The d- and f-Block Elements: NEET notes

1. Position and electronic configuration

NCERT §4.1; §4.2

• The d-block is groups 3 to 12, where the d orbitals of the second-outermost shell fill. The f-block is the separate two-row panel at the foot of the table, where 4f and 5f orbitals fill. The elements are called transition and inner transition metals.
• Four transition series: 3d (Sc to Zn), 4d (Y to Cd), 5d (La, then Hf to Hg) and 6d (Ac, then Rf to Cn). The 4f series (Ce to Lu) are the lanthanoids and the 5f series (Th to Lr) the actinoids.
• IUPAC definition: a transition metal has an incomplete d subshell in the neutral atom or in its ions. Zn, Cd and Hg keep a full d¹⁰ set in the ground state and in their usual oxidation states, so they are not transition metals, though they are studied with them as the last members of each series.
• Outer configuration in general: (n−1)d¹⁻¹⁰ ns¹⁻². Palladium is the exception to the ns part: 4d¹⁰ 5s⁰.
• Because the (n−1)d and ns levels are so close in energy, half-filled and filled d sets win: Cr is 3d⁵ 4s¹ (not 3d⁴ 4s²) and Cu is 3d¹⁰ 4s¹ (not 3d⁹ 4s²).
• Zn, Cd, Hg and Cn share the configuration (n−1)d¹⁰ ns².
• Scandium (3d¹ 4s²) counts as a transition element because its atom has an incomplete 3d set; zinc does not, because it is d¹⁰ both as an atom and as Zn²⁺. Silver is 4d¹⁰ 5s¹ as an atom, but in the +2 state it is 4d⁹, so it is a transition element.
• d orbitals reach further out towards the edge of the atom than s and p orbitals, so they feel the surroundings more and affect neighbouring atoms. Partly filled d orbitals lead to variable oxidation states, coloured ions, complex formation, catalytic activity and paramagnetism.
• Elements along a transition row resemble one another more closely than the members of a main-group period do, though some group similarities exist too.

2. Physical properties and atomic size

NCERT §4.3.1; §4.3.2

• They behave as metals should: they are strong in tension, can be drawn into wires and beaten into sheets, carry heat and electricity well and shine. All of them except Zn, Cd, Hg and Mn adopt one or more of the usual metallic crystal structures at room temperature.
• Except Zn, Cd and Hg they are very hard, of low volatility, and have high melting and boiling points, because (n−1)d electrons join the ns electrons in metallic bonding.
• Along each row the melting point rises to a maximum near d⁵ and then falls steadily, with Mn and Tc as anomalies.
• Enthalpies of atomisation are high and peak near the middle of each series, where one unpaired electron per d orbital gives the strongest interatomic bonding. First-row values run from 326 (Sc) up to 515 kJ mol⁻¹ (V) and fall to 126 kJ mol⁻¹ for Zn, whose 3d electrons take no part in metallic bonding.
• A metal with a very high enthalpy of atomisation tends to be noble in its reactions. The 4d and 5d metals have higher enthalpies of atomisation than the 3d metals, which is one reason metal-metal bonds are common in compounds of the heavy transition metals.
• Ions of the same charge shrink across a series, and atomic radii shrink too, though only slightly: each added d electron shields the growing nuclear charge poorly.
• Radii grow from the 3d to the 4d series, but the 5d radii are almost the same as the 4d ones. The 4f orbitals fill between them, and the steady shrinkage this causes (the lanthanoid contraction) cancels the expected increase. Zr (160 pm) and Hf (159 pm) are the classic pair.
• Smaller radii with rising atomic mass mean density increases across the row: 4.1 g cm⁻³ for Ti up to 8.9 g cm⁻³ for Cu.

