Transition Elements (d-Block): Co, Ni, Cu and Zn Groups – Properties, General Trends and Complex Formation Explained For PGT, JEE Advanced, IIT-JAM, GATE, TGT, BHU, BITSAT, CSIR-NET, NEET Exam Aspirants

Part I — Group 9: The Cobalt Group (Co, Rh, Ir)

Cobalt (Co), Rhodium (Rh), and Iridium (Ir) constitute Group 9 of the d-block. These elements share the characteristic of having odd atomic numbers, which contributes to their relatively low natural abundance. Understanding their electronic configurations, oxidation states, and coordination chemistry is essential for virtually every competitive examination.

1.1 Electronic Configuration and Oxidation States

Element	Symbol	Electronic Configuration	Key Oxidation States
Cobalt	Co	[Ar] 3d⁷ 4s²	−I, 0, I, II, III, (IV)
Rhodium	Rh	[Kr] 4d⁸ 5s¹	−I, 0, I, II, III, IV, (VI)
Iridium	Ir	[Xe] 4f¹⁴ 5d⁹	−I, 0, (I), (II), III, IV, (V), (VI)

🎯 Exam Tip: The most important oxidation states for Co are +II and +III. For Rh and Ir, it is +III. A possible Co(+V) state has been disproved; even Co(+IV) is unstable. Always remember: Co(+II) is stable as simple hydrated ions; Co(+III) is stable only in complexes.

1.2 Occurrence, Extraction, and Uses

All three elements have odd atomic numbers and consequently low crustal abundance. Co occurs at ~23–30 ppm by weight; Rh and Ir are among the rarest stable elements on Earth.

Key Cobalt Ores

Cobaltite: CoAsS
Smaltite: CoAs₂
Linnaeite: Co₃S₄

These ores always occur with Ni ores and often with Cu or Pb ores. Co is extracted as a by-product of these metals. In 1992, world Co ore production was 30,100 tonnes of contained metal; major producers were Zaire (22%), Canada (17%), Zambia and the Soviet Union (15% each), and Australia (9%).

Extraction Steps

Ore is roasted → mixture of oxides (speisses); As₂O₁₀ and SO₂ are recovered as by-products.
Oxides treated with H₂SO₄ → Co and Ni dissolve; Fe (impurity) is precipitated using lime as Fe₂O₃·H₂O.
NaOCl added → precipitates Co(OH)₃.
Hydroxide ignited → Co₃O₄, then reduced with H₂ or charcoal → Co metal.

Co₃O₄ + 4H₂ → 3Co + 4H₂O

Uses of Cobalt

~1/3 of production: High-temperature alloys with steel for gas turbine engines and high-speed cutting tools (e.g., Stellite: 50% Co, 27% Cr, 12% W, 5% Fe, 2.5% C).
~1/3 of production: Pigments for ceramic, glass, and paint industries. CoO historically used as the famous blue pigment ("smalt"). Blue cobalt glass absorbs the yellow Na flame, allowing the K violet colour to be seen — basis of the flame test for potassium.
~1/5 of production: Ferromagnetic alloys; Alnico (Al, Ni, Co) makes permanent magnets 20–30× stronger than Fe magnets.
Co salts of fatty acids (linseed oil) used as "driers" to speed paint drying.
Co is an essential trace element in fertile soil and present in some enzymes and in Vitamin B₁₂.

📌 JEE/NEET Fact: The artificial isotope ⁶⁰Co is radioactive and undergoes β decay (t_½ = 5.2 years), simultaneously emitting intense high-energy γ radiation used in radiotherapy of cancerous tumours. It is prepared by neutron irradiation of ⁵⁹Co (the only naturally occurring isotope) in a nuclear reactor:

⁵⁹₂₇Co → ⁶⁰₂₈Ni + ⁰₋₁e + ν + γ

1.3 General Properties of Cobalt

Co resembles Fe; very tough, harder, and higher tensile strength than steel.
Bluish-white, lustrous, ferromagnetic — loses magnetism above 1000°C.
Relatively unreactive — does NOT react with H₂O, H₂, or N₂ at room temperature.
Reacts with steam → CoO; burns at white heat → Co₃O₄.
Dissolves slowly in dilute acids; passive in concentrated HNO₃.
Combines readily with halogens and, at elevated temperatures, with S, C, P, As, Sb, Sn.

1.4 Oxidation States in Detail

Lower States: −I and 0

These states occur only with π-bonding ligands such as CO, PF₃, NO, and CN⁻.

−I state: Found in tetrahedral complexes:

[Co(CO)₄]⁻, [Rh(CO)₄]⁻, [Co(CO)₃NO], K[Ir(PF₃)₄]

0 state: Exemplified by the famous dicobalt octacarbonyl:

Co₂(CO)₈ — exists in two isomeric forms, both with metal–metal bonding

Co₄(CO)₁₂, Rh₄(CO)₁₂, Ir₄(CO)₁₂ — all contain M–M bonds, cluster of 4 metal atoms

Fig. 1 — Schematic of the bridged isomer of Co₂(CO)₈ showing terminal (red) and bridging (orange) CO groups with a direct Co–Co bond.

+I State

Co(+I) is better developed for Co than for any other first-row transition metal except Cu. Compounds are usually made by reducing CoCl₂ with Zn or N₂H₄ in the presence of a π-bonding ligand. Structures are typically trigonal bipyramidal or tetrahedral.

An important reaction — direct uptake of N₂ gas to form a dinitrogen complex:

Co(acac)₃ + N₂ + 3PPh₃ → [Co^I(H)(N₂)(PPh₃)₃]

(acac = acetylacetonate; the complex has a trigonal bipyramidal structure)

🎯 Key Concept (Dinitrogen Complexes): In [Co^I(H)(N₂)(PPh₃)₃], the N≡N bond length is 1.11 Å vs 1.098 Å in free N₂. The bond is almost unchanged because σ-bonding from N to Co is extremely weak — the N–Co bond is mainly due to π back-bonding from Co to N. This is analogous to [Ru(NH₃)₅N₂]²⁺.

There is also extensive chemistry of Rh(+I) and Ir(+I) with π-bonding ligands (CO, phosphines PR₃, alkenes). These normally adopt square planar or trigonal bipyramidal structures.

