Sciencemadness Discussion Board - The short questions thread (3)

The short questions thread (3)

Pages: 1 2 3 .. 8

Nicodem - 29-7-2010 at 23:49

The old thread has become too long so please continue here. Those who never received a proper reply in the old thread, please feel free to repost the question.

Quote:

This is a thread where you can post all those short questions you always wanted to ask but did not consider them worthy of a new thread. You can ask amateur science related questions of all kind as long as you think they are simple enough to be answered by other forum members in a preferably single post.
Consequently, self discipline in avoiding off topic replies is expected.

Nicodem - 30-7-2010 at 01:08

Quote: Originally posted by mewrox99

Anyone know what the PiHKAL cover is meant to be? Is it like a mountain or something

I guess this was one of the last unanswered questions in the last short questions thread. Having some experience with how cover designers think and work, I would guess that being more or less clueless on the specific content of the book, he/she resorted to use stereotypes and therefore used a stylised picture of Mount Sinai as a symbolism for spiritual revelation or cognitive illumination that the pop culture associates with the use of psychedelics. It is just a guess, but when it comes to book covers you can be pretty sure they try to use stereotypes to have a major impact on the potential book buyer.

psychokinetic - 30-7-2010 at 02:12

I was thinking similarly

Could also be a 'peak', as in 'peaking' and the obvious mountain 'peak'. (Along with a potential pronunciation of the book, 'peak-al').

Picric-A - 1-8-2010 at 10:55

Does anybody know the best solvent for 4-Aminobenzoic acid?

I am looking to purify a load of PABA to get rid of the starch impurity so i want to dissolve as much 4-Aminobenzoic acid in as little solvent as possible.

Thanks

JohnWW - 1-8-2010 at 13:30

It should be soluble in water, and in polar organic solvents such as acetone, diethyl ether, ethanol, isopropanol, tetrahydrofuran, dioxane, pyran, etc.

DJF90 - 1-8-2010 at 13:31

Where the hell did the starch come from?! Purification of Laboratory Chemicals says:

Quote:

p-Aminobenzoic acid [150-13-01] M 137.1, m.p. 187-188*C, pK1 2.45, pK2 4.85.
Purified by dissolving in 4-5% aqueous HCl at 50-60°C, decolorising with charcoal and carefully precipitating with 30% Na2CO3 to pH 3.5-4 in the presence of ascorbic acid. It can be crystd from water, EtOH or EtOH/water
mixtures.

Picric-A - 2-8-2010 at 00:13

Quote: Originally posted by DJF90

Where the hell did the starch come from?! Purification of Laboratory Chemicals says:

Quote:

As i said in my post.... i am purifying the 4-aminobenzoic acid from PABA vitamin supplements and so the starch is present as the filler i am guessing. They were in caps so i meerly pulled the cap open and poured out the powder.

It is only very slightly soluble in water, 4.9g/kg so this is out. I think i will try EtOH or MeOH.

solo - 2-8-2010 at 09:17

Will the amine from alpha -trans-methylcinnamaldehyde produce stereoisomers of the l enantiomer or the d, or a recimic product of l-(+/-) , or d-(+/-)....i"ve read from the links below , but still not clear to me......solo

http://en.wikipedia.org/wiki/Diastereomer

http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/ster...

Red/Yellow benzaldehyde?

turd - 2-8-2010 at 13:43

Someone who is me tried to make 3,4,5-trimethoxybenzaldehyde from 5-bromo-vanillin. After Ullmann and methylation an oil was obtained which distilled in a narrow range of 102-104°C - so far so good. But what was really strange: in the beginning of the distillation the distillate was a pale yellow oil which then crystallized to a distinctly red solid. An absolutely weird sight, more so considering that usually it's the other way round: slightly coloured oil giving colourless crystalls. Question: how comes?

My only explanation, although not very convincing, is that in the beginning of the distillation there was a small amount of polymeric stuff bumping into the receiver. When crystallizing, the red impurity was concentrated on the outside of the crystals, making it much more visible than when mixed with the liquid distillate.

IMG_2669.JPG - 106kB

NH4NO3 Odour?

mewrox99 - 3-8-2010 at 00:50

Today was the first time using some AN I bought a month or so ago.

It claims to be Lab grade prills.

For some reason it seems to have a very slight unusual odor that one my friends was also able to notice.

Anyone have any clues what it could be (my guess is small amounts of NH3 or N2O)

psychokinetic - 3-8-2010 at 22:06

NH3 makes sense, mewrox99.

Nicodem - 4-8-2010 at 14:07

Quote: Originally posted by solo

Will the amine from alpha -trans-methylcinnamaldehyde produce stereoisomers of the l enantiomer or the d, or a recimic product of l-(+/-) , or d-(+/-)....i"ve read from the links below , but still not clear to me......solo

http://en.wikipedia.org/wiki/Diastereomer

http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/ster...

You mean the product of reductive amination? The transformation of CHO into CH2NH2? That would give beta-methyl-cinnamylamine. This compound has no chirality so it has no enantiomers (the -CHO is not a prochiral). It can have two diasteromers though, the E and Z isomers. However, which of the isomers would be the main or sole product depends on the reaction conditions and the configuration of the starting compound. If the conditions are not such as to cause E/Z-isomerization then the product will retain the same configuration around the C=C double bond. For example if you start with the E diastereoisomer you end up with the E diastereoisomer of the product. But only if the reaction conditions are not such as to cause E/Z-isomerization, either in the starting or ending compound.

Also, mind that some methods of reductive amination would also cause the reduction of the double bond. That double bond is also quite electrophilic so it can also alkylate amines via the Michael reaction. In short, the reductive amination of that aldehyde is not trivial and only few methods apply.

turd - 5-8-2010 at 01:43

Quote: Originally posted by Nicodem

Also, mind that some methods of reductive amination would also cause the reduction of the double bond.

I guess that's the product solo is talking about? Evidently you would get the racemate, as it is not a diastereomer. The result would be the same for the cis- and trans-isomers of the reactant.

Picric-A - 5-8-2010 at 12:31

This is probably a very trivial question but here we go;

How does one form formate (methanoate) esters?

I know conc H2SO4 reacts with conc formic acid producing carbon monoxides so i guess if one were to try make the ester the usual way (acid + alcohol + conc H2SO4) you would just be left with a bunch of alcohol and CO. So how does one actually go about it? I would specifically like to make ethyl formate (from conc formic acid and methylated spirit)

Lambda-Eyde - 5-8-2010 at 12:36

Any strong acid will catalyze the Fischer esterification reaction. You can use hydrochloric acid instead of sulfuric acid. Using a Dean-Stark trap or a similar device will improve the yield.

