The d- and f-Block Elements

The d-block sits in the middle of the periodic table, groups 3 to 12, between the s-block and the p-block. In these elements the last electron enters the d orbital of the penultimate shell, so their outer configuration is (n−1)d^1–10 ns^1–2. There are four series: the 3d series (Sc to Zn, period 4), the 4d series (Y to Cd, period 5), the 5d series (La and then Hf to Hg, period 6) and the 6d series (Ac and then Rf to Cn, period 7). The 3d series is the one the board asks about.

A transition element is defined by NCERT as an element that has incompletely filled d orbitals in its ground state or in any of its oxidation states. This definition decides a favourite 1- or 2-mark question: Zn, Cd and Hg have a full d¹⁰ configuration both as atoms and in their common +2 state, so they are d-block elements but are not regarded as transition elements. Copper and silver, on the other hand, have 3d¹⁰ 4s¹ and 4d¹⁰ 5s¹ as atoms, but Cu²⁺ is 3d⁹ and Ag²⁺ is 4d⁹, so they count as transition elements.

The name 'transition' comes from their position: their properties are intermediate between the highly reactive s-block metals and the p-block elements. The f-block elements, in which the last electron enters the f orbital of the ante-penultimate shell, are placed separately at the bottom of the table and are called inner transition elements.

The general configuration is (n−1)d^1–10 ns^1–2. For the 3d series, write [Ar] 3d^x 4s² for most elements, with two exceptions the examiner loves: chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹. The reason is the extra stability of a half-filled and a fully filled set of d orbitals (symmetrical distribution, larger exchange energy), and the small energy gap between 3d and 4s that allows an electron to move across. In the 4d series palladium is even more irregular, 4d¹⁰ 5s⁰.

For ions, remove the ns electrons first, then the (n−1)d electrons. So Fe (3d⁶ 4s²) gives Fe²⁺ = 3d⁶ and Fe³⁺ = 3d⁵; Cu⁺ = 3d¹⁰, Cu²⁺ = 3d⁹; Mn²⁺ = 3d⁵; Cr³⁺ = 3d³; Co²⁺ = 3d⁷; Ni²⁺ = 3d⁸; Zn²⁺ = 3d¹⁰. NCERT exercise 4.1 asks exactly this for eight ions, including the f-block ones: Pm³⁺ = [Xe] 4f⁴, Ce⁴⁺ = [Xe] (4f⁰), Lu²⁺ = [Xe] 4f¹⁴ 5d¹, Th⁴⁺ = [Rn] (5f⁰).

Why do transition elements behave so differently from the s- and p-block metals? Because the (n−1)d and ns orbitals are close in energy, both sets of electrons take part in bonding. That single fact explains variable oxidation states, high enthalpies of atomisation, colour, paramagnetism and catalytic activity, and you will use it again and again in this chapter. The elements within a series also resemble each other much more closely than p-block elements in the same period, since the differentiating electron goes into an inner orbital.

Nearly all transition elements are typical metals: hard, high tensile strength, ductile, malleable, lustrous, good conductors of heat and electricity. Except for Zn, Cd, Hg and Mn, they have one of the typical close-packed metallic structures.

Their melting and boiling points are high. Across a series the melting point rises to a maximum at the d⁵ region and then falls, because the number of unpaired d electrons available for metal–metal bonding is greatest in the middle. Cr, Mo and W have the highest melting points of their series; tungsten melts near 3410 °C. Mn and Tc are anomalous: their d⁵ s² configuration is stable and the d electrons are less available for bonding, so their melting points are unexpectedly low.

The same logic gives the enthalpy of atomisation, which is high for transition metals (often above 400 kJ mol⁻¹) and again maximal in the middle of the series. The board asks 'why is the enthalpy of atomisation of zinc the lowest in the 3d series?'; the answer is that Zn has no unpaired d electrons (3d¹⁰ 4s²), so there is no d–d contribution to metallic bonding. Two more consequences you should link: metals with high enthalpy of atomisation tend to be noble, and the 4d and 5d metals have higher enthalpies of atomisation than the 3d metals because of stronger metal–metal bonding, which is why they form more metal–metal bonded compounds.

Density increases along a series (Sc about 3.4, Cu about 8.9 g cm⁻³) because atomic radius decreases while atomic mass increases. The 5d metals (Os, Ir, Pt) are the densest of all.

