So, overall, my current investigations have led me to learning about oxidation. I'm learning about oxidizing various primary and secondary alcohols to
their respective aldehydes and ketones, in particular. I've read in a number of text books about a fantastic oxidizing agent for primary and secondary
alcohols being potassium dichromate. I had also read about using NH4NO3 and Cupric Acetate in a solvent of Acetic Acid, which is the method I will be
using first. But while waiting for my products to arrive in the mail, and after reading about how stainless steel was a combination of nickel,
chromium, and iron, I decided to attempt to dissolve a bunch of old stainless steel flatware in acid so that I could isolate the chromium, and
possibly prepare potassium dichromate at some point in the future.
I've since read that HCl is a significantly better option for dissolving stainless steel, but I used H2SO4 this time around, which, on a side note,
prompted me to investigate proper handling of this acid, as the stuff is ridiculously corrosive!
I took a 1L pyrex Erlenmeyer flask and placed about 7 forks inside. Unfortunately, I was confined to using the smaller utensils as they were the only
ones that would fit through the neck of the flask which is a standard 1L erlenmeyer of unknown specific width of the neck. I then proceeded to pour
roughly 600mL of 96% H2SO4 into the flask, which came about half way up the utensils. I was under the impression that this would take a minute, so I
decided to jack up the temperature. I was afraid to submerge a thermometer into the conc. acid, so I am not really sure what the temperature rose to,
but there were some *horrible* gas and fumes escaping from the flask, which I imagine was sulfur dioxide and possibly water vapor when the temperature
rose about 100C. The temperature probably raised to at least 130C, at most 160C. I hypothesized this based on the heat setting I was using on my hot
plate.
I let it ride at a fluctuating heat of 130-160 for about two hours where the solution went from transparent and colorless to a nice full-bodied deep
green yet mildly opaque. I then turned off the heat and left for 2 days. When I returned the solution was completely transparent but still deep green.
I turned up the heat again and it started getting more opaque and more and more green. I noticed that the flatware was kind of stripped. Like, all the
pretty, ornate looking work on the flatware was dissolved, the shininess gone leaving a darker silvery metallic colour and the flatware looked like
some basic military shit . Being sort of pleased about this and thinking it was predominately chromium on the outside of the flatware, I took it off
the heat and let it cool down.
Then made a ~25% NaOH solution which I added about 5mL and it bubbled like mad and the solution got hot as hell. Being comfortable in the strength of
Pyrex, I would add the solution in spurts of about 5-10mL, let it bubble a bit and smoke hard, then once the gases reduced a bit and bubbling calmed
down, I would add a little more. It took nearly 30 minutes to add 100mL and I was noticing a lot of precipitation. At first the solution got really
really green, but later on after letting it cool a bit, I noticed the precipitate was kind of like a light grey color and the solution was
consistently more and more green.
Due to how crazy the solution would heat up and bubble, I decided that perhaps I should do this another way. So, I made another 600mL of ~25% NaOH
solution and put it into a 2L Erlenmeyer flask (not Pyrex, I was a little apprehensive about this, but thought it might be okay since I was out of
Pyrex Erlenmeyer flasks). I swirled the 1L flask a few times to mix up the solution well and then started pouring it into the 2L flask with the basic
solution. As the green acid solution hit the aqueous basic solution, it transformed to a brilliant violet/blue color and got really hot, but there was
minimal "out-of-control" bubbling and other craziness, but plenty of gas evolution still. I was trying to neutralize the solution so I kept checking
the pH, but it literally went from being at a pH of 11 to being at pH 9 and then with a little bit more acid straight to a pH of 2! So, I slowly added
more NaOH solution, but then it raised from 2 to 11 again, just like that.
I didn't go about filtering the precipitate because I was fearful of the acid, even once basified. Wasn't sure what kind of damage it would do to my
Teflon Buchner (not to mention the filter paper!). Would it be safe to work with once basified? Would it still burn through filter paper like nuts at
such extreme pHs?
What I am thinking happened is this: The chromium oxide was lining the flatware which came off first, along with a minor amount of nickel and iron. As
it reacted with the H2SO4, I was getting sulfates of these metals. When I added the NaOH solution, it reacted like mad and I got Iron Hydroxide and
Chromium Hydroxide as the NaOH yielded Na2SO4 which collected at the bottom (Sodium Sulfate is a good drying agent, right?)
