Sciencemadness Discussion Board

Quest for the elements

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Upsilon - 28-9-2015 at 16:31

I'm soon to embark on a journey to collect as many of the natural elements as I can by isolating them from compounds containing them. Obviously some will need to be purchased, like the noble gases, but I want to produce as many of them myself as possible. Sure I could just buy samples, but I would find no enjoyment in that.

Needless to say some elements will be more difficult to isolate than others. I will need advice from you all on occasion when I come across a difficult problem. Instead of making a new thread every time I run into an obstacle, I'll just let this thread serve as a sort of journal for my experiences in trying to isolate elements, which I will also use to outsource advice from the community.

And to start it off: I need help with beryllium. Its compounds aren't particularly useful in chemistry so they are virtually impossible to find. However, it does find use as a component in some alloys. It is also found in some fairly common minerals, such as beryl and bertrandite.

That being said I see two viable options. I can try and refine it out of a mineral (research suggests treating crushed bertrandite with sulfuric acid leaves a water-soluble sulfate, but sulfate of what?), or I can try to separate it from an alloy. I have noticed that beryllium-copper alloys are used in certain golf club heads, which aren't terribly expensive. Perhaps treatment with some HCl would tear out the beryllium and leave the copper behind?

elementcollector1 - 28-9-2015 at 16:42

I believe plante1999 posted our only thread on extracting beryllium compounds, and he used crushed raw beryl.
http://www.sciencemadness.org/talk/viewthread.php?tid=23007
Beryllium-copper is expensive and has low beryllium content - best to go with the mineral.

Remember, Be is fairly toxic!

[Edited on 9-29-2015 by elementcollector1]

Velzee - 28-9-2015 at 16:44

I'm no expert, but since beryllium has a slight higher melting point than copper, could you just heat the club up to melt the copper out, leaving Be behind?

Little_Ghost_again - 28-9-2015 at 16:51

I think you can get it from microwave ovens! be careful though as breathing the dust is bad news. Its normally like a ceramic disk in the magnatron, so find an old microwave and bingo



be.jpg - 130kB
found this on google I think its the pink bit :D

[Edited on 29-9-2015 by Little_Ghost_again]

Upsilon - 28-9-2015 at 17:09

Quote: Originally posted by elementcollector1  
I believe plante1999 posted our only thread on extracting beryllium compounds, and he used crushed raw beryl.
http://www.sciencemadness.org/talk/viewthread.php?tid=23007
Beryllium-copper is expensive and has low beryllium content - best to go with the mineral.

Remember, Be is fairly toxic!

[Edited on 9-29-2015 by elementcollector1]


Will definitely be looking into that. Seems like a mammoth task but a fun one as well

Velzee, alloys have one specific melting point. Take solder for electrical circuits, for example. It is composed of tin and lead, leading to a material that melts at a significantly lower temperature than either of the constituent metals.

j_sum1 - 28-9-2015 at 17:34

If you came across an (expensive) alloy of Cu/Be (Used for non-sparking tools, at 4%Be) you could have a go at isolating it from that.
For me, it was an element I bought for a couple of bucks.

blogfast25 - 28-9-2015 at 18:06

Obtaining Be compounds (typically BeSO<sub>4</sub> if you're taking an aqueous route) from Beryl or some other source is only a small part of the problem. Isolating the metal, if that's what you desire is far far harder and possibly very hazardous as a process.
Quote: Originally posted by Little_Ghost_again  
I think you can get it from microwave ovens! be careful though as breathing the dust is bad news. Its normally like a ceramic disk in the magnatron, so find an old microwave and bingo




found this on google I think its the pink bit :D

[Edited on 29-9-2015 by Little_Ghost_again]


Please post the link.

Oh, this one?

https://www.youtube.com/watch?v=4Zev8Ws4P1Y

He couldn't tell Be from his *rse. Low level info that should be ignored.


[Edited on 29-9-2015 by blogfast25]

Upsilon - 28-9-2015 at 19:09

Quote: Originally posted by blogfast25  
Obtaining Be compounds (typically BeSO<sub>4</sub> if you're taking an aqueous route) from Beryl or some other source is only a small part of the problem. Isolating the metal, if that's what you desire is far far harder and possibly very hazardous as a process.

[Edited on 29-9-2015 by blogfast25]


Yes, I would need to learn what to watch out for with isolating it. I'm still quite far off from doing any of this yet, but when I get to it, it'll almost certainly be by electrolysis (either of a molten beryllium salt in an inert atmosphere, or in propylene carbonate if I can ever get that to work).

j_sum1 - 28-9-2015 at 22:47

You could akways practise isolating Mg from epsom salt and when you had perfected that, move onto Be.

Little_Ghost_again - 29-9-2015 at 00:28

All I did was post a pic of the part needed, why is that low level information to be ignored?

[Edited on 29-9-2015 by Little_Ghost_again]

phlogiston - 29-9-2015 at 07:05

BeO ceramics can also be found in high power (RF) transistors. It has excellent thermal conductivity.

Upsilon - 29-9-2015 at 07:59

I found this on eBay:
http://ebay.com/itm/WHOLESALE-2lb-Lot-GREEN-BERYL-AQUAMARINE...

Seems like a really good deal; for this purpose it would be better to buy rough, ugly samples like this instead of the really good-looking crystals.

Magpie - 29-9-2015 at 08:55

I would be really careful handling Be. If Be powder is inhaled it can cause beryllium lung disease if you are susceptible. I have no Be or Be compounds in my lab and don't ever plan to.

Be was once used for nuclear fuel fabrication. Many of the fabrication workers now have this lung disease. Many lawsuits and much money is now involved.

blogfast25 - 29-9-2015 at 09:11

Quote: Originally posted by Little_Ghost_again  
All I did was post a pic of the part needed, why is that low level information to be ignored?



Because the video maker only states that it's Be ceramic but provides nothing to back that up with. Anyone can make such claims.

blogfast25 - 29-9-2015 at 09:16

Quote: Originally posted by Magpie  
I would be really careful handling Be. If Be powder is inhaled it can cause beryllium lung disease if you are susceptible.


Unless you're working with Be (or compounds) based powders, ordinary precautions (don't inhale or ingest) should suffice. We're not talking about mass fabrication here.

Originally the element was called 'glucinium' because its 'salts' taste sweet.

[Edited on 29-9-2015 by blogfast25]

blogfast25 - 29-9-2015 at 09:21

Quote: Originally posted by Upsilon  
I found this on eBay:
http://ebay.com/itm/WHOLESALE-2lb-Lot-GREEN-BERYL-AQUAMARINE...

Seems like a really good deal; for this purpose it would be better to buy rough, ugly samples like this instead of the really good-looking crystals.


Really good looking crystals of this stuff are gems or at least semi-gems and would be seriously expensive. https://en.wikipedia.org/wiki/Emerald

But even for these 'ore grade' pieces I would seek reassurance (like a certificate) from the seller. Then, good luck processing it: it's incredibly hard stuff, see also the link higher up. http://www.sciencemadness.org/talk/viewthread.php?tid=23007&...

[Edited on 29-9-2015 by blogfast25]

[Edited on 29-9-2015 by blogfast25]

Upsilon - 29-9-2015 at 11:37

Quote: Originally posted by blogfast25  
Quote: Originally posted by Upsilon  
I found this on eBay:
http://ebay.com/itm/WHOLESALE-2lb-Lot-GREEN-BERYL-AQUAMARINE...

Seems like a really good deal; for this purpose it would be better to buy rough, ugly samples like this instead of the really good-looking crystals.


Really good looking crystals of this stuff are gems or at least semi-gems and would be seriously expensive. https://en.wikipedia.org/wiki/Emerald

But even for these 'ore grade' pieces I would seek reassurance (like a certificate) from the buyer. Then, good luck processing it: it's incredibly hard stuff, see also the link higher up. http://www.sciencemadness.org/talk/viewthread.php?tid=23007&...

[Edited on 29-9-2015 by blogfast25]


Thanks for the info. I'm going to be making my own little arc furnace at some point; hopefully I can actually get that hot enough to melt the beryl. Would definitely make it much easier to break apart. If not, a good old sledgehammer ought to do the trick.

blogfast25 - 29-9-2015 at 12:58

Quote: Originally posted by Upsilon  


Thanks for the info. I'm going to be making my own little arc furnace at some point; hopefully I can actually get that hot enough to melt the beryl. Would definitely make it much easier to break apart. If not, a good old sledgehammer ought to do the trick.