Atomic size. Across a series, the atomic radius first decreases (Sc to Cr) as nuclear charge grows, then stays nearly constant (Fe, Co, Ni) because the added d electrons shield the outer 4s electrons almost as fast as the nuclear charge increases, and finally rises at Cu and Zn where electron–electron repulsion in the filled d¹⁰ shell dominates. Ionic radii of M²⁺ ions follow the same pattern and decrease with increasing charge on the ion.

Down a group, the radius increases from the 3d to the 4d series as expected, but the 4d and 5d elements of the same group have almost identical radii, for example Zr 160 pm and Hf 159 pm. This is the most-asked consequence of lanthanoid contraction: the 14 lanthanoids sit between La and Hf, and the poor shielding by 4f electrons cancels the size increase that a new shell should have produced.

Ionisation enthalpy. The first ionisation enthalpies of the 3d series increase along the series, but only slightly and irregularly (from about 631 kJ mol⁻¹ for Sc to 906 for Zn), because the increase in nuclear charge is largely offset by the extra shielding from the added 3d electrons. The irregularities come from configuration: removing an electron that leaves behind a stable d⁵ or d¹⁰ is easy, and removing an electron from such a stable arrangement is hard. So the second ionisation enthalpy is unusually high for Cr (3d⁵ 4s¹ → must break d⁵) and Cu (3d¹⁰ 4s¹ → must break d¹⁰); Cu has the highest second ionisation enthalpy of the 3d series. The third ionisation enthalpy is exceptionally high for Mn (Mn²⁺ is d⁵) and for Zn (Zn²⁺ is d¹⁰). This is why Mn²⁺ resists oxidation to Mn³⁺, and why Fe²⁺ (d⁶) is oxidised to Fe³⁺ (d⁵) easily. The 5d series has higher first ionisation enthalpies than 3d and 4d because the poorly shielding 4f electrons leave a greater effective nuclear charge.

Transition elements show a wide variety of oxidation states because the (n−1)d and ns orbitals have nearly equal energies, so electrons from both can be lost or shared. The number of oxidation states rises to a maximum in the middle of the series and then falls: Sc shows only +3, Ti +2 to +4, V +2 to +5, Cr +2 to +6 (common +3 and +6), Mn +2 to +7, Fe +2 and +3 (+6 in ferrates), Co +2 and +3, Ni mainly +2, Cu +1 and +2, Zn only +2. Manganese shows the largest number of oxidation states because it has the maximum number of unpaired electrons (3d⁵ 4s²: seven electrons available). The early elements have too few d electrons and the late ones have too many paired electrons, so both ends show fewer states.

Points the examiner looks for. (1) The oxidation states differ by one unit (Fe²⁺, Fe³⁺), unlike the p-block where they differ by two (Sn²⁺, Sn⁴⁺, because of the inert pair). (2) In the p-block the lower state becomes more stable down the group; in the d-block the higher state becomes more stable down the group (Mo(VI) and W(VI) are stable while Cr(VI) is a strong oxidant). (3) The highest oxidation state of a group-3 to group-7 element equals its group number and is shown only in oxides and fluorides, because O and F are small and highly electronegative. (4) Low oxidation states, including zero, appear in complexes with π-acceptor ligands such as CO: Ni(CO)₄, Fe(CO)₅.

Stability of a state comes back to configuration. The +2 state becomes more stable from Ti to Mn because removing a third electron gets harder as nuclear charge rises, and Mn²⁺ is d⁵. Cu⁺ disproportionates in water, 2Cu⁺ → Cu²⁺ + Cu, because the hydration enthalpy of the small, doubly charged Cu²⁺ more than compensates for the second ionisation enthalpy.

E°(M²⁺/M). The value depends on the sum of the enthalpy of atomisation, the first and second ionisation enthalpies and the hydration enthalpy of M²⁺. Because ionisation enthalpies vary irregularly, E° does not follow a smooth trend. Three anomalies are asked. Copper is the only 3d metal with a positive E° (+0.34 V): its high enthalpy of atomisation and high second ionisation enthalpy are not balanced by the hydration enthalpy, so Cu does not liberate H₂ from dilute acids. Mn (−1.18 V) and Zn (−0.76 V) are more negative than expected from the trend because their M²⁺ ions are stable d⁵ and d¹⁰. Nickel's E° (−0.25 V) is more negative than the trend suggests because of the high hydration enthalpy of the small Ni²⁺ ion.

