Transition elements: general properties, reactions, variable oxidation states, complex formation

Transition Elements – General Properties, Reactions, Variable Oxidation States and Complex Formation

1. Definition and Position in the Periodic Table

• Transition elements are the d‑block elements of groups 3 – 12 (periods 4 – 7).
• In at least one of their common oxidation states they possess a partially filled (n‑1)d subshell.
• Exceptions (often examined): Scandium (Sc) and Zinc (Zn) have a completely filled (n‑1)d subshell in the ground‑state configuration but display typical transition‑metal chemistry in their ions.

2. General Physical and Chemical Properties

2.1 Physical Properties

2.2 Origin of Colour in Transition‑Metal Compounds

• Colour arises from *d‑d transitions*: an electron is promoted from a lower‑energy d‑orbital to a higher‑energy d‑orbital within the same metal centre.
• These transitions are Laporte‑forbidden (no change in parity) but become partially allowed through vibronic coupling, giving rise to the vivid colours observed.
• The energy gap (Δ_oct or Δ_tet) depends on:
  • Oxidation state – higher charge increases Δ.
  • Nature of the ligands – the spectrochemical series (I⁻ < Br⁻ < Cl⁻ < F⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO) orders ligands from weak to strong field.
  • Geometry – octahedral complexes split d‑orbitals into t₂g and e_g sets; tetrahedral splitting is smaller (≈4/9 of octahedral).

2.3 Magnetic Behaviour – d‑Electron Count ↔ Unpaired Electrons ↔ Magnetism

3. Variable Oxidation States

3.1 Why Transition Metals Show Several Oxidation States

• The (n‑1)d and ns orbitals are close in energy; electrons can be removed from either set with comparable ionisation energies.
• The most stable oxidation state is usually the one that yields a half‑filled d⁵ or a completely filled d¹⁰ subshell, because these configurations are especially stable.

3.2 Worked Examples (Predict the Most Stable Oxidation State)

Chromium (Cr)

1. Ground‑state configuration: [Ar] 3d⁵ 4s¹.
2. Removing 1 e⁻ → Cr⁺ (d⁵) – half‑filled but rarely isolated.
3. Removing 3 e⁻ → Cr³⁺ (d³) – common in aqueous solution; balance of ionisation energy and hydration energy.
4. Removing 6 e⁻ → Cr⁶⁺ (d⁰) – very strong oxidising agent (chromate, dichromate).
5. In most chemical environments, +3 is the most stable because it offers a reasonable compromise between lattice/solvation energy and the cost of removing additional electrons.

Manganese (Mn)

1. Ground‑state: [Ar] 3d⁵ 4s².
2. +2 oxidation state → Mn²⁺ (d⁵) – half‑filled, very stable in many salts.
3. Higher states (+3, +4, +6, +7) are also observed but require higher oxidation energy; +7 (MnO₄⁻, d⁰) is a powerful oxidiser.

Iron (Fe)

1. Ground‑state: [Ar] 3d⁶ 4s².
2. +2 oxidation state → Fe²⁺ (d⁶) – moderately stable; commonly found in ferrous compounds.
3. +3 oxidation state → Fe³⁺ (d⁵) – half‑filled, therefore more stable in oxidising environments (e.g., ferric salts, Fe₂O₃).

3.3 Common Oxidation States, d‑Electron Counts and Typical Colours

4. Redox Behaviour and Standard Electrode Potentials

• Transition‑metal ions participate in characteristic redox couples. The standard electrode potential (E°) indicates the tendency of a species to be reduced (more positive E° = stronger oxidising agent).
• Values are given versus the standard hydrogen electrode (SHE) at 25 °C.

Using E° values: A redox reaction will proceed spontaneously in the direction that gives a positive overall cell potential (E°_cell = E°_cathode – E°_anode). Example:

$$\ce{2Fe^{2+} + MnO4^- + 8H^+ -> 2Fe^{3+} + Mn^{2+} + 4H2O}$$

Cell potential = 1.51 V (MnO₄⁻/Mn²⁺) – 0.77 V (Fe³⁺/Fe²⁺) = **+0.74 V**, so the reaction is spontaneous.