An arc furnace would be a good tool here: heat your pieces until they melt (or nearly melt), then quench them in cold water. The vitrification process changes the crystalline structure dramatically and the material then becomes much, much easier to grind. Grind under water to avoid dust.

As regards the sledgehammer, that's probably not a great idea unless you have an extremely strong surface you'll simply mostly be hammering beryl pieces into that surface.

aga - 29-9-2015 at 13:17

Reading your link to the Be experiments in the penultimate post reminds me of why i first came here and stayed.

Science, Chemistry, Facts, Experiments, Results !

There are even chemical equations in there to accompany the photos.

Time to turn SM around and re-take that Fascinating high ground and get back to some actual interesting Science instead of all these random chat-room style random Lizard Alien type comments that appear all the time.

blogfast25 - 29-9-2015 at 14:15

I've long contended that in recent years all that is good about SM, evidence based amateur science in a nutshell, is in decline. Although it's subjective and impossible to prove, that's what my perception tells me.

A turnaround would be great but I'm not holding my breath.

aga: if you're interested in experimentation with Beryl with the aim of preparing some BeO or BeSO<sub>4</sub>, let me know.

[Edited on 29-9-2015 by blogfast25]

Upsilon - 1-10-2015 at 12:34

I will set aside the beryl experiment for later when I have better equipment, then. Next I would like to discuss phosphorus. My only phosphorus-containing compound is trisodium phosphate. Given the ability to make phosphoric acid from it, I have found the following half-reactions in the table of reduction potentials:
H3PO4(aq) + 2H+ + 2e- -> H3PO3(aq) + H2O (-0.276 V)
H3PO3(aq) + 3H+ + 3e- -> P(red) + 3H2O (-0.454 V)
H3PO3(aq) + 2H+ + 2e- -> H3PO2(aq) + H2O (-0.499 V)
H3PO2(aq) + H+ + e- -> P(white) + 2H2O (-0.508 V)
P(red) + 3H+ + 3e- -> PH3(g) (-0.111 V)
P(white) + 3H+ + 3e- -> PH3(g) (-0.063 V)

This suggests that these reactions occur in the presence of an acidic reducing agent, such as oxalic acid:
2CO2(g) + 2H+ + 2e- -> H2C2O4 (-0.430 V)

Based on this information, the first reaction is thermodynamically favorable:
H3PO4(aq) + H2C2O4(aq) -> H3PO3(aq) + 2CO2(g) + H2O

But after this, not even the reaction involving red phosphorus is thermodynamically favorable, let alone the one involving hypophosphorous acid (which is an intermediate to white phosphorus).

I'm not going to be picky about which type of phosphorus is in my collection (though it would be cool to have both), but I won't get either of them at 1M concentration at 25C. The question then becomes, how can I alter the reaction conditions to increase the favorability of, to keep things simple for now, the phosphorous acid to red phosphorus reaction, without going too far causing the formation of phosphine gas? I am aware of phosphine's toxicity and would prepare for its formation regardless, but its formation signals the loss of elemental phosphorus, which I don't want to happen.

softbeard - 1-10-2015 at 14:18

I hate to be a spoilsport, but ASAIK there is no practical way to reduce the phosphate anion to solid phosphorus, or even phosphite, in aqueous solution.
H3PO4(aq) + 2H+ + 2e- -> H3PO3(aq) + H2O (-0.276 V) vs
2CO2(g) + 2H+ + 2e- -> H2C2O4 (-0.430 V) reduction potentials may imply
H3PO4(aq) + H2C2O4(aq) -> H3PO3(aq) + 2CO2(g) + H2O
but in reality it will never work.
Time for the experts in aqueous phosphorus chemistry to weigh in.

blogfast25 - 1-10-2015 at 14:51

Quote: Originally posted by softbeard  
I hate to be a spoilsport, but ASAIK there is no practical way to reduce the phosphate anion to solid phosphorus, or even phosphite, in aqueous solution.
H3PO4(aq) + 2H+ + 2e- -> H3PO3(aq) + H2O (-0.276 V) vs
2CO2(g) + 2H+ + 2e- -> H2C2O4 (-0.430 V) reduction potentials may imply
H3PO4(aq) + H2C2O4(aq) -> H3PO3(aq) + 2CO2(g) + H2O
but in reality it will never work.
Time for the experts in aqueous phosphorus chemistry to weigh in.


I don't know whether the above scheme would work (I'm very sceptical about it) but one should understand that thermodynamics says NOTHING about kinetics. A thermodynamically favourable scheme does not mean that it will proceed at any appreciable rate.

Do you see a lot of coal (carbon) burst spontaneously into flames? If not, ask yourself why! :D:D (Hint: C + O2 ===> CO2 is highly thermodynamically favourable).

[Edited on 1-10-2015 by blogfast25]

Upsilon - 1-10-2015 at 15:12

Quote: Originally posted by blogfast25  


I don't know whether the above scheme would work (I'm very sceptical about it) but one should understand that thermodynamics says NOTHING about kinetics. A thermodynamically favourable scheme does not mean that it will proceed at any appreciable rate.

Do you see a lot of coal (carbon) burst spontaneously into flames? If not, ask yourself why! :D:D (Hint: C + O2 ===> CO2 is highly thermodynamically favourable).

[Edited on 1-10-2015 by blogfast25]


I feel like this is true mainly for these half-reactions involving solids and gases (like your example of C + O2). Like in my H2S experiment in the other thread, the reaction is thermodynamically favorable, but both substances had to be melted to overcome the kinetic barriers (btw, I tried that experiment again and posted some goodies there). Perhaps with aqueous solutions, these barriers are overcome by the act of ionizing the substances? I don't know for sure; perhaps I will conduct this experiment this weekend or sometime during next week.

[Edited on 1-10-2015 by Upsilon]

j_sum1 - 1-10-2015 at 15:58

Well upsilon.
If it is phosphorus you are after then you can always scrape it off the side of matchboxes.
If it is phosphorus chemistry you are interested in then that is also not a bad start point. You can sublime to convert from red to white and this also acts to purify it. There are a few solvents you can play with also. There are a few threads around on oxygen and chlorine compounds of phosphorus.

If (as I suspect) it is the synthesis you want for your element collection and you want a reasonable sized sample, then this thread will be too small to contain the issues. Fortunately, there is a lengthy thread stickied on the production of white phos. There is lots to read there but it is almost guaranteed to answer any question you might have. The alchemists boiled down urine but there are other methods. SM member rogeryermaw is probably the person who had the most success. He also has a youtube clip showing his methodology. I like it. It is OTC (or almost). But there are significant hazards and phosphorus doesn't play fair. Be real careful and read up first.

Actually, let me emphasise that point. Do some reading. For most of the threads you have started, there are already significant threads discussing the question. There is a wealth of information on these boards and it is good to respect the expertise that has already been freely offered. There is also an advantage in having all the related information together in one place.

You and I have similar goals with regards to our element collections -- with the exception that you seem to want to tackle the difficult ones first. I think it is wiser to work up to the difficult/hazardous.
Here are some projects that I have either done or are on my short list -- all inspired by what I have read on SM

At present I have around 45 elements and probably almost 100 different samples. Phosphorus, neodymium and lanthanum are on my list too but they can wait until I upskill and accumulate the necessary equipment. Anyway, we should compare notes. U2U me if you want.

Upsilon - 1-10-2015 at 16:08

Thanks for the information, I'll do some looking around. Also, I can see why you would think that I'm trying to tackle the difficult elements first, but that's not what I'm doing here. Right now I am just in the thinking stage; I am planning out how to isolate each one. The easy ones I don't really need clarification on. Essentially I am asking about these ahead of time; I am not necessarily trying to get them in the order I'm posting them.

Little_Ghost_again - 1-10-2015 at 17:18

Quote: Originally posted by aga  
Reading your link to the Be experiments in the penultimate post reminds me of why i first came here and stayed.

Science, Chemistry, Facts, Experiments, Results !

There are even chemical equations in there to accompany the photos.

Time to turn SM around and re-take that Fascinating high ground and get back to some actual interesting Science instead of all these random chat-room style random Lizard Alien type comments that appear all the time.



