Revision: d-block and f-block Elements Chemistry Science (English Medium) Class 12 CBSE

Elements having partially filled d-orbitals in ground state or in excited state are known as transition elements.

In the lanthanoids, the electrons are filling the 4f-subshell. On moving from left to right, the nuclear charge increases, and this increase is expected to be compensated by the increase in the magnitude of the shielding effect by the 4f-electrons. However, the f-electrons have very poor shielding effects. With an increasing atomic number in the lanthanoid series, there is a progressive decrease in the atomic as well as ionic radii of trivalent ions from La³⁺ to Lu^3+, and this is known as lanthanoid contraction.

A homogeneous mixture of two or more metals or a metal and a non-metal is called alloy.

The property of transition elements to exhibit more than one oxidation state is called variable oxidation state.

The electrode potential of a half-cell measured under standard conditions is called standard electrode potential.

The extra stability associated with parallel spin electrons in degenerate orbitals is called exchange energy.

The elements in which the last electron enters the (n–1)d subshell are called d-block elements.

The elements which have incompletely filled d-orbitals in their ground state or in any of their common oxidation states are called transition elements.

The elements in which electrons are progressively filled in the 4f or 5f orbitals are called inner transition elements.

The series of elements from cerium (Ce) to lutetium (Lu) in which 4f orbitals are progressively filled are called lanthanoids.

The series of elements from thorium (Th) to lawrencium (Lr) in which 5f orbitals are progressively filled are called actinoids.

The gradual decrease in atomic and ionic radii of lanthanoids with increasing atomic number is called lanthanoid contraction.

Compounds formed when small atoms like H, C or N occupy interstitial spaces in the crystal lattice of metals are called interstitial compounds.

The property of a substance to get attracted in a magnetic field due to the presence of unpaired electrons is called paramagnetism.

The property of a substance to get weakly repelled by a magnetic field due to the absence of unpaired electrons is called diamagnetism.

The property of a substance to be strongly attracted by a magnetic field and retain magnetism is called ferromagnetism.

Statement:
The standard electrode potential values of transition metals depend on ionisation enthalpy, enthalpy of atomisation and hydration enthalpy.

Explanation:

• More negative E° → stronger reducing agent.

• Mn and Zn show highly negative E° values.

• Cu shows positive E° value due to high ionisation enthalpy.

Statement:
Ionisation enthalpy generally increases across a transition series due to increase in effective nuclear charge.

Explanation:

• Increase in nuclear charge pulls electrons strongly.

• Variation is not smooth due to extra stability of half-filled and fully filled configurations.

• Removal of 4s electrons occurs before 3d electrons during ion formation.

Statement:
Transition elements exhibit variable oxidation states due to the comparable energies of ns and (n−1)d orbitals.

Explanation:

• Both ns and d electrons participate in bonding.

• Early elements show higher oxidation states.

• Middle elements show maximum oxidation states.

• Later elements prefer lower oxidation states.

Example:
Mn shows +2 to +7 oxidation states.

Statement:
The magnetic behaviour of transition metal ions depends on the number of unpaired electrons present in their d-orbitals.

Explanation:

• Presence of unpaired electrons → Paramagnetic

• Absence of unpaired electrons → Diamagnetic

• Magnetic moment increases with increase in number of unpaired electrons

• Given by formula: \[\mu=\sqrt{n(n+2)}\]

Statement:
Transition metal ions are coloured due to d–d electronic transitions in the presence of ligands.

Explanation:

• Absorption of visible light promotes electron from lower to higher d-orbital.

• Colour observed is complementary to absorbed light.

• d⁰ and d¹⁰ configurations are colourless (e.g., Zn²⁺, Sc³⁺).

Statement:
Transition metals act as catalysts due to their variable oxidation states and ability to form intermediate complexes.

Explanation:

• They provide alternate reaction pathway.

• Lower activation energy.

• Surface adsorption of reactants.

Examples:

• Fe in Haber process
• V₂O₅ in Contact process
• Ni in hydrogenation

Statement:
The most stable oxidation state of lanthanoids is +3 due to the stable 4f electronic configuration.

Explanation:

• Some show +2 and +4 states due to extra stability of empty, half-filled or fully filled f-orbitals.

• Ce⁴⁺ and Eu²⁺ are important examples.

d-Block Elements:

• Elements in which the last electron enters (n−1)d orbital.
• Located in the middle of the periodic table.
• Also called transition elements.
• Show variable oxidation states, colored compounds, complex formation.

General Electronic Configuration: \[(n-1)d^{1-10}ns^{0-2}\]

f-Block Elements:

• Elements in which the last electron enters (n−2)f orbital.
• Placed separately at the bottom of the periodic table.
• Known as inner transition elements.

