[Fe(CN)₆] — potassium hexacyanoferrate(III)
5. [Ni(CO)₄] — tetracarbonylnickel(0)
6. [Fe(CN)₆]⁴⁻ — hexacyanoferrate(II) ion
7. [Co(NH₃)₄Cl₂]⁺ — tetraamminedichloridocobalt(III) ion
8. [Ag(CN)₂]⁻ — dicyanidoargentate(I) ion
9. [CuCl₄]²⁻ — tetrachloridocuprate(II) ion
10. [Co(NH₃)₅(NO₂)]²⁺ — pentaamminenitrito-N-cobalt(III) ion
Mistake: "Ligands are named in order of charge or
priority."
Correction: Ligands are named in
alphabetical order based on their name in the
complex (not their formula, not their charge). Ignore prefixes
(di-, tri-) when alphabetizing. Example: tetraamminedichlorido
(ammine before chlorido, ignoring "tetra" and "di").
IUPAC rules: cation first, anion second. Alphabetical order for ligands. Prefixes: di-, tri-, etc. (or bis-, tris- for complex names). Anionic complex: metal name + -ate. Oxidation state in Roman numerals. No spaces within complex name.
I. Structural isomerism: Different connectivity.
II. Stereoisomerism: Same connectivity, different 3D arrangement.
Optical isomers: Non-superimposable mirror images (enantiomers). Rotate plane-polarized light in opposite directions.
Conditions for optical activity: The complex must be chiral — lacking a plane of symmetry or center of inversion.
Examples:
Structural isomerism: ionization, hydrate, linkage, coordination, ligand. Stereoisomerism: geometrical (cis/trans) and optical (d/l). Cis-platin is anticancer; trans-platin isn't. Optical activity needs chiral complex (no plane of symmetry).
Bonding in VBT: Metal ion uses empty d, s, p orbitals to accept electron pairs from ligands → coordinate bonds form.
Hybridization determines geometry:
Inner vs outer orbital complexes:
| Complex | Config | Hybrid | Geometry | Magnetic | Type |
|---|---|---|---|---|---|
| [Co(NH₃)₆]³⁺ | d⁶ | d²sp³ | Octahedral | Diamagnetic | Inner (low spin) |
| [CoF₆]³⁻ | d⁶ | sp³d² | Octahedral | Paramagnetic (4 unpaired) | Outer (high spin) |
| [Fe(CN)₆]³⁻ | d⁵ | d²sp³ | Octahedral | Paramagnetic (1 unpaired) | Inner (low spin) |
| [FeF₆]³⁻ | d⁵ | sp³d² | Octahedral | Paramagnetic (5 unpaired) | Outer (high spin) |
| [Ni(CN)₄]²⁻ | d⁸ | dsp² | Square planar | Diamagnetic | — |
| [NiCl₄]²⁻ | d⁸ | sp³ | Tetrahedral | Paramagnetic (2 unpaired) | — |
| [Ni(CO)₄] | d¹⁰ (Ni⁰) | sp³ | Tetrahedral | Diamagnetic | — |
VBT is a simple model but has limitations:
→ CFT (next section) overcomes these limitations.
VBT: metal uses empty d/s/p orbitals to accept ligand pairs. Hybridization: sp³ tetrahedral, dsp² square planar, d²sp³/sp³d² octahedral. Inner orbital = low spin (d²sp³) — strong-field ligands. Outer orbital = high spin (sp³d²) — weak-field ligands. Limitations: no explanation for colour, quantitative magnetism, or CFSE.
Core concept: Ligands are point charges (or dipoles). They interact with the d-electrons of the metal ion. This interaction splits the 5 d-orbitals into different energy levels.
