This batch of hydroxide is mostly gray Fe(OH)2, produced by electrolytic corrosion of hot rolled steel in salt solution. It doesn't settle well nor is it easy to filter. The broken orange surface is oxidation from ferrous to ferric (either Fe(OH)3 or FeOOH), broken because I shook it a bit to illustrate the difference. As always, hydroxides can be produced by neutralization of respective salts.
The nice thing about iron is it readily exists in two valence states: +2 and +3. (Although ferrous oxide is unstable, preferring to disproportionate to iron and the iron spinel, ferrous(II) ferrate(III), i.e., FeO.Fe2O3, a.k.a. magnetite.) Salts such as sulfate and chloride can go back and forth freely, and this allows great reduction and oxidation power from these compounds. Ferric chloride for example is used to etch even noble copper, through the reaction 2FeCl3 + Cu = CuCl2 + FeCl2. Likewise, ferrous sulfate is a reducing agent, as: 3FeSO4 + NaClO3 + 1½H2O = NaCl + Fe(OH)3 + Fe2(SO4)3 (chlorate in acidic medium).
Here's some rather nice rust I've produced in various ways, mostly involving water and thus producing the hydroxide. This has been calcined near red heat, losing all water and producing the rater stable reddish-brown (it has an almost purplish tint in person) powder seen here. (I say "stable" because Fe2O3, Al2O3, Cr2O3 and so on, in crystalline form, are, at best, slow to react. Fe2O3 and Al2O3 dissolve slowly in acid (preferrably a complexing one such as hydrochloric), Al2O3 dissolves slowly in strong base, and Cr2O3 really needs an alkaline fusion (such as molten KNO3 and NaOH, producing chromate) to do anything. Other transition metal oxides are worse still!)
(6/8/06) Here's my rust bucket. My bubbler burned out long ago, so I have to stir this periodically to keep it rusting. When you stir, you get a solid murkiness of a drab light brown, ferric hydroxide. The metal (various scrap sheetmetal I have no use for) on top is rusting away nicely, while the metal below remains clean, etched by ferric ions in solution.
I haven't prepared much of these yet but the same old chemical rules apply: dissolve iron (metal or (hydr)oxide) in HCl acid, crystallize if desired. Ferrous chloride is preferred, while ferric chloride is a strong oxidizer and will etch pretty much any metal put into it. To reach FeCl3, you need to add an oxidizer. Chlorine gas is a pretty good way, and since you need excess chloride (to prevent it from precipitating rust, as shown above for the copperas), adding a hypochlorite to the acidic solution would be a great way to go. I'm sure it can also be prepared electrolytically, anodically oxidizing the solution. It'll have to be strongly acidic to prevent iron plating out during the process.
Here is some ferric chloride hydrate I managed to make. The solution was hydrochloric acid pickle, rich with iron from various rust and scale soakings and oxidized from long exposure to air. Due to its high acidity, I decided to distill the acid off in a "solar still" type device. The slow evaporation apparently helped to make suprisingly bulky crystals, acicular though they are. The solution had goop of some sort in it (possibly organic or silicate in nature) and it's hard to see anything in the dark brown liquid, plus there's probably copper and whatnot that found its way into this as well. I would guess this at maybe 90-95% purity. I don't particularly care to recrystallize it!
To obtain anhydrous ferric chloride, you have to dehydrate the crystals in a stream of HCl! I don't even know where to begin with that one on a home scale...
I have however prepared plenty of ferrous sulfate (archaically, copperas or green vitriol). This I made by adding iron metal to sulfuric acid solution. After a week of soaking (and stinking, from impurities in the metal or possibly from reduction of the sulfate ion??), I got a nearly colorless (though murky with dust and rust) solution, with similarly colorless crystals forming on the metal. After keeping warm and recrystallizing the solution during the next week, it make pretty cyan crystals, pictured above. After still longer, I finished evaporation, yielding the bright green crystals [mostly] on the left. I have seen *three* colors of ferrous sulfate, and there is *nothing* else they could possibly be but ferrous sulfate itself. What the heck!?
As for its chemistry, it's just the same as chloride, though ferric sulfate appears to be less favored - I have yet to see any atmospheric oxidation of the solution or the crystals, whereas my iron chloride solution is rusty on the surface. One nice thing about iron salts is they can reduce or oxidize things cyclically: say you put oxygen together with steel. Not much at room temperature, even with water. Add some sulfate and ferrous sulfate forms. This oxidizes to rust (yellow Fe(OH)3), while the sulfate ion keeps some ferric ions in solution. When these return to the metal, some Fe goes into solution, reducing Fe+++ to Fe++ at the same time. Ferrous sulfate goes in and comes back out: a basic catalytic process! After I've finished crystallizing my green vitriol here, I'll hook up a bubbler/circulating pump to a tub with old metal and some FeSO4 and start rusting me some pigment.