The d-Block Elements: Transition Metals

By IUPAC's own precise definition, a transition metal is one with an incompletely filled d subshell, either in its neutral atom or in any of its commonly occurring ions, a definition specific enough to exclude some elements sitting right inside the d-block itself. Zinc, cadmium, and mercury all have a completely filled (n-1)d10 ns2 configuration, both as neutral atoms and in their usual +2 ionic state, so despite occupying group 12 of the d-block, none of them qualify as genuine transition metals. Scandium, by contrast, has a single 3d electron in its ground state, an incomplete d subshell, earning it the label even though it sits at the very start of the row. This isn't a pedantic technicality; it's exactly the property, partly filled d orbitals, that produces every characteristic transition-metal behaviour covered in this chapter, variable oxidation states, coloured compounds, catalytic activity, and complex formation alike.

The outer electronic configuration of most d-block elements follows (n-1)d1-10 ns1-2, the inner (n-1)d orbitals holding anywhere from one to ten electrons while the outer ns orbital typically holds one or two. But the energy gap between (n-1)d and ns orbitals is genuinely small, small enough that the extra stability of a half-filled or completely filled d subshell can tip the balance, producing real, memorable exceptions. Chromium's configuration is 3d5 4s1, not the 'expected' 3d4 4s2, since shifting one electron into 3d completes an exceptionally stable half-filled set. Copper's configuration is 3d10 4s1, not 3d9 4s2, for exactly the same reason, reaching a completely filled, equally stable d10 set instead. These aren't arbitrary quirks to memorise; they're the direct, predictable consequence of the same stability principle that governs electron filling throughout the periodic table.

With few exceptions, transition metals show every classic metallic property in force: high tensile strength, ductility, malleability, and strong thermal and electrical conductivity. Nearly all of them (Zn, Cd, Hg and Mn being the notable exceptions) are very hard, have low volatility, and boast genuinely high melting and boiling points, and the explanation lies in interatomic bonding: a greater number of electrons available for metallic bonding, drawn from (n-1)d orbitals in addition to the usual ns electrons, produces stronger interatomic attraction. Within any row, melting points climb to a maximum around d5 (where the number of unpaired electrons peaks, favouring the strongest interatomic interaction), then fall again as pairing sets in, with manganese and technetium as genuine anomalies. This same enthalpy of atomisation pattern directly explains why heavier transition metals form metal-metal bonds far more readily than lighter ones, and why elements with very high atomisation enthalpy tend toward noble, unreactive behaviour.

Within a given transition series, ionic radii of same-charge ions steadily decrease as atomic number rises, since each additional electron enters a d orbital, whose shielding is genuinely poor, leaving the effective nuclear pull on outer electrons growing stronger and the radius shrinking, though the variation across any single row stays fairly small overall. A genuinely striking pattern emerges comparing series to series: atomic size grows from the 3d to the 4d row as expected, but the 5d row's radii come back almost identical to the 4d row's, rather than growing further, an anomaly explained by lanthanoid contraction, the steady radius shrinkage caused by the 14 elements' worth of poorly-shielding 4f orbitals filling in between the 4d and 5d rows. The net result is that corresponding 4d and 5d elements, zirconium (160 pm) and hafnium (159 pm) being the textbook example, end up with nearly identical size and, consequently, remarkably similar chemical properties, similar enough that separating them from each other in nature is genuinely difficult.

Ionisation enthalpy generally increases across a transition series, but far more gently than across a row of main-group elements, since new electrons enter the genuinely inner (n-1)d orbitals, which shield the outer 4s electrons from the growing nuclear charge relatively effectively, keeping the atomic radius, and therefore the ionisation enthalpy, from climbing steeply. The irregular jumps that do occur trace back to exchange energy: Mn+ (3d5 4s1) has lower ionisation enthalpy than Cr+ (3d5), since removing an electron from Mn+ costs less exchange-energy stabilisation than doing the same from Cr+'s already half-filled set; similarly, Fe2+ (3d6) has lower third ionisation enthalpy than Mn2+ (3d5), since Fe2+'s configuration hasn't yet reached the extra-stable half-filled arrangement Mn2+ already enjoys. This same exchange-energy logic explains why high third ionisation enthalpies for copper, nickel, and zinc make oxidation states above +2 genuinely difficult to reach for these particular elements.

Transition metals show a genuinely remarkable range of oxidation states, and manganese is the extreme case, exhibiting every single state from +2 all the way to +7. The elements with the most oxidation states cluster near the middle of a series, since early elements (Sc, Ti) simply don't have enough electrons to lose, while late elements (Cu, Zn) have too many d electrons occupying orbitals that are hard to free up for higher valence. A genuinely distinctive feature separates transition-metal oxidation-state variability from the p-block's own version: transition metal states typically differ from each other by exactly one (V2+, V3+, V4+, V5+), while non-transition elements' oxidation states normally differ by two instead, a direct consequence of d orbitals filling in individually rather than in pairs. Groups also show a genuinely opposite trend from the p-block: rather than heavier elements favouring lower oxidation states (the inert pair effect), heavier d-block elements favour higher ones instead, which is exactly why molybdenum(VI) and tungsten(VI) are markedly more stable than chromium(VI), making dichromate a strong oxidiser while MoO3 and WO3 are comparatively unreactive.