DJF90 - 5-8-2010 at 12:46

Formic acid is strong enough to catalyse its own esterification; no added catalyst is necessary.

turd - 5-8-2010 at 13:04

Quote: Originally posted by Picric-A

I would specifically like to make ethyl formate (from conc formic acid and methylated spirit)

Why don't you use Google before asking, lazy bastard.

http://pubs.acs.org/doi/abs/10.1021/ed027p245
I cannot upload the PDF since it's larger than 2MB.

Picric-A - 5-8-2010 at 15:12

Quote: Originally posted by turd

Quote: Originally posted by Picric-A

I would specifically like to make ethyl formate (from conc formic acid and methylated spirit)

lazy bastard.
.

Wow calm down! Bit strong!

Thanks for all the proper informative responses.
Will refluxin formic acid and an excess of ethanol produce said ester? Seems very simple.

[Edited on 5-8-2010 by Picric-A]

mnick12 - 5-8-2010 at 15:20

The yield of uncatalized n-propyl formate with 12hrs of reflux is said to be around 65%, so I would expect ethyl formate yields to be similair. Also couldnt one use p-tsa as a catalyst? I have heard of it being used in some esterifications, and it does not have the same dehydrating properties sulfuric acid has.

panziandi - 5-8-2010 at 15:48

I made methyl formate from formic acid (98%) and methanol. Just refluxed and then distilled the very volatile ester. That was a long time ago can't find my notes.

In fact dehydration of formic acid would require a fair quantity of sulphuric acid, unlike the catalytic quantity used for esterification and sulphuric acid diluted in formic acid doesn't dehydrate it, whereas formic acid diluted in sulphuric acid dehydrates on warming. But as we are saying, catalysis is not required for the esterifcation of alcohols with formic acid, maybe zeolites or magnesium sulphate would help by removing water.

gardenvariety - 5-8-2010 at 15:49

@Nicodem: regarding the Pihkal cover, the mountain looks just as much (if not a bit more) like Mt Diablo, which Shulgin's house/lab overlooks.

For my own question, does metabisulfite hydrolyze into two bisulfite ions or one bisulfite ion and one SO2 molecule?

Picric-A - 5-8-2010 at 16:11

Quote: Originally posted by panziandi

I made methyl formate from formic acid (98%) and methanol. Just refluxed and then distilled the very volatile ester. That was a long time ago can't find my notes.

In fact dehydration of formic acid would require a fair quantity of sulphuric acid, unlike the catalytic quantity used for esterification and sulphuric acid diluted in formic acid doesn't dehydrate it, whereas formic acid diluted in sulphuric acid dehydrates on warming. But as we are saying, catalysis is not required for the esterifcation of alcohols with formic acid, maybe zeolites or magnesium sulphate would help by removing water.

Thanks Panziandi, i will try making both the methyl and ethyl ester tommorow. Although maybe a bit trivial, i have 2.5L of formic acid and havnt realy found a use for it. Also making an ester with such a low boiling point will be good tecnique practice.

turd - 6-8-2010 at 00:36

Quote: Originally posted by Picric-A

Wow calm down! Bit strong!

Apparently not strong enough. First of all you refused to search for yourself, and then instead of fetching the reference I posted detailing the synthesis for students (J. Chem. Education), you continue the begging for being spoon fed. If that's not the behaviour of a lazy bastard, I wonder what is.

Hint: the interesting part is not so much the catalyst but the question whether your product forms azeotropes with the reactants and byproducts.

Picric-A - 6-8-2010 at 04:12

Quote: Originally posted by turd

Quote: Originally posted by Picric-A

Wow calm down! Bit strong!

Hint: the interesting part is not so much the catalyst but the question whether your product forms azeotropes with the reactants and byproducts.

I did not continue 'begging' to be spoon fed, i got my question answered and there was no need for you to lash out.
I have done research and yes i know methyl formate forms an azeotrope with both water and diethyl ether. The ether azeotrope boils at 31.5. Enough research for you?!?! now back off.

sakshaug007 - 6-8-2010 at 11:55

does anyone know if potassium iodide is very soluble in glacial acetic acid? i want to try an electrolysis experiment to obtain methyliodide. the iodide shouldn't oxidize to iodine right? because acetic acid is a weak non-oxidizing acid correct?

thanks

bbartlog - 6-8-2010 at 18:43

I would not expect potassium iodide to be very soluble in glacial acetic acid; in fact I'd expect it to be insoluble or nearly so. However if I understand what you are trying to do correctly, you should be able to electrolyze a concentrated aqueous solution of potassium acetate and potassium iodide and (possibly) obtain some methyliodide. This reaction should be mostly about the anions (well, and current density and concentrations and probably anode composition) so potassium acetate rather than acetic acid should actually be an improvement due to the far greater dissociation. Though I suppose I don't know what your experiment is really about...

bbartlog - 6-8-2010 at 18:49

Speaking of solubility questions: can anyone tell me whether anhydrous CuCl2 is slightly soluble in chloroform? I've observed a slight solubility of the hydrate (CuCl2.2H2O) in the azeotrope (97% CHCl3, 3% H2O) (forms a colorless solution, which suggests something other than Cu++, e.g. HCuCl+ and Cl- or ... something), but I can't find any reference data for chloroform in the places I usually look for solubility data. And I thought what I saw might have been due to the presence of water.

Hamilton - 7-8-2010 at 04:09

easy one, or maybe not at home.

i have got braium sulfate and want to make barium nitrate ou chloride with it, since the solubility of barium sulfate is virtually 0 i don't know how to let the Sr free from the SO4.

how would you do it?

thx

Random - 7-8-2010 at 22:27

I took some zinc from the battery, why doesn't it react with 9% vinegar?

bbartlog - 8-8-2010 at 04:45

Quote:

have got barium sulfate and want to make barium nitrate or chloride

The electronegativities are such that simple metathesis is probably not going to do anything (barring access to rare reagents like cesium chloride), barium being very low and sulfate (in the group electronegativity sense) very high on the scale. And it's not like you can distill out the product to push things in the desired direction. Probably your best angle would be reduction to barium sulfide via heating with carbon, followed by leaching the sulfide and reaction with nitric or hydrochloric acid, but I'm sure the carbothermic reduction requires and/or generates glowing hot temperatures.

bbartlog - 8-8-2010 at 04:52

Quote:

I took some zinc from the battery, why doesn't it react with 9% vinegar?

