Sciencemadness Discussion Board - Chromium isolation

Chromium isolation

12AX7 - 27-4-2005 at 08:11

Seewwwwwww.... has anyone been tempted to make some elemental chromium? Say for the hell of it? Not much other use of it afterall..

Any thoughts on purifying:
Dichromate finishes
Stainless steel
???

Tim

Cyrus - 27-4-2005 at 19:24

How about using CrO3 from a pottery store... I know you can make an electroplating soln. with CrO3 and H2SO4. I forget the exact proportions, but it does take about 1 amp per square inch!!!

Electroplating of course isn't the brightest idea to isolate large amounts of anything, but that is one of the main uses of chromium...

edit- oops, you are right! Cr2O3, not CrO3! Sorry.

[Edited on 29-4-2005 by Cyrus]

Polverone - 27-4-2005 at 21:05

You'll get Cr2O3 from a pottery place, not CrO3, and it does not easily dissolve in sulfuric or any otther acid. Fusing it with nitrates (optionally combined with carbonates) produces dichromates and chromates respectively. It may be possible to plate out metal from dichromate solution. Or you could try aluminothermic reduction of Cr2O3. That's likely to be more spectacular and less pure.

12AX7 - 27-4-2005 at 23:17

I'd rather not try thermite, I don't have a good way to melt a relatively reactive high melting point glommed with slag.

Electrowinning is probably the most practical. IIRC, 50% of chromium production never sees a remotely metallic form, it goes right to chromic acid for plating.

I don't have a pottery store within 50 miles that has more than just clay, so that's out.

Tim

Polverone - 28-4-2005 at 00:46

The book <A HREF="http://www.sciencemadness.org/library/books/the_manufacture_of_chemicals_by_electrolysis.pdf">The Manufacture of Chemicals by Electrolysis</A> indicates that dichromates and permanganates can be made by using chromium or manganese-bearing iron alloys as the anode in electrochemical cells with alkali hydroxide solution as the electrolyte.

Quote:

E. Lorenz (2) has shown that it is possible to produce permanganate by electrolysing a solution of caustic potash, if a manganese or ferro-manganese anode be used and a cathode of copper oxide (the positive plate of a cupron cell for example). The same method can be used for preparing potassium bichromate if the anode be of ferro-chromium. In both cases the iron in the anode is converted to ferric hydroxide which collects at the bottom of the cell.

(2) Zeitsch. anorg. Chem., 1896,12, 393, 396.

You might try using a stainless steel as the anode in a cell with NaOH solution and see what you get. The iron may provide plenty of gunk to complicate things. See what goes into solution.

Edit: linkified book and quoted relevant passage.

[Edited on 5-1-2005 by Polverone]

azaleaemerson - 29-4-2005 at 04:41

You could isolate it by what this guy calls "reverse electroplating"

http://yarchive.net/car/reverse_electroplating.html

I looked for this because I remember an uncle doing something like it when I was a kid. It cleaned his old Chevy car part in any case. Oh how that man loved his car.

unionised - 30-4-2005 at 03:30

" IIRC, 50% of chromium production never sees a remotely metallic form, it goes right to chromic acid for plating. "

and then....?

If you use a slight excess of Al when reducing Cr2O3 then the slag should dissolve in NaOH soln leaving the Cr behind.

Esplosivo - 30-4-2005 at 04:11

Remember that Chromium is also amphotheric, like aluminium, and will react with alkalis (although not with the same vigour IIRC). A thermit reaction will be difficult to control. I once carried one out. I must admit it was spectacular, but the whole thing spread out the chromium which evolved around the spot, barely anything remained in the crucible.

[Edited on 30-4-2005 by Esplosivo]

12AX7 - 30-4-2005 at 04:20

Like I said, I'd rather not do thermite. Plus I have no good way of melting the result into a homogenous blob.

Electrowinning seems to be the way to go, just need to get it in solution. I'll try zapping some stainless today.

Tim

a_bab - 30-4-2005 at 04:58

Well, I got once some chromium globules that were definately made via a termite reaction. There was some yellow stuff left on the globules aswell.
Using less Al then necesary is the key to a slower reaction.

