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blogfast25
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Lime green solution...
Below is a photo of a few hundred ml of lime green waste solution, by-product of making (NH4)2SnCl6 from pewter metal dissolved in a mixture of strong
(15 – 22 %) HCl and 38 % HNO3 and adding about a stoichiometric amount of NH4Cl. The solution is then basically simmered down to almost nothing
(during which the green colour clearly intensifies) and then cooled and iced. The hexachlorostannate drops out with yields now of up to 92 % of theo.
Replace the NH4Cl with KCl mole for mole and obtain K2SnCl6, also with the same green by-product.
Using pure tin metal (99.9 % as advertised) no green is obtained.
The pewter should contain some antimony (but I’ve not been able to demonstrate that), no copper (verified) and no appreciable amounts of lead (also
verified). The HCl does contain small amounts of iron even though switching to a cleaner grade didn’t stop the lime green from appearing.
The colour is quite striking and unusual in inorganic chemistry. Its intensity indicates appreciable concentration…
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watson.fawkes
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The easiest OTC source, to my thinking, for a bit of antimony would be in hardened lead, often used for wheel balancing weights. A local tire shop
might spare a few. In the US, bullet casting alloy is readily available in many places and generally has some antimony in it. Along this line, perhaps
harder to find but also possible is "linotype" alloy, which is lead-antimony-tin.
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blogfast25
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Oh but I can get 500 g of > 99 % Sb for 8 quid (plus PP) from a British eBayer (and I probably will) but I’m not sure how this would solve this
mini mystery: I know of no lime green antimony compounds, do you? Antimony from the pewter will almost certainly be found in the (NH4)2SnCl6 (IV) as
NH4SbCl6 (V), unless it is very much more soluble than the Sn homologue. If the solubility is about the same, Sb may stay in the mother liquor as a
minority constituent… There’s a paper freely available on the Net that describes the separation of Sn and Sb using the hexachlorostannates, a
malonated Dowex-1 equivalent exchange resin and malonic acid solution as eluant… Sn and Sb are remarkably similar, chemically.
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DJF90
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How about nickel...? I seem to recall that colour from my school days, [NiCl4]2- IIRC.
[Edited on 23-11-2010 by DJF90]
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watson.fawkes
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Quote: Originally posted by blogfast25  | | Oh but I can get 500 g of > 99 % Sb for 8 quid (plus PP) from a British eBayer (and I probably will) but I’m not sure how this would solve this
mini mystery: I know of no lime green antimony compounds, do you? Antimony from the pewter will almost certainly be found in the (NH4)2SnCl6 (IV) as
NH4SbCl6 (V), unless it is very much more soluble than the Sn homologue. |
I was thinking that you might try
making some NH4SbCl6 directly and see if it's green; nothing particularly sophisticated. If you were in the US I'd recommend Rotometals. They sell Sb in both ingot and shot, but I'm not sure that it's economical overseas.
The nickel suggestion might be it too. I was looking at babbitting metals (for bearings) today. The tin-based babbitts are very much like pewter, and
one of them listed had 0.15% Monel in it, which of the Monel alloys not listed, but they're all high-nickel alloys, with varying amounts of copper,
iron, manganese, and others.
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blogfast25
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Yes, NH4SbCl6 can almost certainly be made as above, substituting Sb for Sn, the former also dissolves in aqua regia.
But if that’s right and the pewter (it’s a food grade –from a drinking tankard) contains, say nominally 5 % Sb, then the hexachloroantimonate
would have to be mightily coloured for the lime colour to be accounted for. It’s not very likely: it has basically the same electronic configuration
as the stannate equivalent: a full d subshell coordinating the chloride ligands..
Still, worth a ponder… The hexachloroantimonate has (or had?) use because the caesium salt is very insoluble.
Nickel is something I'll test for but the colour just seems so ntense...
[Edited on 23-11-2010 by blogfast25]
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MrHomeScientist
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The only thing I've run across in my own experimentation that was similar to that color was copper (II) chloride in HCl - it turns a similar shade of
green when the solution is high on Cl-.
But since you tested negative for copper, my post was useless  Just thought I'd
mention it in case that sparked an idea for you.
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watson.fawkes
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When you put it that way,
you spark a different question: Is the color monochromatic? You don't need a calibrated spectroscope to answer that, just a basic one. A simple prism
or diffraction grating will suffice. If you tilt the optic around, you should see only one color-image of the bottle, not more than one and not a
streaked-across image.