E°(M³⁺/M²⁺). Mn³⁺/Mn²⁺ is high (+1.57 V) because Mn²⁺ is d⁵; Mn³⁺ is therefore a strong oxidant. Cr³⁺/Cr²⁺ is low (−0.41 V) because Cr³⁺ is d³, a half-filled t_2g set; Cr²⁺ is therefore a strong reductant. Fe³⁺/Fe²⁺ is +0.77 V, lower than Mn because Fe³⁺ is d⁵. Co³⁺/Co²⁺ is very high (+1.97 V), so Co(II) is stable in water, but in the presence of complexing ligands it is easily oxidised to Co(III).

Higher oxidation states. The highest oxidation state is shown in fluorides and oxides: TiX₄, VF₅, CrF₆ for halides; the highest fluoride of Mn is MnF₄ but the highest oxide is Mn₂O₇. Oxygen stabilises a higher state than fluorine because it can form multiple bonds to the metal. Beyond group 7 no oxide above M₂O₃ is known for the 3d metals, though Fe(VI) exists in ferrates. Cu(II) halides are known except the iodide: Cu²⁺ oxidises I⁻ to I₂, 2Cu²⁺ + 4I⁻ → Cu₂I₂(s) + I₂.

Reactivity. The 3d metals are fairly reactive; all except Cu dissolve in 1 M H⁺, though many react slowly because of a protective oxide layer.

Magnetism. A substance is diamagnetic if it is repelled by a magnetic field (all electrons paired) and paramagnetic if attracted (unpaired electrons present). Ferromagnetism, shown by Fe, Co and Ni, is an extreme form of paramagnetism where the moments align permanently. Most transition metal ions have unpaired d electrons, so their compounds are paramagnetic. Each unpaired electron has a magnetic moment from its spin, and the spin-only moment is μ = √[n(n+2)] BM, where n is the number of unpaired electrons and BM is the Bohr magneton. Values to memorise: n = 1, 1.73; n = 2, 2.83; n = 3, 3.87; n = 4, 4.90; n = 5, 5.92 BM. The moment increases up to Mn²⁺ (d⁵, five unpaired, 5.92 BM) and then falls as pairing begins. The orbital contribution is quenched in most 3d ions, so the observed values are close to spin-only.

Worked example (intext 4.8): M²⁺ with Z = 27 is Co²⁺, 3d⁷, three unpaired electrons, μ = √(3 × 5) = √15 = 3.87 BM. Similarly Ce³⁺ (4f¹) gives √3 = 1.73 BM.

Colour. In a transition metal ion the five d orbitals are split into two energy sets by the ligands (water in aqueous solution). An electron in the lower set absorbs light of a particular frequency in the visible range and jumps to the upper set; this is a d–d transition. The colour seen is the complementary colour of the light absorbed. Ions with d⁰ (Sc³⁺, Ti⁴⁺) or d¹⁰ (Zn²⁺, Cu⁺) have no d–d transition possible and are colourless. Colours of aqueous ions from the NCERT table: Ti³⁺ purple, V³⁺ green, Cr³⁺ violet, Mn²⁺ pink, Fe²⁺ green, Fe³⁺ yellow, Co²⁺ pink, Ni²⁺ green, Cu²⁺ blue. So Zn²⁺ salts are white while Cu²⁺ salts are coloured, the most repeated 1-mark question in this chapter.

Complexes. Transition metal ions form a large number of complex compounds, for example [Fe(CN)₆]³⁻, [Cu(NH₃)₄]²⁺, [PtCl₄]²⁻, [Fe(CN)₆]⁴⁻. Three reasons, and the examiner wants all three: comparatively small size of the ions, high ionic charge, and the availability of vacant d orbitals to accept lone pairs from ligands.

Catalysis. Transition metals and their compounds are good catalysts because of (i) variable oxidation states, which let them form unstable intermediates with reactants and provide a new path of lower activation energy, and (ii) their ability to adsorb reactants on the surface through partially filled d orbitals. Examples: V₂O₅ in the Contact process, finely divided Fe in the Haber process, Ni in hydrogenation of oils. In solution, Fe(III) catalyses the reaction between iodide and persulphate: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂, then 2Fe²⁺ + S₂O₈²⁻ → 2Fe³⁺ + 2SO₄²⁻.

