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blogfast25
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[*] posted on 28-10-2015 at 19:09


A selenate is by far the best option here. Se is easily reduced to Se(0).



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[*] posted on 28-10-2015 at 19:21


Quote: Originally posted by blogfast25  
A selenate is by far the best option here. Se is easily reduced to Se(0).


Indeed, though there is always the risk of hydrogen selenide, which is apparently even more toxic than phosgene (exposure limit is 50 ppb over an 8 hour period!) It seems quite similar to arsenous acid reduction but with more wiggle room before the undesirable gas is generated. As long as I use the same procedures used for the arsenous acid reduction then it should be safe.
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[*] posted on 28-10-2015 at 19:58


There's not really a risk of hydrogen selenide if the reduction is done properly. I've done this experiment before for a selenium sample for my element collection. It was a ~5g scale; the source of selenium was an old bottle of selenium dioxide that a lab was throwing out and was improperly stored so the compound had absorbed enough water from the air that it completely dissolved in it and thus I couldn't get perfect stoichiometry. To this I added an excess of a solution of sodium bisulfite, with strong stirring. It's quite an interesting reaction, no reaction is immediately evident, but after several minutes the solution starts to redden in color and after stirring for several hours the reaction is complete.

This yields the red allotrope of selenium which is initially quite a brilliant red color, similar to that of cadmium selenide. Over time however, the sample slowly begins to darken. I prepared my sample about 6 months ago and ampouled it immediately after and the sample is currently a reddish brown color instead of the initial brilliant red; it currently looks quite similar to red phosphorus actually.




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[*] posted on 28-10-2015 at 23:55


If sodium selenate can be purchased for $25 and your only goal is making selenium, then I would suggest simply buying selenium. Selenium is not that expensive and you can buy nice samples on eBay for just a few bucks.

If your goal is to make every element yourself, then of course selenate (or selenite) is one possible way to go. But I would say, keep the selenate. That compound is rare and not easy to obtain, while elemental selenium is quite common. I would not use a rare chemical for just making a common chemical.

The risk of producing H2Se from selenates or selenites is nearly zero. I have done the reduction myself quite a few times and I even electrolysed solutions, containing selenites, but never obtained any H2Se.





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[*] posted on 29-10-2015 at 06:07


Quote: Originally posted by woelen  

The risk of producing H2Se from selenates or selenites is nearly zero. I have done the reduction myself quite a few times and I even electrolysed solutions, containing selenites, but never obtained any H2Se.



Is this just due to elemental selenium being a sluggish reactant, much like its cousin sulfur? H2Se production seems feasible based on SRPs but I imagine it is probably similar to why sulfur won't react easily with even strong reductants unless molten.
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[*] posted on 29-10-2015 at 06:54


Quote: Originally posted by Upsilon  
Quote: Originally posted by woelen  

The risk of producing H2Se from selenates or selenites is nearly zero. I have done the reduction myself quite a few times and I even electrolysed solutions, containing selenites, but never obtained any H2Se.



Is this just due to elemental selenium being a sluggish reactant, much like its cousin sulfur? H2Se production seems feasible based on SRPs but I imagine it is probably similar to why sulfur won't react easily with even strong reductants unless molten.


woelen is right.

To obtain H2Se in significant quantities you probably need to start off from selenides (plus acid).




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[*] posted on 29-10-2015 at 16:35


Would it be safe to use oxalic acid as a reductant, then? Since sulfur takes much coaxing to be reduced by oxalic acid, I would assume selenium is very similar?
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[*] posted on 29-10-2015 at 16:45


Quote: Originally posted by Upsilon  
Would it be safe to use oxalic acid as a reductant, then? Since sulfur takes much coaxing to be reduced by oxalic acid, I would assume selenium is very similar?


My broad guess would be 'No'. Oxalic acid isn't much of a reducer, more something that can be oxidised by powerful oxidisers. But I could be wrong on that. Testing is everything... very easy to test in a test tube.

Reduction of selenites (Se(+4)) with SO<sub>2</sub> works. Probably acidified sulphites would too.

