d-Block

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

d-Block — Transition Elements (Groups 3–12)

The d-block contains transition metals with incomplete d-orbitals in their atoms or ions. They bridge the s-block and p-block and exhibit characteristic properties: variable oxidation states, colored compounds, complex formation, and catalytic activity.

Electronic Configuration: General configuration: (n−1)d¹⁻¹⁰ns⁰⁻²

• Sc: [Ar] 3d¹ 4s²
• Ti: [Ar] 3d² 4s²
• V: [Ar] 3d³ 4s²
• Cr: [Ar] 3d⁵ 4s¹ (anomalous — half-filled d subshell)
• Mn: [Ar] 3d⁵ 4s²
• Fe: [Ar] 3d⁶ 4s²
• Co: [Ar] 3d⁷ 4s²
• Ni: [Ar] 3d⁸ 4s²
• Cu: [Ar] 3d¹⁰ 4s¹ (anomalous — fully filled d subshell)
• Zn: [Ar] 3d¹⁰ 4s²

⚡ Exam tip: Zn, Cd, Hg have complete d¹⁰ configuration and are not true transition metals (they only show +2 oxidation state). Technically, a transition metal must have partially filled d-orbitals in at least one oxidation state.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

d-Block — Chemistry Study Guide

General Properties:

1. Atomic and Ionic Radii: Decreases from Sc to Cu across the period (poor shielding by d-electrons). After Cu, radii increase slightly in the next 3d series. Ionic radii: M⁺ > M²⁺ > M³⁺ for the same element.

2. Ionization Enthalpy: Increases from Sc to Zn (generally). Irregularities: Cr (lower IE due to half-filled stability), Cu (higher IE than Zn because removing electron from 3d is easier than from 4s).

3. Oxidation States: All transition metals show variable oxidation states. The common oxidation states:

   • Sc: +3 only
   • Ti: +2, +3, +4 (Ti³⁺ is stable in aqueous; TiO²⁺ in solution)
   • V: +2, +3, +4, +5 (V⁵⁺ is VO₂⁺; vanadyl)
   • Cr: +2, +3, +6 (CrO₄²⁻ yellow, Cr₂O₇²⁻ orange)
   • Mn: +2, +4, +6, +7 (MnO₄⁻ purple, powerful oxidant)
   • Fe: +2, +3 (Fe₂O₃, Fe₃O₄ called magnetite)
   • Co: +2, +3
   • Ni: +2
   • Cu: +1, +2 (Cu²⁺ blue; Cu⁺ unstable in water, disproportionates)
   • Zn: +2 only

4. Magnetic Properties: Paramagnetic if unpaired electrons exist. μ = √(n(n+2)) BM where n = number of unpaired electrons.

   • Sc³⁺: d⁰ → 0 unpaired → diamagnetic
   • Mn²⁺: d⁵ → 5 unpaired → 5.92 BM
   • Fe²⁺: d⁶ → 4 unpaired → 4.90 BM
   • Fe³⁺: d⁵ → 5 unpaired → 5.92 BM
   • Zn²⁺: d¹⁰ → 0 unpaired → diamagnetic

5. Complex Formation: d-block ions form coordination compounds due to small size, high charge, and available d-orbitals. Ligands cause crystal field splitting (Δ₀ for octahedral, Δₜ for tetrahedral).

6. Catalytic Properties: Transition metals and their compounds act as catalysts because they can exhibit variable oxidation states, provide surface area, and form intermediates.

   • Fe in Haber process (N₂ + 3H₂ → 2NH₃)
   • V₂O₅ in Contact process (2SO₂ + O₂ → 2SO₃)
   • Pd/Cu in hydrogenation
   • Ni in hydrogenation of oils
   • Pt in oxidation of SO₂

7. Alloy Formation: Transition metals form alloys with each other (steel = Fe + C + Mn, Cr + Ni = stainless steel).

Colored Compounds: d-d transition causes color in transition metal complexes. The color depends on the metal ion, oxidation state, ligand, and coordination number.

• [Ti(H₂O)₆]³⁺: violet/purple (d¹)
• [V(H₂O)₆]²⁺: violet (d³)
• [Cr(H₂O)₆]³⁺: blue-violet (d³)
• [Mn(H₂O)₆]²⁺: pale pink (d⁵) — weak absorber
• [Fe(H₂O)₆]³⁺: yellow-brown (d⁵, high spin)
• [Fe(H₂O)₆]²⁺: pale green (d⁶)
• [Co(H₂O)₆]²⁺: pink (d⁷)
• [Ni(H₂O)₆]²⁺: green (d⁸)
• [Cu(H₂O)₆]²⁺: blue (d⁹)

⚡ Exam tip: The intensity of color depends on the number of d-d transitions possible. High-spin complexes are often pale-colored because d-d transitions are spin-forbidden in some cases.