Key facts:
Ligands arranged by their ability to cause d-orbital splitting:
I⁻ < Br⁻ < S²⁻ < SCN⁻ < Cl⁻ < NO₃⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < NH₃ < en < NO₂⁻ < CN⁻ < CO
Weak field ligands (left): Small Δ, high spin → outer orbital complexes
Strong field ligands (right): Large Δ, low spin → inner orbital complexes
Memory trick: "I Bring Some Salads, Can Not Feed Our Cats" (I⁻ Br⁻ S²⁻ SCN⁻ Cl⁻ NO₃⁻ F⁻ OH⁻ C₂O₄²⁻ H₂O)
then "Nice Neighbors Eat Nothing Cold" (NCS⁻ NH₃ en NO₂⁻ CN⁻ CO)
CFSE = energy gained by electrons in lower t2g orbitals − energy lost by electrons in higher eg orbitals
For octahedral complex with configuration t2g^x eg^y:
CFSE = [−0.4x + 0.6y] × Δ₀
For tetrahedral complex with configuration e^x t2^y:
CFSE = [−0.6x + 0.4y] × Δ_t
Also, Δ_t = (4/9) × Δ₀ — tetrahedral splitting is smaller.
Enter d-electron count and geometry — calculate CFSE and number of unpaired electrons.
Limitation: Treats ligands as point charges (doesn't consider covalent character). Molecular Orbital Theory (MOT) is more accurate.
CFT: ligand field splits d-orbitals into t2g (lower) and eg (higher) sets. Δ₀ = octahedral splitting, Δ_t = (4/9)Δ₀. Strong field → large Δ → low spin. Weak field → small Δ → high spin. CFSE = stabilization from electron distribution. Spectrochemical series ranks ligands by field strength. CFT explains colour, magnetism, and CFSE quantitatively.
1. Thermodynamic stability: Related to the magnitude of equilibrium constant (K) for formation. Larger K → more stable.
2. Kinetic stability: Related to the rate of ligand exchange. Slow exchange → kinetically inert. Fast exchange → kinetically labile.
Kinetic and thermodynamic stability are independent — a complex can be kinetically stable but thermodynamically unstable (and vice versa).
For [ML_n] formation:
M + L ⇌ ML (K₁)
ML + L ⇌ ML₂ (K₂)
...
Overall formation constant: β_n = K₁ · K₂ · ... · K_n
Larger β → more stable complex.
Stability = thermodynamic (K_f) + kinetic (rate of exchange). Chelate effect → more stable. Higher charge, smaller size → more stable. Macrocyclic ligands → extremely stable. Applications in biology (porphyrin complexes) and industry (EDTA, extraction).
Metal carbonyl: Complex with CO ligand. CO is an ambidentate ligand (can bind via C or O; usually via C).
Examples: Ni(CO)₄, Fe(CO)₅, Cr(CO)₆, Mn₂(CO)₁₀, Co₂(CO)₈
Metal oxidation state: Metals are in zero oxidation state in homoleptic carbonyls (e.g., Ni(CO)₄ has Ni⁰).
The metal-carbonyl bond has two components:
1. σ-bond (donation): CO donates lone pair from carbon to empty d-orbital of metal → σ bond (M←C).
2. π-back bond (back-donation): Filled d-orbital of metal donates electron density back into π* antibonding orbital of CO → π back-bond (M→C).
This synergic bonding strengthens the M-C bond (both directions reinforce each other).
Evidence: M-C bond is shorter than expected. C-O bond is longer than in free CO (due to back-bonding putting electron density into antibonding π*).
18-electron rule: Stable transition metal complexes have 18 valence electrons (like the next noble gas).
Formula: electrons from metal + 2 × (electrons from each ligand) = 18
Many organometallic compounds follow this rule (similar to octet rule for organic).
Metal carbonyls: metals in 0 oxidation state. Synergic bonding: σ donation from CO + π back-bonding from metal → strong M-C bond, longer C-O bond. Ni(CO)₄ tetrahedral, diamagnetic. 18-electron rule: stable complexes have 18 valence electrons. Applications: Mond process for Ni purification.
