From the splitting diagram to the bottle on the shelf
You arrive at this guide already armed. From the earlier rungs you know what a transition metal is, why these elements show many oxidation states, and where their colour and magnetism come from — the d orbitals split in a ligand field, and electrons fall into the resulting t2g and eg sets. This guide spends none of its budget re-deriving that. Instead it does something the theory chapters could not: it walks down the middle of the first row, element by element, and shows you what Cr, Mn, Fe, Co, Ni, and Cu actually look like in a flask.
One organizing idea threads the whole tour. On the left of the row the high oxidation states are easy to reach and even stable — chromium happily reaches +6, manganese +7 — but as you move right toward iron, cobalt, nickel and copper, the top oxidation states become increasingly hard to hold and ferociously oxidizing, while the +2 state quietly takes over as the comfortable, everyday one. By copper at the far end, +2 (and even +1) is the whole story and +3 is an exotic. Keep that left-to-right slide in mind and the catalogue below stops being a list to memorize and becomes a trend you can almost predict.
Chromium and manganese: the high oxidizers
Chromium is the element of two famous oxidation states. As Cr3+ it is a small, hard, kinetically inert d3 ion — recall that a half-filled t2g set is especially well stabilized — and its aqua ion [Cr(H2O)6]3+ is the violet-grey ion of chrome tanning and chrome plating, sluggish to exchange its ligands. Climb to Cr(VI) and the chemistry flips: there are no Cr6+ cations floating around, because such a charge is a bookkeeping fiction, not a bare ion. Instead chromium hides inside oxoanions — yellow chromate CrO4^2- in base and orange dichromate Cr2O7^2- in acid, interconverting as you swing the pH. Dichromate is a workhorse oxidant, and as it does its job the orange Cr(VI) is reduced back down to green Cr3+.
Manganese stretches the oxidation-state range even wider, from Mn2+ all the way to Mn(VII). The deep purple permanganate ion MnO4- is the showpiece — but note that its intense colour is not a d-d transition at all, since Mn(VII) is formally d0 with no d electrons to excite. The colour comes instead from a charge-transfer band, oxygen lone pairs hurling electron density onto the metal, which is why it is so much more vivid than any pale d-d colour. Permanganate is a powerful oxidant; in acid it is reduced all the way to near-colourless Mn2+, the high-spin d5 ion. That d5 ion sits at a special place: five unpaired electrons, zero crystal field stabilization, so Mn2+ complexes are pale and comparatively labile — manganese's quiet, stable resting state.
Iron: rust, blood, and the Fe2+/Fe3+ shuttle
Iron is the hinge of the whole row, and its chemistry lives almost entirely in the easy shuttle between two oxidation states: pale green Fe2+ (d6) and pale yellow-to-brown Fe3+ (d5). In open air, dilute Fe2+ is slowly oxidized to Fe3+ — half of why iron rusts. Both are weak-field aqua ions, high-spin, and their aqua colours are muted; the dramatic colours arrive when the ligand changes. Add thiocyanate to Fe3+ and you get the blood-red [Fe(SCN)]2+ of the classic test; add cyanide and you reach the famously inert hexacyanoferrates, [Fe(CN)6]4- (ferrocyanide) and [Fe(CN)6]3- (ferricyanide).
That same Fe2+/Fe3+ couple, sitting at just the right potential, is exactly why life leans on iron so heavily. The iron at the centre of haem can hold an Fe2+ that binds dioxygen reversibly in haemoglobin and myoglobin, or shuttle between Fe2+ and Fe3+ to ferry electrons down the respiratory chain in the cytochromes. This is a good moment to retire a myth: "inorganic" does not mean lifeless. Inorganic chemistry is the chemistry of all the elements, carbon included, and some of its most beautiful results — the iron in your blood, the cobalt in your vitamin B12 — are squarely biological. Iron is one of the clearest essential metals of life.
