| Pages:
1
..
3
4
5 |
Texium
Administrator
Posts: 4580
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Ohhh, ok. I haven't had my tea yet this morning. Yes, the original grey-black sediment does sound like CuS. The picture that j_sum posted does not
really look like an iodine/KI solution to me. I would expect a dark solution with no precipitate in that case, but I suppose if it was dilute enough a
very fine suspension of I2 in a solution with no excess KI present could appear orange like that.
|
|
|
Melgar
Anti-Spam Agent
Posts: 2004
Registered: 23-2-2010
Location: Connecticut
Member Is Offline
Mood: Estrified
|
|
Quote: Originally posted by zts16  | | Ohhh, ok. I haven't had my tea yet this morning. Yes, the original grey-black sediment does sound like CuS. The picture that j_sum posted does not
really look like an iodine/KI solution to me. I would expect a dark solution with no precipitate in that case, but I suppose if it was dilute enough a
very fine suspension of I2 in a solution with no excess KI present could appear orange like that. |
Why do you say "no excess KI"? Excess KI would make the I2 water-soluble, and in fact the liquid above the precipitate does look a bit yellow. The
precipitate looks like it's probably dilute Lugol's solution mixed with a white substance that'd mostly consist of zinc compounds.
edit: The iodide may have reduced sulfate, precipitating both itself as I2 and zinc as ZnS. But there already would have been quite a bit of ZnS from
the zinc reduction of sulfuric acid. The original light yellow precipitate sounds like ZnS and possibly elemental sulfur. Add HCl and I bet you'd
smell hydrogen sulfide, in any case.
[Edited on 10/25/17 by Melgar]
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
|
|
|
Texium
Administrator
Posts: 4580
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
I say no excess KI because if there was excess KI the I2 would be water soluble, exactly like you said, and that would mean a much darker solution.
It's pretty clear from looking at the color of the solution vs that of the precipitate that the precipitate itself has color to it and it is not just
the pale yellow solution over a white precipitate. The yellow color of the solution is not unlike what you'd see in a saturated iodine/water solution
without KI present to increase its solubility.
To be honest though, I don't think that this sort of speculation is very useful without further experimentation being done to narrow the
possibilities.
[Edited on 10-25-2017 by zts16]
|
|
|
Melgar
Anti-Spam Agent
Posts: 2004
Registered: 23-2-2010
Location: Connecticut
Member Is Offline
Mood: Estrified
|
|
Well, he did rule out copper (and many other metals) via flame spectroscopy. Although he claims he only ruled out copper II, copper I isn't stable
enough to not be oxidized in a flame, and if it was present, it would have been visible as such. Other than contamination, the only colored element
is iodine, and iodine will give that color very easily if a white substance is present. In any case, if KI is added to sulfuric acid, there would
definitely be a change in color. There's probably very little KI though, due to the acidity.
[Edited on 10/25/17 by Melgar]
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
|
|
|
j_sum1
Administrator
Posts: 6320
Registered: 4-10-2014
Location: At home
Member Is Offline
Mood: Most of the ducks are in a row
|
|
Good thought on the CuS coming out as a precipitate. I have it sitting in some filter paper at present and can do some testing.
I have hit it with some HCl and did not observe H2S coming off. (I am sure I would have noticed.)
It gave only faint hints of blue in a flame and that could easily have been some particles of copper present -- after all, that was the aim of the
exercise.
Should I expect CuS to liberate H2S with concentrated HCl?
Should I expect CuS to decompose in a flame and give it some colour?
The orange precipitate is definitely orange. It formed slowly in the test tube over the course of two minutes and the particles gently rained down.
It was quite pretty really.
I am still quizzed by this one.
|
|
|
Texium
Administrator
Posts: 4580
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
| Quote: Originally posted by j_sum1 | I have hit it with some HCl and did not observe H2S coming off. (I am sure I would have noticed.)
It gave only faint hints of blue in a flame and that could easily have been some particles of copper present -- after all, that was the aim of the
exercise.
Should I expect CuS to liberate H2S with concentrated HCl?
Should I expect CuS to decompose in a flame and give it some colour? |
I've only made CuS once before (for the
Copper Carnival) and I found it to be pretty unreactive. IIRC it refused to react with HCl and what I did for the next reaction was dissolve it in hot
nitric acid, which it still only did reluctantly.
|
|
|
ninhydric1
Hazard to Others
Posts: 345
Registered: 21-4-2017
Location: Western US
Member Is Offline
Mood: Bleached
|
|
The only thing that comes to my mind about the orange precipitate is a cadmium salt, a dichromate precipitate, or V2O5. Maybe vanadium contamination?