1. Medicine:
2. Analytical Chemistry:
3. Metallurgy (Extraction & Refining):
4. Photography:
5. Industrial Catalysis:
6. Electroplating:
7. Water Softening:
8. Biological Systems:
9. Pigments and Dyes:
10. Agriculture:
Cisplatin was discovered by accident in 1965 by Barnett Rosenberg, who noticed that platinum electrodes inhibited bacterial growth. It's now one of the most important anticancer drugs — and its cis isomer is effective while the trans isomer is not!
Coordination compounds have widespread applications: medicine (cisplatin, EDTA), analysis (EDTA titration, detection), metallurgy (cyanidation, Mond process), photography (fixer), catalysis (Ziegler-Natta, Wilkinson), electroplating, water softening, biological systems (Hb, chlorophyll, B₁₂), pigments, and agriculture.
"I Bring Some Salads, Can Not Feed Our Cats"
I⁻ < Br⁻ < S²⁻ < SCN⁻ < Cl⁻ < NO₃⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < NH₃ < en < NO₂⁻ < CN⁻ < CO
"Nice Neighbors Eat Nothing Cold" (NCS⁻, NH₃, en, NO₂⁻, CN⁻, CO)
"Chelate = Crab Claw" — polydentate ligand grips the metal tightly, like a crab's claw.
Bidentate en > monodentate NH₃ in stability.
"CAPA" — Cation first, Anion second, Prefixes alphabetical, Anionic complexes end in -ate
"Mono has 1, Bi has 2, Tri has 3, Tetra has 4, Hexa has 6 (EDTA)"
"Strong Field → Low Spin → Paired"
"Weak Field → High Spin → Unpaired"
Strong field ligands (CN⁻, CO) → small size, cause large Δ → electrons pair up in t2g → low spin.
Weak field ligands (I⁻, Br⁻, Cl⁻) → large size, small Δ → electrons spread → high spin.
"t2g lower, eg higher" in octahedral field.
Remember: t2g has 3 orbitals (lower), eg has 2 (higher). Octahedral stabilization comes from filling t2g.
CN=2 → Linear, CN=4 → Tetrahedral or Square Planar, CN=6 → Octahedral
"CIS = Cure, TRANS = Trash"
Cisplatin (cis form) is the anticancer drug. Transplatin (trans form) is inactive.
Fe→ferrate, Cu→cuprate, Ag→argentate, Au→aurate, Pb→plumbate, Sn→stannate, Sb→antimonate
Use "CIS = Cure, TRANS = Trash" for cisplatin. "Strong Field → Low Spin → Paired" for CFT. Spectrochemical series mnemonic helps remember field strength order.
Q1. Give the IUPAC name of [Co(NH₃)₆]Cl₃.
Cation = [Co(NH₃)₆]³⁺ — complex
Let Co be x. x + 6(0) = +3 → x = +3
6 NH₃ = hexaammine
Hexaamminecobalt(III) chloride
Q2. Write the formula for potassium hexacyanoferrate(III).
Potassium = K⁺ (cation outside)
Hexacyanoferrate(III) = anionic complex with Fe³⁺ + 6 CN⁻
[Fe(CN)₆]³⁻ + 3 K⁺ → K₃[Fe(CN)₆]
Q3. Calculate the oxidation state of the central metal in [Cu(NH₃)₄]SO₄.
Overall complex is neutral.
Cation = [Cu(NH₃)₄]ˣ; Anion = SO₄²⁻
Charge: x + (−2) = 0 → x = +2
In [Cu(NH₃)₄]²⁺: Cu + 4(0) = +2 → Cu = +2
NH₃ is neutral ligand.
Q4. Explain why cis-[Pt(NH₃)₂Cl₂] is an effective anticancer drug but trans form is not.
In cis-isomer, the two Cl⁻ ligands are adjacent (90° apart). In trans, they are opposite (180°).
Cisplatin binds to DNA via the two Cl⁻ positions — they must be adjacent (cis) to form the characteristic intrastrand crosslink.