Famous examples:

Vaska's Compound — trans-[Ir(Cl)(CO)(PPh₃)₂]: yellow, readily absorbs O₂ to become orange. The reversible oxygenation has been studied as a model for haemoglobin oxygen-carrying.
Wilkinson's Catalyst — [Rh(Cl)(PPh₃)₃]: red-violet solid, square planar. Highly effective for selective hydrogenation of organic molecules at room temperature and pressure. Alk-1-enes (terminal alkenes) are hydrogenated; internal double bonds are unaffected. Crucial in the pharmaceutical industry.

📌 GATE/CSIR-NET: Wilkinson's catalyst undergoes oxidative addition — a square planar (+I) complex adds a neutral molecule to give an octahedral (+III) complex:
[Ir^I(Cl)(CO)(PPh₃)₂] + HCl → [Ir^III(Cl)₂(CO)(PPh₃)₂H]
Oxidative addition requires: (a) non-bonding d electrons on metal, (b) two vacant coordination sites.

+I State: The OXO Process (Industrial Importance!)

Wilkinson's catalyst and Co compound HCo^I(CO)₄ are catalysts in the OXO process — one of the most industrially significant reactions in the world:

RCH=CH₂ + HCo(CO)₄ → RCH₂CH₂Co(CO)₄

RCH₂CH₂Co(CO)₄ + CO → RCH₂CH₂CO·Co(CO)₄

RCH₂CH₂CO·Co(CO)₄ + H₂ → RCH₂CH₂CHO + HCo(CO)₄

CO and H₂ are added to an alkene forming an aldehyde. Temperature: 150°C, Pressure: 200 atm. About 3 million tonnes of C₆–C₉ alcohols are produced annually — used to make polyvinyl chloride and detergents.

Also: Acetic acid synthesis from methanol:

CH₃OH + CO → CH₃COOH

(catalysed by [Rh(Cl)(CO)(PPh₃)₂] or [Rh(Cl)(CO)₂]₂ in presence of CH₃I, I₂, or HI)

+II State

This is the most important state for simple compounds of Co, though +III dominates in complexes. Rh(+II) and Ir(+II) are of only minor importance.

Simple Co(+II) compounds known: CoO, Co(OH)₂, CoS, CoF₂, CoCl₂, CoBr₂, CoI₂, CoSO₄, Co(NO₃)₂, CoCO₃. Hydrated salts are all pink or red and contain the hexahydrate ion [Co(H₂O)₆]²⁺.

CoCl₂ — The Classic Humidity Indicator

CoCl₂ is used as a test for water (cobalt chloride paper) and as an indicator in silica gel:

[Co(H₂O)₆]²⁺ ⇌ [Co(H₂O)₄]²⁺ + 2H₂O

pink (octahedral) blue (tetrahedral)

When silica gel indicator is blue → drying agent is effective; when it turns pink → needs to be replaced/regenerated.

Fig. 2 — [Co(H₂O)₆]²⁺ (pink, octahedral) ⇌ [CoCl₄]²⁻ (blue, tetrahedral). Co(+II) tetrahedral complexes have more intense colours due to absence of a centre of symmetry (Laporte rule).

🎯 Why Tetrahedral Co(II) complexes are MORE intensely coloured than octahedral ones? A tetrahedron lacks a centre of symmetry (no inversion centre), so it easily overcomes the Laporte selection rule (Δl = 1). Octahedral complexes must rely on asymmetric vibrations to destroy the centre of symmetry. Hence tetrahedral Co(II) complexes show much stronger absorption bands.

+III State — The Most Important for Complexes

Co(+III) forms more coordination complexes than any other element. This is the most extensively studied oxidation state in coordination chemistry — much of our knowledge of stereochemistry and isomerism of octahedral complexes comes from Werner's studies of Co(+III) complexes in the 1890s.

Why so many Co(+III) complexes?

Co³⁺ has a d⁶ configuration. In an octahedral field with strong ligands, the arrangement is (t_2g)⁶(e_g)⁰ — giving very large crystal field stabilization energy (CFSE). The complexes are:

Diamagnetic (all spins paired)
Inert — ligand exchange reactions are very slow (kinetically stable)
Octahedral in almost all cases

⚠️ Contrast: Simple Co³⁺ ions are strongly oxidizing (unstable in simple compounds — they oxidize water). Co(+III) is stable ONLY in complexes. In contrast, Co(+II) ion [Co(H₂O)₆]²⁺ is stable in water. Many Co(+II) complexes are readily oxidized to Co(+III) complexes in the presence of air/O₂.

Common Co(+III) Complexes with Colors:

Complex	Colour
[Co(NH₃)₆]³⁺	Yellow
[Co(NH₃)₅H₂O]³⁺	Pink
[Co(NH₃)₅Cl]²⁺	Purple
[Co(NH₃)₄CO₃]⁺	Purple
[Co(NH₃)₃(NO₂)₃]	Yellow
[Co(CN)₆]³⁻	Violet
[Co(NO₂)₆]³⁻	Orange

[Co(CN)₆]³⁻ is extremely stable — not decomposed even by alkalis. The CN⁻ ligands are very firmly bonded through π back-bonding, and the CFSE is very high. The complex is claimed to be non-toxic.

Peroxo and Superoxo Cobalt Complexes

[Co^II(CN)₅]³⁻ + O₂ (air) → K₆[(CN)₅Co^III–O–O–Co^III(CN)₅]

Brown-coloured peroxo complex (O–O bond = 1.45 Å, cf. 1.48 Å in H₂O₂)

Further oxidation by Br₂: O–O bond shortens to 1.26 Å (superoxo, bond order 1.5)

Cobalt and Vitamin B₁₂ — Biology Meets Inorganic Chemistry

Vitamin B₁₂ (cobalamin) is one of the most important Co complexes biologically. Isolated from liver — large amounts of raw liver cured pernicious anaemia. Now B₁₂ injections are used. It serves as a coenzyme (prosthetic group) tightly bound to several enzymes in the body.

