The d- and f-Block Elements

In this chapter: position and electronic configuration of d-block elements, trends in size, ionisation enthalpy and oxidation states, electrode potentials, magnetic moment, coloured ions, complex formation, catalytic behaviour, interstitial compounds and alloys, potassium dichromate and potassium permanganate, lanthanoids and lanthanoid contraction, actinoids.

Position and electronic configuration

The d-block occupies groups 3 to 12 of the periodic table. There are four transition series: 3d (Sc to Zn), 4d (Y to Cd), 5d (La, Hf to Hg) and 6d (Ac, Rf to Cn). The f-block, placed separately at the bottom, has the lanthanoids (Ce to Lu, filling 4f) and the actinoids (Th to Lr, filling 5f).

A transition element is one that has an incompletely filled d subshell in its ground state or in any of its common oxidation states. Zinc, cadmium and mercury have a full d¹⁰ configuration in the ground state and in their common oxidation states, so they are not regarded as transition metals, although they are studied with them. Copper (3d¹⁰4s¹) counts as a transition metal because Cu²⁺ is 3d⁹.

General configuration: (n−1)d¹⁻¹⁰ns¹⁻². Palladium (4d¹⁰5s⁰) is an exception. In the 3d series, Cr is 3d⁵4s¹ and Cu is 3d¹⁰4s¹, because half-filled and completely filled d subshells are extra stable. When a transition metal forms an ion, the 4s electrons are removed first: Fe is [Ar]3d⁶4s², Fe²⁺ is [Ar]3d⁶ and Fe³⁺ is [Ar]3d⁵.

General properties of transition elements

Physical properties and enthalpy of atomisation

Almost all transition elements are hard, lustrous metals with high melting and boiling points, typical metallic structures and good thermal and electrical conductivity. Zn, Cd and Hg are exceptions with low melting points (Hg is a liquid). The strong metallic bonding comes from the participation of (n−1)d as well as ns electrons, which also gives high enthalpies of atomisation. These peak near the middle of each series, where the number of unpaired d electrons is largest. Metals of the 4d and 5d series have higher enthalpies of atomisation than the 3d metals, so heavy transition metals more often form metal-metal bonds.

Atomic and ionic sizes

Across a series, ions of the same charge become smaller as nuclear charge increases, because a new d electron shields the others poorly. The decrease is small in the middle of the 3d series. The 5d elements are almost the same size as the 4d elements (Zr 160 pm, Hf 159 pm) because of the lanthanoid contraction: the 4f subshell fills before the 5d series begins, and f electrons shield the nucleus very poorly.

Ionisation enthalpy

Ionisation enthalpy increases along each series, but less steeply than for main group elements, and not smoothly. Irregularities come from the extra stability of d⁵ and d¹⁰ configurations. For example, the third ionisation enthalpy of Mn is very high because Mn²⁺ (d⁵) resists losing an electron, while the second ionisation enthalpies of Cr and Cu are high because removing an electron from Cr⁺ (d⁵) or Cu⁺ (d¹⁰) breaks a stable configuration.

Oxidation states

• Variable oxidation states arise because (n−1)d and ns electrons have similar energies. The largest number of states is shown in the middle of the series; Mn shows every state from +2 to +7.
• Early members (Sc, Ti) show fewer states because they have few electrons to lose; late members (Cu, Zn) because they have too many d electrons and few vacant orbitals to use.
• In the p-block the lower oxidation state becomes more stable down a group (inert pair effect). In the d-block the opposite is true: Mo(VI) and W(VI) are more stable than Cr(VI).
• Fluorine and oxygen stabilise the highest oxidation states. The highest fluoride of Mn is MnF₄, but its highest oxide is Mn₂O₇, because oxygen can form multiple bonds to the metal.

Standard electrode potentials

The E°(M²⁺/M) values of the 3d metals are mostly negative and become less negative across the series. Copper is the only 3d metal with a positive E°(M²⁺/M) (+0.34 V), so it does not liberate H₂ from dilute acids. Its high enthalpy of atomisation plus ionisation enthalpies is not balanced by the hydration enthalpy of Cu²⁺.

Among E°(M³⁺/M²⁺) values, Mn³⁺/Mn²⁺ is high (+1.57 V) because Mn²⁺ is a stable d⁵ ion, while Fe³⁺/Fe²⁺ is comparatively low (+0.77 V) because Fe³⁺ is d⁵. Cr²⁺ is a strong reducing agent since it readily becomes Cr³⁺, whose t_2g³ configuration (in water) is especially stable; Mn³⁺ and Co³⁺ are strong oxidising agents in solution.

Magnetic properties

Substances with unpaired electrons are paramagnetic (attracted by a magnetic field); those with all electrons paired are diamagnetic. For 3d ions, the magnetic moment can be estimated from the number of unpaired electrons n using the spin-only formula:

Worked example: Calculate the spin-only magnetic moment of Fe²⁺ (Z = 26). Which 3d M²⁺ ion would have μ = 3.87 BM?
Solution: Fe²⁺ is 3d⁶; in a free ion Hund's rule gives four unpaired electrons, so μ = √(4 × 6) = √24 = 4.90 BM. For μ = 3.87 BM, n(n + 2) = 15, so n = 3: a d³ ion (V²⁺) or a d⁷ ion (Co²⁺).

Coloured ions

In a transition metal ion with a partly filled d subshell, an electron can jump from a lower d orbital to a higher one (a d-d transition) by absorbing a particular wavelength of visible light. The colour seen is the complementary colour of the light absorbed. The ligands decide the size of the gap, so the same ion can have different colours in different complexes. Ions with d⁰ (Sc³⁺, Ti⁴⁺) or d¹⁰ (Zn²⁺, Cu⁺) configurations are colourless.