It should, but slowly slowly. A strip of zinc metal will take days or even weeks to dissolve in dilute room temperature acetic acid. Use zinc powder or heat up the reagents to speed things up.

DJF90 - 8-8-2010 at 16:19

Quote:

Its a Kolbe-type electrolysis experiment, trapping the methyl radical with iodide to yield MeI. Theoretically very clever, but may be difficult in practice. I'm not sure of the solubility of KI in acetic acid but I would expect CRC to tell you or at least give you a vague idea. However the methyl radical is typically nucleophilic in behaviour, leading me to think that elemental Iodine would be preferable. I don't think this will work well though because radical combination (2Me* => EtH) should be much quicker.

Quote:

The electronegativities are such that simple metathesis is probably not going to do anything.

This has very little to do with it. You're not going to metathesise anything if you can't get it into solution first! Solid state may work but again its difficult experimentally. The best route would be the one you post:

Quote:

Probably your best angle would be reduction to barium sulfide via heating with carbon, followed by leaching the sulfide and reaction with nitric or hydrochloric acid, but I'm sure the carbothermic reduction requires and/or generates glowing hot temperatures.

And yep, it is a high temp reaction. Please bear in mind that whilst Barium sulfate is pretty benign (owing to its negligable solubility), other barium salts are pretty toxic. Be careful.

un0me2 - 8-8-2010 at 18:02

Ok, I have a question - we've all (I'm presuming) seen the use of N-(Trifluoroacetyl)-amino acid chlorides/etc. used in Friedel-Crafts Acylation of various phenols, alkoxyphenols, benzene, etc.

The reason for the Trifluoroacetyl Group is that it is easily removed (basic hydrolysis?) and reacts fairly easily and quickly with the amine (especially if one uses the haloacetones - sounds like fun - to do so).

Now, the pKa of Trifluoroacetic acid (TFA) = ~0.00 & that of Trichloroacetic acid (TCA) is ~0.77, so they should be similar insofar as their reactivity, I have also seen several papers where TCA has been used to protect the amino group of various compounds, being equally easily cleaved by base as TFA.

So, having gone through all of that, is there any REAL reason why so many papers use TFA instead of TCA for the F-C Acylation of benzene or alkoxyphenols? Could N-(Trichloroacetyl)-glycine for instance be used instead of the TFA counterpart in the F-C Acylation of 1,4-dimethoxybenzene? I'm searching for it, but cannot find any references, although (given what is said in one paper I've requested) the relative cost of preparing the TCA derivative is a LOT less than the preparation of the TFA derivative (so I'd be expecting industrial scale-up to use it if it worked).

manimal - 9-8-2010 at 13:44

I am curious: is dimethyl sulfide/sulfoxide/sulfone oxidizable to methanol+sulfate ion? The reason I ask is that I read a report in popular media of a possible medical explanation for the death of an individual as in-situ oxidation of dmso to dimethyl sulfate.

Formatik - 9-8-2010 at 20:35

DMSO doesn't convert to dimethyl sulfate in the body, it does convert to the sulfone though ("MSM", which is also sold as a supplement). One person in the world had a magical body that supposedly converted DMSO to dimethyl sulfate. DMSO is used even intravenously as a transporter in medicinal formulations.

Sulfoxide oxidizes to sulfone with strong oxidants, like H2O2, organic peroxides, hydroperoxides. Acidic aq. Cl2 yields the sulfone and methanesulfonyl chloride. Not even ozone (one of the most powerful oxidants there is) apparently oxidizes beyond sulfone. Have a look at this bulletin on page 17: http://www.gaylordchemical.com/bulletins/Bulletin105B/index.... There seem to be a few papers in the net on the oxidation of sulfoxide with ozone.

And dimethyl sulfone is a very stable material, literature reports when it was heated with fuming nitric and fuming sulfuric acid nitrating mixture for several hours, it did not decompose it!

I know myself also that if you store DMSO for a long time (several years), it becomes rancid (worthless medicinally, but still usable for some chemical reactions) and takes on strong dimethyl sulfide odor. I don't recall if I read somewhere, if not some paraformaldehyde might also form in the mix also.

d010060002 - 10-8-2010 at 07:01

The strongest concentration of ammonia that is pure and easy to get was 10%. A lot of the recipes I'm looking at require a more concentrated form. Has anyone concentrated ammonia with good results using an easy method (boiling, freezing at a reasonable temperature).

sakshaug007 - 10-8-2010 at 11:29

Quote:

However if I understand what you are trying to do correctly, you should be able to electrolyze a concentrated aqueous solution of potassium acetate and potassium iodide and (possibly) obtain some methyliodide.

...Though I suppose I don't know what your experiment is really about...

Its a Kolbe-type electrolysis experiment, trapping the methyl radical with iodide to yield MeI. Theoretically very clever, but may be difficult in practice. I'm not sure of the solubility of KI in acetic acid but I would expect CRC to tell you or at least give you a vague idea. However the methyl radical is typically nucleophilic in behaviour, leading me to think that elemental Iodine would be preferable. I don't think this will work well though because radical combination (2Me* => EtH) should be much quicker.

This is exactly right, and the purpose of this experiment in which I seek is to determine the yield and perhaps refine the reaction conditions so as to obtain a primary product of methyliodide versus the ethane Kolbe product. I've done some further research and it looks as though methanol as a solvent using a smooth platinum anode is preferable and of course potassium or sodium acetate would be the better reagent as opposed to acetic acid due to increased dissociation. My thoughts are to mix a saturated solution of potassium acetate and potassium iodide (in the proper molar ratios) in methanol and then attempt electrolysis. I'm not sure about certain details such as stirring or heating the solution while electrolyzing, and also whether the methyliodide would undergo anodic oxidation as well, any suggestions? Does anyone think this will work?