3. Ionisation enthalpies

NCERT §4.3.3

• Ionisation enthalpy rises from left to right along each series because the nuclear charge grows as the inner d set fills.
• The rise is gentle: successive ionisation enthalpies of transition elements do not climb as steeply as those of main-group elements, and the spread across a series is much smaller than across a main-group period. First values for Sc to Cu lie between 631 and 762 kJ mol⁻¹; Zn is 906.
• When d-block atoms form ions, the ns electrons leave before the (n−1)d electrons. So Fe (3d⁶ 4s²) becomes Fe²⁺ 3d⁶, and every M²⁺ and M³⁺ ion of the row is a pure dⁿ ion.
• The irregular first ionisation enthalpies come from the way removing one electron shifts the relative energies of 4s and 3d. The 3d electrons shield the 4s electrons fairly well, so radii and first ionisation enthalpies change only a little.
• Three terms set an ionisation enthalpy: attraction to the nucleus, electron-electron repulsion and exchange energy. Exchange energy grows with the number of parallel-spin pairs in a degenerate set and stabilises it, so a configuration that would lose exchange energy on ionisation holds its electron more tightly.
• The steady rise breaks at d⁵: the second ionisation enthalpy of Mn (making Mn²⁺) and the third of Fe (making Fe³⁺). Mn⁺ is 3d⁵ 4s¹ while Cr⁺ is d⁵, so Mn's second value (1509) is below Cr's (1592); Fe²⁺ is d⁶ while Mn²⁺ is d⁵, so Fe's third value (2962) is below Mn's (3260).
• Second ionisation enthalpies are unusually high for Cr (1592) and Cu (1958), where M⁺ is d⁵ and d¹⁰. Zn's is comparatively low (1734), because it removes a 4s electron and leaves a stable d¹⁰ ion.
• Third ionisation enthalpies are generally high; the very high values for Ni, Cu and Zn explain why these metals rarely go above +2.
• Ionisation enthalpies are a guide to the stability of oxidation states, but the full picture is complex and does not reduce to a simple rule.

4. Oxidation states

NCERT §4.3.4

• Transition elements show many oxidation states. The widest range is in or near the middle of the series: Mn shows every state from +2 to +7.
• Few states appear at the two ends. Sc and Ti have too few electrons to give up or share; Cu and Zn have so many d electrons that little room is left to bond with. So Sc(II) is almost unknown, Ti(IV) outlasts Ti(III) and Ti(II), and zinc sticks to +2.
• Scandium is the transition element that shows no variable oxidation state (only +3).
• As far as Mn, the top stable state is the sum of the 4s and 3d electrons: TiO₂ has Ti at +4, VO₂⁺ has V at +5, CrO₄²⁻ has Cr at +6 and MnO₄⁻ has Mn at +7. Beyond Mn the high states drop off quickly: iron and cobalt show +2 and +3, nickel +2, copper +1 and +2, zinc +2.
• Transition-metal oxidation states usually differ by one (V(II), V(III), V(IV), V(V)); in main-group elements they typically differ by two.
• In groups 4 to 10 the heavier members favour the higher states, the reverse of the inert pair effect in the p-block. Mo(VI) and W(VI) are more stable than Cr(VI), so dichromate in acid is a strong oxidant while MoO₃ and WO₃ are not.
• Ligands that accept π electrons as well as donate σ electrons allow very low oxidation states: Ni in Ni(CO)₄ and Fe in Fe(CO)₅ are in the zero state.