Reading the M2+/M standard electrode potential values across the first transition series reveals a general trend toward less negative values, tracking the rising sum of first and second ionisation enthalpies, but copper stands out as a genuine exception, the only metal in the row with a positive E° (+0.34 V), which directly explains why copper alone resists dissolving in ordinary dilute acids, only reacting with genuinely oxidising acids like nitric acid or hot concentrated sulphuric acid instead. The reason traces to energy balance: the substantial energy needed to transform solid copper into aqueous Cu2+ simply isn't compensated by copper's hydration enthalpy the way it is for other, more reactive metals. Looking instead at the M3+/M2+ couple reveals its own logic: Ti2+, V2+, and Cr2+ are all strong reducing agents, readily liberating hydrogen from dilute acid, precisely because their oxidation to the +3 state moves them toward more stable configurations, exactly the opposite situation from Mn3+ and Co3+, which are the strongest oxidising agents in aqueous solution, since being reduced back to +2 is what carries them toward extra stability instead.

Transition metal ions with unpaired electrons are paramagnetic, weakly attracted into a magnetic field, and for first-row transition metal compounds specifically, the orbital contribution to magnetism is effectively quenched, leaving the observed magnetic moment determined almost entirely by unpaired electron count, calculable directly from the spin-only formula, mu = the square root of n(n+2), where n is the number of unpaired electrons and mu is measured in Bohr magnetons. A single unpaired electron gives 1.73 BM; Mn2+, with its five unpaired 3d electrons, gives 5.92 BM, both matching experimentally observed values closely. This relationship runs both ways, genuinely usefully: measure a compound's magnetic moment experimentally, and the spin-only formula tells you directly how many unpaired electrons that ion actually carries, a real, practical structural clue extracted from nothing more than a magnetic susceptibility measurement.

Just as covered for coordination compounds generally, a transition metal ion's colour comes from d-d transitions, an electron absorbing a specific visible wavelength to jump between split d orbitals, with the observed colour being the complementary wavelength left unabsorbed. Tracking hydrated first-row ions confirms the pattern precisely: Sc3+ (d0, no d electrons to excite) and Zn2+ (d10, no vacant d orbital to jump to) are both genuinely colourless, while every configuration in between, d1 through d9, shows some visible colour, Ti3+ purple, V3+ green, Mn2+ pale pink, Fe3+ yellow, Cu2+ blue. This isn't a coincidence needing a separate explanation each time; it's the exact same underlying mechanism, ligand-induced d orbital splitting enabling a light-absorbing transition, playing out identically across an entire row of otherwise very different elements.

Transition metals and their compounds are prized catalysts precisely because of their ability to adopt multiple oxidation states and form complexes; V2O5 in the Contact Process, iron in the Haber Process, and nickel in catalytic hydrogenation are all classic industrial examples, and even the mechanism of iron(III) catalysing iodide's oxidation by persulphate can be shown explicitly, iron cycling between +3 and +2 to shuttle electrons between the two reactants. Small atoms like hydrogen, carbon, or nitrogen can also become trapped directly inside a transition metal's crystal lattice, forming interstitial compounds like TiC or Fe3H, genuinely non-stoichiometric, harder than the pure metal, and still metallically conductive. And because transition metals share such similar atomic radii, they readily form alloys, homogeneous solid solutions where one metal's atoms are simply distributed among another's, producing everyday materials like brass (copper-zinc), bronze (copper-tin), and the entire family of steels, whose hardness and strength come directly from chromium, vanadium, manganese, and other transition metals mixed into the base iron.

Potassium dichromate is prepared from chromite ore (FeCr2O4), fused with sodium carbonate in air to give sodium chromate, acidified to sodium dichromate, then converted to the less-soluble potassium salt by treatment with potassium chloride, yielding potassium dichromate's characteristic orange crystals. Chromate (tetrahedral) and dichromate (two corner-sharing tetrahedra) genuinely interconvert in solution depending on pH, though chromium's own oxidation state stays identical throughout, and in acidic solution, dichromate is a powerful oxidiser (E° = 1.33 V), oxidising iodide to iodine, sulphide to sulphur, and Fe2+ to Fe3+. Potassium permanganate follows a parallel story: prepared from pyrolusite ore (MnO2) fused with alkali and an oxidiser to give green manganate, which disproportionates in acid or neutral solution into permanganate and MnO2, and commercially finished by electrolytic oxidation. Deep purple, only modestly soluble, and diamagnetic (versus manganate's paramagnetism, from the difference of one unpaired electron), acidified permanganate oxidises iodide, Fe2+, oxalate, sulphite, and nitrite alike, making both compounds genuinely indispensable oxidising reagents in analytical and preparative chemistry.