Thanks

[Edited on 10-8-2010 by sakshaug007]

[Edited on 10-8-2010 by sakshaug007]

Nicodem - 10-8-2010 at 13:46

Quote: Originally posted by sakshaug007

Well, there will not be much acetate anion oxidation at the anode until all the iodide gets oxidized to the iodine (or iodate if you do not use a membrane to separate from the cathode). Check their redox potentials! Of course, if an overvoltage of the iodide is achieved then some acetate will start to get oxidized, but not under normal operating conditions. Anyway, once you have all the iodide oxidized to I3-, or more likely only when oxidized to I2, the acetate will start to oxidise to methyl radicals and CO2. At this point methyl iodide could start forming by methyl radicals being oxidized by I3- and/or I2.
I don't know why you think methanol would be a good solvent for this. Check the C-H dissociation energy for methanol and methane. You might be surprised that methyl radicals easily abstract hydrogens from many substances, especially the alpha-hydrogens of alcohols. Water is obviously an ideal solvent for such an experiment. Triiodide is very soluble and iodine soluble enough in water.
The electrodes should be separated by a semipermeable membrane. If so, I don't think you should worry much about MeI reacting at the electrodes. In an non-separated cell it could react at the cathode and being reduced to methane, but I doubt this happens that easily. Chances are it could also react at the anode to form a iodoso compound, but I doubt. It could however succumb in the reaction with radicals, because iodine is easily abstracted from alkyl iodides. but if a methyl radical abstracts iodine from methyl iodide, guess what it gives... However, there could be other radical species involved, like the acetoxyl radical (which is short lived as it decomposes to methyl radical and CO2). Also, if the formation of MeI is faster than its consumption in putative side reactions, once the aq. phase is saturated with MeI (and this happens fast) this will start depositing in its own phase (where it is protected from the electrodes).

PS: You already asked about this in the previous short questions thred. Perhaps you should consider opening a dedicated thread given that you keep bringing this up. Describe the problem and add all the pertaining references and link back to your question in this thread.

sakshaug007 - 10-8-2010 at 16:17

Quote:

I don't know why you think methanol would be a good solvent for this.

I don't want the H2O -> O2(g) + e- + H+(aq) reaction interfering at the anode.

Quote:

PS: You already asked about this in the previous short questions thred. Perhaps you should consider opening a dedicated thread given that you keep bringing this up. Describe the problem and add all the pertaining references and link back to your question in this thread.

I did already bring this up in an existing thread about iodomethane prep. http://www.sciencemadness.org/talk/viewthread.php?tid=12394&... but nobody provided any input!!!! I seem to be the only one interested in this technique which makes me think its not gonna work.

As always thanks for being the only one to thoroughly answer my questions Nicodem.

[Edited on 11-8-2010 by sakshaug007]

solo - 11-8-2010 at 06:51

Is there a study or does someone know, about this compound (N-methyl-1-phenyl-propan-2-amine) where the there exists a double bond between C1 and C2, effects on the activity of the compound....since the compound is derived from alpha-methyl cinnamaldehyde...(perhaps i could have worded my inquiry better), ...as i seem to recall a studies where the variables are changed with the amine group and the benzene ring but never noted any changes on the variables of the alkyl chain..............solo

Nicodem - 11-8-2010 at 10:16

Quote: Originally posted by sakshaug007

I don't want the H2O -> O2(g) + e- + H+(aq) reaction interfering at the anode.

I suggest you to always check redox potentials, species concentrations (including their pKa and activity) and their electrode overpotentials before worrying about such things. Besides, what makes you think methanol would not get oxidized at the anode? Already intuitively you should know that it is more easily oxidized than H2O.

Quote:

You wanted to start a discussion on electrochemistry in a thread dedicated to the classical MeI preparation and you gave no references whatsoever. Obviously that was doomed to end there. Compile all the references you got on the topic and open a thread in the proper way. That is the only way to open a fruitful discussion topic. Just posting fragments here and there, unsupported by any literature or experiments, is obviously not going to lead anywhere.

Quote: Originally posted by solo

Is there a study or does someone know, about this compound (N-methyl-1-phenyl-propan-2-amine) where the there exists a double bond between C1 and C2, effects on the activity of the compound....since the compound is derived from alpha-methyl cinnamaldehyde...(perhaps i could have worded my inquiry better), ...as i seem to recall a studies where the variables are changed with the amine group and the benzene ring but never noted any changes on the variables of the alkyl chain..............solo

Methamphetamine with "a double bond between C1 and C2" is an enamine so it can not be subject to activity studies (assuming you mean biological studies). How about starting to use IUPAC nomenclature? It is always so confusing to figure out what compound you are talking about. Alternatively post the structure. The direct product of reductive amination of alpha-methylcinnamaldehyde is something else, like I already told you above, and still you keep calling it that way. The enamine compound I assume you ask about is N-methyl-1-phenyl-2-aminopropene and is not something that can be derived directly from alpha-methylcinnamaldehyde without a multistep process, so you can't go on calling it a compound derived from alpha-methyl cinnamaldehyde. IUPAC has given us a nomenclature to be used, so use it.

sakshaug007 - 11-8-2010 at 12:15

Quote:

Compile all the references you got on the topic and open a thread in the proper way.

Sadly, nicodem, the only reference I have on this direct subject (i.e. alkyl halide synthesis using NON-Kolbe electrolysis) is written in German, HENCE ALL OF MY QUESTIONS!! I have the reference attached if anybody would care to translate it for me, but until then the only other material I have come across pertains to Kolbe electrolysis itself (radical dimerization). I simply expanded on this topic to apply it to the synthesis of alkyl halides (methyliodide being a simple one) as they are not easy to obtain for at home experiments. I don't have halide acids, I don't have phosphorus, and I don't have thionyl chloride. What I have are acetates/acetic acid and potassium iodide thats it. So please, bare with me or perhaps provide me with some references on this exact area of research.

On the other note, I can't find standard electrode potentials for either methanol or acetate ion. Again, can you point me in the right direction?

Attachment: electrolytic production of CH3I.pdf (776kB)
This file has been downloaded 954 times

turd - 11-8-2010 at 12:15

Well, there surely has been some interest in rigid amphetamine analogs, if that is the question.
Random list of links:
http://www.erowid.org/library/books_online/pihkal/pihkal056....
http://www.google.com/search?q=rigid+amphetamine
http://en.wikipedia.org/wiki/2-Aminotetralin
http://en.wikipedia.org/wiki/2-Aminoindane
http://en.wikipedia.org/wiki/EXP-561

I can see why - if it's rigid there's no question which conformer resides in the receptor. (?)

manimal - 11-8-2010 at 16:23

Quote: Originally posted by Formatik

DMSO doesn't convert to dimethyl sulfate in the body, it does convert to the sulfone though ("MSM", which is also sold as a supplement). One person in the world had a magical body that supposedly converted DMSO to dimethyl sulfate. DMSO is used even intravenously as a transporter in medicinal formulations.

Sulfoxide oxidizes to sulfone with strong oxidants, like H2O2, organic peroxides, hydroperoxides. Acidic aq. Cl2 yields the sulfone and methanesulfonyl chloride. Not even ozone (one of the most powerful oxidants there is) apparently oxidizes beyond sulfone. Have a look at this bulletin on page 17: http://www.gaylordchemical.com/bulletins/Bulletin105B/index.... There seem to be a few papers in the net on the oxidation of sulfoxide with ozone.