5. Electrode potentials

NCERT §4.3.5; §4.3.6

• E°(M²⁺/M) sums up three steps: atomising the metal, removing two electrons (ΔiH₁ + ΔiH₂), and hydrating the M²⁺ ion. First-row values: Ti −1.63, V −1.18, Cr −0.90, Mn −1.18, Fe −0.44, Co −0.28, Ni −0.25, Cu +0.34 and Zn −0.76 V.
• Copper alone has a positive E°(M²⁺/M), so it cannot release H₂ from acids; only oxidising acids (nitric acid and hot concentrated sulphuric acid) attack it, and the acid itself is reduced. Turning Cu(s) into Cu²⁺(aq) costs more energy than the hydration of Cu²⁺ returns.
• The general drift to less negative values across the row follows the rising sum of the first and second ionisation enthalpies.
• Mn, Ni and Zn are more negative than the trend predicts: Mn²⁺ (d⁵) and Zn²⁺ (d¹⁰) have stable configurations, and Ni²⁺ has the most negative hydration enthalpy.
• E°(M³⁺/M²⁺) values: Ti −0.37, V −0.26, Cr −0.41, Mn +1.57, Fe +0.77, Co +1.97 V. Sc³⁺ is very stable (noble gas core), so Sc's value is low.
• Mn's value is high because Mn²⁺ (d⁵) is especially stable, and it is much higher than Cr's or Fe's because the third ionisation enthalpy of Mn (d⁵ → d⁴) is so large; this is why Mn(III) matters little. Fe's value is comparatively low because Fe³⁺ is d⁵. V's is low because V²⁺ has a half-filled t₂g set (Unit 5). Zn would have the highest value, since it would mean breaking into d¹⁰.
• Cr²⁺ and Mn³⁺ are both d⁴, yet Cr²⁺ is a reducing agent (becoming d³, a half-filled t₂g set) while Mn³⁺ is an oxidising agent (becoming d⁵).

6. Higher oxidation states and reactivity

NCERT §4.3.7; §4.3.8

• The highest oxidation states show up in fluorides and oxides. Among the halides they are TiX₄, VF₅ and CrF₆. Mn(VII) has no simple halide, but MnO₃F exists. After Mn, the only trihalides are FeX₃ and CoF₃.
• Fluorine holds high states through high lattice energy (as in CoF₃) or strong covalent bonds (as in VF₅ and CrF₆). V(V) exists only as VF₅; the other V(V) halides hydrolyse to oxohalides VOX₃. Fluorides are unstable in low states, for example VX₂ (X = Cl, Br, I), and the same holds for CuX.
• Every Cu(II) halide is known except the iodide, because Cu²⁺ oxidises iodide: 2Cu²⁺ + 4I⁻ → Cu₂I₂(s) + I₂.
• Many Cu(I) compounds disproportionate in water: 2Cu⁺ → Cu²⁺ + Cu. Cu²⁺(aq) wins because its hydration enthalpy is so much more negative that it more than pays for copper's second ionisation enthalpy.
• In oxides the top oxidation number equals the group number from Sc₂O₃ to Mn₂O₇. Iron goes no higher than Fe₂O₃, though ferrate(VI), FeO₄²⁻, forms in alkali and soon decomposes to Fe₂O₃ and O₂. Oxocations hold high states too: VO₂⁺ for V(V), VO²⁺ for V(IV), TiO²⁺ for Ti(IV).
• Oxygen holds high oxidation states even more firmly than fluorine does, because it can form multiple bonds to the metal: the top fluoride of Mn is MnF₄ but the top oxide is Mn₂O₇, in which each Mn is tetrahedrally surrounded by O with one Mn-O-Mn bridge. Tetrahedral [MO₄]ⁿ⁻ ions are known for V(V), Cr(VI), Mn(V), Mn(VI) and Mn(VII).
• Oxidising power rises in the order VO₂⁺ < Cr₂O₇²⁻ < MnO₄⁻, because the lower species they are reduced to become more stable in that order.
• Copper aside, the 3d metals are fairly reactive, and 1 M H⁺ oxidises them, sometimes slowly. At room temperature Ti and V resist dilute acids that are not oxidising. The E°(M²⁺/M) values show that the drive to form M²⁺ weakens across the row.
• Mn³⁺ and Co³⁺ are the strongest oxidising agents among the aqueous M³⁺ ions. Ti²⁺, V²⁺ and Cr²⁺ are strong reducing agents that release hydrogen from dilute acid: 2Cr²⁺ + 2H⁺ → 2Cr³⁺ + H₂.