Interstitial compounds. Small atoms such as H, C, N and B get trapped in the voids of the metal lattice. The products are usually non-stoichiometric (TiC, Mn₄N, Fe₃H, VH_0.56, TiH_1.7) and are neither typically ionic nor covalent. Properties: high melting points (higher than the pure metal), very hard (some borides approach diamond), retain metallic conductivity, and are chemically inert. Steel is harder than iron for this reason.

Alloys. Because transition metal atoms have radii within about 15% of each other, they easily replace one another in the lattice to form homogeneous solid solutions. Alloys are hard and have high melting points. Ferrous alloys contain Cr, V, W, Mo and Mn; stainless steel is the best-known one. Non-ferrous examples: brass (Cu–Zn) and bronze (Cu–Sn).

Transition metals form oxides when heated with oxygen at high temperature. All the 3d metals except scandium form the monoxide MO, which is ionic and basic. As the oxidation state of the metal rises the M–O bond becomes more covalent and the oxide more acidic. The best example is manganese: MnO (+2) is basic, Mn₃O₄ and Mn₂O₃ are intermediate, MnO₂ (+4) is amphoteric, and Mn₂O₇ (+7) is a covalent green oil that dissolves in water to give permanganic acid HMnO₄. Similarly CrO (+2) is basic, Cr₂O₃ (+3) is amphoteric and CrO₃ (+6) is acidic, giving chromic acid H₂CrO₄ and dichromic acid H₂Cr₂O₇. Vanadium runs V₂O₃ (basic), V₂O₄ (less basic), V₂O₅ (amphoteric: it gives VO₄³⁻ with alkali and VO₂⁺ salts with acid).

The highest oxidation state of a metal appears in its oxoanions because oxygen is small, highly electronegative and can form multiple bonds. So Cr shows +6 in CrO₄²⁻ and Cr₂O₇²⁻, and Mn shows +7 in MnO₄⁻. The oxometal anions where the metal's oxidation state equals its group number are vanadate VO₄³⁻ (group 5, +5), chromate CrO₄²⁻/dichromate (group 6, +6) and permanganate MnO₄⁻ (group 7, +7); this is NCERT exercise 4.6.

Also note the ferrate ion FeO₄²⁻, where Fe is +6, formed only in alkaline medium; it shows that oxides above Fe₂O₃ do not exist but a higher oxoanion can. From here the chapter turns to the two compounds that carry most of the marks: potassium dichromate and potassium permanganate.

Preparation from chromite ore (FeCr₂O₄), three steps. (1) Fuse the ore with sodium carbonate in excess air: 4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 8Na₂CrO₄ + 2Fe₂O₃ + 8CO₂. The yellow sodium chromate is filtered from the iron oxide. (2) Acidify the filtrate with sulphuric acid to get sodium dichromate: 2Na₂CrO₄ + 2H⁺ → Na₂Cr₂O₇ + 2Na⁺ + H₂O; the less soluble Na₂SO₄ crystallises out first. (3) Treat the concentrated sodium dichromate solution with KCl; potassium dichromate, being less soluble, crystallises as orange crystals: Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇ + 2NaCl.

Effect of pH. Chromate and dichromate are interconvertible: 2CrO₄²⁻ + 2H⁺ → Cr₂O₇²⁻ + H₂O (acid turns yellow to orange) and Cr₂O₇²⁻ + 2OH⁻ → 2CrO₄²⁻ + H₂O (alkali turns orange to yellow). In both ions Cr is +6. Structure: chromate is tetrahedral; dichromate is two tetrahedra sharing one corner oxygen, with a Cr–O–Cr angle of 126°.

Oxidising action in acidic medium: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O, E° = 1.33 V. The orange solution turns green. Ionic equations the board asks: with iodide, Cr₂O₇²⁻ + 14H⁺ + 6I⁻ → 2Cr³⁺ + 7H₂O + 3I₂; with iron(II), Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O; with H₂S, Cr₂O₇²⁻ + 8H⁺ + 3H₂S → 2Cr³⁺ + 7H₂O + 3S; with tin(II), Cr₂O₇²⁻ + 14H⁺ + 3Sn²⁺ → 2Cr³⁺ + 3Sn⁴⁺ + 7H₂O. Sodium dichromate is more soluble than the potassium salt; potassium dichromate is preferred as a primary standard in volumetric analysis because it is non-hygroscopic and can be obtained pure. Heating ammonium dichromate gives the 'volcano' reaction: (NH₄)₂Cr₂O₇ → Cr₂O₃ + N₂ + 4H₂O.