12AX7 - 30-4-2005 at 07:39

I've been electrolzing a strong lye solution between two stainless electrodes (at least that ought to be what they are, too heavy for aluminum, which would be long gone in the electrolyte anyway, and nonmagnetic so aren't plain steel), last check a thin gray deposit was growing on the cathode but the solution is only slightly yellow (probably more from soaking some dichromate finished hardware with the solution first, which only blackened them). So far it seems to be a good Brown's gas production cell.

Edit: the gray layer is up to about 1/8". The anode doesn't appear to be plating out, but its surface is discolored orange/red. The solution is still only very slightly yellow.

Tim

[Edited on 30-4-2005 by 12AX7]

unionised - 30-4-2005 at 13:34

"Remember that Chromium is also amphotheric, like aluminium, and will react with alkalis (although not with the same vigour IIRC)"
That's news to me and to those who wrote the text books I have seen.
Granted, if you were to oxidise it to Cr(VI) it would be, but that's not really relevant to washing the metal with caustic.

Polverone - 30-4-2005 at 13:57

To my surprise, when I tried it with a stainless steel spoon as anode, and a steel nail as cathode, I immediately began to see purple/red swirls coming from the anode. This happened when I switched the polarity too, mostly near the bottom of the cell where some undissolved NaOH remained. It seems that I must have been producing ferrate? I don't know of anything else with that color that could come from a steel nail.

After a couple hours of running the cell (5v, 5 A max (it got pretty hot)), I started pipetting solution out and introducing it to various other compounds. Ascorbic acid, ethanol, salicylic acid, sucrose, and hydrogen peroxide all seemed to reduce (?) it to an unremarkably iron-yellow color. Acids seemed to discharge the color rapidly as well.

I also tried seeing what happened when I ran the cell with an oxalic acid solution, the same spoon at the anode, and the nail at the cathode. I obtained a solution of very interesting color. In small amounts, like in a pipet, it appeared pale green like a solution of iron (II) compounds. In larger amounts or different lighting, it looked red. It seemed there was some very fine dispersion of particles that made the solution appear red in certain lighting conditions while the "true" transmitted color was green.

None of this puts you closer to getting pure chromium from stainless steel, but I still found it interesting.

12AX7 - 30-4-2005 at 17:09

Hmm odd.

Well the cathodic growth is up to 3/16" thick, I'll scrape it off and see if it...uh, well no response to bases obviously, but I can try heating, acid, etc...

I'll try acid and neutral (salt) later.

Tim

neutrino - 30-4-2005 at 17:33

Cr3+ + 4OH- -> Cr(OH)4-

This reaction drives the equilibrium between solid Cr2O3 and the dissolved form to the left, dissolving the solid.

unionised - 1-5-2005 at 06:23

I just wonder if you read my post?
What did you think was the point of adding excess Al? Did it occur to you that there shouldn't be much Cr2O3 left over?

It's generally the case that oxides get more acidic as the oxidation state rises. Cr is a case in point; CrO3 is strongly acid, Cr2O3 rather weakly so CrO isn't and (and this is the important bit) unlike aluminium, chromium as the metal (which is what this thread's about) isn't soluble in alkali.
So, if you did this thermite reaction (and I can understand someone not wishing to, it's quite violent) you would end up with Chromium (goody goody!), Al2O3, Some Al, some Cr2O3 and a few odds and ends from the crucible.
Al2O3 disolves in alkali because it's amphoteric
Cr2O3 disolves in alkali because it's amphoteric

Al disolves in alkali because its one of the amphoteric metals (there aren't many)
And the Cr sits there because it doesn't disolve in alkali. It would do if you left it long enough in the presence of oxygen but, in this (ie the relevent) case its not going to because of the Al disolving and scrubbing any O2 from solution.

Chromium sesquioxide is amphoteric. You could just about argue that the trioxide is too but you would be pushing it to find an instance of it acting as a base. Chromium, a metal, isn't amphoteric.
BTW, you haven't said what drives that eqm to the right.