Structure: Contains a Co(+III) ion at the centre of a corrin ring system (similar to haemoglobin's porphyrin but the corrin ring is less conjugated and rings A and D are joined directly). The Co atom is bonded to four ring N atoms. The fifth position is occupied by another N from a side chain (5,6-dimethylbenzimidazole). The sixth position (active site) is occupied by:

CN⁻ in cyanocobalamin (isolated form — CN is introduced during isolation)
OH in hydroxocobalamin
H₂O in aquocobalamin
CH₃ in methylcobalamin (biologically active)
Adenosine in adenosylcobalamin (biologically active)

📌 Important for CSIR-NET/IIT-JAM: The cobalamins can be reduced from Co(+III) → Co(+II) → Co(+I) in neutral or alkaline solution. The Co(+I) complex is strongly reducing and five-coordinate — the site usually occupied by CN or OH is vacant. Methylcobalamin is important in the metabolism of methane-producing bacteria — these bacteria can also transfer CH₃ to metals like Pt^II, Au, and Hg^II, the latter posing a significant ecological threat (conversion of elemental Hg/inorganic Hg salts into toxic methylmercury CH₃Hg⁺ and dimethylmercury (CH₃)₂Hg in lake sediments).

+IV State

This is the highest oxidation state normally obtained for cobalt. Oxidation of alkaline Co²⁺ solutions gives a product thought to be hydrated CoO₂. A complex Ba₂Co^IVO₄ has been reported.

Part II — Group 10: The Nickel Group (Ni, Pd, Pt)

2.1 Electronic Configuration and Oxidation States

Element	Configuration	Most Stable State	Key Other States
Nickel (Ni)	[Ar] 3d⁸ 4s²	+II	−I, 0, (I), III, (IV)
Palladium (Pd)	[Kr] 4d¹⁰	+II	0, (I), IV
Platinum (Pt)	[Xe] 4f¹⁴ 5d⁹ 6s¹	+II, +IV	0, (I), (III), V, VI

🎯 Key Trend: Ni is predominantly divalent (+II) in simple compounds and complexes. The higher oxidation states of Ni, Pd, Pt are all unstable. The two most prolific complex-forming elements are Co and Pt.

2.2 Occurrence and Extraction

Nickel is the 22nd most abundant element. Commercially important Ni ores include sulphides (pentlandite (Fe,Ni)₉S₈ — the most important), always with Fe or Cu sulphides and alluvial deposits of silicates and oxides/hydroxides.

The Mond Process — A Masterpiece of Chemical Engineering

Ni + 4CO →^50°C Ni(CO)₄ →^230°C Ni + 4CO

(Mond Process — patented by L. Mond, used in South Wales 1899–1970)

NiO and water gas (H₂ and CO) are warmed under atmospheric pressure to 50°C. H₂ reduces NiO to Ni, which reacts with CO to form volatile nickel tetracarbonyl Ni(CO)₄ (highly toxic and flammable!). Impurities remain solid. The gas is then heated to 230°C — it decomposes to give pure Ni metal and CO is recycled.

Uses of Nickel

~2/3 of production: Ferrous and non-ferrous alloys — stainless steel (12–15% Ni), Monel metal (68% Ni, 32% Cu), Nimonic series (75% Ni, Cr, Co, Al, Ti) for gas turbine/jet engines, Hastelloy C (corrosion resistance), Nichrome (60% Ni, 40% Cr — electric radiator wire).
Ni/Fe storage batteries (rapid charging).
Raney Nickel (finely divided Ni) — catalyst for hydrogenation of hexamethylenediamine, anthraquinone → anthraquinol (H₂O₂ production).
Electroplating protective coat on steel.

Platinum and Palladium

Both are rare and expensive, but more abundant than other platinum group metals. Obtained as anode sludge from electrolytic refining of Ni. Uses:

Pt: Jewellery (since several centuries BC), catalyst in oil refining (hydrocarbon reforming), three-way catalytic converters with Pd and Rh (requires lead-free petrol!), Pt crucibles in laboratory.
Pd: Wacker process (PdCl₂ catalyses C₂H₄ → CH₃CHO), hydrogenation catalyst, absorbs 935× its own volume of H₂ at red heat.

2.3 Low Valency States

Ni(−I): Found in carbonyl anion [Ni₂(CO)₆]²⁻.

Ni(0): [Ni⁰(CO)₄] — tetrahedral, volatile, very poisonous, easily oxidized, flammable. Perhaps the best-known metal carbonyl. Stability much lower than carbonyls in earlier transition metal groups.

2K₂[Pt^IICl₄] + N₂H₄ + 8PPh₃ →^EtOH 2[Pt⁰(PPh₃)₄] + 4KCl + 4HCl + N₂

2.4 The +II State — Square Planar Geometry

The +II state is very important for all three elements. A wide variety of simple Ni²⁺ compounds exist.

Nickel(II) Chemistry

The hydrated ion [Ni(H₂O)₆]²⁺ gives rise to the green colour characteristic of many hydrated Ni salts. Many anhydrous Ni salts are yellow. Ni(+II) forms octahedral, square planar, and a few tetrahedral complexes.

🎯 Why does Ni(II) (d⁸) prefer square planar over tetrahedral? In an octahedral field with d⁸, the arrangement is (t_2g)⁶(e_g)². In a strong field, electrons pair in the lower t_2g and one e_g orbital, approaching a square planar geometry. The CFSE gain from going square planar is much larger for d⁸ than for other configurations, especially for heavier metals (Pd²⁺, Pt²⁺ are exclusively square planar).

Fig. 3 — d-orbital energy level splitting for d⁸ Ni²⁺ in octahedral (left) and square planar (right) environments. In strong-field square planar, all 8 electrons are paired; in weak-field octahedral, 2 electrons are unpaired (paramagnetic, μ ≈ 2.8–3.4 BM).

The Dimethylglyoxime (DMG) Test for Nickel

When Ni²⁺ solution is treated with dimethylglyoxime in slightly ammoniacal solution, a bright red precipitate forms — the basis of the most sensitive gravimetric detection of Ni.