I ended up scraping the whole thing after diluting it like mad, but correct me if I'm wrong about what could have been. The chromium and minimal
amount of nickel are pretty much going to be together (not bonded, but reacting the same way). The fairly colorless iron hydroxide and violet chromium
hydroxide (very beautiful color, btw) once precipitated would need to be separated. So, I could start adding acid again, which would encourage the
chromium hydroxide to dissolve back into the solution, yet the iron hydroxide would not. I noticed this happening, as the I was adding the acid to a
NaOH solution, when I was adding the acid, the solution got more violet-blue towards violet-purple and the precipitate that I saw when I was pouring
it out was a light, just light grey metallic looking stuff. So, I guess I could've filtered out the iron hydroxide once the chromium (and nickel?)
re-dissolved and then evaporated off the water (which would've taken forever, but i guess i could have distill it off with vacuum and then evaporated
the last bit) to yield chromium hydroxide. I am not really too worried about what to do with it at that point, I would just like to be able to isolate
the chromium so that I could potentially get to potassium dichromate at some point in the future.
If I was to use HCl to dissolve the stainless steel, how much faster does it break it down? I guess heat would speed it up, how much heat? What
containers are suitable for containing HCl or sulfuric acids full of chromium ions and friends? Any difference from the acids without
chromium/iron/nickel? Is the main reason for not using sulfuric acid with glass that isn't pyrex just because of the hot temperatures when reacting
with a base?
Brief theoretical overview regarding going from chromium (III) chloride to potassium/sodium dichromate. Once I've isolated the chromium(III) ions as
Chromium (III) Hydroxide probably, then it discusses using KOH or NaOH to oxidize to [Cr2(OH)6]3-. If I am isolating the Cr as Chromium (III)
Hydroxide, then would this step be necessary, or would I actually be converting the sodium to sodium chromate(VI) by bonding with the 6+ chromium
ions? I think that's what's going on when reacting with H2O2, correct? Followed by heating to decompose the remaining H2O2 and then adding AcOH to go
from chromate to dichromate.
If when isolating the chromium from the nickel and iron, I didn't remove all of it, would the remaining Ni and Fe oxidize in the same manner as the
chromium altering the stages of oxidation and ultimately creating a problematic contaminant? In fact, I'm wondering if the Ni would actually attempt
to cause reductions, or only in the presence of a co-catalyst?
I know a lot of these are stupid questions and I might be biting off more than I can chew, but any help or arrows in the right direction where I can
read/investigate would be much appreciated! thanks.
[Edited on 2-2-2010 by havarti_gouda]bbartlog - 1-2-2010 at 20:02
Not sure I can address all your points, but a few comments:
- Stainless steel as far as I know does not have a coating per se (other than the passivating oxide one a couple molecules thick). So whatever
dissolves is likely to be a proportional mix of the actual composition, unless you see some sort of leaching that results in a spongy formation
(unlikely here in my opinion). Ideally, you'd want to have an idea of what alloy of stainless you are dealing with so as to be able to calculate the
amount of iron, chromium, and nickel in solution. Conventional cutlery is 18/8 stainless, 18% chromium 8% nickel and the rest iron.
- Teflon is impervious to acids and bases for all practical purposes. Filter paper should resist your solution if it's basic. Strong sulfuric acid
might destroy it though especially if hot.
- I don't know if you're aware, but the pH scale is logarithmic. Which means that if you're dealing with strong acids and bases, it's a little tricky
to hit the middle of the pH range - a few milliliters of NaOH solution can move the pH of a liter of solution from 3 to 11. This is doubly true if you
don't measure anything. If you carefully weigh the acid and the cutlery, then weigh what's left of the cutlery, and do some calculation, you would at
least have a starting point to aim for in terms of how much base has to be added for neutralization.
- I don't have a good recommendation off the top of my head for separating the Na++, Fe+++, Cr+++ and Ni++ ions that are in your solution at this
point. Turning everything into either hydroxide or carbonate and then filtering and washing would let you get rid of the sodium, but for the rest I'm
not sure. It's possible that their reactivity is different enough that you could then (with care) dissolve them in order using acid, but you'd have to
research it.
