Sciencemadness Discussion Board

The short questions thread (1)

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Paddywhacker - 10-3-2009 at 01:56

Just speculating that it could be the difference in solubility between hot and cold for the two substances. NaCl dissolves just a little more when heated wheras NaClO3 dissolves a lot.

Evaporate a mixed solution down to give crystals when hot, then filter and cool. Much more NaClO3 will crash out than NaCl. Make sense?

Ebao-lu - 10-3-2009 at 11:59

2 KClO4: boric acid can help to convert acid to amide. It is used to acylate alkylamines http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=v81...
and probably can be used with ammonium salts, that also not_important suggested once
Quote:

benzoic acid => ammonium benzoate => benzamide => aniline

the conversion to amide being done by heating with B(OH)3 catalyst in xylene, the water formed being azeotroped off.

It can be a problem in your case to get rid of unreacted fat via acid-base extraction, because if you simply basify it to get salt, it will solubilize your amide into emulsion. So better to use an excess of urea. You can get rid of urea/ its decomposion products by dissolving your amide in acetone (other solvents may also suit, but biuret that is one of the decomposion products is comperatively soluble in ethanol).

ps: Urea dissolves it fat because it melts at 133




[Edited on 10-3-2009 by Ebao-lu]

kclo4 - 10-3-2009 at 14:00

Quote:
Originally posted by barbs09
Hi, This will probably be an easy one for some one.. For the purposes of extracting a crop of NaClO3 from a cell containing NaCl and NaClO3, two methods are generally proposed 1, concentrating he solution untill the chlorate will crystallise out and 2, salting the chlorate out by the addition of a salutation solution of NaCl.

My question is: since NaClO3 is over twice the solubility of NaCl for the same volume of water, why does the more soluble of the two salts precipitate out?? I would have thought the least soluble salt (NaCl) would have crystallised out first. I cannot find a satisfactory answer anywhere.

Thanks in advance



I think perhaps you are confusing NaClO3 with KClO3.
They normally add KCl to salt out KClO3 from a NaClO3 solution.
There are a lot of threads on this I believe, and I know there are a few very good websites about this.

kclo4 - 10-3-2009 at 14:11

Quote:
Originally posted by Ebao-lu
2 KClO4: boric acid can help to convert acid to amide. It is used to acylate alkylamines http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=v81...
and probably can be used with ammonium salts, that also not_important suggested once
Quote:

benzoic acid => ammonium benzoate => benzamide => aniline

the conversion to amide being done by heating with B(OH)3 catalyst in xylene, the water formed being azeotroped off.

It can be a problem in your case to get rid of unreacted fat via acid-base extraction, because if you simply basify it to get salt, it will solubilize your amide into emulsion. So better to use an excess of urea. You can get rid of urea/ its decomposion products by dissolving your amide in acetone (other solvents may also suit, but biuret that is one of the decomposion products is comperatively soluble in ethanol).

ps: Urea dissolves it fat because it melts at 133




[Edited on 10-3-2009 by Ebao-lu]


I don't think the Urea dissolved, instead it just melted and remained at the bottom.

I'd still like to find an easier way to get the glycerol from the fatty acids then to react it with NaOH, and then add an acid. I don't have any NaOH at this moment and right now is not the best time for me to be buying it.


Thanks for the info on the boric acid! I didn't know that helped, perhaps that would allow the urea and the triglycerides to react. I wonder if heating up boric acid with olive oil might free up the fatty acids by forming borogylcerin. I doubt it would since the boric acid is so weak but it would be nice if it did.

Is there any other effective method of freeing the fatty acids from trigycerlides? such as something that would increase hydrolysis? boiling in a solution of acid, perhaps? I've tried to make a soap with sodium carbonate but that failed miserably.

dann2 - 10-3-2009 at 16:25

Hello,

Regarding the solubility of NaCl + NaClO3 it will depend on concentration of each salt and there mutual solubility. You can read up on the joys (and joys they are) of mutual solubility here.

Salting out was actually used in industry to obtain a crop of Sodium Chlorate. Adding a concentrated solution of NaCl to a (suitably concentrated and at a suitable temperature) solution of NaCl + NaClO3 caused some NaClO3 to fall out of solution.

Dann2

barbs09 - 11-3-2009 at 00:32

Thanks for all your replies and dann2 special thanks for the link to the mutual solubility curve. This is off your Geocities-Cape Canaveral website I believe? If it is you have added more to it since I last visited-good stuff by the way. Do you regularly update the site?

I thought by asking a chlorate question outside one of the umpteen dozen chlorate production strings/sites I ran the risk of getting flamed but I was genuinely confused as to why the more soluble sodium chlorate could come out of solution before the less soluble sodium chloride. I believe my answer will be found in the complicated looking mutual solubility diagram. It may take a few single malts to get there though :D

Thanks all, AB

not_important - 11-3-2009 at 01:14

Quote:
Originally posted by kclo4(much snipping)
Is there any other effective method of freeing the fatty acids from trigycerlides? such as something that would increase hydrolysis? boiling in a solution of acid, perhaps? I've tried to make a soap with sodium carbonate but that failed miserably.