Ni²⁺ + 2 dmgH₂ → [Ni(dmgH)₂] + 2H⁺

(dmgH₂ = dimethylglyoxime; the red complex has square planar structure)

The complex is stabilized by: (1) two five-membered chelate rings, (2) internal hydrogen bonding (N–H···O), (3) stacking of planar molecules (Ni–Ni ≈ 3.25 Å)

📌 IIT-JAM/GATE: The Ni–Ni distance in solid [Ni(dmgH)₂] is 3.25 Å. This was one of the earliest examples of metal-to-metal interaction through stack bonding (dz² overlap). In the solid, Ni should be regarded as octahedrally coordinated rather than square planar due to this intermolecular interaction.

Pd(+II) and Pt(+II)

Both Pd²⁺ and Pt²⁺ are exclusively square planar and diamagnetic. The Pd²⁺ ion has d⁸ configuration and is paramagnetic as [Pd(H₂O)₄]²⁺ — but this complex is spin paired and presumed square planar.

Wacker Process (Industrial!)

C₂H₄ + PdCl₂ + H₂O → CH₃CHO + Pd + 2HCl

Pd + 2CuCl₂ → PdCl₂ + 2CuCl

2CuCl + 2HCl + ½O₂ → 2CuCl₂ + H₂O

———————————————————————————

H₂C=CH₂ + ½O₂ → CH₃CHO (overall reaction)

Zeise's Salt — The First Alkene Complex

K[Pt(η²-C₂H₄)(Cl)₃]·H₂O — Known since 1825! The [Pt(C₂H₄)(Cl)₃]⁻ ion is essentially square planar with Cl at three corners and the ethylene molecule (H₂C=CH₂) perpendicular to the PtCl₃ plane. The C=C distance in the complex is 1.375 Å vs 1.337 Å in free ethene — the double bond is only slightly lengthened.

Bonding (Dewar-Chatt-Duncanson model, 1951-1953):

σ bond: Filled π orbital of ethene donates to an empty hybrid orbital on the metal.
π back-donation: Filled metal d orbital overlaps with empty antibonding (π*) orbital of ethene.

Cisplatin — Anticancer Drug

The cis isomer of [Pt(NH₃)₂Cl₂] (cisplatin) is an important anti-cancer drug. The trans isomer is ineffective.

cis-[Pt(NH₃)₂Cl₂] injected into bloodstream

→ reactive Cl groups lost → Pt bonds to N atoms in guanosine (part of DNA)

→ bridges between two guanosine units → upsets normal DNA reproduction

→ rapidly dividing cancer cells are attacked

⚠️ Critical balance: Cisplatin is highly toxic — it also attacks bone marrow cells (red and white blood cells) and testes cells. A careful balance must be maintained between giving enough to kill tumours and leaving sufficient white blood cells to protect against bacterial/viral attack.

2.5 Horizontal Comparisons: Fe, Co, Ni vs Ru, Rh, Pd vs Os, Ir, Pt

The ferrous metals (Fe, Co, Ni) and platinum metals differ as follows:

Ferrous metals are much more reactive; reactivity decreases Fe → Co → Ni.
Maximum oxidation states: Fe(+VI), Co(+IV), Ni(+IV) — rarely exceeded +III in practice.
Platinum metals are much more noble, little affected by acids.
Reactivity of platinum metals increases: Ru→Rh→Pd and Os→Ir→Pt (opposite trend to ferrous).
Both groups show: coloured compounds, variable valency, catalytic properties, large number of coordination compounds.
Key differences going down: increased stability of higher oxidation states, disappearance of simple ionic forms, increased nobility.

Part III — Group 11: The Copper Group / Coinage Metals (Cu, Ag, Au)

3.1 Introduction and Electronic Configuration

Element	Configuration	Most Stable State	Other States
Copper (Cu)	[Ar] 3d¹⁰ 4s¹	+II	+I, (+III)
Silver (Ag)	[Kr] 4d¹⁰ 5s¹	+I	(+II), (+III)
Gold (Au)	[Xe] 4f¹⁴ 5d¹⁰ 6s¹	+III	+I, V

All three have one s electron outside a completed d shell. The d electrons are involved in metallic bonding, making melting points and enthalpies of sublimation much higher than for Group 1 metals. The higher ionization energies and enthalpies of sublimation make Cu, Ag, and Au much less reactive — they show noble character.

🎯 Why Cu is less reactive than K (despite similar outer electronic structure)? Cu has a much higher melting point (higher enthalpy of sublimation due to d-electron participation in metallic bonding). The nuclear charge of Cu holds orbital electrons more tightly → higher ionization energy. The enthalpy of hydration is NOT large enough to offset these, so K is much more reactive.

3.2 Standard Reduction Potentials

Fig. 4 — Latimer reduction potential diagrams for Cu, Ag, and Au in acid solution. Note: Cu⁺ and Au⁺ disproportionate in water (marked with * and their E° pattern). Ag⁺ is stable.

📌 JEE Advanced — Disproportionation: Cu⁺ and Au⁺ disproportionate in aqueous solution:
2Cu⁺ ⇌ Cu²⁺ + Cu K = [Cu²⁺]/[Cu⁺]² = 1.6 × 10⁶ (large → equilibrium strongly to right)
3Au⁺ ⇌ Au³⁺ + 2Au K = [Au³⁺]/[Au⁺]³ = 1 × 10¹⁰

Therefore Cu⁺ exists in solution for less than a second! Only Cu(+I) compounds that are insoluble (CuCl, CuCN, CuSCN) or form stable complexes are stable to disproportionation.

3.3 The +I State

Copper(I)

Cu⁺ has a d¹⁰ configuration. Simple compounds and complexes are typically diamagnetic and colourless. Exceptions exist where colour arises from charge-transfer bands: Cu₂O is yellow or red, Cu₂CO₃ is yellow, CuI is brown.

Fehling's Test: Cu²⁺ is reduced to Cu₂O (brick-red precipitate) by mild reducing agents (reducing sugars like glucose). This is the chemical basis of Fehling's test:

Fehling's A = CuSO₄ + sodium potassium tartrate in H₂O (deep blue solution)
Fehling's B = NaOH
Mix immediately before adding the sugar and warming → yellow/red precipitate of Cu₂O if reducing sugar present

Silver(I) — Most Important State for Ag

Practically all simple ionic Ag compounds contain Ag⁺. Most Ag^I salts are insoluble in water. Soluble exceptions: AgNO₃, AgF, AgClO₄.