[Edited on 1-5-2005 by unionised]

neutrino - 1-5-2005 at 06:36

Sorry, I thought the post was referring to dissolving the oxide, not the metal. I agree, the metal shouldn’t dissolve in a base with any significant speed.

12AX7 - 5-5-2005 at 11:52

BUMP!

Am currently electrolyzing stainless (food grade I think, 304 or 316?) in HCl (half dilute muriatic acid) at 5V. Strong green after a minute or two.

Should go check it in case something's plating out and growing towards a shorted cell...

Edit: That went smoothly. I now have a (relatively neutralized with NaOH, no H2 bubbles from Zn metal testing) solution of Fe + Ni + Cr + Cl. Now, to seperate!

Something interesting I noticed was the range of boiling/sublimation points, CrCl3 (- is CrCl2 formed, or is that less preferred?) seems to be the highest. I could do worse than loading the mixed, dried salts in a copper tube, heating just shy of melting the Cu and scraping out the sublimation layers.

Anything else? Precipitate (say with Na2CO3) then basify with lye solution, bubbling air to oxidize to chromate? Then finally of course plating out a nice shiny gob of the stuff.

Tim

[Edited on 6-5-2005 by 12AX7]

12AX7 - 6-5-2005 at 21:23

Say, instead of atmospheric oxidation, would bleach (Ca(OCl)2 to be exact), or sodium chlorate for that matter, oxidize it at all? A small test (without basification) certainly "bleaches" the iron to ferric hydroxide quite nicely.

Oh, was going to add, I precipitated it all with sodium hydroxide and carbonate. So there's a mixture of hydroxides and carbonate precipitates (a shaly, greenish pastel blue), well whatever...

Tim

[Edited on 7-5-2005 by 12AX7]

Esplosivo - 6-5-2005 at 21:48

Quote:

Originally posted by unionised
"Remember that Chromium is also amphotheric, like aluminium, and will react with alkalis (although not with the same vigour IIRC)"
That's news to me and to those who wrote the text books I have seen.
Granted, if you were to oxidise it to Cr(VI) it would be, but that's not really relevant to washing the metal with caustic.

Isn't chromium metal amphoteric? I mean not to the extent of aluminium, but if boiled with a concentrated NaOH solution shouldn't it react? If not then I'm sorry for confusing you, but damn it its written on my class notes. I hate these teachers.

unionised - 7-5-2005 at 02:18

I haven't tried it (though, if I find some chromium about the place I might) so I can't be sure.
Even if the oxidation potential favours it, which I doubt, in order to disolve it would have to get oxidised to Cr(III) and I think it wouldn't get past the insoluble Cr(II) without an oxidant, like air, present.

12AX7 - 9-5-2005 at 12:28

Added lye to the "stainless hydroxide" mixture (which BTW has been producing floating brown material, I'm guessing iron (in preference to Cr(III) and Ni(II)) is oxidizing to ferric oxy/hydroxides.

Anyways, I'll tell you what color the solution is once the Ni & Fe ppt's settle.

In the mean time, can I reduce chromates with sugar, possibly slightly acidified? I don't have any ferrous sulfate on hand, and haven't been able to make any (I made a 50% sulfuric acid (liquid fire) solution and added rust and metal with no reaction!?). Or can I just flush it without guilt?

Tim

12AX7 - 11-5-2005 at 15:48

Solution has remained perfectly colorless after several days of soaking with lye and air exposure (lye's probably weakening from CO2 absorption now! Say, is CrO4 stronger than CO3?).

I am officially out of ideas. Anyone?

Tim

S.C. Wack - 11-5-2005 at 17:52

The only way I know of to separate the metals at this point would be to add NaOH/peroxide to get soluble chromate or filter and roast in air with or without oxidizer with carbonate and/or hydroxide and leach the chromate.

12AX7 - 19-5-2005 at 03:49

I have dried the iron, chromium and nickel hydroxide/carbonate precipitates and calcined to a red heat, about 1400-1500°F (750-800°C). (Geez, now I'm listing temperature in *three* units?!)

