← Chemistry — Std 12
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Transition and Inner Transition Elements

Ch. 8Std 12

Easy Overview

Ever wondered why copper wires conduct electricity so well, why iron rusts, why chromium shines, or why gemstones are brilliantly colored? The answer lies in the transition elements — the d-block elements that sit in the middle of the periodic table. These elements are the workhorses of chemistry, forming colored compounds, variable oxidation states, and complexes that make modern technology possible. Transition elements are defined as elements whose atoms have partially filled d-orbitals (d¹ to d⁹) in the ground state or common oxidation states. This includes Groups 3 to 12 in periods 4, 5, and 6 — with the f-block (lanthanides and actinides) set apart as inner transition elements. They are all metals, they are hard and lustrous, good conductors of heat and electricity, and have high melting and boiling points. Many of the objects around you — from the stainless steel in your kitchen to the electronic devices you use — rely on the unique properties of these elements. Before we dive into each property, one idea ties it all together: the partially filled d-orbitals. These d-orbitals can accept and donate electrons, which explains variable oxidation states. They also split in energy when ligands approach, giving rise to that beautiful rainbow of colored compounds. The d-electrons also contribute to metallic bonding (making them strong and hard) and can align their spins to produce magnetic materials. Mastering the d-block requires understanding one core concept: what the d-orbitals can do — and the answers come down to electron configuration, the (n-1)d¹⁻¹⁰ ns¹⁻² pattern, and how different elements use those d-electrons.

Definition and electron configuration

A transition element (IUPAC) is an element with an atom that has a partially filled d-subshell, or that can give rise to cations with incomplete d-subshells. This includes d-block elements of groups 3-12. Sc (2,8,9,2) with configuration 3d¹4s² forms Sc³⁺ which has 3d⁰ (empty) — strictly not a transition element! Similarly Zn (3d¹⁰4s²), Cd, and Hg (d¹⁰ configuration) are NOT transition elements because they have completely filled d-orbitals in both the atom and common ions (Zn²⁺ is 3d¹⁰, Hg²⁺ is 5d¹⁰). Of the elements from Sc to Zn in period 4, Sc and Zn are excluded — the true transition elements are Ti through Cu. General electronic configuration: (n-1)d¹⁻¹⁰ ns¹⁻². Common anomalies: Cr (3d⁵4s¹) — half-filled d gives extra stability; Cu (3d¹⁰4s¹) — completely filled d gives extra stability. These exceptions arise because half-filled and fully-filled d-orbitals have lower energy (exchange energy and symmetry).

Trends in atomic and ionic radii

Across a period (Sc → Zn), atomic and ionic radii decrease gradually, but much less sharply than in s- or p-block. This is due to the d-orbital shielding effect — d-electrons shield the outer 4s electrons from the nuclear charge less effectively than s or p electrons (d-orbitals are more diffuse). So as nuclear charge increases, the 4s electrons are pulled inward only slightly. The radii of the second and third transition series (periods 5 and 6) are very similar due to lanthanide contraction. The 4f orbitals of lanthanides have poor shielding, so the 14-element increase in nuclear charge from La to Lu strongly contracts the atomic size. This means Zr (period 5) and Hf (period 6) have almost identical radii (160 pm and 159 pm), making them extremely difficult to separate — their chemistry is nearly identical. This is why Zr and Hf occur together in nature and require fractional distillation or ion exchange for separation.

Ionization enthalpies and oxidation states

Successive ionization enthalpies increase gradually across the series. The first ionization energy (I₁) involves removing a 4s electron; I₂ removes the second 4s electron; I₃ begins removing 3d electrons. The sum I₁ + I₂ + I₃ determines the stability of the +3 oxidation state. Variable oxidation states are a hallmark of transition elements. Fe shows +2 (Fe²⁺) and +3 (Fe³⁺); Mn shows +2, +3, +4, +6, +7 — from Mn²⁺ to MnO₄⁻. The highest oxidation state increases across the period up to Mn (Mn₂O₇, +7) and then decreases. The highest oxidation states are achieved when all 3d and 4s electrons are used — for Sc (+3), Ti (+4), V (+5), Cr (+6), Mn (+7). But removing more electrons requires too much energy after Mn. Lower oxidation states are favored by heavier elements (W is more stable in +6 than Cr). Stability of common oxidation states: Fe³⁺ > Fe²⁺ (because 3d⁵ half-filled is extra stable), Mn²⁺ > Mn³⁺ (again 3d⁵). The +2 state for first-row transition elements follows the order: Mn > Fe > Co > Ni > Cu > Zn (this is the Irving-Williams series for stability).