Two series:

• Lanthanides (4f): Ce (58) → Lu (71)
• Actinides (5f): Th (90) → Lr (103)

General Electronic Configuration: \[(n-2)f^{1-14}(n-1)d^{0-1}ns^{2}\]

• Located between s-block and p-block elements
• Occupy Groups 3 to 12
• Present in 4 periods (4th to 7th)

Series of d-block:

Atomic and Ionic Radii:

• Atomic and ionic radii of d-block elements are smaller than s-block but larger than p-block elements.
• Within a 3d series, atomic radii decrease for the first five elements (Sc to Mn), then remain almost constant for the next five (Fe to Zn). This is because the increase in ENC (effective nuclear charge) first causes shrinkage, but additional d-electrons increase shielding and counterbalance further shrinkage.
• The 4d and 5d series elements have larger atomic and ionic radii than 3d series elements (due to more electron shells). However, 4d and 5d elements have nearly the same size — due to lanthanoid contraction.

Atomic Volume and Density:

• Atomic volume decreases along a period (as atomic size decreases).
• Density increases along the period.

Melting and Boiling Points:

• All transition elements have high melting points (typically above 900°C) in their solid state.
• Zn, Cd, Hg have abnormally low melting points because their completely filled d-orbitals prevent strong covalent metallic bonding.
• As unpaired electrons increase, metallic bonding strengthens → higher melting point. Tungsten (W) has the highest melting point of all metals.
• Mn and Tc have abnormally low melting points.

Enthalpies of Atomisation:

• Due to strong interatomic attraction, transition metals have high enthalpies of atomisation.
• Greater the number of valence electrons → stronger metallic bonding → higher enthalpy of atomisation.
• Members of 4d and 5d series have greater enthalpy of atomisation than 3d series.

Ionisation Energies:

• IE values of d-block elements lie between those of s-block and p-block elements.
• IE first increases up to Mn, then becomes irregular or constant due to the irregular trend of atomic size in 3d series.
• IE of Zn, Cd, and Hg are abnormally high due to the greater stability of completely filled d-subshells.
• The first two IE values of Ni are lower than Pt → Ni(II) compounds are more thermodynamically stable than Pt(II).

IE₁ order (important anomalies):

1. Hg > Cd > Zn
2. Au > Cu > Ag
3. Pt > Pd > Ni

Oxidation States:

All transition elements except the first and last of each series show a number (variable) of oxidation states.

• Mn shows the maximum number of oxidation states in the first series (7 states) — because it has 5 unpaired 3d electrons + 2 s-electrons available.
• Higher oxidation states are more stable for heavier members of a group (e.g., Mo(VI) and W(VI) are more stable than Cr(VI)).
• Lower oxidation states are more stable for lighter (3d) members.

Standard Electrode Potential:

• No regular trend exists in E° (M²⁺/M) values because IE and sublimation enthalpies show irregular variation.
• SRP tends to become more positive across a period (left to right) due to increasing IE and decreasing atomic size.
• Within a group, SRP becomes more negative going down.

• E° for Ni²⁺/Ni and Zn²⁺/Zn are more negative than expected. The high negative value of Ni²⁺/Ni stabilises Ni²⁺ ions. The high negative value for Zn²⁺/Zn is due to the stable, completely filled 3d¹⁰ configuration.
• Cr²⁺ is a strong reducing agent (acts as a reducing agent, gets oxidised to Cr³⁺; the d³ configuration = t₂g³ is very stable).
• Mn³⁺ (d⁴) is an oxidising agent — it gets reduced to Mn²⁺ (d⁵), which has an exactly half-filled d-orbital (extra stability).
• E°(Mn²⁺/Mn) is more negative than expected — due to extra stability of the half-filled 3d⁵ (Mn²⁺) ion.

Coloured Ions: Most of the transition metal compounds (ionic as well as covalent) are coloured both in the solid and in aqueous solution, in contrast to the compounds of s and p-block elements.

Magnetic Properties: In the case of transition metals, as they contain unpaired electrons in (n – 1)d orbitals, most of the transition metal ions and their compounds are paramagnetic.

Magnetic moment is calculated by spin only formula viz.

\[\mu=\sqrt{n\left(n+2\right)}\mathrm{~B.M.}\]

where n = number of unpaired electrons

I. K₂Cr₂O₇ (Potassium dichromate)

Preparation: From chromite (FeO·Cr₂O₃)

4FeO·Cr₂O₃ + O₂ → 2Fe₂O₃ + 4Cr₂O₃

4Na₂CO₃ + 2Cr₂O₃ + 3O₂ → 4Na₂Cr₂O₄ + 4CO₂ × 2

4FeO·Cr₂O₃ + 8Na₂CO₃ + 7O₂ → 8Na₂Cr₂O₄ + 2Fe₂O₃ + 8CO₂

2Na₂CrO₄ + 8H₂SO₄ → Na₂Cr₂O₇ + Na₂SO₄ + H₂O

Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇ + 2NaCl

Properties

Oxidising Properties:

1. Liberates I₂ from KI
   K₂Cr₂O₇ + 7H₂SO₄ + 6KI → 4K₂SO₄ + Cr₂(SO₄)₃ + 3I₂ + 7H₂O
2. Oxidises ferrous salts to ferric salts:
   K₂Cr₂O₇ + 7H₂SO₄ + 6FeSO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 3Fe₂(SO₄)₃ + 2H₂O
3. Oxidises H₂S to sulphur:
   K₂Cr₂O₇ + 4H₂SO₄ + 3H₂S → K₂SO₄ + Cr₂(SO₄)₃ + 7H₂O + 3S
4. Oxidises sulphites to sulphates:
   K₂Cr₂O₇ + 4H₂SO₄ + 3Na₂SO₃ → K₂SO₄ + Cr₂(SO₄)₃ + 4H₂O + 3Na₂SO₄
5. Oxidises nitrites to nitrates:
   K₂Cr₂O₇ + 4H₂SO₄ + 3NaNO₂ → K₂SO₄ + Cr₂(SO₄)₃ + 3NaNO₃ + 4H₂O