Rust itself is worth one honest sentence, because it explains a great deal about iron's chemistry and its economic curse. The familiar brown solid is hydrated iron(III) oxide, roughly Fe2O3.xH2O, formed when Fe2+ is oxidized in moist air and the resulting Fe3+ hydrolyses. Unlike the tough, adherent oxide that protects aluminium, iron's oxide is flaky and porous, so it spalls off and lets fresh metal keep corroding underneath — which is why iron must be painted, galvanized, or alloyed to stainless steel to survive. Notice the recurring theme: high-charge Fe3+ is acidic enough in water that its aqua ion hydrolyses, the same reason Fe3+ solutions are slightly acidic and slowly go cloudy.
Cobalt, nickel, copper: the right-hand trio
Cobalt shows the row's headline trend more sharply than any other element. Its Co2+ aqua ion is the pink-to-blue ion you meet in cobalt chloride moisture indicators (octahedral and pink when wet, tetrahedral and blue when dry). But the truly characteristic chemistry sits at Co3+. As a bare aqua ion Co3+ is unstable and tears water apart, yet wrap it in strong-field ligands and it becomes one of inorganic chemistry's most studied centres: low-spin d6 complexes like [Co(NH3)6]3+ are kinetically inert and beautifully coloured. This is exactly the spin-state lesson from earlier guides made flesh — a single ion swinging from violent oxidizer to placid, robust complex purely by changing what surrounds it.
Nickel is the quietest of the six: it is essentially the +2 story and little else. The green [Ni(H2O)6]2+ aqua ion is d8, and d8 octahedral ions have no high-spin/low-spin choice to make — there is only one sensible filling, t2g^6 eg^2. Add a strong-field, chelating ligand and nickel can also go square planar and diamagnetic, as in the cherry-red bis(dimethylglyoximate) complex used to test for Ni2+. Across this whole right-hand side the metals are still happy to act as Lewis acids at the centre of a coordination compound, binding a ring of donor atoms; nickel simply does it with a smaller spread of oxidation states than its neighbours.
Copper closes the tour with two final twists. Its everyday face is Cu2+, the familiar blue of copper sulfate solution — a d9 ion that, with one electron short of a filled eg, shows a textbook Jahn-Teller distortion: the octahedron stretches along one axis so the orbitals split a little further and the ion drops in energy. Copper's other oxidation state, Cu+ (d10), is the second twist. In water Cu+ is unstable and undergoes disproportionation — 2 Cu+ collapses into Cu2+ plus Cu metal — so cuprous ions survive in solution only when locked up by ligands or precipitated as insoluble salts like CuI. By the end of the row, then, the high oxidation states have vanished entirely and copper's chemistry is a contest between just two low ones.
Reading the row as one connected story
Step back and the catalogue resolves into a few clean patterns. The maximum oxidation state rises to a peak at manganese (+7) and then falls away to the right; the highest states always hide inside oxoanions or fluorides, never as bare cations, because a charge like +7 is an oxidation-state bookkeeping label, not a real ionic charge sitting on an atom. High states are strong oxidants (CrO4^2-, MnO4-), low states near the right are reductants or just stable resting points (Fe2+, Cu+ where trapped), and a couple of awkward ions in between collapse by disproportionation. Colour tracks the d-electron count and the field strength; magnetism tracks the number of unpaired electrons.
Common ions across the row (oxidation states + d-count): Cr +3 (d3, inert, violet) ........ +6 CrO4^2-/Cr2O7^2- (oxidant) Mn +2 (d5, HS, near-colourless) .. +7 MnO4- (deep purple, CT band) Fe +2 (d6) <----> +3 (d5) the redox shuttle of life/rust Co +2 (d7, pink/blue) ............ +3 (d6, low-spin, inert) Ni +2 (d8, green; sq-planar = diamagnetic) Cu +1 (d10, disproportionates) ... +2 (d9, blue, Jahn-Teller) max O.S. peaks at Mn (+7), then falls to the right