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Arthur Dent
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Indeed the green color of your solution is intense. I've made some Nickel Chloride by dropping a few early 70's canadian 5 cent pieces (99.9% Ni) in
conc. HCl and almost immedioately, the acid turned light green, and the finish of the nickels wasn't even affected yet. Now the pieces are completely
dissolved and the solution is the deepest dark emerald green you've ever seen... really beautiful.
So I think it doesn't take very much nickel to give an acid solution a green tinge.
Robert
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woelen
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The color of the solution is not really like the color of nickel(II) in solution. That is really green and not yellowish green. The complex NiCl4(2-)
is not very stable and even a fairly strong solution of NiCl2 in HCl has a more green color.
I can imagine it contains nickel, but its not the only contamination. Molybdenum may also be one of the constituents. Molybdenum can give intense
yellow/green solutions.
Can you add some solution of Na2SO3 to the yellow solution and boil? When molybdenum is present, then you get a blue color.
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blogfast25
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@Watson: I do have a primitive spectroscope but more to the point I also have a DVD-R: perfect for this purpose. Will test later on. To the naked eye
this doesn’t look monochromatic to me. I’ve always suspected some complex and they don’t usually generate monochromatic colours…
@woelen: molybdenum, eh? What the hell would that be doing in pewter?? I will certainly test this but with metabisulphate, that also generates SO2
which I’m guessing is the purpose of the sulphite?
************
Well, the green is certainly not monochromatic, rather it’s a wide green streak in my spectroscope.
Here’s some test tube tests with the liquid:
1: control; solution as such.
2: with small amount of sodium metabisulphite after boiling. Slight intensification of colour due to boiling in. No blue.
3: with a larger amount of sodium metabisulphite after boiling. Slight intensification of colour due to boiling in. No blue.
4: with weak NH3 solution: Sn(OH)4 drops out, no blue Cu2+/NH3 complex (as expected).
5: with 5 M NaOH: Sn(OH)4 drops out but the green hue is clearly maintained; so it can exist in alkaline conditions too. That rather rules out Ni2+
too.
The only conclusion so far is that the 'waste' solution still contains appreciable amounts of hexachlorostannate: the ammonium salt appears very
soluble, even in the cold...
[Edited on 24-11-2010 by blogfast25]
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kmno4
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Take sol.1 and make it acidic with HNO3, next add H2O2, boil it - any changes ?
Take sol.1 make it strongly alkaline with NaOH, filter out Sn salts and to clear solution add H2O2 and boil it - any changes ?
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blogfast25
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| Quote: Originally posted by kmno4 | Take sol.1 and make it acidic with HNO3, next add H2O2, boil it - any changes ?
Take sol.1 make it strongly alkaline with NaOH, filter out Sn salts and to clear solution add H2O2 and boil it - any changes ? |
@kmno4: done!
The series above is with addition of reagents but BEFORE boiling:
1: control; solution as such
6: control plus HNO3 plus H2O2. Solution darkens. Possibly typical of Fe3+ or Fe.NO (2+).
7. Test tube 5 filtered and H2O2 added. Same colour as 6.
Above, AFTER boiling 6 and 7:
6’: after boiling 6 the green re-appears. I’ve seen this many times when making the hexachlorostannates (NH4 or K) from pewter: at some point
during the simmering process the colour turns lime green immediately…
7’: the alkaline solution does not change colour upon boiling of 7…
I really suspect iron but what compound?
[Edited on 24-11-2010 by blogfast25]
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kmno4
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I think that the green something could be some (probably chloride) complex of Cr(III).
Such complexes are intesively green colored.
If 7 really contains chromates - additional tests are needed.
[Edited on 24-11-2010 by kmno4]
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not_important
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I'm going for iron, too. Hard to say for certain, almost certain a chlorine containing complex but it may have other things in it. Doesn't take too
much Fe to give colours, and it's everywhere. A complex might interfere with some of the simple tests for iron, and Fe(III) hydrated oxide goes
colloidal fairly easily when the Fe is not too concentrated so simple filtering might not removed it. Indeed, when purifying compounds that might be
contaminated with iron, a common method involves oxidising the iron to Fe(III), then using a low solubility oxide/hydroxide/carbonate of the cation
metal to push the pH alkaline and precipitate the Fe(III) out on the surface of that so as to be more easily removed.