Transplatin cannot form this specific crosslink because its Cl⁻ are too far apart.
Only cisplatin-DNA adduct stops DNA replication → cell death of cancer cells. Transplatin binds DNA differently (interstrand) → less effective.
Q5. Discuss Werner's theory of coordination compounds with examples.
Two types of valency — primary (ionizable) and secondary (non-ionizable).
Every metal has a fixed secondary valency (CN = 2, 4, 6).
Secondary valencies are directional → geometry of complex.
Formulated as [Co(NH₃)₆]Cl₃. All Cl⁻ ionizable → 3 AgCl formed with AgNO₃.
[Co(NH₃)₄Cl₂]Cl. Only 1 Cl⁻ ionizable → 1 AgCl. Confirms theory.
Q6. Using VBT, predict the geometry and magnetic behavior of [Co(NH₃)₆]³⁺ and [CoF₆]³⁻.
Co (Z=27): [Ar] 3d⁷ 4s². Co³⁺ = [Ar] 3d⁶.
NH₃ is strong-field ligand → forces pairing of d electrons.
3d⁶ pairs to 3d₆ (t2g⁶ eg⁰). Two 3d orbitals empty.
Hybridization: d²sp³ → inner orbital complex. Octahedral. Diamagnetic (no unpaired e⁻).
F⁻ is weak-field ligand → doesn't force pairing.
3d⁶ remains as 4 unpaired (t2g⁴ eg²).
Hybridization: sp³d² → outer orbital complex. Octahedral. Paramagnetic (4 unpaired).
Same metal, same ox state, different ligands → different magnetic behaviour. NH₃ strong → inner (low spin). F⁻ weak → outer (high spin).
Q7. Discuss CFT — splitting of d-orbitals in octahedral field, high-spin vs low-spin, and CFSE calculation.
Ligands are point charges. In octahedral field, 6 ligands approach along ±x, ±y, ±z axes.
dx²-y² and dz² (eg) face ligands directly → higher energy (repelled).
dxy, dyz, dxz (t2g) point between ligands → lower energy.
Δ₀ = energy difference (splitting parameter).
Weak field: Δ₀ small → pairing energy > Δ₀ → high spin (electrons spread out).
Strong field: Δ₀ large → pairing energy < Δ₀ → low spin (electrons pair up in t2g).
CFSE = (−0.4 × n_t2g + 0.6 × n_eg) × Δ₀
Example: d⁶ low spin — t2g⁶ eg⁰ → CFSE = −0.4 × 6 = −2.4 Δ₀
Explains colour (d-d transitions between split orbitals), magnetism (unpaired electron count), stability (larger CFSE → more stable), and prefers certain geometries.
Selected exercise questions from NCERT Class 12 Chemistry Chapter 5. Click each to reveal the solution.
Answer: Werner's postulates:
1. Two types of valency: Primary (ionizable — satisfied by anions, corresponds to oxidation state) and secondary (non-ionizable — satisfied by ligands, corresponds to coordination number).
2. Fixed coordination number: Most metals show CN of 2, 4, or 6.
3. Directional secondary valencies: Determine the spatial arrangement (geometry) of the complex.
Example: CoCl₃·6NH₃ has 6 NH₃ in the coordination sphere (secondary valency = 6), and 3 Cl⁻ outside (primary valency = 3). Formula: [Co(NH₃)₆]Cl₃.
(a) [Co(NH₃)₅Cl]Cl₂ → Pentaamminechloridocobalt(III) chloride
Co in +3: x + 5(0) + (−1) = +2 (charge on cation), so x = +3.
(b) [Co(en)₃]₂(SO₄)₃ → Tris(ethane-1,2-diamine)cobalt(III) sulfate
Each [Co(en)₃] has charge +3, so Co is +3.
(c) K₄[Fe(CN)₆] → Potassium hexacyanoferrate(II)
Fe in +2: x + 6(−1) = −4 → x = +2.
