UNIT-4: d- and f-Block Elements

1. General Group Trends of d-Block Elements

Electronic Configuration

around the d-block elements, the valence shell electronic configuration is generally represented as (n-1)d^(1-10) ns^(1-2), where (n-1) represents the inner d-orbital shell and n represents the outermost s-orbital shell.

General Electronic Configuration: (n-1)d^(1-10) ns^(1-2)

Anomalies occur in electronic configurations when half-filled (d⁵) or fully filled (d¹⁰) subshells offer extra stability due to symmetry and high exchange energy.

• Chromium (Cr, Z = 24): Expected [Ar] 3d⁴ 4s², actual configuration is [Ar] 3d⁵ 4s¹.
• Copper (Cu, Z = 29): Expected [Ar] 3d⁹ 4s², actual configuration is [Ar] 3d¹⁰ 4s¹.

Colour of Transition Metal Complexes

Most transition metal compounds are coloured in aqueous solutions or crystalline states. This property arises due to the presence of unpaired d-electrons and the phenomenon of d-d transitions.

• Mechanism: When ligands approach the central metal ion, the degenerate d-orbitals split into different energy levels (e.g., t₂g and eg in octahedral geometry).
• Light Absorption: An electron absorbs energy corresponding to the visible light spectrum and gets promoted from a lower energy d-orbital to a higher energy d-orbital.
• Transmitted Colour: The observed colour corresponds to the complementary colour of the absorbed wavelength.

Ions with d⁰ or d¹⁰ electronic configurations (e.g., Sc³⁺, Ti⁴⁺, Zn²⁺, Cu⁺) are colourless because no d-d transitions can occur due to the absence of unpaired d-electrons or complete filling of d-orbitals.

Variable Valency (Variable Oxidation States)

Transition metals exhibit variable oxidation states in their compounds. This variability occurs because the energy difference between the (n-1)d and ns orbitals is very small, allowing electrons from both orbitals to participate in chemical bond formation.

• Minimum Oxidation State: Generally corresponds to the loss of outermost ns electrons.
• Maximum Oxidation State: Increases across the series up to the middle (e.g., Mn exhibits up to +7 in KMnO₄) and then decreases as d-orbitals become paired.
• Bonding Nature: Lower oxidation states generally form ionic bonds (e.g., MnCl₂), whereas higher oxidation states form covalent bonds or oxoanions (e.g., MnO₄⁻).

Magnetic Properties

Transition metal ions generally exhibit magnetic properties depending on the presence of unpaired electrons.

• Paramagnetism: Observed in species containing one or more unpaired electrons. These substances are attracted by an external magnetic field.
• Diamagnetism: Observed when all electrons are paired. These substances are weakly repelled by an external magnetic field.

The spin-only magnetic moment (μ) can be calculated using the formula:

μ = √(n(n + 2)) BM

Where:

• n = Number of unpaired electrons
• BM = Bohr Magneton (unit of magnetic moment)

Catalytic Properties

Transition metals and their compounds act as effective catalysts in many chemical processes due to:

1. Variable Oxidation States: Ability to change oxidation states readily to form intermediate complexes with reactants.
2. Large Surface Area: Ability to provide a catalytic surface for adsorption of reactant molecules, lowering the activation energy.

• V₂O₅ in the Contact Process for H₂SO₄ manufacture.
• Finely divided Fe in the Haber Process for NH₃ synthesis.
• Ni in the hydrogenation of oils to fats.

Ability to Form Complexes

Transition metals form a vast number of complex compounds (coordination compounds) due to:

• Small size of transition metal cations.
• High effective nuclear charge (high charge-to-radius ratio).
• Availability of vacant d-orbitals of suitable energy to accept electron pairs donated by ligands.

2. Stability of Oxidation States and E.M.F. (Latimer & Ebsworth Diagrams)

Stability Factors of Oxidation States

The relative stability of an oxidation state of a transition metal in solution depends on the net balance of three thermodynamic parameters:

1. Enthalpy of Sublimation (ΔsubH): Energy required to convert solid metal into gaseous atoms.
2. Ionization Enthalpy (ΔiH): Energy required to remove electrons from gaseous atoms to form metal cations.
3. Hydration Enthalpy (ΔhydH): Energy released when metal cations interact with water molecules.

An oxidation state is stable in aqueous solution if the overall standard electrode potential (E°) for the reduction process is favorable, which is governed by ΔG° = -nFE°.

Latimer Diagrams

A Latimer diagram is a compact representation of standard reduction potential data for successive oxidation states of an element.

• The element is written with its highest oxidation state on the far left, and lower oxidation states are arranged sequentially to the right.
• Numerical values above the connecting arrows indicate standard reduction potentials (E° in Volts).

General Latimer Notation: A --(E°1)--> B --(E°2)--> C

Key Rules for Latimer Diagrams:

• To calculate the potential for a non-adjacent conversion (e.g., A to C), potentials cannot be added directly. Convert to free energy changes: ΔG° = -nFE°. Thus, ΔG°(total) = ΔG°1 + ΔG°2, leading to: E°(A->C) = (n1*E°1 + n2*E°2) / (n1 + n2).
• Disproportionation Criterion: A species in an intermediate oxidation state will undergo spontaneous disproportionation if the potential on its right is more positive than the potential on its left: E°(right) > E°(left).