7. Magnetic properties

NCERT §4.3.9

• A diamagnetic substance is pushed out of a magnetic field; a paramagnetic one is pulled in. A ferromagnetic substance is pulled in very strongly, and ferromagnetism is an extreme form of paramagnetism.
• Paramagnetism comes from unpaired electrons: each has a magnetic moment from its spin and its orbital motion. In first-row compounds the orbital part is effectively quenched, so only spin counts.
• Spin-only formula: μ = √[n(n + 2)] BM, where n is the number of unpaired electrons and BM is the Bohr magneton. One unpaired electron gives 1.73 BM.
• The moment rises with n: n = 1, 2, 3, 4, 5 give 1.73, 2.83, 3.87, 4.90 and 5.92 BM (√8 = 2.83; NCERT's Table 4.7 prints 2.84 for the n = 2 ions Ti²⁺ and Ni²⁺). A measured moment therefore tells you how many electrons are unpaired.
• Worked case: a divalent ion of Z = 25 is Mn²⁺, 3d⁵, with five unpaired electrons, so μ = √35 = 5.92 BM. For Z = 27, Co²⁺ is 3d⁷ with three unpaired electrons: μ = √15 = 3.87 BM.
• d⁰ and d¹⁰ ions (Sc³⁺, Ti⁴⁺, Zn²⁺) have no unpaired electrons and are diamagnetic.
• Measured moments of hydrated ions match the spin-only values closely for Ti³⁺ to Mn²⁺ (Mn²⁺: 5.92 calculated, 5.96 observed) but run higher for Fe²⁺ (5.3-5.5 against 4.90), Co²⁺ (4.4-5.2 against 3.87) and Cu²⁺ (1.8-2.2 against 1.73).

8. Coloured ions and complexes

NCERT §4.3.10; §4.3.11

• Colour arises when an electron jumps from a lower-energy d orbital to a higher-energy d orbital. The energy of that jump matches light of a frequency usually in the visible region, which is absorbed; we see the complementary colour.
• Which frequency is absorbed depends on the ligands around the ion. In water, water molecules are the ligands.
• Colours of aquated ions: Ti³⁺ purple (d¹), V⁴⁺ blue (d¹), V³⁺ green (d²), V²⁺ violet (d³), Cr³⁺ violet (d³), Mn³⁺ violet (d⁴), Cr²⁺ blue (d⁴), Mn²⁺ pink (d⁵), Fe³⁺ yellow (d⁵), Fe²⁺ green (d⁶), Co³⁺ blue (d⁶), Co²⁺ pink (d⁷), Ni²⁺ green (d⁸), Cu²⁺ blue (d⁹).
• d⁰ ions (Sc³⁺, Ti⁴⁺) and d¹⁰ ions (Zn²⁺) are colourless: with no d electron to excite, or no empty d orbital to receive one, there is no d-d jump.
• Complex compounds are those in which a metal ion binds several anions or neutral molecules to form a species with its own properties, such as [Fe(CN)₆]³⁻, [Fe(CN)₆]⁴⁻, [Cu(NH₃)₄]²⁺ and [PtCl₄]²⁻.
• Transition metals form a very large number of complexes because their ions are small, carry high charges and have d orbitals available for bonding.

