| Pages:
1
2 |
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
Sodium sulphide
Is there a simple way to make small quantities of sodium sulphide, not via hydrogen sulphide. I can buy kilogram quantities on line but I only need a
few grams.
|
|
|
entropy51
Gone, but not forgotten
Posts: 1612
Registered: 30-5-2009
Member Is Offline
Mood: Fissile
|
|
Quote: Originally posted by nezza  | | Is there a simple way to make small quantities of sodium sulphide, not via hydrogen sulphide. I can buy kilogram quantities on line but I only need a
few grams. |
There have been threads on this in the past. The methods that do not use hydrogen sulfide are
not simple, they are worse than using H2S for the most part.
You can buy it online from photo suppliers. $5 for 100 grams.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is an interesting path (most likely untried). First, an interesting reaction, heat S and Al (avoid moisture):
2 Al + 3 S --> Al2S3
Depending on how fine the Al, this could be very exothermic (think flash powder).
Add an excess of NaOH and heat again:
6 NaOH + Al2S3 --> 3 Na2S + 2 Al(OH)3
More interesting do it all at once by constructing consecutive layers of Aluminum and Sulfur mix, pure sulfur and then NaOH. Apply heat or ignite.
The reason for the pure S layer is as buffer between Al and the NaOH, which is hygroscopic.
[Edited on 11-3-2012 by AJKOER]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Do you really believe this will give a clean product? NaOH and S also react with each other, without the Al. This combination gives a very complicated
mix of sulfide, thiosulfate and polysulfide (the potassium compound made in this way is called liver of sulphur, which is very far away from pure
sulfide).
Making pure Al2S3 seems achievable to me, but it will not be easy. Expect a lot of unreacted aluminium and sulphur and also expect polysulfides in the
mix. This kind of solid-solid reactions hardly ever give pure products and quantitatively well defined reactions. The main reaction is formation of
Al2S3, but there are side reactions.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Yes, I agree, certainly not a pure synthesis, but an OK excuse to react Al and S and amazingly easier (fun) way than the standard preparation for the
anhydrous salt, albeit with an Al2O3 impurity. For example, per Wikipedia:
"Production
Industrially Na2S is produced by reduction of Na2SO4 with carbon, in the form of coal:[3]
Na2SO4 + 4 C → Na2S + 4 CO
In the laboratory, the anhydrous salt can be prepared by reduction of sulfur with sodium in anhydrous ammonia. Alternatively, sulfur can be reduced by
sodium in dry THF with a catalytic amount of naphthalene:[4]
2 Na + S → Na2S"
To my prior Aluminum sulfide suggested route, here is a modification: Add strong NH4OH to any of the Al2S3 that may have formed along with any
possible un-reacted Sulfur and Aluminum. Filter out (unfortunately, the bad smell of H2S from the Ammonium Sulfide has not been avoided). Expected
reaction:
Al2S3 + 2 NH3 + 4 H2O --> (NH4)2S + 2 Al(OH)3
where water is consumed and one can use the flocculation property of Aluminum hydroxide to help clean the solution from particles.
Now add the NaOH (in excess) and mildly heat or, possibly better, add some alcohol in which Na2S is only slightly soluble, to move the reaction to the
right. Then, cool to isolate Na2S.xH20, which is the commercially available hydrate. Reaction:
(NH4)2S + 2 NaOH --> Na2S + 2 NH3 (g) + H2O
A source notes the following unwanted reaction:
Na2S + 2 H2O <----> 2 NaOH + H2S
The result, if successful, is hopefully a somewhat cleaner Na2S product in hydrated form.
[Edited on 13-3-2012 by AJKOER]
|
|
|
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
Thanks for the answers. I have found a supplier I can get 100g off so that will do. One more question - How stable is it in aqueous solution ?.