Complexes, catalysis, interstitial compounds and alloys

• Complex formation: small ions with high charge and vacant d orbitals of suitable energy bind ligands easily, e.g. [Fe(CN)₆]³⁻, [Cu(NH₃)₄]²⁺, [Ni(CN)₄]²⁻.
• Catalytic activity: comes from variable oxidation states and the ability to form complexes and to adsorb reactants on the surface. Examples: V₂O₅ in the contact process, finely divided iron in the Haber process, nickel in hydrogenation. Fe³⁺ catalyses the reaction between iodide and persulphate: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂, then 2Fe²⁺ + S₂O₈²⁻ → 2Fe³⁺ + 2SO₄²⁻.
• Interstitial compounds: small atoms such as H, C or N get trapped in the gaps of the metal lattice, e.g. TiC, Mn₄N, Fe₃H, VH_0.56. They are usually non-stoichiometric, have higher melting points than the pure metals, are very hard, keep metallic conductivity and are chemically inert.
• Alloys: transition metals have similar atomic radii, so one can replace another in the lattice. Examples: brass (Cu and Zn), bronze (Cu and Sn), and the many ferrous alloys used in steel.

Potassium dichromate, K₂Cr₂O₇

Both of the following compounds are listed in the JEE Main and NEET syllabi, so learn their preparation and oxidising reactions.

Chromate and dichromate change into each other with pH: 2CrO₄²⁻ (yellow) + 2H⁺ → Cr₂O₇²⁻ (orange) + H₂O, and Cr₂O₇²⁻ + 2OH⁻ → 2CrO₄²⁻ + H₂O. The chromate ion is tetrahedral; the dichromate ion is two tetrahedra sharing one corner, with a Cr-O-Cr angle of 126°. Chromium is +6 in both.

In acid, dichromate is a strong oxidising agent: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O (E° = 1.33 V). It oxidises I⁻ to I₂, Fe²⁺ to Fe³⁺, Sn²⁺ to Sn⁴⁺ and H₂S to S. Being available pure, K₂Cr₂O₇ is used as a primary standard in volumetric analysis.

Potassium permanganate, KMnO₄

Both ions are tetrahedral. The green manganate ion (Mn +6, d¹) is paramagnetic; the purple permanganate ion (Mn +7, d⁰) is diamagnetic, its colour arising from charge transfer rather than a d-d transition.

• Acidic medium: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (E° = 1.52 V). Oxidises I⁻ to I₂, Fe²⁺ to Fe³⁺, oxalate to CO₂, H₂S to S, sulphite to sulphate and nitrite to nitrate. The acid used is dilute H₂SO₄; HCl is avoided because permanganate oxidises it to Cl₂.
• Neutral or faintly alkaline medium: MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻. Oxidises I⁻ to iodate (IO₃⁻), thiosulphate to sulphate and Mn²⁺ to MnO₂.

The lanthanoids

• Configuration: the 4f subshell fills from Ce to Lu; general configuration [Xe]4f¹⁻¹⁴5d⁰⁻¹6s².
• Oxidation state: +3 is characteristic. Ce⁴⁺ (4f⁰) exists and is a good oxidising agent (E°(Ce⁴⁺/Ce³⁺) = +1.74 V); Eu²⁺ (4f⁷) and Yb²⁺ (4f¹⁴) act as reducing agents; Tb⁴⁺ (4f⁷) is an oxidant.
• Lanthanoid contraction: atomic and ionic (M³⁺) radii decrease steadily from La to Lu, because the added 4f electrons shield each other poorly from the growing nuclear charge. Consequences: 4d and 5d elements of the same group have nearly equal radii (Zr/Hf, Nb/Ta), so they are very hard to separate; the basic strength of the hydroxides decreases from La(OH)₃ to Lu(OH)₃.
• Chemistry: early members are about as reactive as calcium. They burn in O₂ to give Ln₂O₃, liberate H₂ from dilute acids, and form carbides and nitrides on heating with C and N₂.
• Uses: mischmetall (about 95% lanthanoid metal and 5% iron, with traces of S, C, Ca and Al) is used in a magnesium-based alloy for bullets, shells and lighter flints. Mixed lanthanoid oxides are used as catalysts in petroleum cracking.

The actinoids

• All actinoids are radioactive; the later members have very short half-lives and have been made only in nanogram quantities.
• They show a wider range of oxidation states than lanthanoids (up to +7 for Np and Pu), because the 5f, 6d and 7s levels are close in energy. +3 is common, but for the early members higher states such as +4, +5 and +6 are important (e.g. U shows +6).
• Actinoid contraction from element to element is greater than the lanthanoid contraction, because 5f electrons shield even more poorly than 4f electrons.

Common mistakes: (1) Removing 3d electrons before 4s when writing ion configurations. (2) Calling Zn a transition element; it is d-block but not transition. (3) Taking the colour of MnO₄⁻ as a d-d transition; Mn(VII) is d⁰. (4) Forgetting that in acid KMnO₄ gains 5 electrons but K₂Cr₂O₇ gains 6 per formula unit. (5) Explaining the similar size of Zr and Hf by anything other than the lanthanoid contraction.

JEE and NEET focus

• Electronic configurations of 3d atoms and ions, including Cr and Cu, and the definition of a transition element.
• Trends in oxidation states, E° values and stability, with reasons based on d⁰, d⁵, d¹⁰ and t_2g³ configurations.
• Spin-only magnetic moment calculations and the reason for colour; which ions are colourless.
• Preparation of K₂Cr₂O₇ and KMnO₄, chromate-dichromate equilibrium, structures, and balanced oxidising reactions in acidic and neutral media.
• Lanthanoid contraction, its causes and consequences; oxidation states of Ce, Eu and Yb; actinoids compared with lanthanoids.

Practice questions