5. Characteristic Reactions of Transition Metals

1. Formation of coloured aqueous ions (ligand‑field change):
   $$\ce{[Fe(H2O)6]^{2+} + 4Cl^- -> [FeCl4]^{2-} + 6H2O}$$ Colour changes from pale green to yellow‑brown.
2. Redox reactions (example with standard potentials):
   $$\ce{2Fe^{2+} + H2O2 -> 2Fe^{3+} + 2OH^-}$$ Fe²⁺ is oxidised (E° = +0.77 V) while H₂O₂ acts as the oxidising agent.
3. Displacement reactions (activity series):
   $$\ce{Cu + 2AgNO3 -> Cu(NO3)2 + 2Ag}$$
4. Oxide and hydroxide behaviour
   • Basic oxide: $$\ce{Fe2O3 + 3H2O -> 2Fe(OH)3}$$
   • Amphoteric oxide: $$\ce{ZnO + 2NaOH + H2O -> Na2[Zn(OH)4]}$$
5. Complex‑ion formation (ligand exchange):
   $$\ce{[Co(H2O)6]^{3+} + 6NH3 -> [Co(NH3)6]^{3+} + 6H2O}$$

6. Complex Formation (Coordination Chemistry)

6.1 Key Concepts

• Coordination number (CN) – number of donor atoms attached to the metal centre. Most common CN = 4 (tetrahedral or square‑planar) and CN = 6 (octahedral).
• Ligand classification
  • Monodentate – one donor atom (e.g., $\ce{NH3}$, $\ce{Cl^-}$).
  • Bidentate – two donor atoms (e.g., ethylenediamine, $\ce{en}$).
  • Polydentate (chelating) – three or more donor atoms (e.g., $\ce{EDTA^{4-}}$).
• Geometry
  • Tetrahedral (CN = 4) – typical for high‑spin $d^{10}$ and $d^{0}$ metals.
  • Square planar (CN = 4, $d^{8}$) – common for Ni(II), Pd(II), Pt(II).
  • Octahedral (CN = 6) – the most common geometry for transition‑metal complexes.

6.2 Crystal‑Field Theory (CFT) – Qualitative Summary

• In an octahedral field the five d‑orbitals split into a lower‑energy t₂g set (d_xy, d_xz, d_yz) and a higher‑energy e_g set (d_z², d_x²‑y²).
• Δ_oct (the splitting energy) increases with:
  • Higher oxidation state.
  • Stronger‑field ligands (see the spectrochemical series).
  • Shorter metal–ligand distances.
• When Δ_oct > pairing energy, electrons pair in the lower t₂g set → *low‑spin* complexes (e.g., $\ce{[Fe(CN)6]^{3-}}$). Otherwise, electrons occupy higher e_g orbitals → *high‑spin* complexes (e.g., $\ce{[Fe(H2O)6]^{3+}}$).
• Low‑spin complexes are usually less magnetic (fewer unpaired electrons) and often display different colours compared with their high‑spin counterparts.

6.3 Stability of Complexes

• Stability constant (K_f) – quantitative measure of how tightly a ligand binds. Larger K_f ⇒ more stable complex.
• Chelate effect – multidentate ligands form complexes with markedly higher K_f than comparable monodentate ligands because several bonds are formed simultaneously, reducing the entropy loss.
• Example (formation of a hexaaqua ion):
  $$\ce{[Fe(H2O)6]^{2+} <=> Fe^{2+} + 6H2O}\qquad K_{\mathrm f}=10^{4.0}$$

6.4 Representative Complex‑Ion Reactions (Exam‑style)

1. Ligand substitution (labile vs inert):
   $$\ce{[Co(H2O)6]^{2+} + 4Cl^- -> [CoCl4]^{2-} + 6H2O}$$ (labile, fast exchange).
2. Formation of a chelate:
   $$\ce{[Cu(H2O)4]^{2+} + en -> [Cu(en)2]^{2+} + 4H2O}$$
3. Redox coupled with complex formation:
   $$\ce{2[Fe(CN)6]^{4-} + MnO4^- + 8H^+ -> 2[Fe(CN)6]^{3-} + Mn^{2+} + 4H2O}$$

7. Summary Checklist for Cambridge AS & A‑Level

• Know the definition, position (group 3‑12) and the two common exceptions.
• Be able to describe trends in melting point, density, conductivity and explain why transition metals are good conductors.
• Explain colour of compounds via d‑d transitions and ligand‑field splitting.
• Use the d‑electron‑count table to predict magnetic behaviour (high‑spin vs low‑spin).
• Identify the most stable oxidation state using the half‑filled d⁵ / fully‑filled d¹⁰ rule; apply this to Cr, Mn and Fe.
• Read and interpret standard electrode potentials; predict the direction of redox reactions.
• Recall characteristic reactions: coloured ion formation, redox, displacement, oxide/hydroxide behaviour, complex‑ion formation.
• Understand coordination number, ligand classification, common geometries and the basics of crystal‑field theory.
• State the chelate effect and recognise factors that increase complex stability.