[Edited on 30-10-2015 by blogfast25]




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[*] posted on 29-10-2015 at 17:38


SRPs say the reaction is favorable, but then again it says reduction of sulfate by oxalic acid is also favorable. Still, reduction of sulfate is only 0.66V while reduction of selenate is 1.64V. Maybe this is favorable enough that the reaction could yet proceed (by comparison, reduction of MnO2 by OA is only 0.07V greater, which we know for sure works). Definitely something to be testing.
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[*] posted on 29-10-2015 at 17:56


Quote: Originally posted by blogfast25  
Quote: Originally posted by Upsilon  
Would it be safe to use oxalic acid as a reductant, then? Since sulfur takes much coaxing to be reduced by oxalic acid, I would assume selenium is very similar?

Reduction of selenites (Se(+3)) with SO<sub>2</sub> works. Probably acidified sulphites would too.

See a couple posts above. I reduced selenium dioxide(selenous acid) with sodium bisulfite, the free acid provides all the acidity needed. If using sodium selenite, acidifying the solution is likely necessary.




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[*] posted on 29-10-2015 at 18:02


Quote: Originally posted by Upsilon  
SRPs say the reaction is favorable, but then again it says reduction of sulfate by oxalic acid is also favorable.


SRPs say lots of things: thermodynamics is not kinetics.




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[*] posted on 29-10-2015 at 23:59


@blogfast25: Selenite has selenium in oxidation state +4, not +3.

@Upsilon: Oxalic acid probably will do the job of reducing selenites, but only slowly. I also noticed that reduction of selenite is much faster/easier in the presence of chloride ion. I tried reducing a solution of Na2SeO3 in 10% H2SO4 with SO2 and this is rather slow. When a little HCl or NaCl is added, then the reaction becomes much faster, you see the liquid turn orange and then turbid and brick-red in just a few tens of seconds. So, maybe a combination of dilute HCl and oxalic acid?

Also, keep in mind that selenate is not the same as selenite. I am not sure how easily selenate is reduced. The highest oxidation state of many non-metal elements is much less reactive than lower oxidation states. This is most natable with sulphur, phosphorus and chlorine, but it may also be so with selenium. E.g. sulfite can easily be reduced to sulphur or sulfide in acidic solution, while sulfate is very hard to reduce under similar conditions, although it has sulphur in higher oxidation state.




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[*] posted on 30-10-2015 at 05:02


Quote: Originally posted by woelen  

Also, keep in mind that selenate is not the same as selenite. I am not sure how easily selenate is reduced. The highest oxidation state of many non-metal elements is much less reactive than lower oxidation states. This is most natable with sulphur, phosphorus and chlorine, but it may also be so with selenium. E.g. sulfite can easily be reduced to sulphur or sulfide in acidic solution, while sulfate is very hard to reduce under similar conditions, although it has sulphur in higher oxidation state.


Well, that is reflected by SRPs. Reduction of sulfate to sulfurous acid has a potential of a meager 0.17V, while sulfurous acid to thiosulfate or elemental S have potentials of 0.4V and 0.45V respectively - so it makes sense that sulfates are harder to reduce than sulfites. Selenates and selenites, on the other hand, exhibit a reverse relationship - reduction of selenate to selenous acid has a potential of 1.15V while selenous acid to selenium has a potential of 0.74V. So based on this selenate is much easier to reduce than selenite.
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[*] posted on 30-10-2015 at 07:58


Quote: Originally posted by woelen  
@blogfast25: Selenite has selenium in oxidation state +4, not +3.




Oooopsie. Yes. (Blushes badly). :mad:




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[*] posted on 31-10-2015 at 08:41


I am going to try ampouling some chlorine gas at some point not to far off. I have a good idea of what I need to do, but I want to know what the most effective method for neutralizing excess chlorine is. I am thinking of using a sodium hydroxide solution; does anyone have a better idea?
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[*] posted on 2-11-2015 at 18:51


Well if nobody has a better idea then I'll just use the NaOH to deal with the excess chlorine gas.