9. Catalysts, interstitial compounds and alloys

NCERT §4.3.12; §4.3.13; §4.3.14

• Transition metals and their compounds are good catalysts because they can take several oxidation states and form complexes. Examples: V₂O₅ in the Contact process, finely divided iron in the Haber process, nickel in catalytic hydrogenation.
• On a solid catalyst, reactant molecules bond to surface metal atoms (first-row metals use their 3d and 4s electrons). This raises the reactant concentration at the surface and weakens bonds in the reacting molecules, lowering the activation energy.
• A change of oxidation state lets an ion shuttle electrons. Iron(III) catalyses 2I⁻ + S₂O₈²⁻ → I₂ + 2SO₄²⁻ in two steps: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂, then 2Fe²⁺ + S₂O₈²⁻ → 2Fe³⁺ + 2SO₄²⁻. The iron ends as it began.
• Interstitial compounds form when small atoms such as H, C or N are trapped in the holes of a metal lattice. They are usually non-stoichiometric and neither typically ionic nor covalent: TiC, Mn₄N, Fe₃H, VH₀.₅₆ and TiH₁.₇. Their formulas do not correspond to normal oxidation states.
• Interstitial compounds melt higher than the pure metals, are very hard (some borides approach diamond), keep metallic conductivity and are chemically inert.
• An alloy is a blend of metals. In a homogeneous solid-solution alloy one metal's atoms are scattered at random among the other's, which works when the metallic radii are within about 15 percent of each other. Transition metals, with similar radii, alloy readily, and their alloys are hard and often high-melting.
• Cr, V, W, Mo and Mn are used to make many steels and stainless steel. Alloys with non-transition metals matter too: brass (copper-zinc) and bronze (copper-tin).

10. Potassium dichromate and permanganate

NCERT §4.4; §4.4.1

• First-row oxides form when the metals are heated in oxygen. Every metal except Sc forms an ionic MO oxide. As the oxidation number rises, ionic character falls and acidic character rises: Mn₂O₇ is a covalent green oil and gives HMnO₄; CrO₃ gives H₂CrO₄ and H₂Cr₂O₇. Vanadium runs from basic V₂O₃ to less basic V₂O₄ to amphoteric V₂O₅. CrO is basic and Cr₂O₃ amphoteric.
• Disproportionation: when one oxidation state is less stable than the states on either side, it converts into both. In acid, manganate(VI) does this: 3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O.
• K₂Cr₂O₇ from chromite ore: fuse FeCr₂O₄ with sodium carbonate in air, 4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 8Na₂CrO₄ + 2Fe₂O₃ + 8CO₂. Acidify the yellow chromate solution to get orange dichromate, 2Na₂CrO₄ + 2H⁺ → Na₂Cr₂O₇ + 2Na⁺ + H₂O, then add KCl: Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇ + 2NaCl. The potassium salt, being less soluble, crystallises.
• Chromate and dichromate interconvert with pH, and chromium is +6 in both: 2CrO₄²⁻ + 2H⁺ → Cr₂O₇²⁻ + H₂O, and Cr₂O₇²⁻ + 2OH⁻ → 2CrO₄²⁻ + H₂O. CrO₄²⁻ is tetrahedral; Cr₂O₇²⁻ is two tetrahedra sharing one corner, with a Cr-O-Cr angle of 126°.
• Acidified dichromate is a strong oxidant: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O, E° = 1.33 V. It oxidises I⁻ to I₂, H₂S to S, Sn²⁺ to Sn⁴⁺ and Fe²⁺ to Fe³⁺: Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O. K₂Cr₂O₇ is a primary standard in volumetric analysis and is used in the leather industry; the more soluble sodium salt is used in organic chemistry.
• KMnO₄ from pyrolusite: fuse MnO₂ with KOH and an oxidant such as KNO₃ (or air) to get green K₂MnO₄, 2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O, which disproportionates in neutral or acid solution to permanganate. Commercially, manganate(VI) is oxidised electrolytically in alkali. In the laboratory, peroxodisulphate oxidises Mn²⁺: 2Mn²⁺ + 5S₂O₈²⁻ + 8H₂O → 2MnO₄⁻ + 10SO₄²⁻ + 16H⁺.
• KMnO₄ forms dark purple, almost black crystals (isostructural with KClO₄), dissolves only 6.4 g per 100 g of water at 293 K, and decomposes at 513 K: 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂. Manganate and permanganate are tetrahedral, with π bonds from oxygen p to manganese d orbitals; green manganate has one unpaired electron (paramagnetic) and permanganate none (diamagnetic).
• Electrons taken by MnO₄⁻ depend on the medium: to MnO₄²⁻, 1 e⁻ (E° = +0.56 V); to MnO₂ in acid, 3 e⁻ (+1.69 V); to Mn²⁺ in acid, 5 e⁻ (+1.52 V): MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.
• In acid: 5Fe²⁺ + MnO₄⁻ + 8H⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺ (green to yellow); 10I⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 8H₂O + 5I₂; oxalate at 333 K, 5C₂O₄²⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 8H₂O + 10CO₂; also H₂S to S, sulphite to sulphate and nitrite to nitrate. HCl is not used to acidify, because permanganate oxidises it to chlorine.
• In neutral or faintly alkaline solution MnO₄⁻ goes to MnO₂: iodide becomes iodate (2MnO₄⁻ + H₂O + I⁻ → 2MnO₂ + 2OH⁻ + IO₃⁻), thiosulphate becomes sulphate, and Mn²⁺ becomes MnO₂ (catalysed by ZnSO₄ or ZnO). Permanganate is also used to bleach wool, cotton and silk and to decolourise oils.

