vano
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Indium sulfide
Hi. I made indium sulfide. I used indium metal nitric acid and sodium sulfide. It just has a nice color. It is not anhydrous. Anhydrous sulfide has a
reddish-brown color.
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ChemTalk
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That is a very nice compound, thanks for sharing. I may try it sometime, I have all of the necessary ingredients.
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vano
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Good luck ChemTalk.
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Bedlasky
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It's interesting that these soft Lewis acid/soft Lewis base salts have nice colours, even if they harder analogues are white. Al2S3, Ga2S3 are white,
In2S3 yellow and Tl2S black.
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vano
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Gallium sulfide is also yellow. I have quite a lot of gallium, but it is difficult to make sulfide because it is hydrolyzed in water. You are right
thallium sulfide is black, but I remember the crystals have a blue color. Most likely dark blue.
[Edited on 23-1-2021 by vano]
[Edited on 23-1-2021 by vano]
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Bedlasky
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Quote: Originally posted by vano  | Gallium sulfide is also yellow. I have quite a lot of gallium, but it is difficult to make sulfide because it is hydrolyzed in water. You are right
thallium sulfide is black, but I remember the crystals have a blue color. Most likely dark blue.
[Edited on 23-1-2021 by vano]
[Edited on 23-1-2021 by vano] |
I always think that Ga2S3 is white. But according to wiki, Ga2S3 can be yellow or white, depending of process of preparation:
https://en.wikipedia.org/wiki/Gallium(III)_sulfide
[Edited on 23-1-2021 by Bedlasky]
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vano
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No problem 
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Lion850
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Vano your photo inspired me to also give this a go. Details below.
Indium nitrate solution:
First step was to get indium metal in solution, by reacting indium with nitric acid.
I expected the reaction to be:
2In + 6HNO3 = 2In(NO3)3 + 3H2
and accordingly 5.5g indium in 4 lumps was reacted with 11g 70% HNO3. The reaction started quickly with bubbling and release of brown NO2 fumes, see
below.
After some time the reaction slowed down, and the beaker had to be heathed to speed it up. I also added another 6.5g 70% HNO3 as more acid was
required when taking into account this reaction:
In + 4HNO3 = In(NO3)3 + NO + 2H2O
I left the reaction to slowly proceed, and went back every 3 or 4 hours to again heat the solution to near boiling and then let it cool slowly with
the hotplate. The solution always went green to some extend which I first thought was due to impurities in the indium but I realised later the
dissolved NOx caused the green color.
By the next morning all the indium was dissolved and I had a completely clear water colored solution of some 30ml. I then added water to increase the
total volume to 150ml, and slowly boiled this down to around 30ml again. Purpose being to try and expell and acid or NOx that may still have been
present. During the final phase of boiling there was no acid odor.
Next step was to bubble H2S through the indium nitrate. First step was to make iron sulphide. 42g iron filings was mixed with 28g sulphur and poured
into a test tube:
The test tube was then heated in a bunsen burner. The reaction took quite a bit of heating to start; more then I remember from previous occasions.
Maybe the iron filings was old and covered with oxidation? Anyway the expected red glow eventually started:
Once cool, the test tube was broken and a lump of 63g was recovered.
This was very hard but brittle, and was broken into smalled lumps using a hammer and chisel. Some 10g of FeS was placed in an erlenmeyer flask.
Approximately 30ml hydrochloric acid (pool acid, around 320g/L) was diluted with an equal amount of water and added to the flask. Prior to this the
solution of indium nitrate was placed in a large testube and some water added. I had never reacted H2S in solution before, and thought the longer time
the bubbles will remain in the solution the better - thus using a tall test tube with a bit of water added. The simple setup is shown below:
The diluted HCl was added to the flask and the stopper applied. Bubbles was observed in the indium nitrate solution at some 1-2 bubbles per second.
Intially nothing seemed to happen; probably it took some time to displace the air in the flask. Then a yellow suspension started to appear. Although
some portion of the bubbles reached the surface there was no strong H2S smell at the test tube; probably some hydrogen was generated with the H2S?
After 45 minutes there was thick yellow suspension and then the smell of H2S exiting the test tube became noticeable. At this point I removed the tube
from the test tube and at the same time dumped an excess of copper sulphate solution into the Erlenmeyer flask to neutralise any H2S as it was being
generated. The test tube with yellow product:
I left the test tube overnight for the yellow ppt to settle. The next morning it was still yellow, but a few hours later it changed to orange!