Magnetic properties

Paramagnetism: attraction to an external magnetic field, caused by unpaired electrons. Each unpaired electron contributes a magnetic moment. The spin-only formula: μ = √(n(n+2)) Bohr Magnetons (BM), where n = number of unpaired electrons. Diamagnetism: repelled by a magnetic field (all electrons paired). Ti³⁺ (d¹): 1 unpaired e⁻, μ = 1.73 BM. V³⁺ (d²): 2 unpaired e⁻, μ = 2.84 BM. Cr³⁺ (d³): 3 unpaired e⁻, μ = 3.87 BM. Fe²⁺ (d⁶ in high spin): 4 unpaired e⁻, μ = 4.90 BM. Cu²⁺ (d⁹): 1 unpaired e⁻, μ = 1.73 BM. Zn²⁺ (d¹⁰): 0 unpaired e⁻, diamagnetic. Magnetic moments are measured to determine the number of unpaired electrons in a complex, which helps deduce geometry and spin state. Ferromagnetism (permanent magnets) appears in Fe, Co, Ni and some alloys — here, unpaired spins on neighboring atoms align parallel in domains. This is different from paramagnetism, where alignment is temporary and lost when the field is removed.

Catalytic properties

Transition metals and their compounds are excellent catalysts. They provide surfaces for reactant molecules to adsorb (heterogeneous catalysis) and can adopt multiple oxidation states to form intermediate complexes (homogeneous catalysis). Haber process (NH₃ synthesis): Fe with promoters (Al₂O₃, K₂O) — N₂ and H₂ adsorb on Fe surface, N≡N triple bond weakens, NH₃ forms and desorbs. Contact process (H₂SO₄): V₂O₅ — SO₂ + ½O₂ → SO₃ (VO₂ intermediate). Ostwald process (HNO₃): Pt-Rh gauze — 4NH₃ + 5O₂ → 4NO + 6H₂O. Hydrogenation (margarine production): Ni catalyst — adds H₂ across C=C double bonds. Automobile catalytic converters: Pt, Pd, Rh — convert CO (toxic) to CO₂, NOₓ to N₂, and unburnt hydrocarbons to CO₂ + H₂O. The catalyst lowers Ea by providing an alternative reaction pathway. Transition metals are especially good because: (a) they have empty d-orbitals to accept electrons from reactants and d-electrons to donate, (b) they can readily change oxidation states, and (c) they can adsorb reactant molecules on their surface.

Formation of colored compounds

Most transition metal compounds are colored — and this is directly linked to d-orbitals. In a free transition metal ion, all five d-orbitals are degenerate (same energy). When ligands (molecules or ions like H₂O, NH₃, Cl⁻) surround the metal, the d-orbitals are no longer degenerate. In an octahedral complex, the d-orbitals split into two sets: t₂g (lower energy: dxy, dyz, dzx) and eg (higher energy: dx²-y², dz²). The energy difference (Δ, splitting energy) corresponds to the energy of visible light (400-700 nm). When visible light shines on the complex, d-electrons can absorb energy and jump from lower to higher d-orbitals (d-d transition). The color we see is the complementary color to the absorbed wavelength. Ti³⁺ (d¹) in water: ground state t₂g¹, and it absorbs green light (~540 nm) to promote the electron to eg — so [Ti(H₂O)₆]³⁺ appears purple (complement of green). Mn²⁺ (d⁵ high spin, half-filled): t₂g³ eg² — d-d transitions are spin-forbidden (need spin-flip), so Mn²⁺ compounds are very pale pink (almost colorless). Cu²⁺ (d⁹): gives blue color in CuSO₄·5H₂O. The exact color depends on the ligand — NH₃ (a stronger ligand) causes larger splitting than H₂O, so [Cu(NH₃)₄]²⁺ (deep blue) is darker than [Cu(H₂O)₆]²⁺ (pale blue). This is why you see different colors when different ligands coordinate — it is all about how much the d-orbitals split.