11. The lanthanoids

NCERT §4.5

• The lanthanoids are the 14 elements after lanthanum (Ce to Lu); La is usually discussed with them, and Ln stands for any of them. They resemble one another more closely than the members of a transition series do, and have essentially one stable oxidation state, +3.
• Atoms share 6s² with a variable number of 4f electrons (La 5d¹ 6s², Ce 4f¹ 5d¹ 6s², Eu 4f⁷ 6s², Gd 4f⁷ 5d¹ 6s², Yb 4f¹⁴ 6s², Lu 4f¹⁴ 5d¹ 6s²). Every Ln³⁺ ion is simply 4fⁿ, with n = 1 to 14 across the series.
• Lanthanoid contraction: atomic and ionic radii fall steadily from La to Lu because one 4f electron shields another even more poorly than one d electron shields another. The M³⁺ radii fall regularly, from 106 pm for La³⁺ to 86 pm for Yb³⁺; atomic radii fall less regularly (Eu's is 199 pm).
• Its big consequence is in the third transition series: 5d elements end up nearly the same size as their 4d partners. Zr (160 pm) and Hf (159 pm) occur together in nature and are hard to separate.
• +3 dominates, but +2 and +4 appear where they give an empty, half-filled or filled 4f set. Ce⁴⁺ (noble gas configuration) is a strong oxidant returning to +3; E°(Ce⁴⁺/Ce³⁺) = +1.74 V, which could oxidise water, but the reaction is slow, so Ce(IV) is a good analytical reagent. Tb⁴⁺ (f⁷) is also an oxidant; Pr, Nd, Tb and Dy show +4 only in oxides MO₂.
• Eu²⁺ (f⁷) and Yb²⁺ (f¹⁴) are reducing agents that go back to +3; samarium, like europium, shows both +2 and +3.
• The metals are silvery-white, soft and tarnish quickly in air; hardness rises with atomic number (samarium is steel-hard). Melting points are 1000-1200 K, with samarium at 1623 K. Properties change smoothly except for Eu and Yb (and sometimes Sm and Tm).
• Many Ln³⁺ ions are coloured, from f-electron transitions that give narrow absorption bands; La³⁺ and Lu³⁺ are colourless. All the ions are paramagnetic except f⁰ (La³⁺, Ce⁴⁺) and f¹⁴ (Yb²⁺, Lu³⁺).
• First ionisation enthalpies are about 600 kJ mol⁻¹ and second about 1200 kJ mol⁻¹, similar to calcium. Third ionisation enthalpies are unusually low for La, Gd and Lu. Early lanthanoids react like calcium and later ones more like aluminium; E°(Ln³⁺/Ln) is −2.2 to −2.4 V (Eu −2.0 V).
• Reactions: they combine with hydrogen on gentle heating, give carbides (Ln₃C, Ln₂C₃, LnC₂) with carbon, release H₂ from dilute acids and from water (giving Ln(OH)₃), burn in halogens to LnX₃, form Ln₂O₃ in oxygen, Ln₂S₃ with sulphur and LnN with nitrogen. Ln(OH)₃ are true compounds and basic, like alkaline earth hydroxides.
• Uses: alloy steels for plates and pipes; mischmetall (about 95% lanthanoid metal, 5% iron, traces of S, C, Ca, Al) in magnesium-based alloys for bullets, shells and lighter flints; mixed oxides as petroleum-cracking catalysts; some oxides as phosphors in television screens.