I put it back in a lye solution and at last, it is turning yellow.

Any ballpark as to how long I have to wait for it to dissolve? Would hypochlorite speed it up?

Tim

chemoleo - 19-5-2005 at 04:41

Yellow as in Na2CrO4? Try and add some HCl, it'd form dichromate which is red.
Otherwise, it could be the hydrated form of Fe2O3, which is yellow too (albeit insoluble, but maybe it forms a very fine suspension).

S.C. Wack - 19-5-2005 at 08:23

The electric yellow of chromate and sugar drink orange of dichromate are unmistakable but not indicative of the amount in solution. One could compare with the weight before leaching.

edit: obviously this would require knowing the type of steel, to find the Cr %. Perhaps acid/formaldehyde -> Cr+3, would be a better test of Cr content, though this would require one to have a little formaldehyde on hand.

Anyone interested should download/look under chromium in Thorpe's Dictionary of Applied Chemistry.

[Edited on 19-5-2005 by S.C. Wack]

Pyridinium - 19-5-2005 at 10:28

Quote:

Merck Index indicates Cr metal is attacked by caustic alkali and even alkali carbonates.

It didn't say anything about them having to be molten, either.

So I could leave a piece of stainless in conc. NaOH for oh, about 6 months, and I might get some Cr salts even without electrolysis.

I have an old qual. analysis text here that says Cr(OH)3 is an amphoteric hydroxide.

Picture

12AX7 - 19-5-2005 at 17:52

It looks a bit greener, say like dilute ferrous chloride, than the picture shows. Doesn't seem too concentrated to me, although I'm sure it's strongly basic (it's kind of slightly viscous like that).

Dropping HCl didn't cause anything to precipitate or change color in the vicinity of the drop.

Other thing I was thinking was to heat it with a reducer (like some powdered charcoal) to get it down to powdered oxides and magnetically seperate the Fe3O4. That would at least concentrate it to a brittle nichrome.

From what I've read, the stainless should be from 15 to 20% Cr and 8 to 15% Ni.

Tim

Chromate.jpg - 45kB

Pyridinium - 19-5-2005 at 18:28

Quote:

Originally posted by 12AX7
Other thing I was thinking was to heat it with a reducer (like some powdered charcoal) to get it down to powdered oxides and magnetically seperate the Fe3O4. That would at least concentrate it to a brittle nichrome.

That sounds like a good idea; roast everything to dryness (first wash away all soluble alkali) to turn all hydrous Fe oxides to Fe2O3. Then roast again, this time in contact with chunks of carbon (maybe not powder or it could go *poof* when it gets good and hot)... hot Fe2O3 will form Fe3O4 in contact with carbon (I've done this).

Weird reaction?

Jome - 20-5-2005 at 02:49

Yesterday I tried to make elemental cromium from a stochiometric mix of Cr2O3 and Al.

7g initiated by 1,5g of KClO3+sugar burned in about 1,5 seconds, leaving a dark gray, partially molen powder which obviously had some Cr2O3 green in it.

First I was not sure how to "prove" it was cromium, but I've read about this wonderful method of cromate-manufacture from the hive, so I decided to give it a try. I added the dark grey powder to an ammount of 15% muriatic acid.

I didn't measure since this was only a 100ml beaker size experiment, not an actual manufacture. After a few hours the solution was very dark dark green, with some solid pieces still laying in the bottom. Now about 10ml of this solution was added to 20ml of water.
After KOH flakes was added and the solutions stirred. First the excess acid was neutralised, then Cr(OH)3 precipated and at last the precipate dissolved to form some sort of cromite.
To this cromite solution was then added 20ml of 30%hydrogen peroxide solution, which made it bubble about as much as a soda and coloured it dark yellow.

Now I had potassium cromate! So far, nothing weird. The cromate actually looked just like soda, "forrest fruit fanta" to be exact. Oxygen being evolved at a slow but steady rate. Guess I added to much H2O2....

But then I remembered a procedure from frogfots page to form red dicromate from glacial acetic acid and cromate. I decided to add some 25% to the clear yellow solution in a test-tube.