Formation of complex compounds

Transition metals form a vast number of coordination compounds because they have small size, high nuclear charge, and available d-orbitals for accepting lone pairs from ligands. A complex has a central metal ion bonded to ligands (Lewis bases that donate electron pairs). Examples: [Fe(CN)₆]⁴⁻ (ferrocyanide), [Cu(NH₃)₄]²⁺ (tetraamminecopper(II) ion — deep blue), [Ag(NH₃)₂]⁺ (Tollen's reagent), [Co(NH₃)₆]³⁺, K₃[Fe(C₂O₄)₃] (potassium ferrioxalate). Stability of complexes depends on the nature of the metal ion and ligand. Chelating ligands (multidentate ligands that bind through multiple sites — like EDTA with 6 binding sites) form more stable complexes due to chelate effect (entropy-driven: more free particles released when chelating ligand replaces monodentate ones). Complex formation is important in biological systems: hemoglobin (Fe-porphyrin for O₂ transport), chlorophyll (Mg-porphyrin for photosynthesis), vitamin B₁₂ (Co-corrin), and many enzymes contain transition metal ions at their active sites.

Lanthanides and their contraction

The lanthanide series (Ce to Lu, atomic numbers 58-71) follows lanthanum in period 6. These are f-block elements where the 4f orbitals are being filled. Lanthanides all have very similar chemistry — they are +3 ions with gradual decrease in ionic radii from La³⁺ (106 pm) to Lu³⁺ (85 pm). This lanthanide contraction (a steady decrease in size across 14 elements) occurs because 4f electrons shield each other poorly — the increased nuclear charge pulls all electrons inward. Consequences of lanthanide contraction: (1) The second and third transition series elements (like Zr and Hf) have nearly identical radii and properties, making separation difficult. (2) There is a decrease in basic character from La(OH)₃ to Lu(OH)₃. (3) The ionic radii of Y³⁺ (90 pm) is close to Ho³⁺ and Er³⁺ — explaining why Yttrium occurs in rare earth ores. Separation of lanthanides is challenging due to their chemical similarity — methods include ion-exchange chromatography and solvent extraction.

Key Points

  • •Transition elements: (n-1)d¹⁻¹⁰ ns¹⁻²; partially filled d-orbitals in atom or common ion
  • •Sc and Zn are NOT transition elements (Sc³⁺ is d⁰, Zn²⁺ is d¹⁰)
  • •Cr: 3d⁵4s¹ (not 3d⁴4s²); Cu: 3d¹⁰4s¹ (not 3d⁹4s²) — half-filled/filled d stability
  • •Variable oxidation states — hallmark of transition metals
  • •Atomic radii decrease slightly across period due to poor d-orbital shielding
  • •Lanthanide contraction: Zr (160 pm) and Hf (159 pm) have nearly same size
  • •Magnetic moment: μ = √(n(n+2)) BM, where n = number of unpaired electrons
  • •Colored compounds: due to d-d transitions (splitting of d-orbitals in ligand field)
  • •Superior catalysts due to variable oxidation states and ability to adsorb reactants
  • •Fe (Haber), V₂O₅ (Contact), Pt (Ostwald, catalytic converters), Ni (hydrogenation)
  • •Chelate effect: multidentate ligands form more stable complexes (entropy driven)
  • •Lanthanides: 4f¹⁻¹⁴, all form +3 ions, very similar chemistry
  • •Lanthanide contraction: steady decrease in size from La to Lu
  • •f-block elements are still transition elements (inner transition)

Practice Questions

  • What are transition elements? Why are Sc and Zn not considered transition elements?
  • Explain the electronic configuration of Cr and Cu. Why do they differ from expected configurations?
  • Why do transition metals form colored compounds? Explain with reference to d-d transitions using [Ti(H₂O)₆]³⁺ as an example.
  • What is lanthanide contraction? Give its causes and consequences.
  • Explain the catalytic properties of transition metals with one industrial example.
  • Calculate the magnetic moment of Cr³⁺ (d³), Fe²⁺ (d⁶ high spin), and Zn²⁺ (d¹⁰).
  • Why do transition metals show variable oxidation states? Give examples from Mn (oxidation states +2 to +7).
  • Why are Zr and Hf so difficult to separate? Discuss the role of lanthanide contraction.