12. The actinoids and applications

NCERT §4.6; §4.7

• The actinoids are the 14 elements from Th to Lr, discussed together with actinium. All are radioactive. The early members have long half-lives; the later ones last from a day down to 3 minutes (lawrencium, Z = 103) and have been made only in nanogram amounts, which makes them hard to study.
• Configurations share 7s² with variable 5f and 6d occupancy. Thorium has no 5f electron; from Pa onwards 5f fills, completing at element 103. f⁰, f⁷ and f¹⁴ are favoured: Am is [Rn] 5f⁷ 7s² and Cm is [Rn] 5f⁷ 6d¹ 7s².
• 5f orbitals are less buried than 4f orbitals, so 5f electrons take part in bonding far more.
• Actinoid contraction: atoms and M³⁺ ions shrink across the series, and the shrinkage from one element to the next is larger than in the lanthanoids, because 5f electrons shield poorly.
• Because 5f, 6d and 7s are close in energy, actinoids show a wider range of oxidation states. +3 is general, but the first half reaches higher: the maximum rises from +4 (Th) to +5 (Pa), +6 (U) and +7 (Np), then falls. +3 and +4 ions tend to hydrolyse.
• The metals are silvery but have varied structures, since their metallic radii are much more irregular than the lanthanoids'. Finely divided, they are highly reactive: boiling water gives a mix of oxide and hydride, HCl attacks all of them, nitric acid barely affects most (a protective oxide forms), and alkalis have no action.
• Early actinoids have lower ionisation enthalpies than early lanthanoids, because 5f electrons are better shielded from the nucleus and so more available for bonding. Their magnetic behaviour is more complex; susceptibility varies with the number of unpaired 5f electrons roughly as in the lanthanoids, but lanthanoid values are higher.
• Lanthanoid-like behaviour appears only in the second half of the actinoid series. The lanthanoid contraction matters more than the actinoid contraction, because little is known of the elements after the actinoids.
• Applications: iron and steel are the chief construction materials (carbon plus Cr, Mn, Ni); MnO₂ goes into dry cells, and batteries also use Zn and Ni/Cd. Group 11 are the coinage metals.
• Industrial catalysts: V₂O₅ oxidises SO₂ in making sulphuric acid; TiCl₄ with Al(CH₃)₃ is the basis of Ziegler catalysts for polythene; iron in the Haber process; nickel for hydrogenating fats; PdCl₂ in the Wacker process that makes ethanal; nickel complexes for polymerising alkynes. AgBr is prized in photography for its sensitivity to light.