AgNO₃: Most important silver salt — used in qualitative analysis to precipitate Cl⁻, Br⁻, I⁻ as AgCl (white), AgBr (pale yellow), AgI (yellow). Presence confirmed by solubility in NH₄OH: AgCl soluble in dilute NH₄OH, AgBr soluble in strong 0.880 ammonia, AgI insoluble even in 0.880 ammonia.

Gold(I) and Gold Drugs

Au(+I) is less stable and known mainly as oxide Au₂O. It exists in linear complexes [NC→Au←CN]⁻, [Cl→Au←Cl]⁻, [R₃P→Au←Cl].

Au(+I) drugs are used to treat rheumatoid arthritis. The drugs are thought to be linear complexes of type [RS→Au←SR] or [R₃P→Au←PR₃].

3.4 The +II State (Copper)

This is the most stable and important state for Cu. Cu²⁺ has d⁹ configuration — one unpaired electron → paramagnetic. Compounds are typically coloured due to d–d spectra.

Jahn-Teller Distortion in Cu(II)

The octahedral arrangement causes crystal field splitting of d orbitals. With d⁹ configuration, the e_g level has (t_2g)⁶(e_g)³. The e_g is not symmetrically filled → Jahn-Teller distortion occurs → tetragonally distorted octahedron with 4 short bonds and 2 long trans bonds.

Fig. 5 — Jahn-Teller distorted octahedral Cu²⁺ complex. The d⁹ configuration has (e_g)³ which is asymmetrically occupied → the two axial bonds elongate. The d_x²−y² orbital has one unpaired electron (shown in red).

The Copper(II) Acetate Dimer — Cu₂(CH₃COO)₄·2H₂O

Copper(II) acetate is dimeric and hydrated. The structure consists of two Cu atoms each with a roughly octahedral structure. Four acetate groups act as bridging ligands between the two Cu atoms. The fifth coordination position around each Cu is occupied by O from water. The other Cu atom occupies the sixth position. The Cu–Cu distance is 2.64 Å.

⚠️ Key Distinction: Cu does NOT form a metal-metal bond (Cu–Cu = 2.64 Å, significantly longer than Cu–Cu = 2.55 Å in metallic copper). Compare with Cr, Mo, Rh, Ru which DO form M–M bonds in similar carboxylate structures. The Cu^II ions have d⁹ — with one unpaired electron each, if bonded, electrons would be paired → diamagnetic. But measured μ at 25°C is 1.4 BM/Cu, rather than the spin-only 1.73 BM. This suggests weak antiferromagnetic coupling (interaction between unpaired spins) via lateral overlap of 3d_x²−y² orbitals — sometimes called δ bonding.

3.5 Photography — The Chemistry of Silver Halides

Silver halides (AgBr primarily, AgI for fast emulsions, AgCl sometimes) are used as light-sensitive materials in photographic film.

Exposure: Light excites an electron from the halide ion (Br⁻) into the conduction band. This electron moves to the surface of the silver bromide grain and reduces Ag⁺ → Ag metal (10–50 atoms). The film now contains a latent image — invisible to the eye.
Development: A mild reducing agent (quinol) preferentially reduces more AgBr to Ag metal where Ag atoms already exist. This intensifies/amplifies the latent image.
Fixing: Unchanged AgBr is dissolved with sodium thiosulphate (hypo):

AgBr + 2Na₂S₂O₃ → Na₃[Ag(S₂O₃)₂] + NaBr
Printing: Light is passed through the negative onto paper coated with AgBr emulsion → developed and fixed as before.

3.6 The +III State

Gold(III) — Most Common State for Au

Au(+III) is the most important state for gold, unlike Cu or Ag. Au³⁺ has d⁸ configuration (like Pt²⁺), forms square planar complexes, and these decompose to the metal on heating.

Au + HNO₃ + HCl → H₃O⁺[AuCl₄]⁻·3H₂O → AuCl₃

[AuCl₄]⁻ + OH⁻ → Au(OH)₃ →^dehydrate Au₂O₃ →^150°C Au + Au₂O + O₂

Liquid gold (used in decoration of picture frames, glass, ceramic ornaments) is a chloro complex of Au^III dissolved in an organic solvent. When heated, it decomposes to leave a thin film of metallic gold.

Part IV — Group 12: The Zinc Group (Zn, Cd, Hg)

4.1 Introduction — Are They Really Transition Metals?

Zinc, Cadmium, and Mercury all have a d¹⁰s² electronic arrangement and typically form M²⁺ ions. Because they have a complete d shell, they do NOT behave as typical transition metals:

Zn and Cd do not show variable valency.
They have d¹⁰ configuration → cannot produce d–d spectra. Many compounds are white/colourless. (Some Hg(+II) and Cd(+II) compounds are coloured — due to charge transfer from ligands to metal.)
The metals are relatively soft compared to other transition metals — d electrons do not participate in metallic bonding.
Melting and boiling points are very low.

Element	Config.	Oxidation State	Covalent Radius (Å)	Melting Pt (°C)
Zinc (Zn)	[Ar] 3d¹⁰ 4s²	+II only	1.25	420
Cadmium (Cd)	[Kr] 4d¹⁰ 5s²	+II only	1.41	321
Mercury (Hg)	[Xe] 4f¹⁴ 5d¹⁰ 6s²	+I, +II	1.44	−39 (liquid!)

4.2 Occurrence and Extraction

Zinc

24th most abundant element (132 ppm by weight).
Ore: ZnS (sphaelerite/zinc blende in USA; zinc blende in Europe) — structure like diamond, formula may be (Zn,Fe)S.
Also: ZnCO₃ (smithsonite/calamine), Zn₄(OH)₂(Si₂O₇)·H₂O (hemimorphite).

Extraction of Zinc — Two Routes

Route 1 (Pyrometallurgical):

2ZnS + 3O₂ → 2ZnO + 2SO₂ (roasting)

ZnO + CO ⇌ Zn + CO₂ (at 1200°C — equilibrium pushed right by shock cooling)

Route 2 (Electrolytic — expensive but high purity):

ZnS + 2O₂ → ZnSO₄ (at lower temp.)