It wasnt long ago you were a newb, this is the newb section! People have to start somewhere like you did, if there is a lack of science going on then maybe its time attitudes changed to those that are new and of little knowledge.
How many times have people complete debunked things in a high and mighty manner only to be proved wrong later on. I posted the picture because its well known in recycling that you have to be careful of the magnetron and the Be bit on it.
The intention was merely to show a potential source and not provide in depth details of what it was or how to process it. Ok aga you do experiments but before moaning about real science take a look in in pre pub, those are the people who actually do real science and there arnt many names in that place.
If you dont have a section like beginners then you will never grow a community, 99% of people who post in this section wont get passed it, but the 1% that do will end up doing real science.
The danger is discouraging people from posting crap, once you start doing that you are going to miss out on the 1%.
Also some of us do things that fail and dont like posting about it, I asked in a thread on the iodine about getting back some of the acid with vacuum distillation, I thought a quick simple answer would give me a direction.
I didnt think it worth bothering searching for ages and as no one posted anything back I assumed it was ok.
Turned out far from ok.
So yes now I will research first but I am less likely to ask questions in the first place now.

Upsilon - 1-10-2015 at 17:57

I'm going to try the phosphate reduction reaction for fun. However I think it would be reasonably difficult to determine the presence of H3PO3 in a mixture of oxalic and phosphoric acid. Instead I think I'm going to construct an electrochemical cell; that way I will easily be able to tell if a reaction is occurring by measuring the voltage of the cell (theoretically it will be 0.154 V). I don't have much knowledge beyond your basic, easy-to-understand cells, but I'm going to take a shot in the dark here. One of the half-cells will be a 1M phosphoric acid solution, the other a 1M oxalic acid solution. Graphite electrodes. A salt bridge would consist of an acid to allow migration of H+ ions. Am I on target here, or completely off?

woelen - 1-10-2015 at 23:12

Phosphoric acid cannot easily be reduced. Thermodynamically it may seem feasible, but there is no easy mechanism (pathway) from phosphoric acid to lower oxidation state compounds. Only very strong and very reactive reductors are capable of reducing phosphoric acid.

This property sometimes is used to an advantage as well. E.g. from phosphoric acid it is possible to make HBr from NaBr or HI from NaI or KI. With sulphuric acid this is impossible, the iodide and also the bromide are oxidized and you get all kinds of unclean side reactions.

Thermodynamic data only is one part of the equation, a viable pathway for the reaction is the other part. You can compare this to a heavy ball laying on the top of a hill, with the ball having a lot of potential energy, ready to be used to convert to kinetic energy if the ball moves down the hill and gains speed. If the hill has a rounded top and has a smooth surface, then a small push of the ball suffices to set it in motion and then it goes down with ever increasing speed. Now imagine that at the top of the hill (which still has the same height), the ball has a dike around it. The ball still has the same potential energy, but now you have to put more effort in using that potential energy. You first have to push it over the dike. An occasional push does not lead to the ball going down the hill. You first have to add some energy, before you are rewarded with much more kinetic energy. Now you can compare phosphoric acid in aqueous solution to a ball, which is on a hill, but with very high walls around it. The walls may be even higher than the height of the hill, so a LOT of effort must be put into the system before its energy can be released.

Conclusion: A cell, based on phosphoric acid will not work and oxalic acid will not reduce phosphoric acid. Oxalic acid simply does not push hard enough to get the ball over the dike.



[Edited on 2-10-15 by woelen]

Upsilon - 2-10-2015 at 05:06

So I guess that leaves the question, where did they come up with those reduction potentials? Is it based on a prediction, or did they actually make it work somehow?

blogfast25 - 2-10-2015 at 06:33

Quote: Originally posted by Upsilon  
So I guess that leaves the question, where did they come up with those reduction potentials? Is it based on a prediction, or did they actually make it work somehow?


Once you have a set of redox reactions that involve the H<sup>+</sup>/H<sub>2</sub> half-cell (+0 V, by convention) that kinetically work, you can start combining hitherto unknown reductions (or oxidations, that is only a matter of sign!) with known ones in cells that actually work. The reduction potential is always relative to the H<sup>+</sup>/H<sub>2</sub> half-cell.

Today we have large databases of the Gibbs Free Energies of Formation of aqueous species. Combining these allows, with Nernst's theory, to actually calculate standard reduction/oxidation potentials for 'new' systems.

Upsilon - 2-10-2015 at 15:53

Makes sense. Looks like there's still no better route to phosphorus that doesn't involve high heat and/or toxic gases.

I would also like to note that I just received a golfball sized orpiment (arsenic trisulfide) in the mail today. However I'm not going to try anything with it for a while yet. I'll post again when I'm ready to have a go at separating out the arsenic.

This weekend I'm going to try isolating titanium from titanium dioxide with the information in this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=21332#...


[Edited on 3-10-2015 by Upsilon]

blogfast25 - 2-10-2015 at 16:04

Quote: Originally posted by Upsilon  

This weekend I'm going to try isolating titanium from titanium dioxide with the information in this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=21332#...
[/url]


It'd be nice to see yet someone else prepare some rough Ti metal with my method. :cool:

elementcollector1 - 2-10-2015 at 16:33

Quote: Originally posted by Upsilon  
Makes sense. Looks like there's still no better route to phosphorus that doesn't involve high heat and/or toxic gases.

I would also like to note that I just received a golfball sized orpiment (arsenic trisulfide) in the mail today. However I'm not going to try anything with it for a while yet. I'll post again when I'm ready to have a go at separating out the arsenic.

This weekend I'm going to try isolating titanium from titanium dioxide with the information in this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=21332#...


[Edited on 3-10-2015 by Upsilon]


Out of curiosity, where on Earth did you find such a lump of orpiment? Was it cheap?

Upsilon - 2-10-2015 at 16:46

Quote: Originally posted by elementcollector1  

Out of curiosity, where on Earth did you find such a lump of orpiment? Was it cheap?


Oh absolutely, they're all over eBay. Mine specifically is 91g, and it cost me $10 plus $3 for shipping.

Upsilon - 3-10-2015 at 18:31

Was busy today so I didn't get around to the titanium thermite. However I did start the process of making the calcium sulfate for the reaction. I mixed solutions of calcium nitrate and sodium carbonate to precipitate calcium carbonate (I'll also be keeping the sodium nitrate). I'll then react the calcium carbonate with HCl to make calcium chloride, which I'll then mix with ammonium sulfate to make calcium sulfate. Sure I could just mix the calcium nitrate and ammonium sulfate to begin with but I don't really want to deal with ammonium nitrate right now.

j_sum1 - 4-10-2015 at 00:57

Sounds convoluted, but should work. Make sure your CaSO4 is anhydrous before you attempt the thermite.
Do you have any CaF2 for a flux?

Upsilon - 4-10-2015 at 11:39

I do. However it is taking a long time to process all of this calcium carbonate; I probably won't get to the actual thermite today.

For the time being, I've been speculating about easy ways to create lithium metal. I discovered that lithium nitrate has an extremely low melting temperature, and that it could easily be melted on a hotplate. Is electrolysis of this salt viable? The temperatures are so relatively low that it could even be done under mineral oil. I have lithium carbonate on hand and it would be quite easy to make the nitrate out of it.

blogfast25 - 4-10-2015 at 14:15

Quote: Originally posted by Upsilon  
I do. However it is taking a long time to process all of this calcium carbonate; I probably won't get to the actual thermite today.

For the time being, I've been speculating about easy ways to create lithium metal. I discovered that lithium nitrate has an extremely low melting temperature, and that it could easily be melted on a hotplate. Is electrolysis of this salt viable? The temperatures are so relatively low that it could even be done under mineral oil. I have lithium carbonate on hand and it would be quite easy to make the nitrate out of it.


CaSO4? Just get the cheapest, no-frills wall filler (that's the hemi-hydate) and dehydrate it. That's what I used.

I see you're getting NaNO3 as a by-product? Well, that works very well as a heat booster for the TiO2 thermite too! Better even than CaSO4.

Electrolysis of LiNO3? I believe someone reported on that here. Mine was ridiculously deliquescent but maybe that was just that grade, not sure...