I did not clean the 10ml measuring cylinder from cromate residue before measuring up the acetic acid, they were going to meet anyway I figured. So about 0,1-0,2ml of yellow dicromate drops was still in there, but when the 2ml of acetic acid was added something weird happened: It turned blue. Not red as It was supposed to.

About as blue as concentrated copper-sulphate soln. When 2ml of yellow chromate solution was added to 2ml of 25% acetic acid the solution got so blue it was almost black.

Adding the dark-blue solution to 20 parts water made it dissolve in a weird way, leaving a very slight yellow tone.

I also discovered that the 0,1 cromate to 1ml of acetic quickly looses its strong blue color, after 20 minutes (writing this) the solution in the test tube has lost all it's blue color, in fact the blue is very hard to see. What happened?

Only blue Cr-compound that I know of is perchromate, but could this be? I'd think that would be much harder than this to make.
Any idea, anyone?

sparkgap - 20-5-2005 at 03:28

From your description, the transient blue color does seem to be due to peroxychromate production. The peroxide added, with acidification are the conditions, IIRC, for producing the transitory peroxychromate, which lasts only for a short period of time before the solution turns from a pale green to a slight yellow. That much I remember from undergraduate qualitative analysis.

See also this.

sparky (~_~)

[Edited on 21-5-2005 by sparkgap]

Jome - 20-5-2005 at 04:24

That explains it all! Addning it to water raises the pH so that the strong color gets more diluted than one would think it'd be. This seems like viable way to retrieve cromate ifrom Cr2O3, now if only I could get the Al out....Al is also amphoteric forming aluminates floating around the solution.

Btw how would the FFC cambridge process, that new method of making titanium from TiO2 work for chromium?

This page:
http://www.msm.cam.ac.uk/djf/FFC_Process.htm
States that Cr2O3 and lots of other metaloxides could also be reduced to its metal (even SiO2

) but what I dont get is how it works, TiO2 is a semiconductor but Cr2O3 isn't.
CaCl2 melts at 782 C, not at all impossible to get.

12AX7 - 20-5-2005 at 04:39

Quote:

Originally posted by Jome
This seems like viable way to retrieve cromate ifrom Cr2O3, now if only I could get the Al out....Al is also amphoteric forming aluminates floating around the solution.

Could reduce it down to say CrCl3 + AlCl3 and try cementating with zinc or plating out. If you want metal, anyway.

Quote:

Originally posted by Jome
This page:
http://www.msm.cam.ac.uk/djf/FFC_Process.htm
States that Cr2O3 and lots of other metaloxides could also be reduced to its metal (even SiO2

) but what I dont get is how it works, TiO2 is a semiconductor but Cr2O3 isn't.
CaCl2 melts at 782 C, not at all impossible to get.

An eutectic of NaCl, KCl, CaCl2 and/or MgCl2 (supposedly MgCl2 is quite corrosive so may be the best route, in terms of dissolving things) melts as low or lower than red heat, though solubility may suffer from being cooler.

Anyway, the process described is the Hall process, with absolutely no details about how it is different and why it works where Hall doesn't. In summary, the oxides are *dissolved* in the chloride solvent and electrolyzed, as electrowinning metal from an aqueous solution.

Edit: that was weird. I accidentially doubleposted, got the "only your second message was stopped" screen, but both posts appeared anyway... so I deleted one, but that pulled both down! Good thing I copied to clipboard...

Tim

[Edited on 20-5-2005 by 12AX7]

[Edited on 20-5-2005 by 12AX7]

Jome - 20-5-2005 at 05:55

Great idea to use Zn.

However the FFC process does not seem to be similar to the Hall process except the use of elektricity.
http://www.corrosion-doctors.org/Electrowinning/Aluminum.htm
The Al2O3 is dissolved in cryolite and is reduced by a Carbon cathode. The molten Al sinks down to the bottom and is collected.

In the FFC process the carbon instead is the anode, and the oxide is the cathode, which is reduced to the pure metal (sponge) without ever being dissolved. Quite different I´d say!