On that note, I received my potassium bromide today; I will be attempting isolation of bromine from it at some point. I was originally planning on using chlorine gas to oxidize the bromide in the potassium bromide solution, but I'm wondering if I can get a cleaner product using chloric acid as an oxidant instead. It would be nice to not need to deal with a chlorine gas generator, and would also prevent BrCl contamination.
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[*] posted on 2-11-2015 at 19:01


Chlorine:
I bubble through a bucket of NaOH. It seems to work ok.

Bromine:
My last synth was using MnO2. It took a while to clean the RBF and yield was low. I have used other oxidants as well. H2O2 as usual is the cleanest.
My next attempt will be using Cl2. NileRed I think has a YT vid where he makes Cl2 in situ. This has the advantage of high yield and no significant Cl2 wastage. The other method is to bubble Cl2 through the bromide solution. That is likely the method that I will use.
I can't point to links ATM but from reading here and a few videos, I don't believe BrCl contamination is a significant issue. 2BrCl --> Br2 + Cl2 is favoured. woelen will be along some time soon to correct me.
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[*] posted on 2-11-2015 at 19:13


Quote: Originally posted by Upsilon  
Well if nobody has a better idea then I'll just use the NaOH to deal with the excess chlorine gas.

On that note, I received my potassium bromide today; I will be attempting isolation of bromine from it at some point. I was originally planning on using chlorine gas to oxidize the bromide in the potassium bromide solution, but I'm wondering if I can get a cleaner product using chloric acid as an oxidant instead. It would be nice to not need to deal with a chlorine gas generator, and would also prevent BrCl contamination.


NaOH as a chlorine scavenger is fine.

But the displacement of bromide with chlorine leads inevitably to small amounts of BrCl, which is hard to distinguish from Br<sub>2</sub>.

There are plenty threads on 'clean' Br<sub>2</sub> generation on this site.




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[*] posted on 2-11-2015 at 19:17


Ok, blogfast corrected me. How small is small? Does a redistillation fix the issue?

For an element collection, I probably don't care too much anyway. Although I will probably take steps to remove water.
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[*] posted on 2-11-2015 at 19:20


Quote:
How small is small?

Think about 20 mL... from 100g of starting NaBr. After drying. That's what I got with woelen's method, anyway, and by the time I got it all ampouled (under sulfuric acid, to keep things dry), it had dwindled to just over 2 mL. It's a very perfidious element.




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[*] posted on 2-11-2015 at 19:22


On a related note, does anyone have experience in ampouling liquid chlorine? It would seem to me that dry ice and acetone would get to the right temperature range. It would also seem that torching glass in the presence of a flammable liquid is fraught with potential hazards. Also, the thermal stresses on the ampoule might be significant.

It is the sort of task that one would want to get right first time. I don't fancy the combination of flame, flammable liquid, molten glass, broken glass and free and rapidly-dispersing chlorine.
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[*] posted on 2-11-2015 at 19:25


Quote:
about 20 mL... from 100g of starting NaBr

Are you sure?
That would seem to me to be a contamination of nearly 50% (not knowing off the top of my head the density of a BrCl / Br2 mixture).
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[*] posted on 2-11-2015 at 19:28


Quote: Originally posted by j_sum1  
Ok, blogfast corrected me. How small is small? Does a redistillation fix the issue?

For an element collection, I probably don't care too much anyway. Although I will probably take steps to remove water.


Hard to tell. A tall distillation column would separate it but how tall is a tall distillation column? ;)




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[*] posted on 2-11-2015 at 20:01


Quote: Originally posted by j_sum1  
H2O2 as usual is the cleanest.


Wouldn't using H2O2 not be any better than Cl2 gas, since the O2 evolved could form bromine oxides? Even more so than chlorine would form BrCl since O2 is a stronger oxidant than Cl2?

And looks like the chloric acid method wouldn't be very good after all since reduction of chloric acid would likely produce chlorine gas and form BrCl anyway.
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[*] posted on 2-11-2015 at 20:27


H2O2 is the cleanest physically of all the methods I have attempted thus far. Cleanup is swish and rinse. I don't know about Br2O. Most bromine I have produced has been for demonstration purposes. "Look! Oxidation happened!"
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