[Edited on 4-10-2015 by blogfast25]

Upsilon - 4-10-2015 at 15:16

Quote: Originally posted by blogfast25  


CaSO4? Just get the cheapest, no-frills wall filler (that's the hemi-hydate) and dehydrate it. That's what I used.

I see you're getting NaNO3 as a by-product? Well, that works very well as a heat booster for the TiO2 thermite too! Better even than CaSO4.

Electrolysis of LiNO3? I believe someone reported on that here. Mine was ridiculously deliquescent but maybe that was just that grade, not sure...



[Edited on 4-10-2015 by blogfast25]



I may just use the sodium nitrate then, less steps until I'm ready to actually start. I also wanted to make some calcium chloride so I figured I might as well go through with this process.

As for the lithium nitrate, I may do it this week sometime. I could make it by compining it with the calcium nitrate I have but lithium carbonate isn't very soluble, so that may not be the best route. Otherwise I'll have to wait on my cheapie distillation glassware to come in from China so I can get started making nitric acid.

blogfast25 - 4-10-2015 at 16:07

Quote: Originally posted by Upsilon  

I may just use the sodium nitrate then, less steps until I'm ready to actually start. I also wanted to make some calcium chloride so I figured I might as well go through with this process.



In that case, you need to adjust the overall mix a little. Let me know if you need help with that.

Upsilon - 4-10-2015 at 16:21

Quote: Originally posted by blogfast25  

In that case, you need to adjust the overall mix a little. Let me know if you need help with that.


I probably will. But first, what role exactly does the nitrate play? Is it just used for its strong oxidizing properties, creating a hot-burning side reaction that helps the main thermite keep going? I have seen people add sulfur to their thermite mixture to set up such a "helper" reaction with the aluminum - if the nitrate is used for the same thing, then it should work better than sulfur since it's a stronger oxidizer, right?

blogfast25 - 4-10-2015 at 17:14

Quote: Originally posted by Upsilon  

I probably will. But first, what role exactly does the nitrate play? Is it just used for its strong oxidizing properties, creating a hot-burning side reaction that helps the main thermite keep going? I have seen people add sulfur to their thermite mixture to set up such a "helper" reaction with the aluminum - if the nitrate is used for the same thing, then it should work better than sulfur since it's a stronger oxidizer, right?


The main reaction:

TiO2 + 4/3 Al === > Ti + 2/3 Al2O3

... produces little heat, only about 200 kJ/mol of Ti. This is far short of what's needed to obtain the reaction products in molten form and allow good separation between metal and slag.

To boost heat and reach the needed end-temperature a heat boosting system is used: KClO3, KNO3, NaNO3 and CaSO4 are all usable. They rely on the reaction (here for KClO3):

KClO3 + 2 Al === > KCl + Al2O3 which is VERY exothermic. With KClO3 + Al, you need about 1 to 1.2 mol KClO3 per mol of TiO2 to reach between 2600 to 3000 C end-temperature.

I'll post my NaNO3 based formulation tomorrow (I need to look it up).

Sulphur is really only used for the SiO2 thermites, it doesn't work for Ti.

Further reading (my essay on this method):

http://developing-your-web-presence.blogspot.co.uk/2008/10/o...


[Edited on 5-10-2015 by blogfast25]

Upsilon - 5-10-2015 at 09:10

Quote: Originally posted by blogfast25  


The main reaction:

TiO2 + 4/3 Al === > Ti + 2/3 Al2O3

... produces little heat, only about 200 kJ/mol of Ti. This is far short of what's needed to obtain the reaction products in molten form and allow good separation between metal and slag.

To boost heat and reach the needed end-temperature a heat boosting system is used: KClO3, KNO3, NaNO3 and CaSO4 are all usable. They rely on the reaction (here for KClO3):

KClO3 + 2 Al === > KCl + Al2O3 which is VERY exothermic. With KClO3 + Al, you need about 1 to 1.2 mol KClO3 per mol of TiO2 to reach between 2600 to 3000 C end-temperature.

I'll post my NaNO3 based formulation tomorrow (I need to look it up).

Sulphur is really only used for the SiO2 thermites, it doesn't work for Ti.

Further reading (my essay on this method):

http://developing-your-web-presence.blogspot.co.uk/2008/10/o...


[Edited on 5-10-2015 by blogfast25]


Thanks. With regards to the lithium nitrate electrolysis, I think I'm going to run into some density issues trying to do it under mineral oil. First I don't know the density of molten LiNO3; however its density in solid form is 2.38 g/cm3, so I don't imagine its density would drop below the 0.8 g/cm3 of mineral oil. The molten lithium will certainly float, though. I would have to develop some upside down trap to collect it, but it would be hard to keep it under mineral oil after collecting. Perhaps petroleum jelly would work until I can get it sealed under an inert gas.

elementcollector1 - 5-10-2015 at 09:37

Quote: Originally posted by Upsilon  

Thanks. With regards to the lithium nitrate electrolysis, I think I'm going to run into some density issues trying to do it under mineral oil. First I don't know the density of molten LiNO3; however its density in solid form is 2.38 g/cm3, so I don't imagine its density would drop below the 0.8 g/cm3 of mineral oil. The molten lithium will certainly float, though. I would have to develop some upside down trap to collect it, but it would be hard to keep it under mineral oil after collecting. Perhaps petroleum jelly would work until I can get it sealed under an inert gas.


Why not shape the cathode like an upside-down metal bowl, or some such? That way it would be produced and collected at the same time.

Upsilon - 5-10-2015 at 09:42

Quote: Originally posted by elementcollector1  

Why not shape the cathode like an upside-down metal bowl, or some such? That way it would be produced and collected at the same time.


That's what I'm thinking. I don't know what to do with it, though. I won't be able to seal anything under inert gas for a couple of months. There's really just no good way to display it until then that I can think of; smothering in petroleum jelly isn't very good, and if I put it in a vial full of mineral oil, it'll just float to the top and be hidden by the cap.

[Edited on 5-10-2015 by Upsilon]

blogfast25 - 5-10-2015 at 14:23

Quote: Originally posted by Upsilon  
Quote: Originally posted by elementcollector1  

Why not shape the cathode like an upside-down metal bowl, or some such? That way it would be produced and collected at the same time.


That's what I'm thinking. I don't know what to do with it, though. I won't be able to seal anything under inert gas for a couple of months. There's really just no good way to display it until then that I can think of; smothering in petroleum jelly isn't very good, and if I put it in a vial full of mineral oil, it'll just float to the top and be hidden by the cap.

[Edited on 5-10-2015 by Upsilon]


Have look at SM's library:

http://library.sciencemadness.org/library/index.html

I think Georg Brauer's 'Handbook of Preparative Inorganic Chemistry' has a procedure for elemental lithium.

A dome shaped cathode would only work if you can coat the outside with some insulator. Even the there would be issues with current pathways, I think.


[Edited on 5-10-2015 by blogfast25]

Upsilon - 5-10-2015 at 14:26

Good news; I'm getting a lot of extra work hours in the coming month, so I'll have more spending money to get materials sooner than later! Unfortunately this means that I probably won't get around to doing anything until this weekend.

I'm going to be doing a lot in the coming months. I am hopeful that I will find time to build an arc welder/furnace, an induction heater, and a vacuum/inert gas glovebox. I'll also be working on the element collection. If anyone has any ideas for useful projects under around $200-$300 USD, I'll take suggestions.

Upsilon - 6-10-2015 at 14:24

Quote: Originally posted by blogfast25  


Have look at SM's library:

http://library.sciencemadness.org/library/index.html

I think Georg Brauer's 'Handbook of Preparative Inorganic Chemistry' has a procedure for elemental lithium.

A dome shaped cathode would only work if you can coat the outside with some insulator. Even the there would be issues with current pathways, I think.


[Edited on 5-10-2015 by blogfast25]


Holy crap, I didn't even see your post until just now. Must have posted shortly before my last post. I'll definitely have to look into that; I'm having trouble thinking of a practical collection method.

Instead of using valuable nitric acid to make the lithium nitrate, I'll probably instead react the lithium carbonate with citric acid, then combine with calcium nitrate to precipitate out insoluble calcium citrate. The issue is that I don't know how soluble lithium citrate is, but it should at least be more soluble than the lithium carbonate itself.

