iron chemistry
Iron is the workhorse metal of civilization and of life. It frames our buildings as steel, it rusts on our gates, and it carries oxygen in every breath you take. After aluminium it is the most abundant metal in the Earth's crust, and the planet's core is mostly iron. For all that ubiquity, iron's solution chemistry is built around just two oxidation states, +2 and +3, and the easy flip between them is the engine of much of its usefulness.
Iron(II), the pale green 'ferrous' ion [Fe(H2O)6]2+ (a d6 ion), is the lower state; iron(III), the pale violet (often yellow-brown in practice) 'ferric' ion [Fe(H2O)6]3+ (d5), is the higher. In air and water iron(II) is slowly oxidized up to iron(III), and the two are easily interconverted by mild oxidants and reductants — this single-electron Fe2+ / Fe3+ couple is what makes iron a superb electron carrier. Iron(III), being a small highly charged ion, is acidic in water (its aqua ion hydrolyzes, which is why iron(III) solutions are often yellow-brown and slightly acidic). Iron forms a host of complexes: the blood-red [Fe(SCN)]2+ used to test for iron(III), the deep-blue pigment Prussian blue (an iron(II)/iron(III) cyanide framework), and the stable, low-spin [Fe(CN)6]4- and [Fe(CN)6]3- ions. Higher states like iron(IV) and even iron(VI) (the ferrate ion FeO4 2-) exist but are strong oxidizers and uncommon.
Iron's two-state shuffle is central to both industry and biology. Steel is iron's great gift to engineering, while rusting (iron slowly returning to the oxide it came from) is its great curse, fought with paint, galvanizing and stainless alloys. In the body, the iron at the center of haemoglobin grabs and releases oxygen, iron-sulfur clusters and cytochromes shuttle electrons through respiration, and many enzymes use iron's Fe2+ / Fe3+ switch to do chemistry. Iron is the quiet example that 'inorganic' chemistry runs the living world.
A drop of thiocyanate added to a solution containing iron(III) turns it blood-red as [Fe(SCN)]2+ forms — a classic test for Fe3+. If only iron(II) were present, nothing happens; oxidize it first (say with hydrogen peroxide) to iron(III), and the red appears. The test literally shows you the Fe2+ / Fe3+ distinction.
Thiocyanate turns iron(III) blood-red but ignores iron(II) — a simple visual test of the oxidation state.
Iron(III) solutions are acidic not because of any added acid but because the small, highly charged Fe3+ ion polarizes its bound water molecules until they release protons — aqua-ion hydrolysis. This is a general feature of small, highly charged metal ions, not something peculiar to iron.