Cr would seem to be low probability but a check might be in order.
BTW - at lower concentrations Cu and Ni look similar when aq NH3 is added - the blue colour from each is similar.
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DougTheMapper
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It's probably iron. I have obtained similar results while reacting shavings of Mg with HCl to yield MgCl2. An identical brilliant lime green color
formed which I had assumed was dissolved chlorine. I was surprised, however, when it remained after the solution was boiled so I suspected
contamination of some kind - probably from the steel cutting tools used to machine the Mg. The only other known elements in the Mg alloy were trace
amounts of aluminum and manganese.
Victor Grignard is a methylated spirit.
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kmno4
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Very often Mg shavings are coated with Cr (to prevent oxidation).
Besides - appropriate tests for Fe (II/III) are commonly known and very easy to perform.
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woelen
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If you want to confirm that the color is due to iron, then add a solution of K3Fe(CN)6 or K4Fe(CN)6 to a mix of the green and brown solution. A deep
green or blue color confirms iron. This test is very sensitive! The color is due to formation of prussian blue.
Another very sensitive test is to add the brown solution to a solution of a thiocyanate. This will lead to formation of a bloodred complex,
thiocyanatoferrate(III).
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kmno4
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Yes, but one must remember that tested solutions should be acidic if you want to get reliable results.
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blogfast25
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Thanks ALL for your responses.
Tests with KSCN had been carried out before (Fe was always my prime suspect) but were negative, to my surprise. I will now repeat a thiocyanate test
on solution 6. Bear in mind that the FeSCN(2+) complex is a fairly weak one: anything else complexing the Fe more strongly may obscure it.
I’ve unfortunately no ferro or ferri hexacyanides.
Recapping a bit:
Sources of Fe are: HCl grade used, contaminate Fe in pewter, cutting tool (a hacksaw).
Green colour appears only with pewter, not with 99.9 % tin metal. With the latter a slightly yellow solution is obtained, pointing to naked Fe3+.
Switching from a fairly iron rich HCl grade to a much purer one made no difference.
Although this is essentially still nothing more than a hypothesis, it would seem to point to a complex of Fe3+ in the presence of Sb (V). I’ve
ordered some 99.2 % antimony (of course that too maybe Fe contaminated) for some tests.
One way to almost completely eliminate Fe (while also avoiding Sb) would be to make some (NH4)2SnCl6 from pure tin and to recrystallise it carefully,
then treat the dried product with acetone to leach out any FeCl3. The Fe, Sb free hexachlorostannate could then be reduced with Al/HCl and reoxidised
with HNO3 (or H2O2) to check if the lime gree colour appears or not.
Quick update:
Adding 1 M KSCN to test tube 1 (the lime green control) yields a very strong positive for Fe3+, unmistakably so. So there is definitely quite a bit of
Fe (III) in the solution. Question is, what Fe (III) compound or complex has a lime green colour? I still suspect antimony has something to
do with this and will run some tests when I get my consignment of Sb...
[Edited on 25-11-2010 by blogfast25]
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kmno4
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Fe does not have to be the only metal present in this solution.
It also does not explain why solution after precipitation Sn salts is still gren and why this green colour turns to yellow after H2O2 treatment.
I would strongly recommend to do test for Cr (via CrO5) in solution number 7. All tests you performed tell me that it must be Cr.
Besiedes, if you have some fluoride (KF, NaF...) you can mask Fe(III) directly in solution 1. Then you can see if green colur still remains. Fluoride
gives more stable complex with Fe(III) than thiocyanate.
After masking sol.1 with F(-) and adding larger excess of SCN(-), it is quite possible that you will get (after warming) dark violet-bule
Cr(III)-SCN(-) complex.
Of course, I can be wrong but experiment is a king.
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blogfast25
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@KMnO4: yes, chromium is still in the race, I’d say. I have no soluble fluorides right now though. Not sure by what you mean by ‘via CrO5’? You
mean precipitate the Cr as an insoluble chromate (lead for instance)... Please explain… And Cr (III) forms a thiocyanate complex? News to me…
Edit: oh, I see: CrO5: Cr (VI) peroxide. That was kind of tested for in tube 2, wasn't it?
Carrying on:
So, about 6 g (theor.) of (NH4)2SnCl6 were prepared from tin (99.9 %), HCl 15 % and HNO3 38 %, as per usual. The solution assumes a light yellow
colour right from the start and after simmering and collection of the product a yellowish supernatant solution was obtained and tested for iron (III).