EDIT: After looking around I have found this glass cloth electrical tape capable of acting as an electrical insulator as well as withstanding the high temperatures. If it works correctly, I should be able to run a wire (insulated with this tape) into an upside-down test tube with the mouth submerged into the molten LiNO3, filled the rest of the way with mineral oil. The wire would be connected to a short graphite electrode. Here is a poorly-drawn diagram of my thoughts. Dark blue is lithium nitrate solution, light blue is mineral oil, gray rods are graphite electrodes, and red wiring, is, well, wiring. The wire submerged in the molten lithium nitrate will obviously be insulated with the glass fabric electrical tape.

This is of course assuming the glass fabric tape will actually work for this. Since it is apparently a form of fabric insulation, I don't know how well it will keep a watertight seal. I am also slightly weary about using a borosilicate glass test tube for this, but it should be OK at 255C when lithium nitrate melts, though it will be cutting it close (I will be using a porcelain crucible for the actual lithium nitrate container).

[Edited on 7-10-2015 by Upsilon]

Upsilon - 7-10-2015 at 18:57

I had some time today so I started making the lithium citrate as a precursor to the lithium nitrate I need. I reacted lithium carbonate with citric acid and the reaction proceeded as expected, except that the aftermath solution is extremely syrupy. I'm guessing this is just undissolved lithium citrate?

Upsilon - 8-10-2015 at 15:07

Er, I just added the lithium citrate solution to the calcium nitrate solution, and I got absolutely no precipitate. What the heck happened? Is calcium citrate actually appreciably soluble unlike Wikipedia says?

blogfast25 - 8-10-2015 at 15:31

Quote: Originally posted by Upsilon  
Er, I just added the lithium citrate solution to the calcium nitrate solution, and I got absolutely no precipitate. What the heck happened? Is calcium citrate actually appreciably soluble unlike Wikipedia says?


It really is NOT a good idea to try and carry out these displacement reactions between compounds that have some solubility, like calcium citrate.

Citrates are the salts of a weak triprotic acid and their solubility is likely to be very pH dependent. Represent citric acid by H<sub>3</sub>Ct, so Ca citrate would be the Ca<sup>2+</sup> salt of Ct<sup>3-</sup>. As the conjugated base of a weak acid, Ct<sup>3-</sup> is a weak base:

Ct<sup>3-</sup>(aq) + H<sub>2</sub>O(l) < === > HCt<sup>2+(aq)</sup> + OH<sup>-</sup>(aq)

pH affects the form of the citrate ions present, thus also citrate solubility. Check the pH of your mixed solution.



[Edited on 8-10-2015 by blogfast25]

blogfast25 - 8-10-2015 at 15:42

A much simpler way is to take advantage of the modest solubility of LiCO3 and the high insolubility of CaCO3:

Ca(NO3)2(aq) + 2 LiCO3(s,aq) === > CaCO3(s) + 2 LiNO3(aq)

Simply simmer a slurry of LiCO3(s,aq) with a Ca(NO)2 solution (stoichiometric amounts) and the slurry will precipitate the Ca as carbonate, leaving the Li in solution as the highly soluble nitrate. Filter and recrystallise the LiNO3.

[Edited on 8-10-2015 by blogfast25]

Upsilon - 8-10-2015 at 15:45

Well that's an issue, I don't have any pH indicators. Even if I did, I don't know how I would adjust it without screwing with the contents of the final product. Luckily I did this in small scale to test it so it's not a big deal at all. I think I can make the direct reaction between lithium carbonate and calcium nitrate work.

EDIT: Just ordered a pH meter, don't know why I waited so long to get one. I'll save the solution so I can test it when it gets here, but I'm still going to use the other method to make lithium nitrate.

[Edited on 8-10-2015 by Upsilon]

Upsilon - 11-10-2015 at 16:31

Today I did a quick silicon dioxide thermite since I was frustrated with the fact that I have been making almost no progress on any of this. I found some extra time today, so I used the method entailed here:
http://thehomescientist.blogspot.com/2010/03/experiment-sili...

I didn't use very much (4.5g SiO2, 6g S, and 5g Al). I put it in a small aluminum foil boat and then put the whole thing into my makeshift thermite reactor made of bricks. It burned spectacularly and calmly for at least 30 seconds. I was left with a single mass of slag that I chiseled apart when it cooled. I put the pieces in water to get rid of the majority of the slag (Al2S3), and then moved them to conc. HCl for a brief time. I got some decent lumps which I was initially sure were silicon metal, but I have a hunch that they are just more pieces of slag. They don't really split apart like the other slag but they degrade easily - small bits are falling off even from being knocked around the vial I'm storing them in. The lumps are dark gray and quite rough, yet they shimmer in light. They certainly don't look at all like the image on that website. The material I saved weighs 0.35g. Not sure what to make of all this, anyone have any thoughts?

elementcollector1 - 11-10-2015 at 17:19

Sounds like alumina slag to me - are they a cohesive mass, or sort of sintered in appearance? At any rate, your reaction went forward - there's got to be some silicon in there...

Don't worry about losing your silicon - if any formed, it won't be too badly attacked by the HCl. You might get an unappealing gray powder, but it'll still be pure silicon.

I bought some fumed silica as part of a failed shear-thickening fluid experiment last month - I'm tempted to take it home and try out this reaction on it.

blogfast25 - 11-10-2015 at 17:38

@Upsilon:

These very small amounts don't work well. Minimum 100 g of mix (approx.). Small amounts of mix don't allow enough time for metal/slag separation and also lose far too much heat. 'Go large' with this one!

@EC1:

Si isn't attacked by HCl at all. It's not a metal (it's a metalloid) and doesn't have any 'loose' electrons in it's valence shell.

[Edited on 12-10-2015 by blogfast25]

Upsilon - 11-10-2015 at 17:48

Quote: Originally posted by blogfast25  
@Upsilon:

These very small amounts don't work well. Minimum 100 g of mix (approx.). Small amounts of mix don't allow enough time for metal/slag separation and also lose far too much heat. 'Go large' with this one!


I figured as much. I guess the small mixtures just don't have enough mass to retain heat long enough to allow the molten metal to join together? (EDIT: Duh, I didn't even see that this is exactly what yiu said) I did notice a couple of tiny (2mm diameter max) perfect silvery spheres - I suppose these are the nice pure silicon samples I'm after. It seems wasteful to use so much chemical to make a lot of silicon when I only need a gram or two at most, but if that's what must be done then so be it. At any rate, I guess all I managed to do was stink up the property with H2S.

I'll be redoing this experiment as well as the titanium dioxide thermite this weekend (first days off in weeks!). I may also get to the lithium nitrate electrolysis, the design of which I still need confirmation on (scroll up).

[Edited on 12-10-2015 by Upsilon]

blogfast25 - 11-10-2015 at 17:53

Quote: Originally posted by Upsilon  


I figured as much. I guess the small mixtures just don't have enough mass to retain heat long enough to allow the molten metal to join together? I did notice a couple of tiny (2mm diameter max) perfect silvery spheres - I suppose these are the nice pure silicon samples I'm after. It seems wasteful to use so much chemical to make a lot of silicon when I only need a gram or two at most, but if that's what must be done then so be it. At any rate, I guess all I managed to do was stink up the property with H2S.



H2S is the reason I stopped preparing Si this way. Try and use large scrubbers with thin bleach to capture it, it oxidises the H2S to sulphur very effectively.

Si thermites are quite hard, also because the density of molten Si is actually lower than the density of molten alumina!

Upsilon - 11-10-2015 at 18:09

Quote: Originally posted by blogfast25  

H2S is the reason I stopped preparing Si this way. Try and use large scrubbers with thin bleach to capture it, it oxidises the H2S to sulphur very effectively.


The H2S isn't really an issue until you try to dissolve the slag with either acid or water; dissolving the slag can easily be done in a controlled lab environment where it can be easily neutralized (I just didn't for the sake of time). Some sulfur dioxide is produced when the thermite burns, but that's not a big deal. H2S can be produced by the Al2S3 hydrolyzing with the moisture in air, but this doesn't happen at a fast enough rate to warrant a scrubber setup over the thermite reactor itself.