Acid hydrolysis can be done, but is difficult in a non-industrial situation. Simply heating fats with water under pressure works too, and again isn't the easiest at home. Also remember that the acid catalysed hydrolysis is reversible, use an excess of water and all that.

Sodium carbonate will work, but it is slow. The CO2 of the carbonate is slowly released, actual boiling helps sweep it away in the steam. But it takes hours to go to completion, and you do need to use an excess of Na2CO3. You need good mixing, the steam helps with that but mechanical stirring usually is needed.

The direct route to amides should work, although again time and good stirring are likely needed. A small amount of NH4Cl might help, functioning as an acid catalyst.

Ebao-lu - 11-3-2009 at 03:41

why not to purchase a piece of soap? there should be a sodium stearate soap, without fragrances, it is usually brown. Tar soap is also usualy sodium salt. Boric acid can surely catalyse the hydrolysis of fat via same mechanism lined out in the link, but a'm not sure that the rate of reaction will be fine, besides fat is immiscible with water. Glycerol should also be sold in pharmacies.

Sedit - 13-3-2009 at 09:07

I have recently taken a fondness to electrochemistry and since im just starting to learn how to effectively use a divided cell. Well my question is what one of the products may be.
Table salt is the electrolyte, cell is partioned with unglazed clay pot. The electrodes at this point where nothing more then Galvanized steel. PH indicator showed the inside of the pots PH jumped right away. Both solutions remained clear and the cell was only run for maybe a half an hour or so. So after getting a few mesurements on current ect I got ready to clean it and start something else. When the two clear solutions are mixed Something white with a slight green tinge precipitates. Any idea what it is?
Basicly its just two zinc electrodes with a NaCl electrolyte.

Just woundering so i can get my head around the normal workings of a divided cell.

hissingnoise - 13-3-2009 at 10:52

I would suggest some reading on electrolysis---the cell you describe will simply waste your time and dissolve your anode.
The green precipitate is an anode-metal chloride.

Sedit - 13-3-2009 at 11:00

I know a decent amount about electrolysis thank you for the suggestion though.
Waste of time? Nope not at all I got the information I was after about how much current was going through ect. It had no real purpose other then making sure the power source was functioning properly and the electrodes where used for nothing more then opportunity alone. I highly doubt it is any metal chloride being as zinc chlorides would be soluble in the solution. Both solutions where clear befor being mixed and precipitated when mixed.

Its just a matter of curiosity what is happening when they are mixed.

hissingnoise - 13-3-2009 at 11:32

Quote:
Originally posted by Sedit
I know a decent amount about electrolysis thank you for the suggestion though.


Sedit, with all due respect, your questions on the subject suggest otherwise.
The most elementary reaction in brine electrolysis is dissolution of active anodes as their chlorides.
It's the very first thing one learns. . .

Sedit - 13-3-2009 at 11:44

And the solution is clear. There is no doubt that the Chloride was generated but it precipitates when added to the other partion. The Chloride is reacting with what is at the cathode partion and precipitating a very light green fluff. I am under the assumption that at the cathode is NaOH and Zn(OH)2. This is reacting with the chloride to produce ...? NaCl + ??

You see what Im getting at here?

hissingnoise - 13-3-2009 at 12:26

In a partitioned cell, metal ions will accumulate in the cathode compartment raising pH, while chloride and hypochlorous ions in the anode compartment will lower pH.
Mixing the basic and acidic solutions will precipitate the metal chloride + some hypochlorite. . .

Sedit - 13-3-2009 at 13:14

Im leaning toward hypochlorites and that was my original hunch thats why I was asking. All ZnCl i have ever made from reactions with hydrochloric have been soluble in H2O yet the precipitate was compleatly non soluble so there is no chance of it being metal chloride because that is highly soluable in H2O.

hissingnoise - 13-3-2009 at 13:49

Hypochlorite formation would be negligible because once formed the greater part of the chlorine would react directly with the active anode.
As for salt solubilities, your electrolyte is already saturated with NaCl. . .

Sedit - 13-3-2009 at 14:12

True being that ZnCl2 is more soluble then NaCl its possible that the precipitate is just salt with some hypochlorite contamination possibly explaining the pale green color. Only way to be 100% possitive would be to reproduce it under more formal conditions and test what I precipitates but its not to important to me so it may be a while.

497 - 18-3-2009 at 15:14

Are lactones able to form bisulfite adducts? Somehow I doubt it because I've searched and haven't found anything mentioning it.. But I wanted to make sure.

not_important - 18-3-2009 at 15:34

Lactones are esters, not noted for form bisulfite adducts.