ZnSO₄ solution → electrolysis → pure Zn

Mercury

Mined as bright red ore cinnabar (HgS) mainly in USSR, Spain, Mexico, and Algeria.
Extraction:

HgS + O₂ →^600°C Hg + SO₂

OR: 4HgS + CaO → 4Hg + CaSO₄ + 3CaS (with quicklime)

Hg vapour is condensed and collected.

4.3 Oxidation States

Mercury(I) — The Unique Dinuclear Ion

Mercury is unique in the Hg(+I) state in that it consists of two directly linked metal atoms. The mercury(I) ion has the structure [Hg—Hg]²⁺, not Hg⁺.

🎯 Why is Mercury(I) written as Hg₂²⁺ and NOT Hg⁺? Multiple lines of evidence:
(1) X-ray diffraction shows a linear Cl–Hg–Hg–Cl structure in Hg^ICl (not alternating Hg⁺ and Cl⁻).
(2) Magnetic properties: Hg⁺ would have one unpaired electron (paramagnetic). But all Hg(I) compounds are diamagnetic — in [Hg–Hg]²⁺ electrons are all paired.
(3) Raman spectra: An extra line at 171.7 cm⁻¹ attributed to the Hg–Hg bond stretching (Raman active, IR inactive — this is a homonuclear bond).
(4) Cryoscopic measurements: Depression of freezing point fits Hg₂²⁺ + 2NO₃⁻ (3 particles), NOT Hg⁺ + NO₃⁻ (2 particles).
(5) EMF of concentration cell: Calculated n = 2, consistent with Hg₂²⁺ carrying two positive charges.

Fig. 6 — Linear structure of Hg₂Cl₂ (calomel). The Hg–Hg bond (2.53 Å) is a covalent bond formed by overlap of 6s orbitals. The Hg–Cl distance is 2.25 Å. The two Hg atoms are bonded together, forming the dinuclear [Hg–Hg]²⁺ unit, NOT discrete Hg⁺ ions.

Disproportionation of Hg(I)

Addition of OH⁻ or S²⁻ ions removes Hg²⁺ from equilibrium, causing Hg(I) to disproportionate completely:

Hg²⁺ + Hg ⇌ Hg₂²⁺ E° = +0.13 V K ≈ 170

(solutions of Hg(I) contain one Hg²⁺ for every 170 Hg₂²⁺)

If OH⁻ or S²⁻ added → HgO or HgS precipitate Hg²⁺ → equilibrium shifts left → complete disproportionation:

Hg₂²⁺ → Hg²⁺ + Hg

This explains the absence of Hg(I) hydroxides, sulphides, and cyanides — they disproportionate instantly.

4.4 General Properties

Why Mercury is Liquid at Room Temperature

Mercury is the only metal which is liquid at room temperature (melting point −39°C). The reason:

The very high first ionization energy of Hg (due to: filled 4f shell poor shielding → contracted 6s orbital, plus relativistic effects) makes it difficult for electrons to participate in metallic bonding.
Hg has an appreciable vapour pressure at room temperature → exposed mercury surfaces should always be covered to prevent vaporization and poisoning.
The Hg vapour is unusual as it is monatomic like noble gases.

Reactivity Trends

Zn and Cd dissolve in dilute non-oxidizing acids, liberating H₂; Hg does not.
All three react with concentrated HNO₃ and H₂SO₄.
Zn is amphoteric — dissolves in alkalis forming zincates:

Zn + 2NaOH + H₂O → Na₂[Zn(OH)₄] + H₂

(or: Na₂[Zn(OH)₄] is formulated as Na₂ZnO₂·2H₂O or Na[Zn(OH)₃·H₂O])

4.5 Oxides and Hydroxides

ZnO — Industrially the Most Important

White when cold, yellow on heating (returns to white on cooling) — colour due to defects in solid structure (see defect chemistry).
ZnO is amphoteric: dissolves in both acids and strong alkalis.
Main use: production of rubber (shortens vulcanization time); also white pigment in paint (lower refractive index than TiO₂ but absorbs UV light and re-emits as white light).
World production (1991): 366,500 tonnes.

4.6 Dihalides — Structural Diversity

Compound	Character	Melting Point	Structure
ZnF₂	Most ionic of Zn halides	872°C	Rutile (TiO₂) — Zn octahedrally surrounded by 6F⁻
ZnCl₂	Partly covalent	283°C	Close-packed Cl with Zn in tetrahedral holes; highly soluble (432 g/100 g H₂O at 25°C)
HgCl₂	Covalent (linear molecules)	276°C	Linear Cl–Hg–Cl molecules, Hg–Cl = 2.25 Å
HgF₂	Ionic	645°C (decomp.)	Fluorite (CaF₂) structure; hydrolyzed by water

4.7 Complexes of Zn, Cd, and Hg

Zn²⁺ and Cd²⁺ form complexes with O, N, S donor ligands and halide ions. Hg(+II) forms complexes with N, P, S donor ligands but is reluctant to bond to O. No complexes are known with π-bonding ligands such as CO or alkenes. Since elements have d¹⁰ configuration, there is no crystal field stabilization energy.

🎯 Nessler's Reagent: The complex K₂[HgI₄] (Nessler's reagent) gives a yellow colour or brown precipitate with concentrations of NH₃ as low as 1 part per million. This test is used on drinking water to detect NH₄⁺ ions which may indicate sewage contamination.

4.8 Organometallic Chemistry

Zinc and Cadmium Alkyls — Historical Significance

The first useful organometallic compounds were zinc alkyls ZnR₂ and alkyl zinc halides RZnX, prepared by Sir Edward Frankland in 1849. They were originally used in organic synthesis before Grignard reagents were discovered.

EtI + Zn →^{inert atm, N₂} EtZnI →^heat Et₂Zn + ZnI₂

Mercury Alkyls — Environmental Hazard!

Organomercury compounds R₂Hg and RHgX are extremely poisonous and have caused several environmental disasters:

Minamata Disease (Japan, 1952): 52 people died from eating fish contaminated by mercury from a factory using Hg^II salts to catalyse the production of ethanal from ethyne. HgCl₂ was converted by anaerobic bacteria to MeHgSMe, concentrated in the food chain → fish → humans. Outbreaks also in 1960 and 1965.
Iraq (1956, 1960): Corn seeds treated with organomercurial fungicide were eaten as food — many deaths.
Amazon River contamination: Mercury used to extract gold in Brazil has poisoned the Amazon river ecosystem.