Other methods of producing silicon metal may have advantages, but I'm trying to work strictly with SiO2 and don't really want to buy another silicon compound. SiO2 is obviously quite difficult to work with chemically.

woelen - 12-10-2015 at 03:19

Quote: Originally posted by blogfast25  
A much simpler way is to take advantage of the modest solubility of LiCO3 and the high insolubility of CaCO3:

Ca(NO3)2(aq) + 2 LiCO3(s,aq) === > CaCO3(s) + 2 LiNO3(aq)

Simply simmer a slurry of LiCO3(s,aq) with a Ca(NO)2 solution (stoichiometric amounts) and the slurry will precipitate the Ca as carbonate, leaving the Li in solution as the highly soluble nitrate. Filter and recrystallise the LiNO3.

[Edited on 8-10-2015 by blogfast25]

I do not want to nitpick, but I see a disturbing error in your post. From blogfast I expect better :P

LiCO3 is a very special salt, which only exists in a parallel universe, yet to be discovered. In our universe we do have Li2CO3.

blogfast25 - 12-10-2015 at 06:17

Quote: Originally posted by woelen  


LiCO3 is a very special salt, which only exists in a parallel universe, yet to be discovered. In our universe we do have Li2CO3.


I never get those memos, woelen!

Oooopsie, Li<sub>2</sub>CO<sub>3</sub>, my very bad :mad:. Thank you.

woelen - 12-10-2015 at 07:03

Even the best sometimes has a slip of the brain ;)

Upsilon - 12-10-2015 at 19:58

I attempted the thermite again today - much better results with the larger amounts (I used roughly 100g of mixture). I got similar pieces to what I had before, though (definitely more of them however). I think I didn't allow them enough time in the HCl last time. I'll let them soak overnight and whatnot and report back tomorrow.

I don't think I'll be doing this again though. The sulfur dioxide produced this time was much more significant, as was the H2S involved (it quickly overpowered the bleach I used to neutralize it, I should have used a more concentrated oxidizer). I'm surprised I didn't incite any reaction from the neighbors with the billowing SO2 cloud and rotten egg smell. I didn't even dissolve most of the slag; I picked out what I wanted and left the rest alone. I bagged it up and I may use it as a crude H2S source in the future.

j_sum1 - 12-10-2015 at 21:32

Thanks for reminding me how smelly this one gets. I am due for another attempt on a larger scale but have put it aside because, of all things, my lab has lacked... sand. I now have some nice silica sand and so am ready for silicon extraction. I might want to choose a windy day.

blogfast25 - 13-10-2015 at 08:06

Quote: Originally posted by j_sum1  
Thanks for reminding me how smelly this one gets. I am due for another attempt on a larger scale but have put it aside because, of all things, my lab has lacked... sand. I now have some nice silica sand and so am ready for silicon extraction. I might want to choose a windy day.


In the quest to eliminate sulphur (and thus H2S) from the SiO2 thermite I tried replacing the S+Al booster system with a KClO3 + Al booster system several times (years ago). Although that worked well unfortunately the silicon got caught into extremely hard alumina from which it was impossible to remove it. Very frustrating.

But I wonder if it's worth trying this with the CaSO4 + Al booster method. That would simply mean taking the TiO2/CaSO4/Al/CaF2 formulation and substituting the TiO2 with SiO2, mol for mol. Worth a shot, methinks!

[Edited on 13-10-2015 by blogfast25]

Upsilon - 13-10-2015 at 15:55

So I recovered the pieces a little while ago. Upon applying even slight pressure they crumbled into a bunch of tiny useless shiny flakes (it is literally glitter). I dumped this in frustration and have since tried again with a couple more pieces I pulled out of the slag I saved. It's strange; these pieces that I pulled out looked exactly like what's in the picture on the website I used, and that was before HCl treatment. I put them in HCl and I'll let them simmer a while.

blogfast25 - 13-10-2015 at 17:33

Quote: Originally posted by Upsilon  
So I recovered the pieces a little while ago. Upon applying even slight pressure they crumbled into a bunch of tiny useless shiny flakes (it is literally glitter). I dumped this in frustration and have since tried again with a couple more pieces I pulled out of the slag I saved. It's strange; these pieces that I pulled out looked exactly like what's in the picture on the website I used, and that was before HCl treatment. I put them in HCl and I'll let them simmer a while.


Please bear in mind that silicon prepared this way is not particularly pure: apart from alloyed aluminium it probably also contains micro-occlusions of slag, which would explain the poor mechanical properties.

To increase size of globules obtained, you might want to 'stoke up the heat' by increasing both S and Al content. A higher end-temperature allows more cooling time and thus more longer time for the molten Si metal to coalesce into larger units.

Upsilon - 13-10-2015 at 18:02

Quote: Originally posted by blogfast25  

Please bear in mind that silicon prepared this way is not particularly pure: apart from alloyed aluminium it probably also contains micro-occlusions of slag, which would explain the poor mechanical properties.

To increase size of globules obtained, you might want to 'stoke up the heat' by increasing both S and Al content. A higher end-temperature allows more cooling time and thus more longer time for the molten Si metal to coalesce into larger units.


That's probably it, then. I don't mind impurities so much, but if I can't even get a cohesive sample then I don't really want it. Like I said I don't think I'll be doing this again just because of the nasty sulfur-based gas byproducts. Perhaps the calcium sulfate booster you mentioned would work, but after using so much aluminum powder on this already I'm not really too keen on using some more unless I know it'll work.

That being said, I honestly have no idea what else I would do to prepare elemental silicon. Perhaps the most promising compound is silicon tetraiodide (which I could melt down and electrolyze) but I don't know how I would get here from silicon dioxide. Maybe I could dissolve it in molten phosphoric acid, then react this silicon phosphate (which I cannot find information on surprisingly) with hydroiodic acid?

Either way this is certainly not the quick and easy lunch-break experiment I was hoping for! :P Goes to show that things are so much easier on paper in chemistry.

EDIT: You know, I might as well try carbothermic reduction since I'm planning on building an arc furnace anyway. This silicon can then be purified using a silicon tetrachloride method. I suppose I could have done this already but I didn't really want to try and scrape together a bunch of this fine silicon powder.

[Edited on 14-10-2015 by Upsilon]

blogfast25 - 13-10-2015 at 18:37

Quote: Originally posted by Upsilon  


That being said, I honestly have no idea what else I would do to prepare elemental silicon. Perhaps the most promising compound is silicon tetraiodide (which I could melt down and electrolyze) but I don't know how I would get here from silicon dioxide. Maybe I could dissolve it in molten phosphoric acid, then react this silicon phosphate (which I cannot find information on surprisingly) with hydroiodic acid?



Nah, SiI<sub>4</sub> is a purely covalent compound and can't be electrolysed. It's also hard to prepare at hobby level. Silicon phosphate doesn't even exist: Si is not a real metal, remember?

Industrially, reduction of SiCl<sub>4</sub> with Mg metal is used but forget about trying that in your garage or potting shed lab! :(

Upsilon - 13-10-2015 at 19:07

Quote: Originally posted by blogfast25  

Industrially, reduction of SiCl<sub>4</sub> with Mg metal is used but forget about trying that in your garage or potting shed lab! :(


Does it occur with explosive force or something? What exactly makes it impractical for the home lab?

Also like I said I may try the carbothermic reduction method for producing silicon when I get around to building my arc furnace.

...hell, with an arc furnace, I could probably just run the SiO2 + All thermite on its own, right? And avoid toxic gases altogether? Since the problem with this reaction normally is just not enough heat to sustain the reaction? In an arc furnace this would be no issue, plus I could keep the heat going as long as I want to allow the silicon metal to coalesce as much as I want it to.

[Edited on 14-10-2015 by Upsilon]

j_sum1 - 13-10-2015 at 21:12

Hmm. Sorry it didn't work out for you, Upsilon.
Last time I did this reaction was over a year ago. From memory it was somewhere between 50 and 80 grams of mix in total. IO used the same ratios as you. I got a few pebbles of silicon -- the largest of which would have been 8mm diameter. It seemed to have pretty good mechanical properties.
I can't remember how fine my aluminium was. That might make a difference. I am definitely looking forward to doing it again. I was thinking of perhaps a half kilo batch.

blogfast25 - 14-10-2015 at 07:03

Quote: Originally posted by Upsilon  

Does it occur with explosive force or something? What exactly makes it impractical for the home lab?

Also like I said I may try the carbothermic reduction method for producing silicon when I get around to building my arc furnace.