Structures:

II. Potassium Permanganate (KMnO₄)

Preparation: From Pyrolusite (MnO₂)

2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O

2MnO₂ + 2K₂CO₃ + O₂ → 2K₂MnO₄ + 2CO₂

MnO₂ + 2KOH + KNO₃ → K₂MnO₄ + KNO₂ + H₂O

3K₂MnO₄ + 2CO₂ → 2KMnO₄ + MnO₂↓ + 2K₂CO₃

Properties

Oxidising Properties:

1. Oxidising H₂S to S
   2KMnO₄ + 3H₂SO₄ + 5H₂S → K₂SO₄ + 2MnSO₄ + 3H₂O + 5S
2. Oxidises sulphur dioxide to sulphuric acid
   2KMnO₄ + 5SO₂ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄
3. Oxidises oxalates to CO₂
   2KMnO₄ + 3H₂SO₄ + 5C₂H₂O₄ → K₂SO₄ + 2MnSO₄ + 8H₂O + 10CO₂
4. Oxidises HX to X₂, where X = Cl, Br, I
   2KMnO₄ + 3H₂SO₄ + 10HX → K₂SO₄ + 2MnSO₄ + 8H₂O + 5X₂

Structure:

Position and Introduction:

The f-block consists of two series:

1. Lanthanoids — fourteen elements from Ce (58) to Lu (71), following Lanthanum (La, 57)
2. Actinoids — fourteen elements from Th (90) to Lr (103), following Actinium (Ac, 89)

In lanthanides, electrons enter the penultimate (4f) and pre-penultimate subshells.

General configuration of lanthanoids: [Xe] 4f¹⁻¹⁴ 5d⁰⁻¹ 6s²

Complete Lanthanoid Table:

La (5d¹6s²), Gd (4f⁷5d¹6s²), and Lu (4f¹⁴5d¹6s²) have a 5d¹ electron — they fill 5d before filling 4f again, due to the stability of half-filled (4f⁷) configuration.

Physical State:

• All are silvery white metals with tensile strength; good conductors of heat and electricity.
• Density ranges from 6.77 to 9.74 g/cm³ and increases with atomic number.
• They readily form alloys with other metals, especially iron.

Oxidation States

• Most common and stable OS = +3
• Some exist in +2 (Sm²⁺, Eu²⁺, Tm²⁺, Yb²⁺) — because they achieve stability trying to reach +3 OS.
• Some exist in +4 (Ce⁴⁺, Pr⁴⁺, Tb⁴⁺, Dy⁴⁺) — because they try to approach +3 from +4; hence, these are good oxidising agents.
• Elements in +2 OS act as reducing agents; in +4 OS act as oxidising agents.

Chemical Behaviour:

Lanthanoids (Ln) react with:

• Lanthanoids react with boiling water to give a mixture of oxide and hydride.
• They combine with most non-metals at moderate temperatures.
• Alkalis have no action on them.

Physical Properties:

• Silvery white, soft metals with moderate density
• Good conductors of heat and electricity

Basic Character:

• Hydroxides are ionic and basic
• Basicity decreases from La(OH)₃ to Lu(OH)₃

Ionisation Enthalpy:

• Gradually decreases from La to Lu with irregularities

Oxidation States:

• The most common oxidation state is +3
• Also show +2 and +4

Colour and Spectra:

• Colour due to f–f transitions

Atomic and Ionic Radii (Lanthanoid Contraction):

• Radii decrease from La to Lu
• Due to the poor shielding of the 4f electrons

Magnetic Property:

• Most lanthanoids are paramagnetic

Chemical Reactivity:

• Form:
  • Oxides: Ln₂O₃
  • Hydroxides: Ln(OH)₃
  • Halides: LnX₃
  • Nitrides: LnN
  • Carbides: LnC₂

General electronic configuration: \[[Xe]4f^{0-14}5d^{0-2}6s^{2}\]

Actinoids are the 14 elements from Th (90) to Lr (103). General configuration: [Rn] 5f¹⁻¹⁴ 6d⁰⁻² 7s²

Complete Actinoid Table:

• Silvery white metals with high density and high melting/boiling points
• All are radioactive
• Show variable oxidation states (common: +3; higher states in early elements)
• Strong reducing agents and highly reactive
• React with O₂, halogens, H₂, S (similar to lanthanoids)
• React with hot water to form hydroxides and H₂ gas
• Atomic and ionic radii decrease from Ac to Lr (actinoid contraction)
• 5f orbitals show poor shielding → irregular electronic configuration
• Electron distribution in 5f orbitals is less certain than 4f (lanthanoids)