And dimethyl sulfone is a very stable material, literature reports when it was heated with fuming nitric and fuming sulfuric acid nitrating mixture for several hours, it did not decompose it!

I know myself also that if you store DMSO for a long time (several years), it becomes rancid (worthless medicinally, but still usable for some chemical reactions) and takes on strong dimethyl sulfide odor. I don't recall if I read somewhere, if not some paraformaldehyde might also form in the mix also.

That was pretty much my analysis. Here's the article I was refering to: http://discovermagazine.com/1995/apr/analysisofatoxic493

Formula409 - 11-8-2010 at 18:23

Has anyone had any experience cleaning up the formaldehyde you can buy OTC which has a coloured dye in it (normally for camping toilets). I'm thinking just adding a bit of activated charcoal should be sufficient.

Formula409.

bbartlog - 11-8-2010 at 18:32

Quote:

I have the reference attached if anybody would care to translate it for me

Not going to translate the whole thing, but some points I gleaned:

- they claim great difficulty in condensing the methyl iodide due to it being diluted and (I guess) entrained by the ethane and CO2 which are produced in greater quantity. So for their analysis they lead the gas into an alcoholic solution of dimethylaniline, which apparently reacts with the methyliodide. You likely wouldn't want to take this route even if you had dimethylaniline handy, but it sounds like you'll have to worry about cooling. Or condensing. Obtaining the methyliodide as a separate phase may not be practical. Given that it boils at 42C I'm a little surprised at the claimed difficulty, but I guess it's very volatile.

- in general they appear to have been interested in elucidating the mechanism rather than finding some way of maximizing production, which makes their paper a bit tangential to your purpose.

- they used a cell with a clay membrane, platinum anode, and lead or copper cathode.

- for anode solution they used both potassium acetate and elemental iodine (65g potassium acetate, 160cc water, 10g iodine), and sodium acetate plus potassium iodide (40g and 10g plus 100cc water). One trial was at 2.5 amps for 6 hours, the other at 1.6 amps for 12 hours. In both cases it sounds like they obtained less than a gram of iodine (as methyl iodide) on analysis.

- unfortunately, no information is given on the anode dimensions or the applied voltage.

mewrox99 - 14-8-2010 at 20:37

Is dry calcium carbide flammable

2 CaC2 + 5 O2 = 2 CaO + 4 CO2

stygian - 16-8-2010 at 19:02

Will Na2C2 hydrolyse in alcohols?

12332123 - 21-8-2010 at 05:36

Is it possible to synthesise alkoxides by electrolysis of either halides or hydroxides in alcohol?

Nicodem - 21-8-2010 at 13:49

Quote: Originally posted by stygian

Will Na2C2 hydrolyse in alcohols?

The pKa of most alcohols is rougly 15-16 while that of acetylene is about 25. That pKa values are for water, but what is importance is the difference, and ten magnitudes is one hell of a difference. BTW, it took me exactly seven seconds of googling (seriously, I counted!).

Quote: Originally posted by 12332123

Is it possible to synthesise alkoxides by electrolysis of either halides or hydroxides in alcohol?

If you use an alkali salts for the electrolyte, then yes. There are also a few other electrolytes that reduce the protons of the alcohol instead of their cationic part, but I won't bother searching given you don't bother specifying which alcohol and alkoxide of what you are interested in. For example, if the metal cation is easily reducible (e.g. Cu2+, Pb2+...) then it will simply get reduced and deposit on the cathode instead of forming an alkoxide. Of course, the electrodes must be separated by a semi-permeable membrane and the alcohol must be anhydrous. You might find it a problem to isolate the metal alkoxide from the used electrolyte though. A long enough electrolysis to assure complete oxidation of the original anion of the electrolyte could remove most of it, but that does not sound that easy unless the anion has a very low redox potential. Though for simple alcohols like ethanol, propanol, butanol, etc. and their alkali alkoxides there are a few considerably simpler and faster methods to prepare them (UTFSE for more info).

dann2 - 28-8-2010 at 14:46

Hello,

I have SnCl4:5H2O (white solid) and I want to get to SnCl4 (clear liquid).

Can this be done?

TIA,
Dann2

chemrox - 28-8-2010 at 22:31

If it's soluble in toluene you could reflux it with a dean-stark until 5 eq of water are collected and distill to separate. The latter step under N2

kuro96inlaila - 31-8-2010 at 19:24

Hello,
Anyone have information(density,melting point,solubility etc) on silver(I) dichromate?

jokull - 31-8-2010 at 20:15

Hope these data would be helpful.

Silver dichromate:

Formula: Ag2Cr2O7
Solubility in water (15ºC): 0.083g/L
Density: 4.77 g/cm3

kuro96inlaila - 1-9-2010 at 15:27

Thanks.It do help.but I would love to receive any other data like its toxicity and uses.....

hector2000 - 2-9-2010 at 01:40

We can make KİO3 by Oxidation of Kİ using KClO3 or KMnO4?

Picric-A - 2-9-2010 at 09:13

Does anybody know of an OTC source of methyl acrylate?

I am attempting to perform a Michael addition and i need this compound. If not, does anybody know of a synthesis for it that can be performed at lab scale?

crazyboy - 5-9-2010 at 17:22

I am looking for a vacuum controller so I can accurately measure and control the vacuum of my pump. There is a vacuum controller at the lab I worked at which connected the vacuum to the rotovap. It was small about 6" by 8" by 2". It had an simple digital display and several buttons, two arrows one up one down, two set point buttons, a run/stop button, a max vac button and a vent button.

The vacuum connected directly to the back of the controller and then went to the rotovap. I have looked on eBay but despite having several pages of vacuum controllers none have a vacuum port at the back, only electronic ports. Any advice? I can't have a computer in my lab.

rrkss - 12-9-2010 at 14:50

I use the jkem dvr to regulate my vacuum. Bought one off of ebay. It has two ports on the back. A vacuum in and the out port. It connects to the vacuum source and cycles a solenoid valve to control the vacuum precisely based on the setting.

woelen - 16-9-2010 at 01:26

Making chloro-phenols can be done by simply bubbling Cl2 through a solution of phenol in water. I am only interested in making 2,4,6 trichlorophenol and don't want the others. Is this last compound the main product, or even the only product when Cl2 is passed through an aqueous solution of phenol?