Intergalactic_Captain - 19-3-2009 at 14:00

Are there any uses for carbodiimide peptide coupling reagents aside from making LSD? My combichem group is currently working towards a small tripeptide library and I've been stuck with all the research - From what I've read, it seems that activation of carboxylic acids is quite general (with carbodiimides, DIC in our case). In particular, our scheme for coupling the first AA (fmoc-ala) to the (wang) resin involves DIC/DIPEA (conditions currently being optimized).

Considering this is basically just an accelerated esterification reaction (wang resin is a polystyrene/divinylbenzene polymer functionalized with a benzyl alcohol linker), I'm wondering about the general applicability of coupling a carboxylic acid to other substrates, in particular the preparation of methyl esters for use as methylating agents (dimethyloxalate in particular, PTSA if it's even possible, ...) avoiding the necessity of forming the acid chloride. DIC is way out of my price range, but its preparation by way of diisopropylurea and HgO is not.

fractional distillation question

marksev1 - 20-3-2009 at 11:24

Hello everybody! Which fractional column would you advise me (vigreux or something different), and more importantly how long, 300-400mm or more, for effectivly distilling ethanol from denaturated alcohol (don't know yet which "contaminants" are present in the mixture with ethanol). Cause i would like it pretty nontoxic for some extractions of some medicinal plants and so on..
There would be an easier way of just buying some vodka and distilling it but it wouldn't be such a challange:D. Thanks for the answers

DJF90 - 20-3-2009 at 11:38

Perhaps there is another solvent you could used for your extractions. If you provide details of what you intend to extract and from what, then we may be able to tell you which solvent(s) are best. It's possible that IPA will be suitable for your extraction, and this is available in 70% and 99% concentrations OTC. Again, I cant be sure unless you tell me what you intend to extract.

marksev1 - 21-3-2009 at 02:22

Yes, IPA is also a good choice, but i also don't know which impurities are in the more cheaper technical IPA in my country, if methanol is the case, distilling would be an easier task because of the larger difference in boiling points. Here is a monograph on residual solvents in pharmaceutical products:

http://lib.njutcm.edu.cn/yaodian/ep/EP5.0/05_general_texts/5...

(class 3 solvents would probably be good)

Btw, are spiral fractional coulumns more effective than vigreux...i've read somewhere that for the vigreux the HETP is about 10cm and for spiral column it would be about 3cm, making it more effective...but i also read that metal spirals are even more efective, so which material could be used for the spiral, perhaps stainless steel (just how inert is the chromium oxide coating present on stainless steel)? Some inert metal like platinum would probably be the best choice, but it is costly big time, so i would need a cheaper alternative that would also be quite sufficient.

Does anybody maybe know of a book of qualitative analysis of natural products?

Concerning what would i like to extracted, different things from different plants, starting from caffeine,piperine, atropine, quinine, nicotine, carvone,limolene,safrol,amygdalin,ginkolides,coumarines,menthol ...for more volitale things i would use steam distillation.

querjek - 23-3-2009 at 06:24

I've got some sodium formate and H2SO4 and would like some concentrated formic acid. Is it enough to simply distill a mixture of the two around 101 degrees C or should I add water?

not_important - 23-3-2009 at 07:24

In the presence of strong acids and/or heat formic acid decomposes into water and carbon monoxide; even pure +90% formic acid slowly decomposes so storage containers need pressure relief to prevent explosions. You must use dilute acid, slowly add it to a solution of the formate, always have an excess of format, and keep it cool. When finished, chill and filter off the hydrated sodium sulfate, the fractionate to get about 70% formic acid.

Supposedly it can also be had by mixing sodium formate and bisulfate and distilling under vacuum, but I've no details on that.

Jor - 23-3-2009 at 13:45

Quote: Originally posted by not_important  
In the presence of strong acids and/or heat formic acid decomposes into water and carbon monoxide; even pure +90% formic acid slowly decomposes so storage containers need pressure relief to prevent explosions. You must use dilute acid, slowly add it to a solution of the formate, always have an excess of format, and keep it cool. When finished, chill and filter off the hydrated sodium sulfate, the fractionate to get about 70% formic acid.

Supposedly it can also be had by mixing sodium formate and bisulfate and distilling under vacuum, but I've no details on that.


Aha, so that is why we have PLASTIC Acros bottled with HCOOH at uni. That was the first plastic Acros bottle I have ever seen.

querjek - 25-3-2009 at 04:35

To prepare ~30% HBr in acetic acid, could I saturate some acetic acid with HBr gas generated by dripping H2SO4 onto NaBr?

sakshaug007 - 26-3-2009 at 09:11

Quote: Originally posted by querjek  
To prepare ~30% HBr in acetic acid, could I saturate some acetic acid with HBr gas generated by dripping H2SO4 onto NaBr?