Stability: Resistance to dissociation in solution. Measured by formation constant K_f. Larger K_f → more stable.
Two types:
1. Thermodynamic — related to K_f (equilibrium).
2. Kinetic — related to rate of ligand exchange.
Factors:
(a) Charge on metal — higher charge → more stable (Fe³⁺ > Fe²⁺)
(b) Size of metal — smaller → more stable
(c) Nature of ligand — chelating ligands → more stable (chelate effect)
(d) Number of chelate rings — more rings → more stable
(e) Macrocyclic effect — porphyrin-type ligands → extremely stable
Two components:
1. σ donation: CO donates lone pair from carbon to empty metal d-orbital → M←C σ bond.
2. π back-bonding: Metal donates electron density from filled d-orbital back into π* antibonding orbital of CO → M→C π bond.
This is called synergic bonding — the two components reinforce each other.
Evidence: M-C bond is shorter than expected. C-O bond is longer (weaker) than in free CO (bond order decreases from 3 to ~2.5).
Example: Ni(CO)₄ — Ni⁰, tetrahedral, diamagnetic, 18-electron rule.
(a) Fe in both is +3 (d⁵) and +2 (d⁶) respectively.
[Fe(CN)₆]³⁻ — Fe³⁺ (d⁵), CN⁻ strong field → low spin d⁵ → 1 unpaired electron → weakly paramagnetic (μ = 1.73 BM).
[Fe(CN)₆]⁴⁻ — Fe²⁺ (d⁶), CN⁻ strong field → low spin d⁶ → all paired → diamagnetic.
(b) [Ni(CO)₄] — Ni⁰ (d¹⁰). sp³ hybrid → tetrahedral. Diamagnetic.
[Ni(CN)₄]²⁻ — Ni²⁺ (d⁸). CN⁻ strong field → dsp² hybrid → square planar. Diamagnetic.
Square planar preferred for d⁸ low-spin because CFSE is higher in square planar.
Two geometrical isomers:
1. fac-isomer (facial): Three Cl⁻ occupy one triangular face; three NH₃ occupy the opposite face. Non-polar. Has plane of symmetry.
2. mer-isomer (meridional): Three Cl⁻ occupy a meridian (plane passing through metal); three NH₃ also occupy a meridian. Polar. Also has plane of symmetry.
Neither fac nor mer is optically active (both have planes of symmetry).
(a) Tetrahydroxozincate(II): [Zn(OH)₄]²⁻
Zn is +2, 4 OH⁻ each −1 → charge = −4, so complex has 2− charge.
(b) Potassium tetrachloridopalladate(II): K₂[PdCl₄]
Pd is +2, 4 Cl⁻ = −4, complex charge = 2−, needs 2 K⁺.
(c) Diamminedichloridoplatinum(II): [Pt(NH₃)₂Cl₂]
Pt is +2, 2 NH₃ = 0, 2 Cl⁻ = −2 → total = 0.
(d) Potassium tetracyanonickelate(II): K₂[Ni(CN)₄]
Ni is +2, 4 CN⁻ = −4, complex charge = 2−, needs 2 K⁺.
1. Ionization isomerism: Exchange of ligand inside/outside sphere. [Co(NH₃)₅Br]SO₄ vs [Co(NH₃)₅(SO₄)]Br.
2. Hydrate isomerism: Water in/out of sphere. [Cr(H₂O)₆]Cl₃ vs [Cr(H₂O)₅Cl]Cl₂·H₂O.
3. Linkage isomerism: Ambidentate ligand binds via different atoms. [Co(NH₃)₅(NO₂)]²⁺ vs [Co(NH₃)₅(ONO)]²⁺.
4. Coordination isomerism: Ligand exchange between cation/anion complexes. [Co(NH₃)₆][Cr(CN)₆] vs [Cr(NH₃)₆][Co(CN)₆].
5. Ligand isomerism: Different ligands with same formula (e.g., n-propyl vs isopropyl).
6. Geometrical isomerism: cis/trans arrangements. cis- vs trans-[Pt(NH₃)₂Cl₂].
7. Optical isomerism: d and l forms. [Co(en)₃]³⁺ has d and l isomers.
Structural shape: All three are octahedral (CN = 6).