- Even if you have chromium, turning it into chromate may be a project.
(edit) - oh, and heating hydrochloric acid is even fumier than H2SO4, generally. If you're going to heat/boil concentrated HCl you need to arrange for
reflux and have a trap, or better have a fume hood. If the H2SO4 is working I'd stick with it.
[Edited on 2-2-2010 by bbartlog]havarti_gouda - 1-2-2010 at 20:52
Oh my. Even fumier sounds horrible. I don't currently have a fume hood, but I've prepared schematics for one. I need to break the gumption barrier and
build it. One possible problem in my design would be that it's going to be constructed of a wooden frame tightly wrapped with thick plastic on all
sides, top, and back with a lighter plastic on the front to easily draw back for access to the work space. Then overhead a vacuum source (tweaked
vacuum cleaner modified to reduce noise and control suction) with tubing that runs out the window. I am just wondering about solvent/acid fumes
degrading the plastic or vacuum over time which leaves me with two options for the plastic: a.) make it easy to replace the plastic, b.) treat the
plastic with spray-on teflon or something similar. But I'm still facing the vacuum problem. But that's for another post.. :-)
Regarding the spongy formation, it's funny you mention that because I forgot to. Remember how I said the acid only covered half of the forks, well,
there was a lot of light green spongy looking gunk that had formed on various parts of the top portion of the forks. Probably exactly what you're
talking about.
I definitely should have taken proper measurements, but I guess I was assuming that I would try to isolate as much as possible and whatever was
generated was fine. I will certainly take all weights when I get the next one started. I am fairly positive that the composition of the flatware is
18-8. My general estimation would be that in my last batch I shuld've had about 600-700 grams of chromium.
I need to get an electronic pH meter. I was using litmus paper, which is a hassle to work with when required to check between drops of solution.
From what I was reading, if I can isolate Chromium (III) Hydroxide or Chromium (III) Chloride from the stainless steel, then going to Sodium or
Potassium Dichromate doesn't seem very difficult at all. The processs is outlined very well on that site I updated my last post with. It's just
isolating the Chromium 3+- that's going to be a major bitch. Maybe there's an easier way! :-)12AX7 - 1-2-2010 at 23:26
Here's what I would do:
Since you have sulfuric, that's fine, stick with that. Dilute it to 30% and boil the stainless in it. Hydrogen gas will be released instead of
noxious SO2 gas, because water is being reduced instead of sulfuric acid. You also keep the sulfur in solution, instead of losing it as gas. (BTW,
if you've survived SO2 from hot sulfuric acid, 30% HCl at a somewhat lower temperature will not be any worse.) Use enough acid to dissolve all the
metal (assume Fe(II), Ni(II) and Cr(III) are formed) plus enough excess to leave the solution at 10% acidity.
Once all the metal is dissolved, cool the solution. Iron is predominant (being >70% pure iron), so you'll probably get a lovely deposit of
FeSO4.7H2O. As air exposure and crystallization proceed, the crystals may become darker, due to Fe(III) and nickel impurities. (Seperating Ni from
Fe is good enough for another thread, and it has value since nickel is moderately expensive.) Don't worry about crystallizing everything, you'll
eventually be left with an indeterminate mush anyway. You can simply remove a big chunk of iron this way, reducing the metal content by maybe half,
raising your chrome content from 18% to 36% roughly.
Next, take your solution and neutralize with something cheap (sodium carbonate comes to mind, but if you have lots of NaOH, whatever floats your
boat). You should dilute the solution first, so it's not a syrupy consistency that forms partly neutralized globs. You should get a similarly gray
precipitate which, on exposure to air, turns black, then brown: this is the Fe(II) forming very reactive Fe(OH)2, which oxidizes to Fe3O4 and Fe(OH)3.