The sulfuric acid/NaBr reaction will work to produce HBr gas but is often not preferred since it is an oxidizing acid and therefore will inevitably form bromine, sulfur dioxide, and H2O impurities. Its actually best to use non-oxidizing acids such as phosphoric acid or even acetic acid. I'm also not sure how soluble HBr is in acetic acid I assume a little would dissolve but I don't know about 30%. Lastly, I could be wrong but what might also happen is the acid could protonate the hydroxyl group and form acetic bromide compounds as by products. I hope this was helpful.





[Edited on 26-3-2009 by sakshaug007]

sakshaug007 - 26-3-2009 at 09:25

Hello everyone,

I am interested in synthesizing methyl iodide and was wondering if the following reaction would work:

CH3OH + H2SO4 + KI ---> CH3I + K2SO4 + H2O (heat solution @ ~60°C to distill CH3I)

My HSC calculations show it to be favorable, but again I don't anything about the rate. Has anyone tried this? Or does anyone suspect it to work?

Thanks a lot

Jor - 26-3-2009 at 10:12

Yes, HBr can dissolve to certainly 30% in GAA.

Conc. H2SO4 on KBr will produce lots of bromine. Use phosphoric acid.

No you can't make MeI this way. Even cold, sulfuric acid will oxidises iodides to iodine, itself being reduced to sulfur dioxide, sulfur, hydrogen sulfide, and possibly other sulfur-compounds.

The prefferred method uses red P/I2. These two react to form PI3, wich reacts with water forming phosphorous acid (H3PO3) and HI, and HI will react with methanol, forming MeI.

I also recall (not sure) that if highly concentrated (it was more than 85%, i think 95%) is heated with methanol/KI, MeI is formed, but I'm really not sure, and I can't remember where I read it.

Be careful, MeI is a nasty carcinogen.

[Edited on 26-3-2009 by Jor]

sakshaug007 - 26-3-2009 at 11:02

Quote: Originally posted by Jor  

No you can't make MeI this way. Even cold, sulfuric acid will oxidises iodides to iodine, itself being reduced to sulfur dioxide, sulfur, hydrogen sulfide, and possibly other sulfur-compounds.
[Edited on 26-3-2009 by Jor]


Would my reaction work substituting sulfuric acid for other non oxidizing acids? What are some other strong non-oxidizing acid besides H3PO3?

Thanks

Paddywhacker - 26-3-2009 at 12:31

Quote: Originally posted by sakshaug007  
Quote: Originally posted by Jor  

No you can't make MeI this way. Even cold, sulfuric acid will oxidises iodides to iodine, itself being reduced to sulfur dioxide, sulfur, hydrogen sulfide, and possibly other sulfur-compounds.
[Edited on 26-3-2009 by Jor]


Would my reaction work substituting sulfuric acid for other non oxidizing acids? What are some other strong non-oxidizing acid besides H3PO3?

Thanks

Yeah, what about toluene sulphonic acid, or even a big handful of acid exchange resin that could be recovered afterwards?

User - 27-3-2009 at 03:51

I was wondering about boiling down hydrogen nitrate.
Years ago my only source for nitric acid was " Ph- " about 30% HNO3.
Well according to Wikipedia the water/nitric azeotrope boils at 121,8 degrees Celsius.
So a quite easy method of determining whether the concentrating has reached it's azeotrope or 68% would be watching the temperature stabilize at 121.8 degrees.


Well my question is how could a determine the concentration by an alternate way.

( has anyone a table of concentration/density versus mass )



[Edited on 27-3-2009 by User]

UnintentionalChaos - 27-3-2009 at 09:40

Quote: Originally posted by User  
I was wondering about boiling down hydrogen nitrate.
Years ago my only source for nitric acid was " Ph- " about 30% HNO3.
Well according to Wikipedia the water/nitric azeotrope boils at 121,8 degrees Celsius.
So a quite easy method of determining whether the concentrating has reached it's azeotrope or 68% would be watching the temperature stabilize at 121.8 degrees.


Well my question is how could a determine the concentration by an alternate way.

( has anyone a table of concentration/density versus mass )



[Edited on 27-3-2009 by User]


titrate. It answers everything :P

User - 27-3-2009 at 09:48

Well just in case iam not in the possession of a buret any alternate ideas.

DJF90 - 27-3-2009 at 10:44

You could do a crude titration using syringes.

Paddywhacker - 27-3-2009 at 11:43

Quote: Originally posted by User  
Well just in case iam not in the possession of a buret any alternate ideas.


With a cheap 0.01g balance you could do weight/weight titrations. Just weigh everything before and after and use a dropper to do the titration.

sakshaug007 - 27-3-2009 at 11:59

Does anyone know where to get chloroform? (I don't want to make it). I want to electrolyze it for poly(hydridocarbyne) synthesis.