Must-know facts

1. d-block = groups 3-12; transition metal = incomplete d subshell in the atom or a common ion.
2. Zn, Cd, Hg (and Cn) are (n−1)d¹⁰ ns² and are not transition elements; Sc (3d¹) is, and Ag is (Ag²⁺ is 4d⁹).
3. Cr = [Ar] 3d⁵ 4s¹; Cu = [Ar] 3d¹⁰ 4s¹; Pd = [Kr] 4d¹⁰ 5s⁰.
4. Ions lose ns electrons before (n−1)d: Fe²⁺ = 3d⁶, Fe³⁺ = 3d⁵, Mn²⁺ = 3d⁵, Cu²⁺ = 3d⁹.
5. Melting points peak near d⁵; enthalpy of atomisation peaks mid-series; Zn has the lowest (126 kJ mol⁻¹).
6. Lanthanoid contraction makes Zr (160 pm) and Hf (159 pm) almost the same size.
7. Mn shows +2 to +7, the most states in the 3d row; Sc shows only +3; Zn only +2.
8. Heavier members of d-groups prefer higher oxidation states: W(VI) and Mo(VI) are more stable than Cr(VI).
9. Cu is the only first-row metal with positive E°(M²⁺/M) = +0.34 V; it does not release H₂ from dilute acids.
10. E°(M³⁺/M²⁺): Co +1.97 and Mn +1.57 V (strong oxidants); Cr²⁺, V²⁺, Ti²⁺ are strong reductants.
11. Oxygen stabilises higher states than fluorine: MnF₄ but Mn₂O₇.
12. Spin-only μ = √[n(n+2)] BM: 1.73, 2.83, 3.87, 4.90, 5.92 for n = 1 to 5 (Table 4.7 prints 2.84 for n = 2).
13. d⁰ and d¹⁰ ions (Sc³⁺, Ti⁴⁺, Zn²⁺, Cu⁺) are colourless and diamagnetic.
14. Aquated colours: Cu²⁺ blue, Ni²⁺ green, Fe²⁺ green, Fe³⁺ yellow, Mn²⁺ pink, Co²⁺ pink, Ti³⁺ purple.
15. Catalysts: V₂O₅ (Contact), Fe (Haber), Ni (hydrogenation), PdCl₂ (Wacker), TiCl₄ + Al(CH₃)₃ (Ziegler).
16. Chromate (yellow) ⇌ dichromate (orange) with pH; Cr is +6 in both; Cr-O-Cr angle 126°.
17. Acid permanganate takes 5 e⁻ (to Mn²⁺); neutral or faintly alkaline takes 3 e⁻ (to MnO₂); dichromate takes 6 e⁻ per ion.
18. Ce⁴⁺ is an oxidant (E° = +1.74 V); Eu²⁺ and Yb²⁺ are reductants; common lanthanoid state is +3.
19. Mischmetall ≈ 95% lanthanoid metal + 5% Fe; used in lighter flints.
20. Actinoid maximum oxidation state: Th +4, Pa +5, U +6, Np +7; actinoid contraction is larger per element than lanthanoid contraction.

Common traps

Formulas

Key terms

Transition element

A metal whose atom or common ion has an incomplete d subshell.

Inner transition element

An f-block element, in which 4f (lanthanoids) or 5f (actinoids) orbitals fill.

Enthalpy of atomisation

Energy needed to turn one mole of the solid metal into free gaseous atoms; a measure of metallic bond strength.

Lanthanoid contraction

The steady fall in atomic and ionic size from La to Lu, caused by poor shielding by 4f electrons.

Exchange energy

Stabilisation that grows with the number of parallel-spin electron pairs in a degenerate set.

Disproportionation

A reaction in which one oxidation state converts into a higher and a lower one at the same time.

Paramagnetic

Drawn into a magnetic field, because of unpaired electrons.

Diamagnetic

Pushed out of a magnetic field; all electrons paired.

Ferromagnetic

Very strongly attracted to a magnet; an extreme form of paramagnetism.

Bohr magneton (BM)

The unit of magnetic moment used for atoms and ions.

Spin-only moment

Magnetic moment counting electron spin alone, √[n(n+2)] BM.

Complex compound

A compound in which a metal ion binds several ions or neutral molecules into one species with its own properties.

Interstitial compound

A compound formed when small atoms (H, C, N) sit in the holes of a metal lattice; usually non-stoichiometric.

Alloy

A blend of metals; solid solutions form when metallic radii are within about 15%.

Mischmetall

An alloy of about 95% lanthanoid metal and 5% iron with traces of S, C, Ca and Al.

Oxocation

A cation containing metal and oxygen, such as VO₂⁺, VO²⁺ or TiO²⁺.

Primary standard

A pure, stable substance weighed directly to make a solution of exactly known concentration, as K₂Cr₂O₇ is.