Ebsworth Diagrams

An Ebsworth diagram (or Frost diagram) is a graphical plot of standard free energy function relative to the element in its zero oxidation state versus its oxidation state (N).

• Y-axis: -nE° or ΔG° / F (where n is the number of electrons involved relative to elemental state).
• X-axis: Oxidation State (N).

Interpretation of Ebsworth Diagrams:

• Thermodynamic Stability: The lower a species lies on the diagram, the more thermodynamically stable it is. The lowest point on the plot corresponds to the most stable oxidation state.
• Disproportionation: A species lies above the straight line connecting its two surrounding oxidation states will tend to undergo disproportionation into those two states.
• Comproportionation: If a line connecting two species passes above an intermediate species, those two outer species will react to form the intermediate state.
• Oxidizing/Reducing Strength: Slopes between points reflect standard reduction potentials. A steep positive slope going right indicates a strong oxidizing agent.

3. Difference Between First, Second, and Third Transition Series

The d-block consists of three main transition series: 3d (First, Sc to Zn), 4d (Second, Y to Cd), and 5d (Third, La to Hg).

4. Chemistry of Chromium (Cr) and Manganese (Mn) in Various Oxidation States

Chemistry of Chromium (Cr)

Chromium (atomic number 24, configuration [Ar] 3d⁵ 4s¹) exhibits key oxidation states of +2, +3, and +6.

1. Chromium (+2) State:

• Species: Cr²⁺ ion ([Ar] 3d⁴).
• Properties: Strongly reducing agent in aqueous solution. It converts to the more stable Cr³⁺ state by losing an electron.
• Example reaction: 2 Cr²⁺ + 2 H⁺ -> 2 Cr³⁺ + H₂

2. Chromium (+3) State:

• Species: Cr³⁺ ion ([Ar] 3d³).
• Properties: Most stable oxidation state of chromium. The d³ configuration gives extra stability in octahedral fields (half-filled t₂g subshell).
• Forms stable complexes such as [Cr(H₂O)₆]³⁺ (violet) and hydrated salts. Cr(OH)₃ is amphoteric, dissolving in both acids and strong bases.

3. Chromium (+6) State:

• Species: Chromate (CrO₄²⁻, yellow) and Dichromate (Cr₂O₇²⁻, orange).
• Properties: Highly oxidizing in acidic media.
• pH Dependent Equilibrium: Chromate and dichromate ions interconvert in aqueous solution depending on pH:

2 CrO₄²⁻ + 2 H⁺ ⇌ Cr₂O₇²⁻ + H₂O

In acidic solution (low pH), dichromate (Cr₂O₇²⁻) predominates (orange). In basic solution (high pH), chromate (CrO₄²⁻) decision predominates (yellow).

Chemistry of Manganese (Mn)

Manganese (atomic number 25, configuration [Ar] 3d⁵ 4s²) exhibits oxidation states ranging from +2 to +7.

1. Manganese (+2) State:

• Species: Mn²⁺ ([Ar] 3d⁵).
• Properties: Exceptionally stable in aqueous solution due to the extra stability of the half-filled d⁵ configuration. High-spin Mn²⁺ complexes are pale pink or colourless. Resistant to oxidation.

2. Manganese (+3) State:

• Species: Mn³⁺ ([Ar] 3d⁴).
• Properties: Unstable in aqueous solution; acts as a strong oxidizing agent or undergoes disproportionation:

2 Mn³⁺ + 2 H₂O -> Mn²⁺ + MnO₂ + 4 H⁺

3. Manganese (+4) State:

• Species: Manganese dioxide (MnO₂).
• Properties: Stable brown/black insoluble solid. Strong oxidizing agent in acidic solution; acts as a catalyst in the decomposition of H₂O₂.

4. Manganese (+6) State:

• Species: Manganate ion (MnO₄²⁻, green).
• Properties: Stable only in strongly alkaline solutions. Disproportionates rapidly in neutral or acidic media:

3 MnO₄²⁻ + 4 H⁺ -> 2 MnO₄⁻ + MnO₂ + 2 H₂O

5. Manganese (+7) State:

• Species: Permanganate ion (MnO₄⁻, intense dark purple).
• Properties: Powerful oxidizing agent in acidic, neutral, and alkaline media. The deep purple colour is caused by Ligand-to-Metal Charge Transfer (LMCT) rather than d-d transition (since Mn⁷⁺ is d⁰).

5. Lanthanoids and Actinoids

Lanthanoids (4f-Block Elements)

The series consists of 14 elements following Lanthanum, from Cerium (Z = 58) to Lutetium (Z = 71), where the 4f subshell is progressively filled.