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
The sulphide anion hydrolizes and forms hydrosulfuric acid which is very weak, so most of the H<sub>2</sub>S molecules exist as hydrogen
sulphide which slowly evaporates from the solution. Generally speaking, it takes over 15 minutes for the stench to fully develop after the solid has
been dissolved in water.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Endimion17 | | The sulphide anion hydrolizes and forms hydrosulfuric acid which is very weak, so most of the H<sub>2</sub>S molecules exist as hydrogen
sulphide which slowly evaporates from the solution. Generally speaking, it takes over 15 minutes for the stench to fully develop after the solid has
been dissolved in water. |
Per my cited equation:
Na2S + 2 H2O <----> 2 NaOH + H2S
In an open vessel, the formation of H2S gas means that the equilibrium is moving to the right. But, in a closed container with a little added NaOH,
the smell should be less noticeable upon opening assuming the vessel is not heated. Also, upon cooling, the formation of the hydrate moving the
reaction to the left.
[Edited on 16-3-2012 by AJKOER]
|
|
|
BauArf56
Hazard to Self
Posts: 68
Registered: 22-8-2019
Location: between the moon and the sun
Member Is Offline
Mood: energetic
|
|
i know thread is old, but heating sulfur with an excess of sodium carbonate gives sodium sulfide and carbon dioxide (if an excess of sulfur, it will
react to form polysulfides).
|
|
|
Bedlasky
International Hazard
Posts: 1241
Registered: 15-4-2019
Location: Period 5, group 6
Member Is Offline
Mood: Volatile
|
|
There will be thiosulfate as impurity.
|
|
|
unionised
International Hazard
Posts: 5128
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Heating sulphur with sodium carbonate will give a mixture of sulphide and thiosulphate.
Which is probably still better than heating a mixture of aluminium, sulphur and sodium hydroxide which will give you a shower of molten alkali (with
some sulphide, for added toxicity)
|
|
|
RU_KLO
Hazard to Others
Posts: 226
Registered: 12-10-2022
Location: Argentina
Member Is Offline
|
|
I need some sulphide for qualitative chemistry (sulphide precipitation). Never used this procedure.
I do not own a fumehood, and know that H2S is toxic. This is why I try to avoid it.
I read that Na2S could be used instead. After reading this post and others from wiki (SM), tried this process:
An arbitrary ammount of Sulfur (powered) and Sodium carbonate was heated in a test tube (-TT) (this was done outside). The TT was heated with an
alcohol burner till the sulfur melted. It became a black liquid (some sulfur deposited in the neck of the TT)
After 5 minutes or so, the TT was removed from heat. Cooled (it solidifies in a red-orange rock) and breaked. The red-orange rock was crushed.
I read that this is a mixture of sulphides, polysulphides, thiosulfates, and probably sulfur and sodium carbonate.
Cold (1-2° H2O was poured. the idea was that Na2S is less soluble in cold H2O than thiosulfate.
but everything was dissolved. Giving a orange solution.
A few drops of this solution was dropped in a solution of Lead nitrate. Black/Brown precipitate immediately appeared. So probably some sulphide is in
there.
Questions:
1) Is there a way to purify this solution?
2) Could be used as it is for qualitative chemistry? (or the thio - poly sulfides/sulfates will interfere?)
3) Currently is stored in a pyrex capped bottle (80 ml), and as I do not own a fumehood, how dangerous is to work with this solution inside my house?
The idea is to use 3-5 drops each time.
4) Read that Na2S hydrolyses, this means that as time pases, the solution will be useless?
Thanks,
[Edited on 2-4-2023 by RU_KLO]
Go SAFE, because stupidity and bad Luck exist.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
1) I see no easy way to do that. Soooo many impurities!
2) I would not use it for any analytical purposes. As I said, soooo many impurities!
3) Yes, you can safely use this inside the house. Things may be somewhat smelly, but using a few drops of solution will not lead to dangerous amounts
of H2S in the air. Even on addition of acid to these few drops you will only get a bad smell, but no dangerous concentration of H2S.