The philosophy of one century is the common sense of the next.
|
|
|
j_sum1
Administrator
Posts: 6320
Registered: 4-10-2014
Location: At home
Member Is Offline
Mood: Most of the ducks are in a row
|
|
| Quote: Originally posted by zts16 | | Quote: Originally posted by j_sum1 | I have hit it with some HCl and did not observe H2S coming off. (I am sure I would have noticed.)
It gave only faint hints of blue in a flame and that could easily have been some particles of copper present -- after all, that was the aim of the
exercise.
Should I expect CuS to liberate H2S with concentrated HCl?
Should I expect CuS to decompose in a flame and give it some colour? |
I've only made CuS once before (for the
Copper Carnival) and I found it to be pretty unreactive. IIRC it refused to react with HCl and what I did for the next reaction was dissolve it in hot
nitric acid, which it still only did reluctantly. |
Ok. That is consistent with what I observed about the solid material. I will drop it in some nitric to confirm.
That leaves the yellow solution and the orange precipitate (which presumably is an iodide salt of some kind). We have acidic reducing conditions.
The only known species to enter the mix are zinc, copper, sulfate ions, potassium ions and iodide ions. no obvious contaminants that would produce
these observations spring to mind.
|
|
|
j_sum1
Administrator
Posts: 6320
Registered: 4-10-2014
Location: At home
Member Is Offline
Mood: Most of the ducks are in a row
|
|
Thanks to all. Here's an update.
Today I threw out my lesson plan for my chem students, put this thread up on the screen and discussed each contribution and the reasoning and chemical
theory behind each suggestion made.
We also did a few extra tests as they were suggested.
We are doing a unit of work on qualitative and quantitative testing and so this is completely on topic. It is refreshing to be doing something
"live", so to speak, and to tackle something that is a genuine mystery.
The sediment first:
This actually shows some reddish hues of elemental copper but is mostly a dark and still damp sludge.
1. Testing with azeotropic nitric acid:
Unmistakable thick red NO2 gas given off. Blue substance appeared in the bottom of the test tube. This is consistent with copper but there was also
some cloudiness to the mixture. None of the original material remained -- nothing settled in teh botom of the test tube.
2. Testing with dilute nitric acid:
No visible reaction even when heated to boiling.
3.Adding concentrated nitric acid dropwise to the test tube in (2):
Reaction occurred. Some NO2 given off. No evidence of H2S (unsurprising with an oxidising acid). Formation of a yellow suspension of sulfur in a
blue liquid.
From this we concluded that the material was a mixture of elemental copper and copper sulfide.
We also reasoned that there was unlikely to be any zinc remaining -- this would have been detected in earlier acid tests.
This lead us to conclude that the remaining yellow filtrate was largely an acidified and concentrated solution of zinc sulfate. I am not sure what
caused the yellow colour -- it might be worth running the iron tests again.
We ran out of time but were able to duplicate the orange precipitate with KI. We set some up for filtering. I will do a third one and video it i
think. It really is a nice precipitation.
We will also see if anything of the yellow solution precipitates when made alkaline -- with NaOH and also with ammonia. Going slowly of course to see
if anything redissolves. Not expecting much here except confirmation of the presence of zinc ions.
Any further suggestions?
My best guess for the precipitate is ummmm I keep drawing blanks here. Zinc triiodide does not fit but I feel like we are getting closer.
Video coming.
|
|
|
Melgar
Anti-Spam Agent
Posts: 2004
Registered: 23-2-2010
Location: Connecticut
Member Is Offline
Mood: Estrified
|
|
The orange color, I don't think is a salt, I think it's elemental iodine that's precipitating out alongside what I assume is white or yellow zinc
sulfide. You say you have acidic reducing conditions, but depending on the substance, it could either be oxidized (by sulfate or sulfuric acid) or
reduced (by zinc). Iodide cannot be reduced, but can be oxidized very easily, especially by sulfuric acid. It's the reason that you can't distill
alkyl iodides from sulfuric acid, KI, and an alcohol; you'd just have all the iodide oxidized to iodine. Also, as sulfate is reduced, it forms SO2 or
bisulfite, which can be either an oxidizing or reducing agent, sort of like -NO2.
You definitely have zinc SULFIDE (as well as sulfate) in there. Zinc sulfide is much less soluble in water, but can probably be oxidized to zinc
sulfate, which is water-soluble. It is probably forming as the sulfate oxidizes iodide to iodine, then falling out of solution and taking the iodine
with it. Thus, you have a homogeneous mixture of I2 and ZnS. I meant earlier that if you add HCl to the orange/white precipitate, you should smell
H2S.
ZnS + 2HCl = ZnCl2 + H2S
I'm like 99% sure that works with zinc.