Don't bother washing or filtering the precipitate, it will be quite horrible. Instead, let the entire mess dry out. Big sodium sulfate crystals
will form (which is indeed a drying agent, and it will have done its job forming the decahydrate from solution), don't mind them. When everything is
more-or-less solid, heat it up to drive out moisture. I would guess 300C is a good start. This will firm up the precipitate so it can be filtered,
while the soluble materials are unaffected (Na2SO4 melts at ~800C). Don't go too hot or the oxides will turn into refractory sand that you'll never*
get anything out of. Beware of SOx fumes from unreacted metal salts: one shortcoming of calcination. There may be better ways, but it's hard to
argue with raw heat. Notice the junk will turn brown, if it hasn't already.
Once everything is cooled down, suspend it in water again. Wash and filter the precipitate; it should now be relatively free of sodium and sulfate
ions.
At this point, you can put the goo back in water with NaOH and H2O2 to see if the chrome can be oxidized, or...
*My favorite method for forming potassium dichromate directly is by fusing the dry solids with potassium chlorate. Combine in (ironically) a
stainless steel crucible and apply heat cautiously with stirring. Where the potassium chlorate melts (~400C), it starts bubbling, oxidizing any
chromium present, forming an evil dried-blood-red goo. The reaction is exothermic and will run away (foaming and overflowing your crucible) if
unwatched. Chlorine gas is released.
I haven't decided if iron and nickel are poisons to this process, if they catalyze the decomposition without making dichromate. I do know when it is
carried out on ordinary pottery grade Cr2O3 -- a material renouned for being the chemical basis of the indestructible gates of hell, it proceeds quite
nicely, ultimately resulting in a deep red molten solution which freezes into an orange cake. Dissolve, recrystallize (seperate K2Cr2O7 from
remaining KCl) and you've got a load of orange death to enjoy!
Timhavarti_gouda - 4-2-2010 at 11:26
Okay, I switched over to HCl. So, put several pieces of flatware into the flask with about 600mL of HCl. Start heating it up. *Damn* HCl dissolves
metals fast!
So, it got super super green to almost pitch black. Lots of black particles started collecting JUST above the bubbles, which I originally thought was
a sign the HCl was saturated, but realized in the second run that this was not the case. Anyhow, I took out the remaining undissolved flatware and
moved the solution to another, larger Erlenmeyer. Threw strong (25-30%) NaOH solution into an unjointed sep funnel and dripped it in. Much much nicer
than with the conc. H2SO4. Moderate bubbling, but nothing totally insane like the other day. At one point (probably around pH of 9 or 10), goes from
the black to a deep green and tons of precipitate forms which I am assuming is the Iron Hydroxide, Nickel Hydroxide, Chromium Hydroxide. Perfect!
This is when I realized that I had run out of NaOH! So, mixed it really well to see if that would spread out any areas of conc. base and let the
chromium dissolve, but apparently, one needs a decent bit of excess base for this. So, I just let it cool and decided to stay and babysit for a while
so I watched a movie. After the movie finished, I came and took a look at it, and just like should happen, the precipitate had sunk quite a bit to the
bottom and the solution was fairly clear. Checked pH and it was at 11. There were patches throughout of green, blue, and gray. I poured a little bit
out and let air get to it, and it turned to brown, just like Iron Oxide should. So, I'm going to add some extra base today, let the Cr dissolve,
decant solution and then was going to add some acid to bring pH to 9 ish, to bring the Chromium back out and then isolate that to concentrate it more.
Going to use the H2O2 method, coz it would be more readily available to me. Any suggestion on conc. of H2O2? OTC 3% work okay?
Just got my KNO3 today too. Oxidation abound! :-)12AX7 - 4-2-2010 at 12:51
You'll get a small amount of chromite (species between Cr(H2O)2(OH)4(-) and Cr(OH)6(3-)) by dumping NaOH, but I can't find the acidity of them.
Apparently the cumulative log Kf's for Cr(III)(OH)n are 10.1, 17.8, -, 29.9, -, -. Which means:
Cr(H2O)6(3+) + OH(-) <--> Cr(H2O)5OH(2+) + H2O log Kf = 10.1
Cr(H2O)5OH(2+) + OH(-) <--> Cr(H2O)4(OH)2(+) + H2O log Kf = 7.7
Cr(HO)3 isn't listed, probably because it's insoluble, and therefore listed under the Ks's instead. log Ksp = -30.2, which I believe is from Cr(3+),
so I think the proper log Kf against the (OH)2 species should be 12.4.