⚠️ Alkyl vs Aryl mercury toxicity: Alkyl mercury compounds (dimethylmercury Hg(Me)₂, methylmercury MeHg⁺) are much more toxic than inorganic Hg salts. Aryl mercury compounds are even more dangerous — they cause brain damage, numbness, loss of vision, deafness, madness, and death. They have a very strong Hg–C bond and persist for a long time.

4.9 Biological Roles

Zinc — Biologically Essential

Zinc has an important biological role — it is the second most important transition metal (after Fe). Humans contain about 2g Zn (Fe: 4g). There are about 20 enzymes containing Zn:

Carbonic anhydrase — in red blood cells; speeds CO₂ absorption by blood in muscles/tissues and CO₂ release in lungs; regulates pH. Reaction: CO + OH⁻ ⇌ HCO₃⁻
Carboxypeptidase — in pancreatic juice; digestion of proteins; hydrolyses the terminal peptide (amide) link at the carboxyl end of the peptide chain (selective — only works when the terminal amino acid is aromatic or branched aliphatic with L configuration).
Alkaline phosphatase (energy release)
Dehydrogenases and aldolases (sugar metabolism)
Alcohol dehydrogenase (metabolism of alcohol)

Cadmium and Mercury — Biologically Toxic

It is striking that Zn is essential for life but Cd and Hg are both extremely toxic:

Cd: Accumulates in kidneys; causes kidney malfunction and also replaces Zn in some enzymes, preventing their function. Main threat: near Zn smelters (Cd escapes as dust with flue gases). Also from Ni/Cd batteries — manufacturing problems in Sweden and Japan.
Hg vapour: Toxic if inhaled → giddiness, tremors, lung and brain damage. In the laboratory, Hg should be covered with oil or toluene; spillages treated with flowers of sulphur → HgS.

Exam Tips, Tricks, and Frequently Tested Concepts

📌 JEE Advanced / NEET Hot Topics:

Co(+II) vs Co(+III): Simple compounds prefer +II; complexes strongly prefer +III (d⁶ high CFSE). The [Co(NH₃)₆]²⁺ is oxidized by air to [Co(NH₃)₆]³⁺ in the presence of activated charcoal.
Wilkinson's Catalyst: [RhCl(PPh₃)₃] — selective hydrogenation at room temperature and pressure; alk-1-enes only.
Cisplatin vs Transplatin: Only the cis isomer is the anticancer drug. The trans isomer is completely inactive.
Disproportionation: Cu⁺ and Au⁺ disproportionate; Ag⁺ does NOT disproportionate in water.
Jahn-Teller distortion: Cu²⁺ (d⁹) shows strong tetragonal distortion — 4 short + 2 long bonds.
Photography: AgBr (mainly) used; latent image = few Ag atoms; developed with quinol; fixed with sodium thiosulphate.
Hg(I) ion: Written as Hg₂²⁺ (dinuclear), NOT Hg⁺. Evidence: X-ray, magnetic, Raman, EMF, cryoscopic.
Vitamin B₁₂: Contains Co(+III) at centre of corrin ring; sixth position (active site) has OH/CH₃/adenosine in vivo.
OXO Process: HCo(CO)₄ catalyses alkene + CO + H₂ → aldehyde at 150°C, 200 atm.
Wacker Process: PdCl₂ catalyses C₂H₄ + ½O₂ → CH₃CHO (acetaldehyde).
Mond Process: Ni + 4CO →⁵⁰°C Ni(CO)₄ →²³⁰°C Ni (pure). Ni(CO)₄ is tetrahedral, volatile, very toxic.
ZnO: Amphoteric; white pigment; used in rubber production.
Fehling's test: Cu²⁺ (blue) → Cu₂O (brick-red) with reducing sugars.
Nessler's Reagent: K₂[HgI₄] detects NH₃ at 1 ppm.

📌 CSIR-NET / IIT-JAM / GATE Advanced Concepts:

Oxidative addition: Square planar (+I) Ir or Rh complex + XY → octahedral (+III) complex. Requires d⁸ or d¹⁰ metal, non-bonding d electrons, AND two vacant coordination sites.
Vaska's compound: trans-[Ir(Cl)(CO)(PPh₃)₂] — yellow → orange on O₂ absorption (reversible oxygenation). Model for Hb oxygen binding.
Zeise's salt: K[Pt(η²-C₂H₄)(Cl)₃] — first alkene complex (1825). Bonding: σ donation (ethylene → Pt) + π backdonation (Pt → ethylene π*).
Magnus' green salt: [Pt(NH₃)₄]²⁺[PtCl₄]²⁻ — square planar cation and anion stacked; iridescent green colour from metal–metal interaction; increased electrical conductivity in one dimension.
Polycations of Hg: Hg₃²⁺ in Hg₃(AlCl₄)₂, Hg₄²⁺ in Hg₄(AsF₆)₂ — linear Hg chains. [Hg₂.₈₅(AsF₆)]ₙ is a superconductor at low temperatures.
Crystal field stabilization and spin crossover: Co(+II) d⁷ in octahedral field: [Co(H₂O)₆]²⁺ is high spin (t₂g)⁵(eg)²; in strong field (CN⁻): [Co(CN)₅]³⁻ is low spin (paramagnetic, 1 unpaired e⁻, square pyramidal).
Cobaltocene: [Co^II(η⁵-C₅H₅)₂] — dark purple, air sensitive; easily oxidized to stable yellow [Co^III(η⁵-C₅H₅)₂]⁺ (not oxidized by conc. HNO₃ — like ferrocene the rings are attacked by nucleophiles). Rhodocene [Rh^II(η⁵-C₅H₄)₂] is less stable and tends to dimerize.