Electronic Configuration

General Configuration: [Xe] 4f^(1-14) 5d^(0-1) 6s²

Extra stability is observed for empty (4f⁰), half-filled (4f⁷), and completely filled (4f¹⁴) subshells (e.g., Eu is [Xe] 4f⁷ 6s², Gd is [Xe] 4f⁷ 5d¹ 6s², Yb is [Xe] 4f¹⁴ 6s², Lu is [Xe] 4f¹⁴ 5d¹ 6s²).

Oxidation States

• Common State: +3 is the principal and most stable oxidation state for all lanthanoids.
• Variable States (+2, +4): Occur when ions achieve stable 4f⁰, 4f⁷, or 4f¹⁴ configurations:
  • Ce⁴⁺ ([Xe] 4f⁰): Stable, strong oxidizing agent converting to Ce³⁺.
  • Eu²⁺ ([Xe] 4f⁷): Stable, strong reducing agent converting to Eu³⁺.
  • Tb⁴⁺ ([Xe] 4f⁷): Exists due to half-filled stability.
  • Yb²⁺ ([Xe] 4f¹⁴): Stable, strong reducing agent.

Colour and Spectral Properties

• Tripositive lanthanoid ions (Ln³⁺) are often coloured in solid and solution states due to f-f transitions.
• Sharpness of Absorption Bands: Unlike d-d transition bands, f-f absorption spectra consist of extremely sharp, line-like bands. This occurs because 4f orbitals are deeply buried beneath the outer 5s and 5p shells, effectively shielding 4f electrons from external crystal field perturbations.
• Ions with 4f⁰ (La³⁺, Ce⁴⁺) and 4f¹⁴ (Yb²⁺, Lu³⁺) configurations are colourless.

Magnetic Properties

• Lanthanoid ions (except 4f⁰ and 4f¹⁴) exhibit paramagnetism due to unpaired 4f electrons.
• Orbital Contribution: Unlike transition metals where orbital angular momentum is quenched by crystal fields, 4f orbitals are shielded. Therefore, orbital contribution is not quenched.
• Magnetic moment is calculated using total angular momentum quantum number J (where J = L - S for less than half-filled, and J = L + S for more than half-filled):

μ = g * √(J(J + 1)) BM

Where g is the Landé g-factor.

Lanthanide Contraction

The steady and continuous decrease in atomic and ionic radii of lanthanoid elements with increasing atomic number from La³⁺ to Lu³⁺ is known as Lanthanide Contraction.

Cause: Poor shielding effect of 4f electrons. As nuclear charge increases by one unit at each step, the addition of a 4f electron fails to compensate for the increased nuclear pull due to the diffused shape of f-orbitals.

Consequences of Lanthanide Contraction:

1. Similarity of 4d and 5d Elements: The atomic/ionic radii of elements of the second (4d) and third (5d) transition series belonging to the same group are almost identical (e.g., Zr radius = 160 pm, Hf radius = 159 pm). This makes their chemical properties almost identical.
2. Difficulty in Separation: Due to nearly identical ionic radii and chemical properties, separation of pure lanthanoids from natural mixtures is exceptionally difficult.
3. Trend in Basicity of Hydroxides: Basicity of Ln(OH)₃ decreases from La(OH)₃ to Lu(OH)₃. As ionic radius decreases, covalent character of Ln-OH bond increases, reducing hydroxyl ion release.

Separation of Lanthanoids: Ion-Exchange Method

The ion-exchange method is the most modern, efficient, and reliable method for separating individual lanthanoid ions.

• Principle: Based on slight differences in hydrated ionic sizes and complex formation stabilities of Ln³⁺ ions with chelating agents.
• Procedure:
  1. A aqueous solution containing Ln³⁺ ions is passed through a column filled with a cation-exchange resin containing sulfonated polystyrene groups (-SO₃⁻ H⁺).
  2. Ln³⁺ ions displace H⁺ ions and bind to the resin: 3 Resin-H + Ln³⁺ ⇌ Resin₃-Ln + 3 H⁺
  3. The column is then eluted (washed) with a complexing agent such as an aqueous buffer solution of citric acid / ammonium citrate or EDTA.
• Elution Sequence: Smaller Ln³⁺ ions (e.g., Lu³⁺) have higher charge density, form more stable soluble complexes with the chelating ligand, and are washed out (eluted) first. Larger Ln³⁺ ions (e.g., La³⁺) form weaker complexes and remain bound longer, eluting last.

Actinoids (5f-Block Elements)

The series consists of 14 elements following Actinium, from Thorium (Z = 90) to Lawrencium (Z = 103), where the 5f subshell is progressively filled.

• Electronic Configuration: [Rn] 5f^(0-14) 6d^(0-2) 7s²
• Oxidation States: Exhibit a much wider range of oxidation states (+3, +4, +5, +6, +7) than lanthanoids because 5f, 6d, and 7s energy levels are very close in energy, allowing participation of all these electrons. +3 is common, but +4 is most stable for Th, and +6 for U (e.g., UO₂²⁺).
• Spectral & Magnetic Properties: Absorption spectra show f-f transitions but bands are broader and more intense than lanthanoids due to less shielding of 5f electrons. Magnetic properties are structurally more complex.