Preparation from pyrolusite (MnO₂). Fuse MnO₂ with KOH in the presence of air or an oxidant like KNO₃ to get green potassium manganate: 2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O. The manganate then disproportionates in neutral or acidic solution: 3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O. Commercially the manganate is oxidised electrolytically instead: MnO₄²⁻ → MnO₄⁻ + e⁻. In the laboratory a Mn(II) salt is oxidised by peroxodisulphate: 2Mn²⁺ + 5S₂O₈²⁻ + 8H₂O → 2MnO₄⁻ + 10SO₄²⁻ + 16H⁺.

Properties. Dark purple crystals, isostructural with KClO₄; the intense colour is due to charge transfer, not d–d transition (Mn(VII) is d⁰). Manganate (green) has one unpaired electron and is paramagnetic; permanganate is diamagnetic. Both are tetrahedral. On heating to 513 K: 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂.

Oxidising action depends on the medium. Acidic: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O, E° = +1.52 V. Neutral or faintly alkaline: MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻. Strongly alkaline: MnO₄⁻ + e⁻ → MnO₄²⁻. In acid medium: with iodide, 2MnO₄⁻ + 10I⁻ + 16H⁺ → 2Mn²⁺ + 5I₂ + 8H₂O; with Fe²⁺, MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O; with oxalate, 2MnO₄⁻ + 5C₂O₄²⁻ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O; with sulphite, 2MnO₄⁻ + 5SO₃²⁻ + 6H⁺ → 2Mn²⁺ + 5SO₄²⁻ + 3H₂O; with H₂S, 2MnO₄⁻ + 5S²⁻ + 16H⁺ → 2Mn²⁺ + 5S + 8H₂O; with nitrite, 2MnO₄⁻ + 5NO₂⁻ + 6H⁺ → 2Mn²⁺ + 5NO₃⁻ + 3H₂O. In neutral or faintly alkaline medium: iodide is oxidised to iodate, 2MnO₄⁻ + H₂O + I⁻ → 2MnO₂ + 2OH⁻ + IO₃⁻; thiosulphate to sulphate, 8MnO₄⁻ + 3S₂O₃²⁻ + H₂O → 8MnO₂ + 6SO₄²⁻ + 2OH⁻; and Mn²⁺ to MnO₂, 2MnO₄⁻ + 3Mn²⁺ + 2H₂O → 5MnO₂ + 4H⁺.

Uses: volumetric estimation of Fe²⁺, oxalates, nitrites; an oxidant in organic chemistry; bleaching of wool, cotton and silk; decolourising oils. Acidify with dilute H₂SO₄, never HCl (Cl⁻ would be oxidised to Cl₂) or HNO₃ (itself an oxidant).

The 14 elements following lanthanum, cerium (Z = 58) to lutetium (Z = 71), are the lanthanoids; NCERT includes La with them because of its similar properties. The general configuration is [Xe] 4f^1–14 5d^0–1 6s², and the tripositive ions run from 4f¹ (Ce³⁺) to 4f¹⁴ (Lu³⁺), so the +3 state is the characteristic one.

Lanthanoid contraction is the steady decrease in atomic and ionic (Ln³⁺) radii from La to Lu. Cause: as atomic number increases, each added electron enters a 4f orbital, and 4f electrons shield each other very imperfectly from the nuclear charge, so the effective nuclear charge felt by the outer electrons increases and the electron cloud is pulled in. Consequences the examiner expects: (i) the 4d and 5d elements of a group have almost identical radii (Zr 160 pm, Hf 159 pm) and are hard to separate; (ii) the lanthanoids resemble one another so closely that their separation needs ion-exchange methods; (iii) the basic strength of the hydroxides decreases from La(OH)₃ to Lu(OH)₃ as the ion gets smaller and the M–OH bond more covalent; (iv) the ionisation enthalpies of the 5d transition metals are higher than those of 4d.

Oxidation states. +3 dominates. +4 is shown by Ce (4f⁰ after losing four electrons; Ce⁴⁺ is a good oxidant, E° = +1.74 V, and is reduced to Ce³⁺ even by water slowly), and also by Pr, Nd, Tb and Dy (only in their oxides, MO₂). +2 is shown by Eu (Eu²⁺ is 4f⁷, a strong reductant), Yb (Yb²⁺ is 4f¹⁴) and Sm. The pattern follows the extra stability of empty, half-filled and full f shells.