📌 TGT / PGT Level — Important Facts to Remember:

Co is ferromagnetic (like Fe and Ni) but loses magnetism above 1000°C.
Ir has the highest density of any element: 22.61 g cm⁻³.
Rh is an important catalyst in car exhaust control systems (three-way catalytic converters with Pt and Pd).
Smalt (ground blue cobalt glass) was known to ancient Egyptians and Romans.
Zn is the 24th most abundant element; Hg is very scarce; Cd is found as traces in Zn ores.
Group 11 elements (Cu, Ag, Au) have the highest electrical and thermal conductivities of all metals.
Au dissolves in aqua regia (3:1 HCl:HNO₃) — HNO₃ oxidizes, Cl⁻ complexes as [AuCl₄]⁻.
Mercury is the only metal liquid at room temperature; its vapour is monatomic (like noble gases).
Wilson's disease: hereditary shortage of ceruloplasmin → Cu accumulation in liver, kidneys, brain. Treated with EDTA chelation therapy.
Haemocyanin (Cu-containing protein) is the oxygen carrier in snails, crabs, lobsters, octopuses, scorpions — oxygenated = blue (unlike human blood).

Quick Reference: Key Reactions Summary

Group 9 — Key Reactions:

Co(acac)₃ + N₂ + 3PPh₃ → [Co^I(H)(N₂)(PPh₃)₃] (dinitrogen complex)

[Co^II(NH₃)₆]²⁺ →^{air, O₂} [Co^III(NH₃)₆]³⁺ (oxidation in complexes)

4Co²⁺ + 12en + 4H⁺ + O₂ → 4[Co^III(en)₃]³⁺ + 2H₂O

[Ir^I(Cl)(CO)(PPh₃)₂] + HCl → [Ir^III(Cl)₂(CO)(PPh₃)₂H] (oxidative addition)

CH₃OH + CO → CH₃COOH (Rh complex catalysis — acetic acid synthesis)

Group 10 — Key Reactions:

Ni + 4CO →^50°C Ni(CO)₄ →^230°C Ni + 4CO (Mond process)

C₂H₄ + PdCl₂ + H₂O → CH₃CHO + Pd + 2HCl (Wacker process)

Ni²⁺ + 2 dmgH₂ → [Ni(dmgH)₂]↓ (red) + 2H⁺ (DMG test)

Group 11 — Key Reactions:

2Cu⁺ ⇌ Cu²⁺ + Cu (disproportionation of Cu⁺)

[Co(H₂O)₆]²⁺ + 4Cl⁻ → [CoCl₄]²⁻ + 6H₂O (pink → blue)

3Cu + 8HNO₃(dilute) → 3Cu(NO₃)₂ + 2NO + 4H₂O

AgBr + 2Na₂S₂O₃ → Na₃[Ag(S₂O₃)₂] + NaBr (fixing in photography)

4Au + 8NaCN + 2H₂O + O₂ → 4Na[Au(CN)₂] + 4NaOH (cyanide extraction)

Group 12 — Key Reactions:

HgS + O₂ →^600°C Hg + SO₂ (mercury extraction)

Hg²⁺ + Hg ⇌ Hg₂²⁺ (E° = +0.13 V)

ZnO + CO ⇌ Zn + CO₂ (Zn extraction)

Zn + 2NaOH + H₂O → Na₂[Zn(OH)₄] + H₂ (Zn is amphoteric)

CH≡CH →^{Hg²⁺/H₂O} CH₂=CHOH →^H₂O CH₃CHO (Hg²⁺ catalysis)

Graphical Trends — Melting Points Across Groups 9–12

Fig. 7 — Melting points across Groups 9–12 (first, second, third row elements). Note the dramatic drop from Group 11 to Group 12 (especially Hg at −39°C), reflecting the complete d¹⁰ shell not participating in metallic bonding in Group 12.

Summary: The Big Picture

Across Groups 9–12, a beautiful pattern emerges that cuts across all competitive examinations:

Oxidation state trends: Moving from Group 9 → 12, the maximum useful oxidation state decreases. Group 9 reaches +IV (Ir), Group 10 reaches +VI (Pt only), Group 11 reaches +III (Au only), Group 12 is essentially locked at +II (with Hg also forming +I as dinuclear).
Complex formation: Groups 9 and 10 form the most varied and important coordination compounds. Co(+III) and Pt(+II, +IV) are the champions of complex chemistry.
Biological roles: Co (Vitamin B₁₂), Ni (urease enzyme), Cu (haemocyanin, ceruloplasmin, many oxidases), Zn (carbonic anhydrase, carboxypeptidase). Cd and Hg are biologically toxic with no beneficial role.
Industrial catalysis: Rh/Pt/Pd (three-way catalytic converters), Rh (OXO process, acetic acid), Pd (Wacker process), Ni (Mond process, hydrogenation), Pt (cisplatin drug, oil reforming).
Group 12 anomaly: These elements do not behave as true transition elements because of their complete d¹⁰ shell — no variable valency, no d–d spectra, no CFSE, no paramagnetism from d electrons. Mercury's unique liquid state at room temperature is a consequence of high ionization energy (relativistic effects) preventing d¹⁰ electrons from fully participating in metallic bonding.

🎯 Final Exam Strategy: When answering questions on these groups, always think in terms of:

Electronic configuration → predicts geometry, spin state, colour, magnetism.
Oxidation state stability → determines compound type (simple ionic vs complex).
Crystal Field Theory → CFSE determines which spin state/geometry is preferred.
Periodic trends vertically: Heavier elements (2nd and 3rd row) prefer higher oxidation states, are more noble, have stronger M–M bonds (lanthanide contraction makes 2nd and 3rd row atoms nearly the same size).

References & Further Reading: J.D. Lee, Concise Inorganic Chemistry, 5th Ed., Blackwell Science; Atkins, Shriver & Atkins' Inorganic Chemistry; IUPAC Recommendations for Nomenclature of Inorganic Chemistry (2005). All molecular geometries, bond angles, and oxidation states are consistent with IUPAC guidelines and experimental crystallographic data.

Transition Elements (d-Block): Co, Ni, Cu and Zn Groups – Properties, General Trends and Complex Formation Explained For PGT, JEE Advanced, IIT-JAM, GATE, TGT, BHU, BITSAT, CSIR-NET, NEET Exam Aspirants

⚗️ Transition Metals: Groups 9–12 (Co, Rh, Ir | Ni, Pd, Pt | Cu, Ag, Au | Zn, Cd, Hg)