Thanks

DJF90 - 27-3-2009 at 13:10

Try a chemical supplier. I dont think there are any OTC sources of chloroform, although I have seen DCM (at a ridiculous price, about 8 quid for 250mls! It was some kind of hob cleaner IIRC).

Aubrey - 27-3-2009 at 16:04

I attempted the synth of Cyclopentanol from cyclopentanone using Al Isopropoxide. I kept distilling until the vapor temp was 84.5 degrees and have about 300ml of clear liquid in my receiving flask (now in a bottle).
The method calls for reacting the prcpt in the flask with H2SO4 until acidic. now the problem is that if they are talking of the reaction flask rather than the receiving flask, i have a black mess in there from which i cannot even measure the ph because it tuns the litmus black! Should I filter this fist or distill the lot across? I guess i could dip in my digitla ph meter into this mess but im not too keen on that idea.

User - 28-3-2009 at 00:03


For the record.
I don't know if ur purposely giving this guy incorrect info but DCM is not chloroform.
That would be TCM (trichloromethane) not DCM ( Dichloromethane ).

[Edited on 28-3-2009 by User]

DJF90 - 28-3-2009 at 05:26

I know the difference between chloroform and DCM. I was merely stating that I have not seen chloroform OTC, but I have seen DCM albeit at a ridiculous price. In some cases it may be possible to use DCM in place of chloroform.

kclo4 - 28-3-2009 at 09:05

Quote:
I want to electrolyze it for poly(hydridocarbyne) synthesis.


wow that stuff sounds really interesting! From wiki: Upon thermolysis in argon at atmospheric pressure and temperatures of 110°C to 1000°C, decomposition of poly(hydridocarbyne) results in hexagonal diamond (Lonsdaleite).

panziandi - 28-3-2009 at 15:54

I just thought I'd share a nice photo of NaK with you all. Hope you like it. The second photo is a closeup of one of the globules. The NaK was prepared by melting potassium and sodium under pre-dried paraffin, it was transfered to a vial, extruded into paraffin under argon from a pipette to remove surface oxide contamination.



CaptureNaK.JPG - 23kB CaptureNaKcloseup.JPG - 22kB

sakshaug007 - 28-3-2009 at 20:28

Quote: Originally posted by kclo4  
Quote:
I want to electrolyze it for poly(hydridocarbyne) synthesis.


wow that stuff sounds really interesting! From wiki: Upon thermolysis in argon at atmospheric pressure and temperatures of 110°C to 1000°C, decomposition of poly(hydridocarbyne) results in hexagonal diamond (Lonsdaleite).


Yep that's what I plan to do, I want to diamond coat carbon electrodes for electrolysis experiments. But I need to figure out a way to dope the material to make it conductive. Perhaps in situ during electrolysis using some kind of boron compound but I'm not sure which compound to use.

dann2 - 29-3-2009 at 06:59

Hello,
Can anyone suggest another indicator, (more easy to get) other than Sodium Diphenylamine Sulphonate as a redox indicator.
Thanks, Dann2

Paddywhacker - 29-3-2009 at 11:25

Quote: Originally posted by dann2  
Hello,
Can anyone suggest another indicator, (more easy to get) other than Sodium Diphenylamine Sulphonate as a redox indicator.
Thanks, Dann2


Fe[III]/thiocyanate, KI/starch

DJF90 - 29-3-2009 at 17:26

Nice NaK panziandi! It does look very shiny...

sakshaug007 - 30-3-2009 at 23:08

Hello everyone,

I am going to attempt the electrolysis of chloroform in acetone solvent, and I wanted to ask what ionic salts you would suggest for an electrolyte, (i.e., readily soluble in acetone). I understand that I can just use a small amount of NaCl but if I end up doing a lot of this in the future I would prefer a salt that can dissolve in higher concentrations in acetone.

Thanks a lot.

Nicodem - 31-3-2009 at 00:17

Lithium perchlorate can and is commonly used as electrolyte in organic solvents. It is soluble in all solvents that are at least slightly basic, including acetone, and it does not get destroyed by oxidation at the anode like NaCl (which is not soluble in acetone anyway). But where did you saw acetone can be used as solvent? I would imagine methanol or something like that, but not acetone which reacts at both electrodes.

barbs09 - 31-3-2009 at 01:13

Hi all, a simple question that probably belongs in “Beginnings” but is not worthy of a new thread..

Stupidly I have been using stichometric calculations for chemical synthesises without taking the water molecules on some hydrates into consideration. This suddenly explains my low yields :(

When working out a theoretical yield of a reaction using, for example, copper sulphate pentahydrate as a reactant, should I do my molar calculations based on the anhydrous sulphate (159.62 g/mol) then correct for the extra H2O or start out using a combined molar mass of the pentahydrate which is 249.62 g/mol (159.62 g/mol CuSO4 + 90.08 g/mol H2O (5 x 18.016)).

