I came up with the following idea for making powdered Iron:
First a sodium chloride solution is electrolyzed with iron electrodes, the following reaction will take place:
Anode: Fe ==> Fe2+ + 2e-
Cathode: 2H2O + 2e- ==> H2 + 2OH-
The dissolved Iron will directly be converted into the insoluble hydroxide by the following equation: Fe2+ + 2OH- ==> Fe(OH)2
This Iron hydroxide is then reacted with oxalic acid to make Iron Oxalate: Fe(OH)2 + H2C2O4 --> FeC2O4 + 2H2O
The Iron oxide is then heated until it decomposes by the following equation: FeC2O4 --> Fe + CO2
The Iron is released as a very fine powder.The_Davster - 13-4-2006 at 13:24
The iron produced this way is pyrophoric....It catches fire when exposed to air.12AX7 - 13-4-2006 at 13:49
I bet you could also reduce it with carbon, instead of forming an organic salt. I don't know how much it would sinter at that temperature, though, or
how much oxide would be left.
Timgarage chemist - 13-4-2006 at 14:12
Can ferric oxide or hydroxide be reduced with hydrogen at elevated temperature?
That process would directly give dry and very fine iron powder.
I believe it is called "ferrum reductum".
EDIT: Yes it works, very fine iron powder is actually prepared by heating Fe2O3 in a stream of hydrogen.
It is a dark grey powder and of a higher purity than any other grade of iron powder (it contains no carbon).
It is also flammable and easily iginted.
[Edited on 13-4-2006 by garage chemist]mrjeffy321 - 13-4-2006 at 16:56
You could make some type of Iron solution (possible by electrolysis of an Iron electrode, or even easier, steel wool + HCl). The precipitate out Iron
by adding a more reactive metal, displacing the Iron from solution.
For example, say you has an Fe+2 solution, you could then add Aluminum powder which will displace the Iron and make AlCl3 which is "Soluble in water
and alcohol." according to the MSDS. Zinc powder would work really well, or even Magnesium, all form soluble metal chloride salts.
You could use any metal above Iron on the activity series. A powder would work better than chunks due to its increased surface area.
Carbon will reduce Iron Oxide to Iron metal, producing CO2 gas.
Quote:
Originally posted by rogue chemist
The iron produced this way is pyrophoric....It catches fire when exposed to air.
? Please explain? I am interested.
Is this because it is just such a fine powder with such a high surface area / particle that it will so readily react?
[Edited on 14-4-2006 by mrjeffy321]guy - 13-4-2006 at 16:59
Quote:
??? Please explain? I am interested.
It is because the iron has such a high surface area that the reaction with oxigen is very rapid. Oxidation of iron is a very exothermic process, and
the reaction is used in portable heat packs. Since the reaction is now much faster, the heat can make the iron ignite.kclo4 - 13-4-2006 at 21:31
Hey guys! Just wondering, sense I don't really want to "waist" any of my chemicals making hydrogen. I was thinking that propane might work to reduce
the oxide, it would wouldn't it? I am pretty sure it will, I remember heating a copper wire with a torch and then blowing out the torch, the copper
wire remained a nice pure looking copper red. When I took away the stream of propane, it would oxidize into a multicolored wire, when I put back the
propane stream it would go back to the nice copper color. I am pretty sure it was propane it was a longish time ago.
But anyway, propane should work correct?guy - 13-4-2006 at 22:20
That's becuase some CO is produced in the flame and it reduces the CuO to Cu.woelen - 13-4-2006 at 23:06
The original idea of rot probably will not work. When you electrolyse this solution, then you don't get Fe(OH)2. Fe(OH)2 is oxidized VERY easily by
oxygen from air. It turns brown quickly, due to formation of ferric oxide/hydroxide. This compound dissolves with formation of a very stable complex,
the trisoxalatoferrate (III) ion, [Fe(C2O4)3](3-).
I have made this trisoxalatoferrate (III) myself. It is a nice apple-green ion. With ferric hydroxide, ammonia and oxalic acid you can make this
complex yourself.
This compound, however, does not yield iron on decomposition. The potassium salt also can be made easily, it crystallizes well. I also tried to make
the iron (III) salt of this, effectively making ferric trisoxalatoferrate (III) (this only contains iron and oxalate), but I did not succeed. This
salt simply does not crystallize. I obtained a syruppy liquid, which decomposes on further heating, giving basic iron salts and a lot of brown crap.
[Edited on 14-4-06 by woelen]unionised - 14-4-2006 at 10:55
"That's becuase some CO is produced in the flame and it reduces the CuO to Cu"
Er, how exactly does a flame produce CO after it has been blown out?
I have certainly reduced copper oxide with methane as a school experiment. On the other hand, copper oxide is much easier to reduce than iron oxide.
I think propane should work but I'm not sure if iron carbides would also be produced.
As Woelen points out, Fe(OH)2 is very readily oxidised- you would do better making FeCl2 from steel wool and HCl or FeCl3 then adding oxalic acid then
adding NaOH to precipitate the Fe(II) oxalatekclo4 - 14-4-2006 at 11:06
"That's becuase some CO is produced in the flame and it reduces the CuO to Cu. "
I had put out the flame so only propane was passing by the hot copper wire, so there was no flame to produce CO.I am a fish - 14-4-2006 at 11:10
If I were you, I would skip the electrolysis step, and buy iron(II) sulphate from a gardening store. Here in the UK (where most single chemical
fertilizers are difficult to buy) FeSO4 can be bought cheaply from most garden centers and DIY stores.
[Edited on 14-4-2006 by I am a fish]kclo4 - 14-4-2006 at 12:21
"I think propane should work but I'm not sure if iron carbides would also be produced."
I don't think carbides would form, I think it would only oxides the propane into a mixture of alcohol's and aldehydes, and maybe some acids. But even
if some carbides did form, wouldn't they be destroyed from the water, Iron and iron oxides formed in there at those temps?12AX7 - 14-4-2006 at 14:13
Iron doesn't form a carbide under equilibrium conditions.
Impurities and, especially, rapid cooling, can produce iron carbide however.
Only an issue if you have free carbon able to diffuse into the metal (i.e. above red heat).