It was filtered off (and rinsed in the filter) and dried in the sun (under a steel dish) for a few hours, recovery was 3.5g of dry orange powder. This
was less than 50% of the theoretical, based on having started with 5.5g indium. Thus, there had to be a significant amount of indium nitrate left in
the filtrate. Photo of first crop below:
4.5g of sodium sulphide Na2S was dissolved in 30ml water and added to the filtrate. This caused another crop of yellow ppt to form. This was filtered
off, rinsed with water in the filter, and placed in the sun under a steel dish to dry. As it dried this second crop also became orange! The second
crop was 4.1g:
The first crop eventually became as orange as the 2nd crop. Total recovery was this 7.6g, close to the theoretical yield. But why was the reaction in
aquas solution in the test tube so inefficient? Was the 45 minutes simply too short? Was the indium nitrate solution too dilute, or perhaps too
concentrated? Should I have bubbled the H2S into a beaker with magnetic stirring to get better 'contact'?
Atomistry mentions yellow and red as color for In2S3 depending on method of preparation; looks like to got somewhere inbetween with my orange?
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vano
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Thanks for sharing. Very nice. your sulfide is hydrated. I wanted it to be anhydrous. You know waterless red - brownish color, I think it has a more
beautiful color, although when I dried the sulfide it was red, but when I made the powder it still remained yellow, but it looked well dried.
[Edited on 31-1-2021 by vano]
[Edited on 31-1-2021 by vano]
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vano
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It is interesting that it produces thiosalts.
In2S3 + 3Na2S = 2Na3[InS3]
In2S3 + Na2S = 2Na[InS2] (at 120°C)
As a rule, hydrogen sulfide should also form similar compounds, although I do not know how stable it will be. I do not know the stability of these
sodium salts, but if it did not decomposed I would personally make salts of copper, nickel and similar metals, I think they would have very nice
colors, the second does not contain too much sulfide in anion, so I think it will not have black or any other dark color.
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Bedlasky
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Lion850: When nitric acid react with some metal, it's quite rare that hydrogen is evolved (and you need for this quite dilute acid). Most times you
obtain mix of NO and NO2 (but you can get other products depending on concentration of acid, look at this).
So when your In react with azeotropic nitric acid, this equation is closest to reality:
In + 6 HNO3 = In(NO3)3 + 3 NO2 + 3 H2O
Vano: Only cations from 2B class (Sn2+;4+, As3+; 5+, Sb3+; 5+) form soluble thiosalts in aqueous solution. Also some
anions form soluble thiosalts (molybdate, tungstate, rhenate, vanadate). HgS and CuS are partially soluble just in concentrated suflides/polysulfides.
Freshly precipitated CdS is partially soluble in (NH4)2S. Other thiosalts are prepared in solid phase.
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Bezaleel
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| Quote: Originally posted by Lion850 | (snip)
The first crop eventually became as orange as the 2nd crop. Total recovery was this 7.6g, close to the theoretical yield. But why was the reaction in
aquas solution in the test tube so inefficient? Was the 45 minutes simply too short? Was the indium nitrate solution too dilute, or perhaps too
concentrated? Should I have bubbled the H2S into a beaker with magnetic stirring to get better 'contact'?
Atomistry mentions yellow and red as color for In2S3 depending on method of preparation; looks like to got somewhere inbetween with my orange?
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I think that initially 3 H2S + 2 In(NO3) --> In2S3 + 6 HNO3, so the solution turns acid.
Later, HNO3 + In2S3 <--> H2S + In(NO3)3.
I noticed the same behaviour when I tried to scavenge H2S with a solution of CuSO4. This gave some black precipitate but long before the solution had
lost its blue colour, the presence of H2S escaping from it became noticeable. That's why I prefer to use an oxidiser to convert H2S to sulphur.
When you used Na2S, you did not produce any H+, so you could precipitate your In2S3.
Thanks for the extensive write-up, btw. Makes a great read!
| Quote: Originally posted by Bedlasky | (snip)
Vano: Only cations from 2B class (Sn2+;4+, As3+; 5+, Sb3+; 5+) form soluble thiosalts in aqueous solution. Also some
anions form soluble thiosalts (molybdate, tungstate, rhenate, vanadate). HgS and CuS are partially soluble just in concentrated suflides/polysulfides.
Freshly precipitated CdS is partially soluble in (NH4)2S. Other thiosalts are prepared in solid phase. |
So the great orange flakes that Lion850 produced are hydrates of In2S3, and no sodium salts thereof?
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Bedlasky
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Bazaleel: I think that yes, it is hydrous In2S3. Colour may differ depending on reaction conditions and water content (look at CdS).
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woelen
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I also think that they are hydrous In2S3, but there may be some sodium sticking to it as well. That's the reason why precipitating sulfides with Na2S
is not the preferred method of preparation. You get the sulfide, but some sodium ions (and corresponding thio anions) may precipitate as well and
these may be very hard to remove from the solid material you get.
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vano
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Of course both salts are hydrates.
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