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Author: Subject: The short questions thread (3)
Nicodem
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 Quote: Originally posted by solo In the reduction of carbamates.....in the article, Preparation of Methyl N-Substituted Carbamates from Amides through N-Chloroamides Gene A. Hiegel and Tyrone J. Hogenauer http://www.chembbs.com.cn/bbs/File/UserFiles/UpLoad/20090924... .....the Chlorinated amine is treated with sodium methoxide to make the methyl carbamate..... my question, i have sodium ethoxide, can this be used and have the same result?

Many bases can be used for the Hofmann rearrangement. Alkali hydroxides are most commonly used when the isocyanate is to be quenched directly to the amine. If carbamates are the target, then you would want to avoid bases that could potentially give mixtures of products. For example, if you use sodium ethoxide in methanol, then the main product should be the methyl carbamate, but there will be some contamination with ethyl carbamate as well. Depending on the nature of the product (is it recrystallizable?) it might be tedious to remove this side product. In such a case, it would be better to use sodium hydroxide in methanol. The main product would still be the methyl carbamate, but there would also be some contamination with the amine (and potentially of the N,N'-dialkylurea). However, the amine can easily be washed away during the extractive work-up.

Another solution is to prepare sodium methoxide from your sodium ethoxide by dissolving it in absolute methanol and rotavaping to dryness.

Perhaps the simplest solution of them all is to prepare sodium methoxide from sodium or sodium hydroxide. There are methods described on the forum. Ullmann posted what is probably the simplest of them all, which is just drying the methanolic NaOH with molecular sieves.

…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)

Boron Trioxide
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I have a quick question:

I understand that mixing chlorate containing compositions with sulphur is dangerous, though is there any danger with mixing it with a sulfate, namely copper sulfate, for a blue flame?

[Edited on 13-7-2012 by Boron Trioxide]
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 Quote: Originally posted by Boron Trioxide I have a quick question: I understand that mixing chlorate containing compositions with sulphur is dangerous, though is there any danger with mixing it with a sulfate, namely copper sulfate, for a blue flame? [Edited on 13-7-2012 by Boron Trioxide]

Sulfate is not a reducing agent like elemental sulfur. It would not burn at all by itself. Chlorate is a very unstable oxidiser, so pretty much any reducer/chlorate mixture will be dangerously unstable. Copper chloride mixtures would probably be the best blue colourant.

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Boron Trioxide
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Thanks for the response, though another question just occurred to me.

For pyrotechnics would it be possible to use Potassium peroxymonosulfate as an oxidizer? Wikipedia doesnt list its decomposition temperature, while mentioning that is an oxidizer.

Possible Decomposition:
KHSO5 => KOH + SO2 + O2
This seems a bit unlikely but I am unsure of any other decomposition routes more likely.

Sorry if this question is ridiculous in some way, I checked nobody has mentioned this before.
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 Quote: Originally posted by Boron Trioxide Thanks for the response, though another question just occurred to me. For pyrotechnics would it be possible to use Potassium peroxymonosulfate as an oxidizer? Wikipedia doesnt list its decomposition temperature, while mentioning that is an oxidizer. Possible Decomposition: KHSO5 => KOH + SO2 + O2 This seems a bit unlikely but I am unsure of any other decomposition routes more likely. Sorry if this question is ridiculous in some way, I checked nobody has mentioned this before.

It is a good question. I'm guessing peroxymonosulfate has approximately the same oxidising power as peroxydisulfate. A mixture of Ammonium peroxydisulfate and powdered aluminium (by my own experimentation) did not ignite easily but had a similar burn rate to very good potassium nitrate/sugar (and a lot hotter). Try and make the tetraamminecopper(II) salt (if you need help with this just email or U2U me). How pure is your peroxymonosulfate? Are you using "Oxone"?

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Boron Trioxide
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Excellent, I was hoping that potassium peroxymonosulfate would be as good as potassium nitrate, because in Canada nitrates are becoming hard to get, while I have seen lots of peroxymonosulfate available. I do not have any at the moment but I plan to get some soon, I haven't seen Oxone brand around but am sure I could find it.

Right now I have limited lab capability so I doubt I could make the tetraamminecopper(II) salt, though thanks for the advice.

plante1999
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 Quote: Originally posted by Boron Trioxide because in Canada nitrates are becoming hard to get.

You mean near impossible! Thats the reason why I bought a platinum anode for electrolysis, with this I can make oxidizer that I can use to dissolve copper or as general oxidizers in the lab. I'm not in energetic chemistry, but I suppose one could use them for pyrotechnics.

Chemistry Alchemist
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What colour is Lead's tarnish? i was in K-mart today in the fishing section and saw a sinker pretty tarnished, but it was a rainbow colour, most sinkers are lead but some are made from Bismuth... would this sinker i saw the a bismuth sinker if it has a rainbow colour effect starting to form?

99chemicals
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 Quote: Originally posted by Chemistry Alchemist What colour is Lead's tarnish? i was in K-mart today in the fishing section and saw a sinker pretty tarnished, but it was a rainbow colour, most sinkers are lead but some are made from Bismuth... would this sinker i saw the a bismuth sinker if it has a rainbow colour effect starting to form?