Magnetic behaviour:
[Co(CN)₆]³⁻ — Co³⁺ d⁶, CN⁻ strong field → low spin, 0 unpaired → diamagnetic.
[Co(NH₃)₆]³⁺ — Co³⁺ d⁶, NH₃ strong field (but not as strong as CN⁻) → low spin, 0 unpaired → diamagnetic.
[CoF₆]³⁻ — Co³⁺ d⁶, F⁻ weak field → high spin, 4 unpaired → paramagnetic.
Oxidising property: Co³⁺ is a strong oxidizer in aqueous solution (gets reduced to Co²⁺).
Co³⁺ in [Co(CN)₆]³⁻ and [Co(NH₃)₆]³⁺ is stabilized by strong-field ligands → less tendency to reduce → less oxidizing.
In [CoF₆]³⁻ (weak field), Co³⁺ is less stabilized → more tendency to reduce to Co²⁺ → more oxidizing.
(a) Chelate effect: The enhanced stability of complexes with chelating (polydentate) ligands compared to similar complexes with monodentate ligands. Due to entropy — fewer particles needed to encapsulate the metal. Example: [Ni(en)₃]²⁺ is more stable than [Ni(NH₃)₆]²⁺.
(b) Ambidentate ligand: A ligand that has two potential donor atoms but can bind through only one at a time. Examples: NO₂⁻ (via N — nitrito-N, or via O — nitrito-O), SCN⁻ (via S or via N).
(c) Coordination number: Number of donor atoms bonded directly to the central metal atom. Common values: 2 (linear), 4 (tetrahedral/square planar), 6 (octahedral).
(a) [CoCl₂(en)₂]Cl:
Structure: Co³⁺ center, 2 bidentate en ligands + 2 Cl⁻ in coordination sphere + 1 Cl⁻ outside.
IUPAC: Dichloridobis(ethane-1,2-diamine)cobalt(III) chloride
Co ox state: x + 2(0) + 2(−1) = +1 → x = +3.
(b) [Co(NH₃)₄(H₂O)Cl]Cl₂:
Structure: Co³⁺ center, 4 NH₃ + 1 H₂O + 1 Cl⁻ inside; 2 Cl⁻ outside.
IUPAC: Aquatetraamminechloridocobalt(III) chloride
Co ox state: x + 4(0) + 0 + (−1) = +2 → x = +3.
| CN = 2 | Linear | sp hybrid |
| CN = 4 | Tetrahedral | sp³ hybrid |
| CN = 4 | Square planar | dsp² hybrid |
| CN = 6 | Octahedral (inner) | d²sp³ hybrid (low spin) |
| CN = 6 | Octahedral (outer) | sp³d² hybrid (high spin) |
| Octahedral eg | dx²-y², dz² | +0.6 Δ₀ energy |
| Octahedral t2g | dxy, dyz, dxz | −0.4 Δ₀ energy |
| Tetrahedral e | dx²-y², dz² | −0.6 Δ_t |
| Tetrahedral t2 | dxy, dyz, dxz | +0.4 Δ_t |
| CN | Coordination number |
| Δ₀ | Octahedral splitting energy |
| Δ_t | Tetrahedral splitting energy |
| CFSE | Crystal field stabilization energy |
| μ | Magnetic moment (BM) |
| n | Number of unpaired electrons |
| en | Ethane-1,2-diamine (bidentate) |
| EDTA | Ethylenediaminetetraacetate (hexadentate) |
| VBT | Valence bond theory |
| CFT | Crystal field theory |
(A) Both true, R explains A. (B) Both true, R does NOT explain A. (C) A true, R false. (D) A false, R true.
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