Standard Electrode Potentials (E° values — important for stability):

Mn²⁺/Mn: -1.18 V
Fe²⁺/Fe: -0.44 V
Co²⁺/Co: -0.28 V
Ni²⁺/Ni: -0.25 V
Cu²⁺/Cu: +0.34 V
Zn²⁺/Zn: -0.76 V

More negative E° means stronger reducing agent. Mn²⁺ is the most stable d² ion in aqueous solution (has half-filled d⁵ configuration). Cr³⁺ is stable because of half-filled t₂g in octahedral field.

Potassium Dichromate (K₂Cr₂O₇):

• Orange-red crystals, soluble in water
• In acidic medium: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O
• Acts as oxidizing agent in acidic medium:
  • oxidizes Fe²⁺ to Fe³⁺
  • oxidizes I⁻ to I₂
  • oxidizes SO₃²⁻ to SO₄²⁻
• Yellow in alkaline medium (chromate: CrO₄²⁻)

Potassium Permanganate (KMnO₄):

• Deep purple crystals
• In acidic medium: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
• In neutral/alkaline: MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻
• In strongly alkaline: MnO₄⁻ + e⁻ → MnO₄²⁻ (green manganate)
• Preparation: 2MnO₂ + KClO₃ + 2KOH → 2KMnO₄ + KCl + H₂O
• Commonly used titrant in redox titrations

Inner Transition Metals:

• Lanthanides (4f series): Ce to Lu — atomic radii decrease (lanthanide contraction)
• Actinides (5f series): Th to Lr — all radioactive, show variable oxidation states

⚡ Exam tip: Lanthanide contraction causes the 5d series (Hf, Ta, W…) to have nearly the same atomic/ionic radii as the 4d series (Zr, Nb, Mo…), leading to similar properties. This is why Zr and Hf have similar chemical properties — difficult to separate by fractional crystallization.

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

d-Block — Comprehensive Chemistry Notes

Crystal Field Theory (CFT):

In octahedral complexes, the d-orbitals split into two higher energy orbitals (e_g: dₓ²₋ᵧ², dᵧ²) and three lower energy orbitals (t₂g: dₓᵧ, dᵧᵤ, dᵤₓ). The energy gap is Δ₀ (crystal field splitting energy).

For octahedral: e_g are ↑↑ and t₂g are ↑↑↑ In tetrahedral: the splitting is reversed (t₂ higher, e lower) with Δₜ = (4/9)Δ₀.

Spectrochemical Series (ordering of ligand field strength): I⁻ < Br⁻ < SCN⁻ < Cl⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < NH₃ < en < bipy < phen < NO₂⁻ < PR₃ < CN⁻ < CO

Strong field ligands (right side) cause large Δ, low spin complexes. Weak field ligands (left side) cause small Δ, high spin complexes.

Pairing Energy vs Crystal Field Splitting:

• If Δ₀ > P (pairing energy): low spin complex (favored)
• If Δ₀ < P: high spin complex (favored)

Examples:

• [Co(CN)₆]⁴⁻: low spin (CN⁻ is strong field)
• [CoF₆]³⁻: high spin (F⁻ is weak field)

Coordination Number and Geometry:

• CN = 2: Linear (e.g., [Ag(NH₃)₂]⁺)
• CN = 4: Tetrahedral or Square planar (Ni(CN)₄²⁻ is square planar; NiCl₄²⁻ is tetrahedral)
• CN = 6: Octahedral (most common)

Jahn-Teller Distortion: When a degenerate d-orbital is unequally occupied, the complex distorts to remove degeneracy.

• Strong JT effect in d⁹ (Cu²⁺): tetragonally elongated octahedron (4 short bonds, 2 long bonds)
• Moderate JT effect in high-spin d⁴ (Cr²⁺, Mn³⁺)
• Weak/no JT for d⁰, d³, d⁵, d¹⁰

Stability of Coordination Compounds: Stability constant (Kf) = [MLₙ]/[M][L]ⁿ Stepwise formation constants: K₁, K₂, K₃… Overall βₙ = K₁ × K₂ × K₃ × … × Kₙ

Chelate effect: Complexes with polydentate ligands (chelating agents) are more stable than those with monodentate ligands. Example: [Ni(en)₃]²⁺ is more stable than [Ni(NH₃)₆]²⁺.

Isomerism in Coordination Compounds:

1. Geometrical Isomerism:

   • cis/trans in square planar [MA₂B₂] type (Pt(NH₃)₂Cl₂ exists as cis and trans)
   • fac/mer in octahedral [MA₃B₃] type

2. Optical Isomerism:

   • [Co(en)₃]³⁺ is optically active (tris-chelate)
   • [Co(ox)₃]³⁻ is optically active
   • [Cr(C₂O₄)₃]³⁻ is optically active
   • Meso forms exist for some compounds

3. Linkage Isomerism:

   • SCN⁻ can bind through S or N
   • [Co(NH₃)₅SCN]²⁺ vs [Co(NH₃)₅NCS]²⁺

4. Ionization Isomerism:

   • [Co(NH₃)₅SO₄]Br vs [Co(NH₃)₅Br]SO₄

5. Hydrate Isomerism:

   • [Cr(H₂O)₆]Cl₃ vs [Cr(H₂O)₅Cl]Cl₂·H₂O vs [Cr(H₂O)₄Cl₂]Cl·2H₂O

Metal Carbonyls: CO is a π-acceptor ligand (back-bonding from metal d-orbitals to empty π* orbitals on CO).