Another more convenient method of making 2,4,6 trichlorophenol could be adding a solution of NaOCl or Ca(OCl)2 to a solution of phenol and then adding slight excess of HCl. The chlorine, formed in the solution, then immediately is used up by making chlorophenol. I tried this and the clear colorless solution becomes turbid on adding HCl and after a few minutes a big oily yellow/orange blob collects at the bottom. I used excess hypochlorite and after the reaction there is a faint smell of chlorine and some other somewhat "burned phenolic" smell. This smell is VERY persistent. Is this oily blob the desired 2,4,6 trichlorophenol?

[Edited on 16-9-10 by woelen]

Jor - 16-9-2010 at 06:10

You want the 2,4,6-trichlorophenol to make TCPO, the oxalate ester? If so, I would extract with DCM, wash the DCM with water, and distill off the DCM. This should give reasonably pure 2,4,6-trichlorophenol if this is your purpose, and some di- or mono-chlorinated impurities shouldn't really be a problem for the chemiluminescent experiment purpose.
You could use other solvents, but because it's so smelly I would want a solvent that boils of at an as low temperature as possible to prevent excessive evaporation of the compound.

I think it is trichlorphenol, because I have read somewhere that it indeed has a very persistent smell, and can be detected at around 1ppb. A reason why I would never make it, because I live near neighbours. And AFAIK, phenol and aniline are so reactive that they are directly chlorinated to the trichloro-compound with chlorine.

Maybe you could test it's acidity. It have not searched it, but it should be a quite acidic compound.

[Edited on 16-9-2010 by Jor]

Panache - 16-9-2010 at 11:27

is there anyway of estimating the volume of an underground tank without emptying or filling it? The only viewing point i have for it is the standard 3 inch petroleum pipe coming up out of the floor, its full of diesel and is about 2 metres deep but have no idea of the girth.

manimal - 18-9-2010 at 23:30

I am going to try making a sealed stirrer out of a blender blade unit, tentatively by placing the housing over an appropriately-sized pvc pipe and tightening a hollowed-out endcap over it to seal the gasket, then by manipulating the blades so as to affix a stirring shaft.

I guess my question is, can anyone think of a better way in which to make a diy sealed stirrer?

Panache - 19-9-2010 at 16:38

Quote: Originally posted by manimal

I am going to try making a sealed stirrer out of a blender blade unit, tentatively by placing the housing over an appropriately-sized pvc pipe and tightening a hollowed-out endcap over it to seal the gasket, then by manipulating the blades so as to affix a stirring shaft.

I guess my question is, can anyone think of a better way in which to make a diy sealed stirrer?

i would answer differently depending upon whether your system in under vacuum or pressure and if it is imperative for nothing to ingress or exit.

However i have designed and made and used over a couple of years now a 'sealed' stirrer that can be used for either with decent efficacy. There are some limitations however there are some ways those also.
I need to draw a diagram, i'll edit this post later after i have drawn and scanned it.

Formula409 - 19-9-2010 at 19:45

Does anyone have any solubility data for n,n-dimethyltryptophan? Basically I'm looking for a way to separate it from tryptophan. I'm thinking the addition of the methyl groups should make it even less soluble in water (tryptophan is 11.4g/L (http://www.drugs.com/mmx/l-trytophan.html)).

Cheers!

stygian - 20-9-2010 at 12:49

Would phenyl-2-propanol-1-one if formed possibly rearrange to phenylacetylcarbinol?

ayush - 25-9-2010 at 06:29

Will any one write what is the software used to generate this image?

1281371269 - 25-9-2010 at 07:29

Can I extract Ethane1,2,diol from antifreeze safely by distillation?

crazyboy - 25-9-2010 at 07:41

Quote: Originally posted by ayush

Will any one write what is the software used to generate this image?

I don't know the exact program used to create those images but a program called Mercury is apparently quite good and free.

Search for "Mercury - Crystal Structure Visualisation and Exploration Made Easy" It should be the first result.

ayush - 25-9-2010 at 09:13

Thank you crazyboy.

entropy51 - 25-9-2010 at 11:11

Quote: Originally posted by Mossydie

Can I extract Ethane1,2,diol from antifreeze safely by distillation?

I don't know if you can or not, but I have done it without problems. I used vacuum distillation since the BP is nearly 200 C at atmospheric pressure, which is quite a bit above the flash point. I used magnetic stirring because a solid separates out of the antifreeze that I distilled. Don't pour the waste out where animals could drink it. Dogs like it because it tastes sweet, but ingestion causes kidney failure.

1281371269 - 25-9-2010 at 11:42

Would you consider it too dangerous at 1atm then? I have a mantle, so there won't be any sparks / flames anywhere around.

not_important - 25-9-2010 at 12:11

You may get some decomposition depending on the additives. You can try it on a smaller scale with the brand you're planning to use; worth doing anyhow just to determine the rough amount of di- and tri ethylene glycols and other low volatility components.

Quote:

is there anyway of estimating the volume of an underground tank without emptying or filling it?

There's a method that's been used for spacecraft and such; basically you use a piston - a loudspeaker even - to change the air pressure, the pressure change in the system vs the piston excursion allows you to calculate the void space (air). To get the total volume you need to know the tank geometry, or do this when the tank is nearly empty, or do it at several differing depths of fill and estimate the geometry from the void/depth data.

Another possibility is to stick a bright LED and a video camera (webcam grade) in something like a large test tube, and lower that into the tank to image the interior. Only give you a rough idea, but perhaps good enough for corporation work.

entropy51 - 25-9-2010 at 14:10

Quote: Originally posted by Mossydie

Would you consider it too dangerous at 1atm then? I have a mantle, so there won't be any sparks / flames anywhere around.

I don't like to heat distillations to high temperatures if it can be avoided. I haven't looked up the autoignition temperature, but I have a feeling that it's too close for (my) comfort. If the glassware should break when the vapors are above the autoignition temperature it will flash without an ignition source.

If distilled at atmospheric pressure, don't run cooling water through the condenser, lest it break. Better to cool it with the air from a small fan.

Better to set yourself up for vacuum distillations.

Edit: After reading bbartlog's post, I looked up the autoignition temperature and found that it's 400 C. So perhaps an atmospheric distillation is safe after all. But I still just don't like to heat chemicals to a higher temperature than I have to. I'm getting too old to enjoy the excitement of a fire in the basement!

[Edited on 26-9-2010 by entropy51]

[Edited on 26-9-2010 by entropy51]

bbartlog - 25-9-2010 at 17:30

Quote:

Would you consider it too dangerous at 1atm then? I have a mantle, so there won't be any sparks / flames anywhere around.