If you dry out the precipitate and warm it up, elemental iodine would be driven off, and if I'm correct, it'd turn white.
Another experiment that I'm not sure would work, but would be really impressive if it does, is to look at the precipitate under a black light. Zinc
sulfide with a few ppm of copper is what's used to make glow-in-the-dark stuff, and if you happen to have this, it would glow blue-green. If there
happen to be some trace copper ions hanging around when the ZnS precipitates, it might just work. I'm pretty sure iodine would absorb UV, so you'd
have to heat the precipitate to drive off any iodine first.
[Edited on 10/26/17 by Melgar]
[Edited on 10/26/17 by Melgar]
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
|
|
|
Σldritch
Hazard to Others
Posts: 309
Registered: 22-3-2016
Member Is Offline
Mood: No Mood
|
|
I made a really wierd compund...
The compound is black, maybe it is crystalline, it is in the form of a powder and is filtered very easily.. It appears after a few minutes when a
mixture of ferrous sulfate, sodium dithionite and acetone is mixed. It does not appear when only sodium dithionite and ferrous sulfate is mixed and
seems to be proportional to the amount of acetone added to a mixed solution of ferrous sulfate and sodium dithionite.
I left a very small amount on a filter paper and it seemed to disapear into nothing. I have no idea what it could be but this might help: http://pubs.acs.org/doi/abs/10.1021/jo00339a057
Sorry i can not get any more details: i will be away from my lab a few days.
[Edited on 27-10-2017 by Σldritch]
|
|
|
Melgar
Anti-Spam Agent
Posts: 2004
Registered: 23-2-2010
Location: Connecticut
Member Is Offline
Mood: Estrified
|
|
| Quote: Originally posted by Σldritch | I made a really wierd compund...
The compound is black, maybe it is crystalline, it is in the form of a powder and is filtered very easily.. It appears after a few minutes when a
mixture of ferrous sulfate, sodium dithionite and acetone is mixed. It does not appear when only sodium dithionite and ferrous sulfate is mixed and
seems to be proportional to the amount of acetone added to a mixed solution of ferrous sulfate and sodium dithionite.
I left a very small amount on a filter paper and it seemed to disapear into nothing. I have no idea what it could be but this might help: http://pubs.acs.org/doi/abs/10.1021/jo00339a057
Sorry i can not get any more details: i will be away from my lab a few days.
[Edited on 27-10-2017 by Σldritch] |
I think I may have found a winner:
https://en.wikipedia.org/wiki/Iron(II)_sulfide
Metal sulfides seem to be a common thing to accidentally make, apparently.
[Edited on 10/28/17 by Melgar]
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
|
|
|
Σldritch
Hazard to Others
Posts: 309
Registered: 22-3-2016
Member Is Offline
Mood: No Mood
|
|
That kind of makes sense. I guess it disapearing from the filter paper can be explained by left over dithionite being oxidized by air to some sulfur
acid and dissolving the iron sulfide but why would it only form when acetone is added? Acetone should oxidize the dithionite to sodium bisulfite which
would dissolve any formed sulfides from the decomposition of dithionite. Why would acetone even affect it at all?
2 H2S2O4 → 3 SO2 + S + 2 H2O
3 H2S2O4 → 5 SO2 + H2S + 2 H2O
All that sulfur dioxide should dissolve any iron sulfide formed right? Still it seems most likely.
[Edited on 28-10-2017 by Σldritch]
|
|
|
Melgar
Anti-Spam Agent
Posts: 2004
Registered: 23-2-2010
Location: Connecticut
Member Is Offline
Mood: Estrified
|
|
| Quote: Originally posted by Σldritch | That kind of makes sense. I guess it disapearing from the filter paper can be explained by left over dithionite being oxidized by air to some sulfur
acid and dissolving the iron sulfide but why would it only form when acetone is added? Acetone should oxidize the dithionite to sodium bisulfite which
would dissolve any formed sulfides from the decomposition of dithionite. Why would acetone even affect it at all?
2 H2S2O4 → 3 SO2 + S + 2 H2O
3 H2S2O4 → 5 SO2 + H2S + 2 H2O
All that sulfur dioxide should dissolve any iron sulfide formed right? Still it seems most likely. |
Did you see the part on the wikipedia page where it says that iron sulfide will oxidize to iron (ii) sulfate (which is pale blue) in moist air? That
was the kicker for me.
Sodium dithionite is a strong reducing agent, and I'd guess it would be capable of reducing one or more of the products from the reaction with acetone
into H2S. Of course, that's just a guess, since I've never used sodium dithionite for anything. How hard is it to get?
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
|
|
|
| Pages:
1
..
3
4
5 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