• Ni(CO)₄: tetrahedral (Ni⁰, d¹⁰)
• Fe(CO)₅: trigonal bipyramidal (Fe⁰, d¹⁰)
• Cr(CO)₆: octahedral (Cr⁰, d⁶)

17-electron rule applies to these metal carbonyls.

The 18-Electron Rule: Most stable complexes have 18 valence electrons (metal d-electrons + ligand electrons).

• Fe: d⁶ → needs 12 from ligands for 18e
• CO donates 2e each (weak field, good π-acceptor)
• NO donates 3e (nitrosyl)

Krogstad’s Salt and Other Important Compounds:

Sodium Cobaltinitrite: Na₃[Co(NO₂)₆] — used to test K⁺ (yellow precipitate)

Prussian Blue: Fe₄[Fe(CN)₆]₃ — deep blue pigment, formed when Fe³⁺ reacts with [Fe(CN)₆]⁴⁻ Turnbull’s Blue: Fe₃[Fe(CN)₆]₂ — formed when Fe²⁺ reacts with [Fe(CN)₆]³⁻

Tetraamminecopper(II) Sulfate: [Cu(NH₃)₄]SO₄·H₂O — deep blue complex, formed when excess NH₃ is added to Cu²⁺

Hexaamminecobalt(III) Chloride: [Co(NH₃)₆]Cl₃ — yellow, inert complex (low spin d⁶, very stable)

Potassium Hexacyanoferrate(III): K₃[Fe(CN)₆] — used to detect Fe²⁺ (forms Turnbull’s blue)

Disproportionation Reactions: 2Cu⁺ → Cu²⁺ + Cu (in aqueous solution, Cu⁺ is unstable) 3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O (manganate disproportionation)

Chromyl Chloride Test (CrO₂Cl₂): When chloride is heated with K₂Cr₂O₇ and conc. H₂SO₄, orange-red vapors of chromyl chloride are evolved: K₂Cr₂O₇ + 4KCl + 6H₂SO₄ → 2CrO₂Cl₂ + 6KHSO₄ + 3H₂O CrO₂Cl₂ hydrolyzes: CrO₂Cl₂ + 2H₂O → H₂CrO₄ + 2HCl

This test is used to confirm presence of Cl⁻ in a mixture.

Fajan’s Rule for Ionic/Covalent Character: For transition metal compounds:

• Smaller cation → more covalent character
• Higher charge on cation → more covalent character
• Cation with d-electrons → more covalent character (poor shielding)

Thus: Hg²⁺ compounds are more covalent than Zn²⁺ compounds. This is why HgS is insoluble in HCl but soluble in aqua regia.

Interstitial Compounds: d-block metals can form interstitial compounds with small non-metals (C, H, N, B) that occupy lattice voids.

• TiC, ZrC, WC — extremely hard, high melting points
• Fe₃C (cementite) — component of steel
• TiH₂ — hydrogen storage

Non-Stoichiometric Compounds: Compounds where metal-to-nonmetal ratio is not exactly whole number due to variable oxidation states or lattice defects.

• Fe₃O₄ = FeO·Fe₂O₃
• FeS₂ (pyrite): actual composition is Fe₁₋ₓS
• PdHₓ (x varies with pressure)

Comparative Study — First Row vs Second Row vs Third Row:

Property	3d	4d	5d
Size	Smallest	Medium	Largest
M—M bonding	Weak	Stronger	Strongest
Stability of +2 state	Decreases	—	—
Stability of +4 state	Fe << Ru ~ Os
Spin-orbit coupling	Weak	Strong	Very strong
Complex inertia	Low	High	Very high

The 5d metals form the most inert complexes (slow ligand substitution rates).

Eleanor’s Salt and Other Named Compounds:

• Mohr’s salt: (NH₄)₂Fe(SO₄)₂·6H₂O — stable Fe²⁺ source
• Tutton’s salts: (NH₄)₂M(SO₄)₂·6H₂O
• Carnallite: KCl·MgCl₂·6H₂O (double salt)
• Epsom salt: MgSO₄·7H₂O

⚡ Exam tip: In JEE, questions on d-block often ask about color of complex ions, magnetic moment calculation, or disproportionation conditions. Remember: Cu²⁺ is d⁹ — if you calculate magnetic moment using μ = √(n(n+2)), you get √(4×6) = √24 ≈ 4.9 BM for 1 unpaired electron — the formula gives spin-only magnetic moment which is approximate. Actual values may differ due to orbital contributions.

⚡ Exam tip: For KMnO₄ titration in acidic medium, the equivalent weight = M/5 because the change in oxidation state is 5. For K₂Cr₂O₇ in acidic medium, equivalent weight = M/6.

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