...hell, with an arc furnace, I could probably just run the SiO2 + All thermite on its own, right? And avoid toxic gases altogether? Since the problem with this reaction normally is just not enough heat to sustain the reaction? In an arc furnace this would be no issue, plus I could keep the heat going as long as I want to allow the silicon metal to coalesce as much as I want it to.


It's highly exothermic and with a volatile reactant (SiCl<sub>4</sub>;) that means pressurised reactors.

Arc furnace for carbothermic or aluminothemic reduction of sand should work. The carbothermic version is of course used industrially.

[Edited on 14-10-2015 by blogfast25]

blogfast25 - 14-10-2015 at 07:06

Quote: Originally posted by j_sum1  
I was thinking of perhaps a half kilo batch.


Why not try a small batch with CaSO<sub>4</sub> + Al, instead of sulphur?

Upsilon - 14-10-2015 at 08:54

Quote: Originally posted by blogfast25  


Arc furnace for carbothermic or aluminothemic reduction of sand should work. The carbothermic version is of course used industrially.

[Edited on 14-10-2015 by blogfast25]


Honestly I think I'll go with the carbothermic. I was worried about carbon monoxide but then I remembered that it would simply burn immediately after being produced to form carbon dioxide. Besides, the alumina would probably be too hard to melt in the arc furnace (especially since I'll probably be using alumina brick), plus the carbothermic reduction should leave nothing but Si metal behind so I don't have to deal with slag (in reality though, I expect there to be a noticeable amount of carbon left over).

blogfast25 - 14-10-2015 at 10:28

Quote: Originally posted by Upsilon  

Honestly I think I'll go with the carbothermic. I was worried about carbon monoxide but then I remembered that it would simply burn immediately after being produced to form carbon dioxide. Besides, the alumina would probably be too hard to melt in the arc furnace (especially since I'll probably be using alumina brick), plus the carbothermic reduction should leave nothing but Si metal behind so I don't have to deal with slag (in reality though, I expect there to be a noticeable amount of carbon left over).


On of the problems you could be having is to protect the Si from air oxygen since as you need very high temperature to conduct that reduction.

[Edited on 15-10-2015 by blogfast25]

Upsilon - 14-10-2015 at 13:46

Quote: Originally posted by blogfast25  

On of the problems could be having to protect the Si from air oxygen sinceas you need very high temperature to conduct that reduction.


That too; I'll try it out in air first to see just how bad the silicon is attacked. If it's too bad then I'll carry it out in a vacuum or inert gas environment.

But I'm getting ahead of myself. I won't be able to make the arc furnace for a while yet and the vacuum chamber is even further off (probably around December-January). For now I'll see what else I can do. I'll be trying the TiO2 thermite this weekend and possibly the lithium nitrate electrolysis. I think I have a pretty solid plan for lithium; my diagram is still floating around up there somewhere. I found this Kapton tape stuff that's designed for high temperature so I'll be using that to insulate my wires. As long as borosilicate glass can withstand the temperature of lithium nitrate, I'll be good to go.

blogfast25 - 14-10-2015 at 15:32

Electrolysis of a LiCl/KCl eutectic (MP about 300 C), not very well executed:

https://www.youtube.com/watch?v=lGs-4bUM5Gk

The best source of information on lab based preparation of Li is found in "Small-Scale Synthesis of Laboratory Reagents with Reaction Modeling" by Leonid Lerner:

http://www.twirpx.com/file/1618910/

From which it can be downloaded, I think.

Leonid Lerner was a prolific poster here, as 'Len1', before he left for greener pastures.

His method uses LiCl and a graphite crucible as cathode.

Several members here have a free version of the book (as *.pdf) so it shouldn't be too hard to get if you want it.


[Edited on 14-10-2015 by blogfast25]

Upsilon - 14-10-2015 at 15:46

Quote: Originally posted by blogfast25  
Electrolysis of a LiCl/KCl eutectic (MP about 300 C), not very well executed:

https://www.youtube.com/watch?v=lGs-4bUM5Gk


That's pretty interesting, I guess that works because the potassium metal produced is so soluble in its molten salt? Also, is there a way to predict the melting point of mixtures like there?

Though I still think I'll use lithium nitrate since it melts even lower than that mixture (enough so to be conducted under mineral oil where the lithium won't be exposed to air as well as being safe for regular glassware).

blogfast25 - 14-10-2015 at 15:50

See also my edit to my last post.

j_sum1 - 14-10-2015 at 16:40

Quote: Originally posted by blogfast25  
Quote: Originally posted by j_sum1  
I was thinking of perhaps a half kilo batch.


Why not try a small batch with CaSO<sub>4</sub> + Al, instead of sulphur?

I could. I do need to get some CaSO4. Last several visits to the hardware store have had me looking closely at the plaster section wondering how much CaSO4 is present and how much carbonate, hydroxide, magnesium and iron compounds and others -- and how much do these impurities and additions matter. With multiple products and no detailed list of ingredients I have walked away without buying anything.
I already have sulfur. (Or sulphur!)

blogfast25 - 14-10-2015 at 17:06

Quote: Originally posted by j_sum1  
Last several visits to the hardware store have had me looking closely at the plaster section wondering how much CaSO4 is present and how much carbonate, hydroxide, magnesium and iron compounds and others -- and how much do these impurities and additions matter.


I believe no-frills wall filler plaster is almost 100 % CaSO4 hydrate. More sophisticated brands may contain some CaO/Ca(OH)2 because its affects setting behaviour, 'allegedly'. A friend of mine used waste plaster boards from a skip (building site), ground them down and semi-calcined it, no problem.

Carbonates, if any, are of course easy to detect.

I've also used black board chalk sticks, some of which are CaSO4, some CaCO3, not hard to tell which is which, of course.

[Edited on 15-10-2015 by blogfast25]

blogfast25 - 14-10-2015 at 17:12

Quote: Originally posted by Upsilon  


That's pretty interesting, I guess that works because the potassium metal produced is so soluble in its molten salt? Also, is there a way to predict the melting point of mixtures like there?



No, the K<sup>+</sup> ions don't get reduced at all.

Eutectics? Just look for the various eutectic databases, e.g.:

http://www.crct.polymtl.ca/fact/documentation/FTsalt/FTsalt_...

[Edited on 15-10-2015 by blogfast25]

Upsilon - 14-10-2015 at 17:55

Quote: Originally posted by blogfast25  


No, the K<sup>+</sup> ions don't get reduced at all.

[Edited on 15-10-2015 by blogfast25]


Why is that? Is it because the lithium reduction is more favorable? If so, then woudn't K+ start to be reduced at some point when the Li+ concentration drops low enough to where K+ reduction is more favorable? Or is this not what happens at all?

Also, I would like to go off on a tangent about arc furnace design for a little bit, if anyone would like to help me out. I want to have an extremely high-heat design capable of 2500C+. Thus I decided not to use alumina brick and instead will use a silicon carbide crucible. I would like to be able to create a small pour spout in the crucible to easily pour molten material, but given the hardness of silicon carbide mechanical means of doing this are probably not ideal. Perhaps I could use an acetylene torch to do this? Also, how well would tungsten wire work for arc electrodes as opposed to graphite? Tungsten wire would be easier to work with than graphite, plus graphite electrodes burn out quickly due to reaction with oxygen. I feel like tungsten would have a similar problem though, oxidizing too quickly in air at extremely high temperatures

[Edited on 15-10-2015 by Upsilon]

[Edited on 15-10-2015 by Upsilon]

blogfast25 - 15-10-2015 at 06:20

Quote: Originally posted by Upsilon  

Why is that? Is it because the lithium reduction is more favorable? If so, then woudn't K+ start to be reduced at some point when the Li+ concentration drops low enough to where K+ reduction is more favorable? Or is this not what happens at all?

Also, I would like to go off on a tangent about arc furnace design for a little bit, if anyone would like to help me out. I want to have an extremely high-heat design capable of 2500C+. Thus I decided not to use alumina brick and instead will use a silicon carbide crucible. I would like to be able to create a small pour spout in the crucible to easily pour molten material, but given the hardness of silicon carbide mechanical means of doing this are probably not ideal. Perhaps I could use an acetylene torch to do this? Also, how well would tungsten wire work for arc electrodes as opposed to graphite? Tungsten wire would be easier to work with than graphite, plus graphite electrodes burn out quickly due to reaction with oxygen. I feel like tungsten would have a similar problem though, oxidizing too quickly in air at extremely high temperatures



Yes. Li is reduced first. On prolonged operation you'd need to top up the melt with LiCl, if only to maintain the eutectic composition.