Any help would be appreciated.

Cheers



sparkgap - 31-3-2009 at 02:29

Quote: Originally posted by barbs09  

When working out a theoretical yield of a reaction using, for example, copper sulphate pentahydrate as a reactant, should I do my molar calculations based on the anhydrous sulphate (159.62 g/mol) then correct for the extra H2O or start out using a combined molar mass of the pentahydrate which is 249.62 g/mol (159.62 g/mol CuSO4 + 90.08 g/mol H2O (5 x 18.016))


Well, you really should just use the molar mass of the hydrate of whatever salt you're interested right from the start. The "correcting" you speak of looks to be a more error-prone procedure to me.

sparky (~_~)

sakshaug007 - 31-3-2009 at 07:46

Quote: Originally posted by Nicodem  
Lithium perchlorate can and is commonly used as electrolyte in organic solvents. It is soluble in all solvents that are at least slightly basic, including acetone, and it does not get destroyed by oxidation at the anode like NaCl (which is not soluble in acetone anyway). But where did you saw acetone can be used as solvent? I would imagine methanol or something like that, but not acetone which reacts at both electrodes.


I actually didn't see that it can be used as a solvent in this case I'm simply trying it because I don't have any other polar aprotic solvents at the moment (I am going to get DMSO pretty soon though). You said that acetone reacts at both electrodes what does it produce? I was able to dissolve a significant quantity of LiCl in acetone (0.35g/100ml) when I attempted the electrolytic deposition of lithium. It did work, lithium was deposited on the cathode as black amorphous powder which reacted violently when dipped into water, and likewise chlorine was produced at the anode (graphite electrodes were used). By the way would other perchlorates dissolve just as well such as ammonium perchlorate, or sodium perchlorate?

Thanks for the reply.

bquirky - 31-3-2009 at 08:24

Quote: Originally posted by sakshaug007  
Hello everyone,

I am going to attempt the electrolysis of chloroform in acetone solvent, and I wanted to ask what ionic salts you would suggest for an electrolyte, (i.e., readily soluble in acetone). I understand that I can just use a small amount of NaCl but if I end up doing a lot of this in the future I would prefer a salt that can dissolve in higher concentrations in acetone.

Thanks a lot.



Hi ive done some playing around with zinc chloride in acetone
it plates out but a zinc anode should replace the zinc in solution

I have also used aluminium chloride in ethanol which plated out aluminium on the cathode as well (it fizzed in NaOH)

Sodium chloride dosnt seem to dissolve well in any organic solvent that i have been able to try (not many)

Ive messed around with oxalate salts but my results where highly varied i think the oxalate may have been braking down

NaOH seems to dissolve in ethanol there is some conductivity im not sure if that might be usefull for you

On a separate note does anyone know if electrolisis can be used to strip water from organics to make esters ? perhaps in the presence a very small amount of h2so4 with electrolisis simply removing the water from the sulfuric acid ?


marksev1 - 31-3-2009 at 10:54

Hello i would need to know where to download some handbooks useful for searching physico-chemical data..i already have CRC handbook of chemistry and physics, Lange's handbook,International critical tables... i would especially like to have Atlas of spectral data and physical constants for organic compounds, if somebody knows where to find it....
What would you suggest for searching data such as solubility, i would need the most maybe for various natural products.. thanks for the answers!

barbs09 - 31-3-2009 at 11:02

Thanks sparky, sounds good to me

kclo4 - 31-3-2009 at 14:37

sakshaug007, perhaps Sodium Iodide? I remember that it is more soluble in acetone then the other sodium halide compounds.

Sodium Iodate

Panache - 31-3-2009 at 15:50

3 I2 + 6 NaOH → NaIO3 + 5 NaI + 3 H2O

Dehydrate this and heat to around 500C to decompose the iodate to NaI, i guess it would evolve oxygen (wow i'm really insightful today).

Question. Is NaIO3 'explosive' like the MSDS's and similar such information sources proclaim or is this process rather benign.

Is there a more practical route to NaI from KI than liberating the element and doing as i propose?


Picric-A - 1-4-2009 at 12:03

@ Panache- Release HI from the KI using conc phosphoric acid, and run the vapours through a NaOH or Na2CO3 soloution.

Now my question, im having trouble making 4-Aminobenzoic acid from polystyrene. I have managed to nitrate the polystyrene (i think) by dissolving the polystyrene in dichloromethane then nitrating this mix using the normal HNO3/H2SO4 mix.
What im left with is a yellowish powder which is insoluble in water. The next step requires me to oxidise the polymer bond forming a methyl group. I plan on oxidising this with with alkaline KMnO4. The problem i am facing is the nitropolystyrene does not dissolve in the oxidising mix and so nothing happens.
Does anyone know what i can do to mend this problem, eg, what solvent can i dissolve it in to oxidise it?
thanks,


DJF90 - 1-4-2009 at 13:36

I would suggest acetone. PS dissolves in it like a treat so I would expect nitrated polystyrene to also. However you may have compatibility problems with the alkaline permanganate.

panziandi - 1-4-2009 at 13:46

Prolonged refluxing! As you begin to break the C-C polystyrene backbone to form the carboxylic acid it will dissolve in the alkaline solution. It will be a slow reaction requiring prolonged and vigorous conditions, go figure!