4) It hydrolyses, but that is an almost immediate reaction. As soon as you add Na2S to water, the sulfide is almost completely hydrolysed to SH(-)
ions: S(2-) + H2O --> SH(-) + OH(-). This reaction is nearly instantaneous, and nearly 100%. The sulfide, however, still remains sulfide in this
reaction. If you have a sulfide solution and no air can get in, then the sulfide will remain good (as SH(-) ions). If air is allowed to reach the
solution, then it quickly degrades. Sulfide is oxidized to sulfur (and OH(-) is formed in the process as well), and this reacts with more sulfide to
yellow polysulfides. If a lot of air is allowed to come in contact with the solution, it will become turbid and impure sulfur will precipitate. You
also will get all kinds of oxo-compounds of sulfur in solution, it will become a very complicated (and smelly) mix.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Consider making FeS instead. Treated with strong acid this then generates the desired H2S.
Simply combine Fe and S and heat to get the party started:
Fe(s) + S(l) --> FeS(s) (+ much heat)
The FeS obtained is relatively pure.
[Edited on 4-4-2023 by blogfast25]
|
|
|
Cathoderay
Hazard to Self
Posts: 57
Registered: 29-1-2023
Location: US-Texas
Member Is Offline
|
|
Here is a link to Doug's Lab making iron sulfide.
https://www.youtube.com/watch?v=CLBdVM9_FWg
For the small amount you may need there could be another way.
I haven't tried it but I do think that if you mix paraffin wax and sulfur and heat it in a test tube it will produce H2S gas.
You can fit a one hole stopper on the test tube with a bent glass tube that is inserted into the solution you are adding the H2S to.
Some paraffin wax might evaporate and condense in cooler areas and possibly SO2 might be generated but those should not interfere with the process.
|
|
|
Admagistr
Hazard to Others
Posts: 372
Registered: 4-11-2021
Location: Central Europe
Member Is Offline
Mood: The dreaming alchemist
|
|
For the small amount you may need there could be another way.
I haven't tried it but I do think that if you mix paraffin wax and sulfur and heat it in a test tube it will produce H2S gas.
You can fit a one hole stopper on the test tube with a bent glass tube that is inserted into the solution you are adding the H2S to.
Some paraffin wax might evaporate and condense in cooler areas and possibly SO2 might be generated but those should not interfere with the process.
[/rquote]
The reaction of paraffin with sulphur to produce H2S is a known reaction and it works, often with the addition of Al2O3 as a catalyst, the reaction
has the advantage that the development of H2S stops immediately after stopping the heating.
[Edited on 4-4-2023 by Admagistr]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Cathoderay |
I haven't tried it but I do think that if you mix paraffin wax and sulfur and heat it in a test tube it will produce H2S gas.
|
Well, live and learn: never previously heard of that one...
[Edited on 4-4-2023 by blogfast25]
|
|
|
Admagistr
Hazard to Others
Posts: 372
Registered: 4-11-2021
Location: Central Europe
Member Is Offline
Mood: The dreaming alchemist
|
|
The reaction is often reported in the literature and works,but I can't remember the citations at the moment.I read it in about five books on
preparative inorganics,it's many years ago...
|
|
|
B(a)P
International Hazard
Posts: 1139
Registered: 29-9-2019
Member Is Offline
Mood: Festive
|
|
| Quote: Originally posted by Admagistr |
The reaction is often reported in the literature and works,but I can't remember the citations at the moment.I read it in about five books on
preparative inorganics,it's many years ago... |
Here is a good demo by Tdep https://www.youtube.com/watch?v=1hfC_ggJEhs
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Boy, was that longwinded or was it longwinded?  Not to mention a
superfluousness of glass porn...
You could also use scrap PE bags as a hydrocarbon, I guess.