12AX7 - 20-5-2005 at 06:21

Quote:

Originally posted by Jome
However the FFC process does not seem to be similar to the Hall process except the use of elektricity.
http://www.corrosion-doctors.org/Electrowinning/Aluminum.htm
The Al2O3 is dissolved in cryolite and is reduced by a Carbon cathode. The molten Al sinks down to the bottom and is collected.

Well no, the electrolysis directly reduces the aluminum to the cathode, the combustion reaction is incident. Would be great if aluminum didn't form carbides. (I know, you can smelt aluminum (and silicon) just fine, but practically, it requires an arc furnace.)

Quote:

In the FFC process the carbon instead is the anode, and the oxide is the cathode, which is reduced to the pure metal (sponge) without ever being dissolved. Quite different I´d say!

Oops, right you are. I shouldn't read this early in the morning.

Tim

Jome - 21-5-2005 at 01:11

Seems like only half of the chromium pieces / powder really got dissolved in the hydrochloric.... Could it be that this reaction is usually slow but proceeded faster because some of the Cr pieces has Al in it?

Im going to try purifying the Cr today, precipate cromium metal powder by addidion of zink, then rinse this. The thermite method could be a quite useful route to small quantities of elemental cromium powder. Something like:
Cr2O3+Al---->(impure) Cr blobs
Impure Cr is dissolved in hydrochloric acid. The solution filtered and dried down to impure CrCl3 chrystalls, these are dissolved in water and solid Zn is added which precipates Cr and not Al. Powder is rinsed and can then be used to whatever use Cr may have (Im trying to make ammonium dicromate myself)

Just gotta find a good ratio of Cr2O3 to Al, the reaction with 25micron flake Al stoich mix was to fast, most of the powder flew away.

Edit: OR of course I could use NaOH as suggested earlier in this thread to remove Al and get fine Cr powder from the start. Only problem would be Cr2O3 residue, this seems hard to dissolve in KOH at least (I've tried).

[Edited on 21-5-2005 by Jome]

neutrino - 21-5-2005 at 05:27

Or, Cr2O3 + 6HCl -> 2CrCl3 + 3 H2O.

Jome - 21-5-2005 at 07:11

That doesn't seem to occur, neither does KOH affect Cr2O3. I've had them sitting in test tubes now for over a week, and if Cr has been dissolved its not much, in fact I can hardly see any green color. And since I just prepared CrCl3 I've seen how powerful that green color is.

I use 30% HCL and Cr2O3 from ceramic "chrome green".

I tried CuSO4 too but nothing has happened there either.

Aluminotermic reduction of Cr2O3 seems to be the only way if one starts from Cr2O3!

chemoleo - 21-5-2005 at 08:07

Well, there is always fusion of KNO3 with Cr2O3 to yield chromate (iirc). But to get Cr metal - I agree, thermite seems the easiest route from Cr2O3.

As to chrome oxide reacting with HCl - indeed this is NOT the case. Cr2O3 is very unreactive, and (anhydrous) CrCl3 is either made by direct chlorination of Cr metal (forming red-violet needle crystals when sublimed at 945 deg C) or by reacting Cr2O3, coal at red heat with chlorine gas.
Aqueous solutions thereof are the hexahydrate, which is green, and crystallises nicely + is highly soluble (50g CrCl3 per 100g H2O).

Also, I am not sure if Zn is able to reduce Cr3+ to the metal, did you check the reduction potentials? A strongly acidic solution of Cr3+ is reduced by zinc to Cr2+, i.e. the dichloride, which is blue.

CrCl2 is made by reacting HCl gas with Cr metal directly, it is colourless.

[Edited on 21-5-2005 by chemoleo]

neutrino - 21-5-2005 at 09:44

Oh, it's one of those oxides (like calcined MgO).

Jome - 21-5-2005 at 10:29

Zn+2 + 2e- ---> Zn is -0,76
Cr+3 + 3e- ---> Cr is -0,74

So in theory it should work. But perhaps it'd be better to use al (from foil) followed by rinsing the powder with NaOH-soln to remove Al(OH)3 and other shit that could form.