Panache - 2-4-2009 at 01:04
Quote:

I just seen on How its Made of them using some chemical to etch the stainless steel with a picture of the maker, Any idea of what was used? He put the chemical on a sponge and just pressed if for a few seconds and it etched the SS. Im thinking acid but im woundering if anyone has a more definitive answer.



HBr etches stainless quite effectively and quickly and is much safer than HF.

Does anyone know why the density vs concentration tables for h3PO4 solutions stop at 40% concentration of the acid?
Can tungsten be considered the equivalent to tantalum wrt using it for the wire in a whip stirrer or is it far less inert?

[Edited on 2-4-2009 by Panache]

sakshaug007 - 2-4-2009 at 13:00

Does anyone know if acetone is a suitable solvent for carrying out non aqueous electrolysis reactions? i.e. doesn't oxidize or reduce. My experience with it thus far has shown it to be pretty stable, I electrolyzed LiCl in acetone to produce black amorphous lithium metal and chlorine using graphite electrodes then again I'm not sure if I produced lithium-acetone compounds as well. If acetone does undergo oxidation/reduction in an electrolytic cell what does it form? Is it dependent on the electrolytes? Any advice would be great.

Thank you very much

sparkgap - 2-4-2009 at 17:30

Two questions from me:

1. A Fries arrangement will work for diphenyl carbonate and diphenyl oxalate, correct?

2. Would the conditions for hydrogenolysis of a benzyl group reduce double and triple bonds too?

sparky (~_~)

Sedit - 2-4-2009 at 18:59

I have been looking into Phosphorus preperation and I happen to come across this while looking for something else..

How valid is this?
Quote:

Copper Carbonate, artificial malachite.
Cu2(OH)2CO3, toxic, green powder which is soluble in acids and decomposes at 200 C; used in pigments and pyrotechnics and as a fungicide and feed additive; antidote for phosphorous poisoning.


If true and it will stop phosphorous poisoning then how would one administer it?
I dont plan on ever getting phosphorous poisoning but its always good to have a backup plan, I didnt plan on choping my finger tip off either but shit happens.

Sedit - 4-4-2009 at 18:07

LOL for reasons such as this.....

I recently made sodium nitrite from Pb and sodium nitrate. Thru out the whole time around it I was wearing latex gloves and being extreamly careful and yet I started to get a metalic taste in my mouth. I dont really know how I contacted the Pb but the only thing I can think of is that the vapors that came off where more then organic. The lead has been setting outside for years now and was once used as old plumbing pipe so was naturaly quite cruded up before melting and skimming.

My question is what are the detection limits of the PbO or elemental Pb. If a metallic taste is had is that a sure fire sign that Im screwed or should I relax a little bit more?
I was being very careful in ever aspect I could think of from gloves to working on wet news paper to keep dust levels down. I did not have a mask on but I would next time. Should I be very concerned or give it some time. No ill effects are felt just an annoying slight metallic taste.

And what if any Pb could have volatilized from the dirty lead?

not_important - 4-4-2009 at 21:31

Quote: Originally posted by Sedit  
...

If true and it will stop phosphorous poisoning then how would one administer it?
....


See http://tinyurl.com/cawym4 for a 1908 take on this, from Google books.

Usually a solution of CuSO4 was used, but a suspension of the 'carbonate' would work. The stomach is washed out with the copper solution/suspension, insoluble and fairly unreactive copper phosphide being formed. Similar solution can be used to treat phosphorous burns, preventing bits of phosphorous from remaining in the wound to reignite on exposure to air, or to diffuse into the body. Again, just the first step in the treatment, good cleaning of the wound is needed as well as other treatment.

Sedit - 4-4-2009 at 23:07

I like the idea of the turpentine antidote since it stated that some one was able to drink phosphorous followed with turpentine.

Thanks for that paper not_ since then I started drinking alot of milk as it talk about for heavy metals and metallic taste is pretty much gone. Still a little nervous for good reason but should be a problem.

Preparative TLC Plates for Analytical Work

smuv - 5-4-2009 at 09:07

For a good price I can get some 2000 micron silica TLC plates; I have no desire to use these plates preparatively, but wonder if they would work out ok for analytical work? I assume I would just need to spot a little bit more material onto the plate? Also is the thicker silica layer more prone to flaking and cracking?

Thanks in advance; sorry if my questions are a little mundane; I just don't want to waste my money.