Before addition of KSCN 1 M:
8: supernatant liquor from an (NH4)2SnCl6 preparation from tin (99.9 %) metal: slightly yellow with a greenish tinge…
9: HCl 15 w% grade used in the preparation: clear and colourless.
After addition of KSCN 1 M:
8’: tube 8 reacting strongly positive for Fe3+
9’: tube 9 reacting weakly positive for Fe3+
In short, there’s Fe3+ in there too. Due to the iron contamination of the HCl used it’s impossible to tell what the main source is but the HCl is
a clear suspect…
The (NH4)2SnCl6 was rinsed thoroughly with iced DIW and turned snow white. I saw no good reason to recrystallise it as it appears to be iron free. A
KSCN test confirmed that.
When I’ve made some Sb compound, I should be able to test that the complex may form only in the presence of both Sn and Sb.
[Edited on 25-11-2010 by blogfast25]
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woelen
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Chromium does form a complex with thiocyanate, but it is not red. It has a deep purple color, but one needs quite some chromium and thiocyanate at
high concentration to see this complex. It hardly interferes with the very sensitive test for iron(III).
'Naked' iron(III) is not yellow, but it is (nearly) colorless. Only at very high concentration, iron(III) has a weak brown/pinkish color. Only when
chloride is added, a yellow color appears and due to hydrolysis in aqueous solution the color turns more like brown.
http://woelen.homescience.net/science/chem/solutions/fe.html
If you add either phosphate or fluoride to your yellow solutions, then you'll completely mask the iron(III). Both phosphate and fluoride form a rather
stable and colorless complex. If you have phosphoric acid or some soluble phosphate, just add some of this to your solutions. The color of the
iron(III) then is masked and if you still have some color after addition of that, then this is due to some other metal.
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blogfast25
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@ woelen: sorry, I did mean FeCl3 which does have a distinct, almost glowing, yellow colour. All my solutions here are born in a combination of HCl
and HNO3 and are naturally high in Cl-. Solid FeCl3.6H2O is yellow too.
As regards kmno4’s idea of testing for Cr as CrO5 (VI – peroxide), I doubt very much if Cr (III) – that is the most likely oxidation state of Cr
if there really is any present – can be oxidised in aqueous solution with HNO3/H2O2 in acid conditions as he suggested. A couple of quick tests
yesterday seemed to suggest that.
It maybe possible to oxidise chromite – Cr (III) in alkaline conditions – with H2O2 and heat to Cr (VI), I’ll check that this
afternoon. But industrially chromates are the result of oxidising Cr (III) with alkali fusion and in the presence of an oxidiser (although air oxygen
can suffice).
Another reason why I now believe Cr is far less likely than an Fe complex as the cause of the colour lays in test tube 5: control plus strong alkali
gives white precipitate of Sn(OH)4 and lime green supernatant liquid. While ‘in theory’ it’s possible that the suspected Cr (III) had gone into
solution as chromite, in reality that’s not very likely: the pH in solution wasn’t measured but it can’t be that high otherwise the Sn(OH)4
would have redissolved as stannate (Sn(OH)6 (2-)) too. By rights the Cr (III) should have dropped out as Cr(OH)3.nH2O.
So I think it’s probably possible to eliminate Cr from the list of suspects with some more tests along those lines…
[Edited on 26-11-2010 by blogfast25]
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woelen
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The CrO5 test is not easily performed here. CrO5 (actually, CrO(O2)2, the diperoxo complex of hexavalent chromium) is not easily made in your
solutions. H2O2 does not oxidize chromium(III), not even in the presence of a strong oxidizer like HNO3. The only oxidizer which can do the job is
peroxodisulfate, S2O8(2-), with silver ions as a catalyst. This combo quickly oxidizes chromium(III) to chromium(VI) in acidic solution. Next, you
have to be sure that all peroxodisulfate is destroyed before any H2O2 is added, otherwise the peroxodisulfate also oxidizes H2O2, giving sulfate and
oxygen.
It indeed is surprising to see such a strange lime green color. The only iron(III) complex I know which has such a lime green color is the trisoxalato
complex. I have some of that material, the ammonium salt:
http://woelen.homescience.net/science/chem/compounds/ferric_...
It can be made by dissolving iron(III) hydroxide in a solution of excess oxalic acid and adding ammonia to neutralize excess oxalic acid.
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