Yes the sinker is definitively bismuth. Bismuth has a very pretty oxide layer that can be many colors. The sinker in question should have been almost twice as expensive as lead too.

http://en.wikipedia.org/wiki/File:Bi-crystal.jpg

I am pretty sure that the thickness of the oxide is what affects the color

Do you have mole problems? If so, call Avogadro at 602-1023

Sublimatus
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Yeah, the color is mostly dependent on the thickness, based on my experience.

If you melt up some bismuth and skim off the oxidized skin on the top, you'll see a very silvery liquid reminiscent of mercury. Then just let it go, and you'll see a progression through a series of interference colors over the next couple of seconds until the oxide gets so thick that it's just a crusty matte gray.
Chemistry Alchemist
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Is the sinker safe to melt in a pot?

Chemistry Alchemist
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Also, is there ways to test if it is Bismuth with acids? Any acids dissolve lead but not bismuth... HCl for example?

99chemicals
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 Quote: Originally posted by Chemistry Alchemist Is the sinker safe to melt in a pot?

Yes you can safely melt bismuth at home with out any worry of it being toxic. Gloves are highly recommended when working with molten bismuth. Even though it has a low melting point it still hurts like hell when it hits your skin.

As to your other question conc. nitric acid should dissolve bismuth with heating.

Do you have mole problems? If so, call Avogadro at 602-1023

liquidlightning
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Can you heat glassware in a microwave oven? (obviously not one that will ever again be used for food)
Boron Trioxide
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Could someone please tell me what I am doing wrong with my chlorate cell?

I am using a large carbon anode, that when using a current density of 35 mA per cm2 can be run at 1.7 amps, so I ran it at this amperage for 3 hours in potassium chloride, with a stainless steel spoon as my cathode. No chlorate was produced, only a small amount of iron hydroxide. So I switched the spoon for a tube of chrome, this also didn't work at all.

The final reaction mixture smelled slightly of chlorine, a very slight yellow with a pH of 9-10.

Did I just not run it long enough or is 1.7 amps too low to produce chlorates after any amount of time?
plante1999
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In the small scale electrosynthesis of it you will find the info you need.

http://hclo3chem.weebly.com/pdf.html

Hop this helped you, and after the reading please respond to the poll on the pdf tab.

99chemicals
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 Quote: Originally posted by liquidlightning Can you heat glassware in a microwave oven? (obviously not one that will ever again be used for food)

I have done that several times. As far as I know it is safe. If you are doing that use a beaker or a thin walled piece of glass.

Don't use thick glass(like votive candle holders) as there is a higher likelihood that it will crack.

Do you have mole problems? If so, call Avogadro at 602-1023

zoombafu
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Does anyone know where an individual can get suba seals (rubber septa) ? Ive been looking on ebay for a while as well as various other hobby chemistry sites and can't find them.

mnick12
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IIRC they quite available, http://www.ebay.com/sch/i.html?_nkw=rubber+septum&_sacat...

Which ebay were you looking at?
zoombafu
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 Quote: Originally posted by mnick12 IIRC they quite available, http://www.ebay.com/sch/i.html?_nkw=rubber+septum&_sacat... Which ebay were you looking at?

Well hot damn, I must be suffering from a mental deficiency and smoking dope and hitting my head with a hammer wile sniffing glue. Thanks

Lambda-Eyde
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Laboy glass also has rubber septa, as well as cannulae and needles.

This just in: 95,5 % of the world population lives outside the USA
You should really listen to ABBA
liquidlightning
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How well do retorts work for distilling liquids? Would it be feasible to distill about about 500ml hydrochloric acid (not including added water to dilute to 20%) in a few hours with one, or do you have to distill ultra slow due to the air cooled condenser?

[Edited on 2-8-2012 by liquidlightning]

[Edited on 2-8-2012 by liquidlightning]
solo
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C6H5CH2CH(CHNH2)CH3.....not sure what this is called, my guess would be 2-dimethylamine-1- phenylpropane.... Solo

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manimal
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My question has to do with 37% formalin inhibited with methanol. How difficult would it to be to remove this methanol to prevent side reactions, e.g. by heating on a water bath? How much of the methanol exists in equilibrium as methylal, and would boiling this off result in a substantial loss of formaldehyde?
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I've personally rotavapped some formalin to the point that it becomes much more viscous and begins to cloud. It's poured out on a hard surface in a thin layer (read: pyrex baking dish) and as it cools it solidifies. Scraping this solid into flakes or "shavings" will allow proper desiccation over anhydrous calcium chloride. The solid is removed and ground in a pestle and mortar before being returned to the desicator. The fine, dry powder that is obtained is paraformaldehyde, (CH2O)n.
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 Sciencemadness Discussion Board » Fundamentals » Miscellaneous » The short questions thread (3) Select A Forum Fundamentals   » Chemistry in General   » Organic Chemistry   » Reagents and Apparatus Acquisition   » Beginnings   » Responsible Practices   » Miscellaneous   » The Wiki Special topics   » Technochemistry   » Energetic Materials   » Biochemistry   » Radiochemistry   » Computational Models and Techniques   » Prepublication Non-chemistry   » Forum Matters   » Legal and Societal Issues