Nicodem - 5-4-2009 at 23:28

Preparative plates are not really suitable for your needs. I would really think twice before buying them unless they are really, really cheap. It is not so much in that they are less suitable for analytical size of samples, but in that they are most commonly made on glass plates and cutting them without damaging is a challenge in itself. They are used as a whole plate with the sample distributed along the entire starting line and are therefore not made on a thick aluminium foil support like the analytical ones which can be cut with a knife. Also, a package of preparative TLC plates contains much less plates (maybe 10 utmost) than the normal TLC package which usually contains 25 foils (20×20cm), where each foil can be cut to 1/3 lines (so if you use ~1cm wide TLCs you can make approximately 3×25×20=1500 TLCs!).

Aubrey - 6-4-2009 at 10:19

noob question (as alwys)
I attempted the reduction of cylopentanone using Al Isopropoxide. I have a black tarry substance in my flask which stinks on warming and i am afraid of reheating it up until i buy a fume hood to recover my product, because last time my neighbours smelt gas and called out the gas men, this is even wih placing a tube from the vacuum outlet into liquid.
Having said that, my question is, which other reducing agents may I use that are more pleasurable to work with?

Would I do better to look at synthing LaH?
many thanks

Paddywhacker - 6-4-2009 at 11:36

Did you have a reasonable excess of isopropyl alcohol?
Not too much water?
Not too high a pH? Ketones are notorious for giving tarry gunk with alkali.

Bad neighbours are a real nuisance, best dealt with by educating them. Maybe tell them that you are running an alcohol still. That will cover for the occasional smell.

dann2 - 6-4-2009 at 16:09

Hello,

Sorry to butt in (but you always have to butt in when asking a question in ' The Short questions' thread), but can someone sort out the following for me.

Plating equations for Lead Dioxide (Lead Nitrate bath)are:

[1] H<sub>2</sub>O --> OH<sub>ads</sub> + H<sup>+</sup> + e<sup>-</sup><br>
[2] Pb<sup>2+</sup> + OH<sub>ads</sub> --> Pb(OH)<sup>2+</sup> <br>
[3] Pb(OH)<sup>2+</sup> + H<sub>2</sub>O --> PbO<sub>2</sub> + 3H<sup>+</sup> + e<sup>-</sup>;<p>

giving an overall equation of:
Pb(NO<small>3</small>;)<small>2</small> + 2H<small>2</small>O =====>> PbO<small>2</small> + 2HNO<small>3</small> + H<small>2</small>

How many electrons in involved in the overall equation.
It has to be two (surely) for the two Hydrogens that are escaping at the cathode?
So I should put "+ 2e<sup>-</sup> " at the end of the overall equation.

My electrific chemistry knowledge is a bit horrific.

TIA,
Dann2



[Edited on 7-4-2009 by dann2]

sparkgap - 6-4-2009 at 19:32

Quote:

  1. H<sub>2</sub>O → OH<sub>ads</sub> + H<sup>+</sup> + e<sup>-</sup>
  2. Pb<sup>2+</sup> + OH<sub>ads</sub> → Pb(OH)<sup>2+</sup>
  3. Pb(OH)<sup>2+</sup> + H<sub>2</sub>O → PbO<sub>2</sub> + 3H<sup>+</sup> + e<sup>-</sup>


(wait, does the "ads" stand for adsorbed?)

Well, if you put all three together and remove common species from the left and right sides, you should be getting...
Quote:

Pb<sup>2+</sup> + 2H<sub>2</sub>O → PbO<sub>2</sub> + 4H<sup>+</sup> + 2e<sup>-</sup>


Here's your kicker: will nitrate ion be reduced under the conditions? ;)

sparky (~_~)

dann2 - 7-4-2009 at 04:47

Thanks for that Sparkgap,

The original stuff can be seen attached.

Two moles of Nitric acid are formed and two moles of H are released (I presume).

I have been kicked by your kicker..........
Nitrites (by reduction of Nitrates I presume) do build up in the plating tank. How much I do not know. Would small Cathode area's help to stop this happening?

Cheers,
Dann2

Attachment: Journal of The Electrochemical Society, 149 (9) C445-C449 (2002).mht (9kB)
This file has been downloaded 975 times


[Edited on 7-4-2009 by dann2]

dann2 - 7-4-2009 at 04:49

Hello,

And another one.
How easy/difficult is it to reduce Nitrites back to Nitrates by bubbling air though a solution of the salts that contain a small amount of Nitrites and a much larger amount of Nitrates. What other conditions would suit the bubbling operation.

Hydrogen Peroxide can be used as well.
Dann 2

Nicodem - 10-4-2009 at 00:17

After a complain that this thread grew too long to be manageable, I decided that it is indeed so and opened a new one. Please continue there.

The old thread is dead, long life to the new thread!
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