General characteristics. Silvery white soft metals that tarnish in air; melting points 1000–1200 K (Sm 1623 K). Most Ln³⁺ ions are coloured and paramagnetic because of unpaired 4f electrons; La³⁺ (4f⁰) and Lu³⁺ (4f¹⁴) are colourless and diamagnetic. Chemically they resemble the alkaline earths: they burn in O₂ to Ln₂O₃, react with halogens to LnX₃, with water (slowly cold, fast hot) to Ln(OH)₃ + H₂, with N₂ to LnN, with S to Ln₂S₃, with carbon at 2773 K to carbides, and with dilute acids to liberate H₂. Mischmetall, about 95% lanthanoid metal with about 5% iron and traces of S, C, Ca and Al, is used in Mg-based alloys for bullets, shells and lighter flints.

The 14 elements after actinium, thorium (Z = 90) to lawrencium (Z = 103), are the actinoids. All are radioactive; those after uranium are man-made transuranium elements with short half-lives, so their chemistry is less studied. General configuration: [Rn] 5f^1–14 6d^0–1 7s². The 5f, 6d and 7s orbitals are close in energy, so the configurations are irregular.

Actinoid contraction is the gradual decrease in the size of the atoms and M³⁺ ions across the series. It is greater from element to element than the lanthanoid contraction because 5f electrons shield the nuclear charge even more poorly than 4f electrons (5f orbitals are more diffuse and penetrate less).

Oxidation states. +3 is the most common, as in lanthanoids, but the actinoids show a much wider range: Th shows +4 (Th⁴⁺ is 5f⁰), Pa, U and Np show +5, U and Pu show +6, and Np shows +7. This is because the 5f, 6d and 7s orbitals have comparable energies and the 5f electrons are less tightly held. The higher states become less stable as the series progresses; +4 is common in the early members and +3 in the later ones.

General properties. Silvery, highly reactive metals, especially when finely divided; boiling water on them gives a mixture of oxide and hydride, and they combine with most non-metals at moderate temperature. Hydrochloric acid attacks all the actinoid metals, but most are only slightly affected by nitric acid because it forms a protective oxide layer; alkalies have no action. The early actinoids have lower ionisation enthalpies than the early lanthanoids because 5f electrons are less penetrating.

Comparison, the way the board wants it. Similarities: a predominant +3 state, a contraction in size, electropositive and paramagnetic. Differences: actinoids show a wider range of oxidation states; actinoid contraction is greater per element; actinoids are more basic, form complexes more readily, are radioactive and have more complex magnetic behaviour; lanthanoid chemistry is 'smoother' because 4f electrons are well shielded and take little part in bonding.

This short section is the source of case-based passages and 1-mark 'name the catalyst' questions, so learn the list. Iron and steel are the most important construction materials; steel is made by adding controlled amounts of carbon and metals such as Cr, Mn and Ni to iron. TiO₂ is the pigment industry's white; MnO₂ goes into dry cells; Zn is the anode of dry cells and Ni and Cd are used in rechargeable Ni–Cd batteries. Cu, Ag and Au (Group 11) are the coinage metals, though contemporary UK 'copper' coins are copper-coated steel and 'silver' coins a Cu/Ni alloy; Pt and Ag are also used in jewellery. Silver bromide is the light-sensitive salt of black-and-white photography.

Catalysts: V₂O₅ converts SO₂ to SO₃ in the Contact process for sulphuric acid; finely divided iron catalyses the Haber process for ammonia; nickel catalyses the hydrogenation of oils to fats; TiCl₄ with Al(CH₃)₃ is the Ziegler catalyst for polyethylene; PdCl₂ catalyses the Wacker process, the oxidation of ethyne to ethanal; nickel complexes are used in the polymerisation of alkynes. The reason all of these work goes back to variable oxidation states and surface adsorption.

f-block uses: mischmetall in lighter flints and alloys for bullets and shells; mixed lanthanoid oxides as catalysts in petroleum cracking; some lanthanoid oxides as phosphors in television screens; and the actinoids uranium and plutonium as nuclear fuels. If a question asks for 'an alloy containing lanthanoid metals and its use', the answer is mischmetall with its composition and use.