I've distilled ethylene glycol from commercial antifreeze at regular atmospheric pressure a number of times. Seems to work fine. I'm not sure why the flash point should be of particular interest; I mean people distill ethanol on a regular basis and it has a flash point below room temperature. In fact, based on a messy incident I had back before I even had (all) glass equipment (ethylene glycol and HCl spouting out of a flask which was above an open flame) I'd say that ethylene glycol is simply not very flammable at all.

But by all means distill at reduced pressure if you have the gear for it. Maintaining 197C (at least in my case) required using a fairly short path and insulating the neck of the flask. Lower temperature seems preferable if it's convenient.

[Edited on 26-9-2010 by bbartlog]

bahamuth - 29-9-2010 at 13:45

Blue compound from nitration attempt of isopentanol or n-pentanol.

Several years ago I tried to nitrate isopentanol, or n-pentanol with nitric acid and sulfuric acid and remember I got a beautiful blue compound in solution, no attempts were done to try to isolate this.

Of what i think I remember I had the nitric acid mixed with the pentanol with stirring and added sulfuric acid WAY to quick, scorching and boiling of the mix followed, and then it took on a deep but clear blue color.

IIRC I found something on the net then of a blue short chain nitroso compound that maybe was the culprit that formed under such conditions.

So my question is, do anyone know what it maybe was?

stygian - 29-9-2010 at 17:22

Which was it, iso or n-pentanol? I do recall reading of some secondary pseudonitrosite or nitroles or something that are blue.

The pseudo-nitrols, RR':C(NO)(NO 2), may be obtained by the action of nitrous acid on the secondary nitroparaffins; by the action of silver nitrite on such bromnitrosoparaffins as contain the bromine and the nitroso group united to the same carbon atom (0. Piloty, Ber., 1902, 35, p. 3 0 93); and by the action of nitrogen peroxide on ethereal solutions of ketoximes (R. Scholl, Ber., 1888, 21, p. 508; G. Born, Ber. 1896, 2 9, p. 93). They exhibit an intense blue colour when in the liquid condition or dissolved in alkali and possess a very sharp smell. On oxidation with chromic acid they yield dinitrohydrocarbons, and on reduction with hydroxylamine (in alkaline solution) or with potassium sulphydrate give ketoximes, RR': C: NOH (R. Scholl and K. Landsteiner, Ber., 1896, 29, p. 87).

Perhaps some nitrous acid was formed to cause this reaction with the ketone also formed?

[Edited on 30-9-2010 by stygian]

Rogeryermaw - 29-9-2010 at 18:45

does freezing your benzene negatively affect it? i store most of my volatiles in a refrigerator and i guess it was a bit too low...

bahamuth - 30-9-2010 at 10:28

Quote:

Originally posted by stygian

Which was it, iso or n-pentanol? I do recall reading of some secondary pseudonitrosite or nitroles or something that are blue.

Think it was iso-pentanol, but as I vaguely mentioned earlier, do not really remember...

Never tried to reproduce that peculiar reaction but could give it a try, as I now have an idea of what it might be...

1281371269 - 3-10-2010 at 08:29

I'm planning a titration against molar H2SO4 to find out the concentration of an ammonia solution.
According to Wiki, Phenolphthalein should be used for strong acid / strong base neutralisations. Both Phenolphthalein and an indicator called Neutral Red are available from eBay. Whilst the former changes colour around pH 8.3-10, Neutral Red changes between 6.8 and 8, suggesting it would be more accurate.
Which would be better to use?

not_important - 3-10-2010 at 11:15

H2SO4 has two protons, the titration end point will depend on if you intent to titrate to bisulfate or sulfate. Ammonium sulfate when dilute has a pH around 5.5 or so, the bisulfate obviously is lower.

crazyboy - 3-10-2010 at 16:39

Does anyone have experience using aqueous sodium metabisulfite to remove free iodine from compounds? I know sodium thiosulfate is often used for this purpose but I don't have any so I was wondering if I could use sodium metabisulfite. I tested it by several milliliters of aqueous sodium metabisulfite to some methanol with a pellet of iodine and it completely neutralized the color. Is there any reason why it would be a poor choice?

On a separate note, can welding gas be used as an inert gas in reactions of it is comprised of 75% argon 25% CO2? I am most concerned about the CO2, I realize I *might* be able to scrub it with NaOH or KOH but is this necessary?

[Edited on 4-10-2010 by crazyboy]

Oxygen Tanks

smuv - 3-10-2010 at 21:53

Will welding gas suppliers fill rusty oxygen tanks? Pic of tanks in question attached.

@crazyboy: metabisulfite will work fine. About the inert gas, CO2 really limits the reactions that you can use the inert gas for. It seems you have overcome the hurdle of obtaining a regulator and tank, can you not find a place that will fill with N2? As i recall both use the same regulator.

pic.jpg - 7kB

[Edited on 10-4-2010 by smuv]

Rogeryermaw - 4-10-2010 at 00:06

Quote: Originally posted by smuv

Will welding gas suppliers fill rusty oxygen tanks? Pic of tanks in question attached.
[Edited on 10-4-2010 by smuv]

in my experience they will fill a tank as long as it is certified and within the certification date barring any obvious physical damage.

watson.fawkes - 4-10-2010 at 07:49

Quote: Originally posted by Rogeryermaw

in my experience they will fill a tank as long as it is certified and within the certification date barring any obvious physical damage.

Agreed. I want to add that there do exist companies that will do hydrostatic tests on privately-owned cylinders. Mostly this is done in-house by the big gas suppliers (e.g. Praxair), but they don't have a completely exclusive lock on that service.

crazyboy - 4-10-2010 at 15:53

Quote: Originally posted by smuv

It seems you have overcome the hurdle of obtaining a regulator and tank, can you not find a place that will fill with N2? As i recall both use the same regulator.
[Edited on 10-4-2010 by smuv]

Unfortunately, no. The tank contains shielding gas for welding, and while I could get it filled with pure nitrogen or argon, I wouldn't be able to use it for welding.

Sedit - 4-10-2010 at 16:55

Sorry this isn't chemistry related but I didn't know where else to place it and did not want to start a new thread.

Does anyone have a good link for old newspapers?

I wish to read a few from the day of my birth but haven't come across anyplace that has them online yet, so if you know of any online microfilm so to speak could you shoot me a link please, Thanks.

mewrox99 - 5-10-2010 at 04:07

Silly question. Is it possible to make hexamine perchlorate from hexamine and NH4ClO4

smuv - 5-10-2010 at 07:40

Quote: Originally posted by crazyboy

Unfortunately, no. The tank contains shielding gas for welding, and while I could get it filled with pure nitrogen or argon, I wouldn't be able to use it for welding.

