Iron (II) hydroxide and green rust
Isolation of iron (II) hydroxide should be simple, right? We need an iron salt solution and some alkali. Here you can see three test tubes with an iron (II) sulfate solution after adding of: 1. Diluted KOH, 2. 25% ammonia solution in water 3. Hot concentrated KOH. Notice the difference in what is supposed to be the same substance.
What is going on? Iron (II) hydroxide is susceptible to oxidation. Normally it’s a white substance, but in the presence of oxygen (both from air and dissolved in water) it easily oxidizes to the brownish iron (III) oxide. You can see it in the left test tube. Adding potassium hydroxide to iron (II) sulfate resulted in a greenish precipitate which quickly turned brown after shaking. As expected. But what’s about all those shades of green and blue in other test tubes? It doesn’t look like a plain mix of iron (II) and iron (III) hydroxides.
This substance is also known as green rust, which forms naturally during the iron oxidation (mostly under water). Iron (II) hydroxide is a layered structure where iron (II) cations oxidize to iron (III). Water and anions (present in water) come between the layers to compensate the positive charge created by oxidation. This results in a complex formation that can be described by the following formula (in case of iron sulfate like I used) - Fe4Fe2(HO)12SO4·nH2O or even more precise: [Fe2+4Fe3+2(HO−)12]2+ x [SO2−4·2H2O]2− . You might notice the presence of iron in both oxidation states like in Prussian blue. There is a Wikipedia article on Green Rust. An attempt to filter and dry “green rust” will fail as it easily oxidizes to the brown iron (III) hydroxide:
Once we omit the “green rust” stage, we can describe the reaction in the left test tube as a two stage process:
FeSO4 + 2NaOH → Fe(OH)2 (white precipitate) + Na2SO4
4Fe(OH)2 + O2 → 4FeO(OH) (brownish precipitate) + 2H2O
In the middle test tube I used a 25% ammonia solution added to iron (II) sulfate. Some metal hydroxides are dissolved in ammonia solution and form complex compounds. But in case of iron it’s not happening. I found a profound article on this subject here.
It’s still possible to make an iron ammine, but under the different circumstances. Like blowing an excess of ammonia gas through a freshly prepared iron (II) salt solution (made by dissolving iron in an acid) under inert atmosphere. For iron bromide (freshly made by dissolving iron in hydrobromic acid) the process will look as followed:
FeBr2 + 6NH3 (gas) → [Fe(NH3)6]Br2 (Inert atmosphere!)
For the right test tube I added an excess of hot potassium hydroxide to an iron (II) sulfate solution. Amphoteric hydroxides would dissolve under this circumstances. But there are no visible changes. Iron (II) hydroxide is slightly amphoteric (showing weak acidic properties) and slowly (and partly) soluble in concentrated alkali. Again, in inert atmosphere. Like boiling in a strong (> 50%) solution of sodium hydroxide in a nitrogene atmosphere it would go as followed:
Fe(OH)2 + 2NaOH → Na2[Fe(OH)4] (Inert atmosphere!)
Those hydroxocomplex compounds can be isolated as bluish-green crystals. Mostly they are made by boiling of fine iron filings in a strong (>50%) sodium hydroxide solution.
Fe + 2H2O + 4NaOH → Na4[Fe(OH)6] + 2H2 (Inert atmosphere!)
I also made a video of this experiment which has shown slightly different result than on the picture.