Cr(H2O)4(OH)2(+) + 2 OH(-) <--> Cr(H2O)2(OH)4(-) + 2 H2O log Kf = 0.3
So if I have that right, the solubility of Cr in base should be fairly low, requiring a substantial concentration of NaOH (pH > 13) to get it to
dissolve.
The table also shows Fe(II) forming Fe(OH)4(2-) log K = 8.58 (not as stable as Fe(OH)2), so make sure the iron is oxidized (Fe(III) shows no anionic
species), and Ni(II) forming Ni(OH)3(-) log K = 11.33. http://www.scribd.com/doc/6792576/638478
Page 8.85
Damn is Scribd ever hard to navigate. :puke:
Anyway, keep in mind also that, using hydrochloric acid, you get the chloride ions, which means your precipitate still contains a lot of chloride as
chlorohydroxides.
3% H2O2 is most definitely not suitable. At pH = 14, the sheer dilution will precipitate Cr(OH)4 without oxidizing it. You'll need something
stronger, and an extra supply of NaOH to keep things basic. Hmm, Oxiclean is a possibility, carbonate isn't as basic as NaOH but it's not neutral
either, and it has a higher density of H2O2 than the 3% topical stuff.
TimRandom - 3-11-2010 at 01:43
Can we reduce solution of stainless steel with zinc and separate iron and nickel from chromium powder with magnet?Sedit - 3-11-2010 at 12:09
Not without bringing along the non magnetic powders along for the ride in great abundence. This was attempted seperating Nickle and Copper powder but
it proved useless. The Ni did stick to a magnet but Cu came along for the ride as well since its had to seperate particals this small from each other
presumably from electrostatic attraction or something.Random - 18-11-2010 at 01:35
That's bad, but I found one way to separate nickel and copper though.
1. Mix cupronickel in aqua regia to dissolve it.
2. Add aluminium to reduce copper chlorides and nitrates to elemental copper. It will also destroy any excess nitric and hydrochloric acid. (Copper
chloride dissolves aluminium alone without acid).
3. Now we have copper powder and nickel and aluminium in solution. Copper powder should be mixed with hcl to dissolve nickel that has formed too.
Nickel is put into nickel and al solution.
4. Now we can reduce nickel in solution with zinc to metallic nickel.
This should separate them I hope.Sedit - 18-11-2010 at 05:44
Nickle comes along as you try to reduce the Cu with Aluminum as well.
Iv been toying with the use of Ammonia hydroxide solution for the seperation of Ni and Cu and it appears to have some success by precipitating the Ni
as the hydroxide(oxide?) and complexing with the copper keeping it in solution but the ratios are touchy to say the lest. To little hydroxide and
nothing happens, to much and the Nickle is also complexes keeping it in solution.
In between you will find a spot where the Nickle precipitates as a light green hydroxide that can be decanted (filtering is a bitch), washed
repeatedly with fresh water and processed further into the chloride for precipitation with Aluminum or crystalized and processed for storage.tetrahedron - 16-10-2012 at 18:28
3. Now we have copper powder and nickel and aluminium in solution.
can't the nickel simply be plated out (given the right concentration/current/electrodes etc)? or would the aluminum disturb/contaminate the
deposition?12AX7 - 16-10-2012 at 22:32
Yes. The reaction is something like,
6 Cr2O3 + 8 KClO3 = 6 K2Cr2O7 + 2 KCl + 3 Cl2
So you can see it's quite acidic and has extra oxidation power that could be harnessed with a mild base. In practice, the loss due to decomposition,
and the relative cheapness of the reactants (but then, I'm spoiled when it comes to chlorates) means you can use an excess and come out fine.
I don't know if perchlorate will perform the same reaction; if so, it should take a higher temperature, both to melt the stuff and to induce
decomposition. Also I don't think I ever ran this with sodium chlorate, which should work as well. Which would be quite handy given the usefully
high solubility of the product.
TimDaffodile - 23-3-2016 at 15:15
Okay so I've been going over the method in which stainless steel is dissolved, carbonates and hypochlorites are added, etc. Something that gets to me
though is after the hypochlorite addition, when ammonia is added to reduce ferrates to insoluble shit. Would the ammonia mess with the newly formed
chromates as well?