12AX7 - 21-5-2005 at 12:11

Hm, it is quite close. Worse comes to worse you could electroplate it out, possibly with a sulfuric electrolyte instead (being that electrolytically chlorinated heay metal solutions may not be too fun

). Might end up with a sponge (not a problem if you can melt it), but at least it'll be more efficient than hexavalent plating.

Tim

Orange soda! ...Don't drink it!

12AX7 - 24-6-2005 at 00:06

Maybe 50g (damn, how did I neglect to weigh it

) of stainless steel, hydroxydized electrolytically in sodium chloride solution. Suspension washed, dried and calcined to an orange/brown (Fe2O3 primarily). Powdered and added to a Ca(OCl)2 solution (an uh, excessive amount...oops). Solution was warmed briefly, accelerating the already somewhat exothermic reaction. Let sit overnight, yielding an orange solution decanted from brown sludge (which I might as well process further to yield the nickel, and iron oxide for thermite or something). Sulfuric acid added, later precipitating calcium sulfate and forming chlorine and chloric acid (oops, remind me not to heat and acidify hypochlorite solutions) as evidenced by the noxious ClO2 odor.

So, I'm basically ready to plate out chromium from this solution, eh?

Tim

P.S. Oh, and the subject? It DOES look like orange soda!

Pyridinium - 24-6-2005 at 11:15

Quote:

Originally posted by 12AX7
Powdered and added to a Ca(OCl)2 solution (an uh, excessive amount...oops). Solution was warmed briefly, accelerating the already somewhat exothermic reaction. Let sit overnight, yielding an orange solution decanted from brown sludge (which I might as well process further to yield the nickel, and iron oxide for thermite or something). Sulfuric acid added, later precipitating calcium sulfate and forming chlorine and chloric acid (oops, remind me not to heat and acidify hypochlorite solutions) as evidenced by the noxious ClO2 odor.

So, I'm basically ready to plate out chromium from this solution, eh?

I'm not going to try coming up with a balanced equation for that first part with the Ca(OCl)2 but I'm guessing it first makes some calcium chromate, chlorine, and HCl which then form some HOCl, ClO2, and dichromate, maybe with some residual chromate left over

...then the sulfuric acid really pushed it in the direction of dichromate, ClO2, etc.

but at least you got
2H(+) + 2CrO4(2-) <----> Cr2O7(2-) + H2O

I know you knew that I just felt like posting it :-P

So you're going to try plating from dichromate solution... super high current density I guess? Or can you plate out some mossy / granular Cr even at low current?

12AX7 - 24-6-2005 at 12:03

Quote:

Originally posted by Pyridinium
but I'm guessing it first makes some calcium chromate

Ya know, the funny thing about that- data states it is mildly soluble, yet I never saw any precipitate. It was an orange solution I poured off the brown sludge before adding sulfuric acid.

Quote:

So you're going to try plating from dichromate solution... super high current density I guess? Or can you plate out some mossy / granular Cr even at low current?

I forget what they do to get the infamous shiny smooth deposit, I'll have to look it up. I could also add a reducing agent, precipitate Cr(OH)3, cook to Cr2O3 and thermite it out.

Tim

chloric1 - 25-6-2005 at 09:54

Quote:

Well, I had an accident about a year ago. I had made a 29% KOH solution that I was going to use for fuel cells and some how(Don't rember) I mistook this solution for water and mixed some of it into another half filled water bottle. ARRG! I realized that if I boil it back down so it would all fit in the 2 liter bottle I could approximate the 28-30% solution again.

Well, to make a long story short I used a stainless steel pale to due the boiling. To my dismay the alklali was eating the pan! I was floored that it was doing this without an oxidizer and aqueous conditions none the less! Turned out the the attack was mild and superficial. And today I still use the pail without incident. Unfortunately for my carefully prepared KOH electrolyte it has some hydrated ferric oxide contamination floating around in it.

Will have to filter it someday when I can commit an afternoon. Maybe tommorrow?