Most use carbon I think because of it's high MP and because although it does burn off, at least it leaves no trace. W would give WO3. I'm sure in some instances W electrodes are used though.

MrHomeScientist - 15-10-2015 at 08:16

Back to the subject of silicon thermites, I've posted a nicer video of the method that produced larger amounts (though not in nice little spheres like my blog post you linked to). The video link is at the top of the post, but in case you missed it: https://www.youtube.com/watch?v=73YmP_JSrlU
It's my second most popular video, but not really the best quality IMO.

Also re: plaster quality, I had some experience with that in another of my blog posts, in the pursuit of titanium: http://thehomescientist.blogspot.com/2012/09/titanium-thermi...

That brand (DAP) had a 15% - 25% impurity of CaCO<sub>3</sub>, but that's easily remedied by adding sulfuric acid until the bubbling stops. Perhaps cheapo 'wal-mart' brand plaster would be more pure for our purposes.

I'm interested to try the silicon thermite sans sulfur. That one's always nasty.

blogfast25 - 15-10-2015 at 08:42

Quote: Originally posted by MrHomeScientist  


I'm interested to try the silicon thermite sans sulfur. That one's always nasty.


My formulation for TiO2/Al/KClO3/CaF2 was the direct result of research into a sulphur-free SiO2 thermite. I tried replacing the S + Al system by KClO3 + Al system.

Although that worked well, there where problems with coalescence of the Si, of which I found nice chunks completely stuck in rock-hard alumina. Impossible to remove them from there. This coalescence problem is largely due to the similar densities of molten Si and molten alumina.

I then used that same formulation with great success by simply replacing, mol for mol the SiO2 with TiO2.

I would therefore suggest to try the TiO2/Al/CaSO4/CaF2 formulation but substituting TiO2 with SiO2 (mol for mol). That should work because the Heats of Formation of TiO2 and SiO2 are very similar.

Using CaSO4 has one advantage: the slag is a mixture of CaS and Al2O3 and a lot softer than pure Al2O3.

[Edited on 15-10-2015 by blogfast25]

Upsilon - 15-10-2015 at 14:25

Well since I plan on carrying out these arc furnace experiments in an inert atmosphere, I think I'll use tungsten for that anyway. Of course graphite electrodes wouldn't decompose in this situation either, but like I said tungsten wire would be easier to work with.

The issue with this is that the ionization voltage for argon, for example, is much lower than that of regular air. Does this translate to a colder arc, though?

blogfast25 - 15-10-2015 at 16:58

Quote: Originally posted by Upsilon  

The issue with this is that the ionization voltage for argon, for example, is much lower than that of regular air. Does this translate to a colder arc, though?


What do you mean by ionization voltage here?

Upsilon - 15-10-2015 at 17:49

Quote: Originally posted by blogfast25  


What do you mean by ionization voltage here?


There's a real term for it but it has slipped my mind. It's the voltage required to break down the substance into an ionized plasma state that can conduct electricity.

EDIT: "Dielectric strength" is the term I'm referring to.

[Edited on 16-10-2015 by Upsilon]

blogfast25 - 15-10-2015 at 18:28

Quote: Originally posted by Upsilon  

EDIT: "Dielectric strength" is the term I'm referring to.

[Edited on 16-10-2015 by Upsilon]


Yes, that's the right term. And Ar has lower dielectric strength than air? Well, I never knew that!

Upsilon - 15-10-2015 at 18:51

Quote: Originally posted by blogfast25  

Yes, that's the right term. And Ar has lower dielectric strength than air? Well, I never knew that!


Apparently so; search "Dielectric gas" on Wikipedia and you'll get a nice table of dielectric strengths of various gases proportionate to air. Argon has 1/5 the dielectric strength of air, while Neon is only 1/50! Makes sense why they use neon as opposed to other noble gases in signs. On the other end of the spectrum is diamond - 1333 times more difficult to break down than air, at a whopping 4 gigavolts per meter!

[Edited on 16-10-2015 by Upsilon]

Upsilon - 16-10-2015 at 11:33

Just checked on that lithium nitrate/citrate disaster again; the calcium citrate had precipitated out and formed a beaker full of something frighteningly similar to sour cream. I could not resist playing with it before dumping it out :P

[Edited on 16-10-2015 by Upsilon]

blogfast25 - 16-10-2015 at 12:50

Quote: Originally posted by Upsilon  
Just checked on that lithium nitrate/citrate disaster again; the calcium citrate had precipitated out and formed a beaker full of something frighteningly similar to sour cream. I could not resist playing with it before dumping it out :P

[Edited on 16-10-2015 by Upsilon]


That's your calcium citrate: compounds like that often don't neatly crystallise.

aga - 16-10-2015 at 13:12

Quote: Originally posted by Upsilon  
frighteningly similar to sour cream

Sounds horribly familiar.

So much can 'go wrong' in even the simplest syntheses.

Upsilon - 16-10-2015 at 13:49

Quote: Originally posted by blogfast25  

That's your calcium citrate: compounds like that often don't neatly crystallise.


Yup, even though it did eventually precipitate, it would have been a nightmare to try and extract any lithium nitrate from that mess. I tried again with direct reaction of calcium nitrate and lithium carbonate and it worked much better.

Upsilon - 17-10-2015 at 09:47

Ok, so I did more research and learned that the LiNO3 electrolysis will not work because the nitrate ions in the melt will react with lithium metal. However, I also found that a LiNO3+KNO3 eutectic melts at around 120C! This means that solid lithium will be evolved; this solid lithium will form a protective layer insoluble in the melt that will protect it from further oxidation.

blogfast25 - 17-10-2015 at 09:50

Quote: Originally posted by Upsilon  
Ok, so I did more research and learned that the LiNO3 electrolysis will not work because the nitrate ions in the melt will react with lithium metal. However, I also found that a LiNO3+KNO3 eutectic melts at around 120C! This means that solid lithium will be evolved; this solid lithium will form a protective layer insoluble in the melt that will protect it from further oxidation.


What makes you think solid Li will not be attacked by nitrate ions?

Upsilon - 17-10-2015 at 13:22

Quote: Originally posted by blogfast25  

What makes you think solid Li will not be attacked by nitrate ions?


Take a look here:
http://www.sciencemadness.org/talk/viewthread.php?tid=26524&...

Someone on that thread actually did this. Apparently because the protective oxidized layer of lithium is insoluble in the molten mixture and prevents further attack.

EDIT: Sorry, linked the wrong page. It's actually the page before that one. Here is what it says:

Even if lithium/sodium/potassium/etc reacts with propylene carbonate (or any other electrolyte/solvent), it may still be useful for isolating these alkali metals. For example, I've electrolyzed a eutectic mix of LiNO3 and KNO3 at 125-150C, and isolated small amounts of lithium that way. You'd think that lithium would react vigorously with nitrates, and it does. However, lithium is a solid below 180C. The oxide formed on the surface of the lithium is a fast ion conductor for Li+, and is insoluble in the nitrate melt. This protects the bulk of the lithium from the electrolyte. The problem that I had with this setup was that I didn't have a platinum anode. Pretty much anything else gets destroyed by the nitrates. Also, the molten salt has to be very anhydrous, or the oxide layer weakens.


What I don't understand, though, is that even though the nitrate attacks the metal, why does it matter in the long run? Since nitrate is constantly being removed at the anode, eventually it should run out of nitrate ions to react with the lithium, right?
[Edited on 17-10-2015 by Upsilon]

[Edited on 17-10-2015 by Upsilon]

blogfast25 - 17-10-2015 at 17:30

Quote: Originally posted by Upsilon  

What I don't understand, though, is that even though the nitrate attacks the metal, why does it matter in the long run? Since nitrate is constantly being removed at the anode, eventually it should run out of nitrate ions to react with the lithium, right?


Yeah, like when the beaker is empty: for each reduced Li<sup>+</sup> a nitrate ion must be oxidised! :D

I did tell you someone had tried this before, I just couldn't find the reference.

Search for related posts with 'WGTR' as author and you might find what he did precisely.

Remember: molten nitrates will oxidise anything that's oxidisable.
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