I know welding gas suppliers sell argon/helium mixes. If you can use this for the welding you do and chemistry, it would be a good compromise, but filling the tank will be a bit more expensive.

mnick12 - 5-10-2010 at 11:52

Does anyone know if 5-hydroxyvanillin has a distinct odor?
The reason why I ask is I attempted the hydrolysis of 5-bromovanillin and the end product which is crystallizing from ethyl acetate has an odor similar to vanillin (which it is contaminated with).

So after crystallizing from ethyl acetate the suppose-ed hydroxyvanillin was removed from the flask and placed into a small glass dish, the flask was rinsed with 2x5ml of dcm which was then poured into the glass dish. Only some of the material dissolved in the dcm. Next I waited for the dcm to evaporate, and I was rewarded with two distinct crystalline masses. The first being flaky cream colored crystals which seemed to be less soluble in dcm, and the second being sticky yellow crystals which were stuck to the side of the dish. The odor is a pleasant one kind of like a cross between vanilla and bubblegum.The mass of the two combine was 1.83gr from 2.00gr of bromovanillin

thanks

[Edited on 6-10-2010 by mnick12]

Pope - 6-10-2010 at 21:19

How would one approach trying to separate an azeotropic mixture of Toluene/Methyl Ethyl Ketone, my goal is to completely separate and purify the existing Toluene.

The components exist as - 430g Toluene: 410g MEK.

Thank you for your assistance, I have tried the search engine with little success.

woelen - 6-10-2010 at 22:45

Quote: Originally posted by mewrox99

Silly question. Is it possible to make hexamine perchlorate from hexamine and NH4ClO4

Is hexamine a stronger base than NH3? I don't think so. So, I do not expect this to work. The only thing is that NH3 is a gas and that may be advantageous for you, but I definitely do not expect a smooth reaction in which hexamine reacts with NH4ClO4 while NH3 escapes from the mix.

Something you could try is making the sulfate salt of hexamine (does that work, I don't know, try it) and mix a solution of that with a solution of NH4ClO4. Many perchlorates of large cations have low solubility and could be precipitated out.

All of what I have written above is not well-proven, it is just some educated guessing, but you could give it a try.

mewrox99 - 7-10-2010 at 00:59

Found an ancient thread from 2002. This guy mixes HCl with NH4ClO4 and reacts hexamine with it. A precipitate forms that is highly flammable.

mnick12 - 7-10-2010 at 11:42

Yeah I have made the suppose-ed hexamine perchlorate, when I made it was a fluffy white material which upon contact with fire burned with a luminous blue-ish flame. The smoke was rather nasty lots HCl, and kinda stinky. I did this a few years ago and if I remember correctly I used NaClO4 and HCl to generate perchloric acid, but I'm not sure it could have been NH4ClO4 and HNO3. I also made urea perchlorate in a similar manner its kinda boring stuff, in my opionion it is about as interesting as urea nitrate.

Sedit - 7-10-2010 at 16:13

Quote: Originally posted by Pope

How would one approach trying to separate an azeotropic mixture of Toluene/Methyl Ethyl Ketone, my goal is to completely separate and purify the existing Toluene.

The components exist as - 430g Toluene: 410g MEK.

Thank you for your assistance, I have tried the search engine with little success.

Sodium Metabisulfite will form a complex precipitating the MEK as an adduct. Repeated washes of the MeBn with Bisulfite solution should clear up your problem.

dann2 - 7-10-2010 at 17:45

Hello,

Can anyone tell me how to make Aqua Regia.
I have been unable to find a definate answer with some saying it is 3:1 molar or 4:1 by weigh of 'concentrated' HCl and 'concentrate' Nitric acid.
They never specify what concentrated Nitric means and I gather it means the 70% stuff and not the 100% stuff. Also when moles of 'concentrate' HCl are being referred to does it mean actual moles of HCl (one mole = 36.45 grams) and you must use the appropriate amount of 36% solution-in-water of HCl or do you go with one mole of the actual solution in water (molecular weight I will have to figure out.

What ratio (by weight or volume) of 70% Nitric acid 37% Hydrochloric acid?
Or do you need 100% Nitric acid?

TIA,
Dann2

Picric-A - 8-10-2010 at 00:43

Quote: Originally posted by dann2

Hello,

Can anyone tell me how to make Aqua Regia.
I have been unable to find a definate answer with some saying it is 3:1 molar or 4:1 by weigh of 'concentrated' HCl and 'concentrate' Nitric acid.
They never specify what concentrated Nitric means and I gather it means the 70% stuff and not the 100% stuff. Also when moles of 'concentrate' HCl are being referred to does it mean actual moles of HCl (one mole = 36.45 grams) and you must use the appropriate amount of 36% solution-in-water of HCl or do you go with one mole of the actual solution in water (molecular weight I will have to figure out.

What ratio (by weight or volume) of 70% Nitric acid 37% Hydrochloric acid?
Or do you need 100% Nitric acid?

TIA,
Dann2

Aqua Regia does not need to be made perfectly, any mix of conc nitric and hydrochloric acids will dissolve Au and Pt to some extent.

The perfect mix, as you say however, is 3:1, conc (36%) HCl to conc Nitric (70%) and the ratio is done by the solution in water, nothing to do with moles, eg, 30ml HCl:10ml Nitric.

[Edited on 8-10-2010 by Picric-A]

dann2 - 8-10-2010 at 11:53

Thanks for reply.
I am dissolving Bismuth to make BiCl3:2H2O
Simply dissolve Bi metal and evaporate to obtain BiCl3.
Is it fair to assum that all Nitrate will evaporate (go away) and leave the Chloride?
Thats why I though I might need to get the ratio (of HCl to Nitric) close to correct.

Dann2

Picric-A - 8-10-2010 at 22:59

Quote: Originally posted by dann2

Thanks for reply.
I am dissolving Bismuth to make BiCl3:2H2O
Simply dissolve Bi metal and evaporate to obtain BiCl3.
Is it fair to assum that all Nitrate will evaporate (go away) and leave the Chloride?
Thats why I though I might need to get the ratio (of HCl to Nitric) close to correct.

Dann2

As long as you keep to the 3:1 ratio (maybe add a little bit extra HCl to be sure) no nitrate will form as it is completly consumed according to;
HNO3 + 3 HCl → NOCl + 2 H2O + Cl2
You will only be left with the chloride.

solo - 13-10-2010 at 06:25

what is this compound called without the double bond between c1-c2

alpha-methyl cinnamaldehyde

Uploaded with ImageShack.us

[Edited on 13-10-2010 by solo]

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