[Edited on 4-4-2023 by blogfast25]
|
|
|
Cathoderay
Hazard to Self
Posts: 57
Registered: 29-1-2023
Location: US-Texas
Member Is Offline
|
|
I think he got a little too ambitious and made too much and got it too hot. I had seen a analysis video somewhere that showed a dark substance in a
test tube heated with a Bunsen burner, the test tube having a stopper and bent glass tubing. The substance was not identified. I think the apparatus
was kept as is, you don't reuse the test tube. It is meant for qualitative analysis of solutions in test tubes not for synthesis.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Cathoderay |
I think he got a little too ambitious and made too much and got it too hot. I had seen a analysis video somewhere that showed a dark substance in a
test tube heated with a Bunsen burner, the test tube having a stopper and bent glass tubing. The substance was not identified. I think the apparatus
was kept as is, you don't reuse the test tube. It is meant for qualitative analysis of solutions in test tubes not for synthesis.
|
I'll defo try this one day but much simplified. Considering the chemical inertia of paraffins (olefins, see etymology!) I'm surprised this reaction
proceeds at all!
The FeS route has the benefit of leaving you with a bit of a portable H2S source, to be deployed when needed.
|
|
|
teodor
National Hazard
Posts: 922
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by RU_KLO | I need some sulphide for qualitative chemistry (sulphide precipitation). Never used this procedure.
I do not own a fumehood, and know that H2S is toxic. This is why I try to avoid it.
I read that Na2S could be used instead. |
Indeed, there are alternatives for H2S and depending on what you try to analyze different reagents could be used: thiosulfate, thioacetate,
thioacetamide, impure Na2S or even S- free reagents like solvents or chelating agents.
So, main question is: which ions you try to separate and whether you need qualitative or quantitative results.
Making home chemistry last few years I learned that there is an alternative to almost every reaction but there is no alternative to a fume hood. Start
to make it step by step, starting from the suction part.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I tried the reaction of paraffins (from candles) with sulfur. It indeed works, and you can make a fairly constant stream of H2S. The biggest
disadvantage is that there are impurities in the H2S, which are annoying when you want it to make a solution of sulfide (e.g. by bubbling it through
water). I obtained a yellow liquid, because vapor of sulfur goes with the H2S, and although most of it will condense, some sulfur makes it into the
H2S, going into the solution. The reaction also produces some annoying hydrocarbons, which may make it into the solution.
So, I prefer the FeS-route. FeS, made at home, also is quite impure in my experience, but it can be used for making nice pure alkaline solutions of
sulfide. If you mix the impure FeS with acid, then you get H2S and H2 (from unreacted iron), but that H2 does not contaminate your target solution.
The FeS also may contain unreacted sulfur, but at the low temperatures in the aqueous acid, this does not escape and move with the H2S. You just don't
get a clear solution in the flask to which the FeS is added, but the gas mix is clean. The only problem with production of H2S from FeS and dilute
acid, is that droplets of solution can be taken with the H2S, but that problem can be reduced a lot, by using a long tube, and by using a tube, filled
with very loosely packed wadding or glass wool. Any droplets you have remain behind.
The H2S can be bubbled through an aqueous solution of NaOH, which leads to an alkaline solution of a mix of NaHS and NaOH. But a solution of Na2S also
is a mix of these two.
|
|
|
teodor
National Hazard
Posts: 922
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by woelen |
The H2S can be bubbled through an aqueous solution of NaOH, which leads to an alkaline solution of a mix of NaHS and NaOH. But a solution of Na2S also
is a mix of these two. |
The classical method of preparing ammonium sulfide as an analytical reagent includes saturating ammonium solution with H2S and adding ammonium
solution equal to initial volume after that. Based on this I suspect when H2S is absorbed the final result could be mostly HS-.
Update: the complexity of using Na2S of (NH4)2S as an analitical reagents is that it is possible only in alkaline to netral condition, but H2S
systematic analysis is started with acidic solutions.
[Edited on 5-4-2023 by teodor]
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
