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Author: Subject: The Short Questions Thread (4)
DrMario
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[*] posted on 25-10-2014 at 13:36


This should be really short and easy: I prepared a few mL of Lugol's solution in a test tube, and then added a few mg of KOH in the form of small pellets. The yellow color of iodine in the Lugol's solution disappeared. Did I make potassium hypoiodite (KOI)?

If yes, can anyone tell me any peculiar property of KOI? Is it similar to sodium hypochlorite in its oxidizing strength?
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Metacelsus
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[*] posted on 25-10-2014 at 14:08


It's probably potassium iodide and potassium iodate (through disproportionation).



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[*] posted on 25-10-2014 at 14:34


Quote: Originally posted by Cheddite Cheese  
It's probably potassium iodide and potassium iodate (through disproportionation).


Are you saying that I didn't make potassium hypoiodite? That was the question I was asking.
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Metacelsus
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[*] posted on 25-10-2014 at 15:33


You probably didn't.



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[*] posted on 25-10-2014 at 15:35


Quote: Originally posted by Cheddite Cheese  
You probably didn't.


I very much appreciate your direct answer.
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[*] posted on 3-11-2014 at 05:06


Can Sulphur (S8) in water be flushed down the drain or should it be stored and given to a special waste processing facility?



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[*] posted on 3-11-2014 at 05:35


Elemental sulfur should be fine to flush down the drain. It is used in large amounts as a fertilizer for plants mixed with some bentonite clay.
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[*] posted on 3-11-2014 at 05:38


It should not be flushed down the drain (it remains solid and does not dissolve and may clog your tubing). It can be put as solid waste in normal household waste. Sulphur is only marginally toxic and when the waste is burned, it produces SO2, which is formed anyway from other waste and will be scrubbed before any waste gas is released into the air.



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Chemosynthesis
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[*] posted on 3-11-2014 at 05:55


Quote: Originally posted by Jylliana  
Can Sulphur (S8) in water be flushed down the drain or should it be stored and given to a special waste processing facility?


Not sure if there are local regulations, but sulfur doesn't violate the clean water act... however, due to it not being water soluble, I would spread it on your lawn as garden sulfur.
http://avogadro.chem.iastate.edu/MSDS/sulfur.htm
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DubaiAmateurRocketry
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[*] posted on 3-11-2014 at 20:02


when I referred P2O5 or P4O10 as phosphorus pentoxide my professor marked it as wrong however its actually one of the most used name for this chemical, shall i argue with him?
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DrMario
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[*] posted on 3-11-2014 at 20:34


Yes, argue with him by pointing him to:
http://pubs.acs.org/doi/abs/10.1021/jo00987a028
http://pubs.acs.org/doi/abs/10.1021/ja00948a050
http://scitation.aip.org/content/aip/journal/jap/34/8/10.106...
or even
http://pubs.acs.org/doi/abs/10.1021/ja01275a030
(1938)

There's a f###ton of peer-reviewed articles where "phosphorus pentoxide" is used to denote P2O5 as well as P4O10.
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[*] posted on 3-11-2014 at 20:40


And if he's the "corporations are always right"-kind-of-guy, then point him at
http://www.sigmaaldrich.com/catalog/product/sial/214701?lang...
http://www.sandhya-group.com/pdf/phosphoruspentoxide-P2O5.pd...
http://datasheets.scbt.com/sc-203187.pdf
http://www.wuzhouchem.com/cataloged/inor/phosphorus_pentoxid...

And for the grand finale:
http://www.phosphoruspentoxide.net/

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[*] posted on 9-11-2014 at 09:47


Is it possible to produce nitrate salts through electrolysis of a ammonium nitrate solution?
For example, using copper electrodes to produce copper nitrate?

I read the thread Electrolysis of Nitrates
Quote: Originally posted by hodges  
...I've also heard that it is possible to get nitrates (with small yield) by electrolysing ammonium salts...


My idea was to use some silver/copper coins as anodes, then precipitate the silver with copper sheet.
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[*] posted on 9-11-2014 at 11:11


You might end up forming tetraamine copper nitrate, which is explosive. Be careful.

[Edited on 9-11-2014 by Cheddite Cheese]




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[*] posted on 10-11-2014 at 11:28


Hello,

I'm looking to build a recirculating cooling pump to prevent water waste.
I want to add glycol to the coolant, so I can lower the temperature below zero.

I also want to add a valve so I can switch to my vacuum aspirator. But I've no idea how it will perform when glycol is added to the working fluid.

Has anybody an idea what it will do to my vacuum?

[Edited on 10-11-2014 by DutchChemistryBox]
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[*] posted on 10-11-2014 at 12:18


This and this might be useful. You can use Raoult's law to calculate the vapor pressure of the desired mixture of glycerin and water. This vapor pressure will be lower than that of pure water and thus, the vacuum produced will be weaker, though it may not be significantly weaker depending on the amount of glycerin present.
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Metacelsus
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[*] posted on 21-11-2014 at 09:53
Zn amalgam questions


I want to prepare a Zn amalgam for a reduction of a ketone to an alkane. Most references call for reacting Zn metal with an Hg(ii) salt, usually mercuric chloride. I have elemental mercury, but no salts as yet.
Could I use mercurous nitrate (the easiest salt for me to make)?
Better yet, could I just use Hg metal?
How much Hg do I need to activate the Zn?

After searching more, I found that mercurous nitrate works. I've also found references of Al amalgams made with Hg metal. However, I'm still not sure about Zn.

[Edited on 21-11-2014 by Cheddite Cheese]




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Brain&Force
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[*] posted on 21-11-2014 at 09:55


I think you can just dissolve zinc into mercury, if I'm not mistaken.

[edit] happy 210 posts, me!

[Edited on 22.11.2014 by Brain&Force]




Raney nickel can't hydrogenate dank memes.
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[*] posted on 21-11-2014 at 10:05


The mercury is just to activate the zinc; dissolving it would be overkill.



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[*] posted on 21-11-2014 at 10:07


I'm sure that the mercury will work, and you won't need much. The mercury salts are generally more convenient, because it allows a very small amount of the mercury to be deposited evenly over the surface of the zinc, whereas elemental mercury won't spread out so much.



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Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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[*] posted on 21-11-2014 at 22:36


Please someone help me with this one:

How would I make acetamide from acetonitrile and water? Is it possible? If it is, what are conditions? I've tried to just mix them together, but nothing seemed to happen - would I need to boil off water from this mixture to precipitate acetamide?

Thanks!




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[*] posted on 21-11-2014 at 23:25


I think you need a catalyst to aid the reaction.
Here's a link to a patent using Zinc oxide as a catalyst:

http://www.google.com/patents/US3040095

Also, here's an old thread that has a synthesis of Acetamide from Ethyl acetate and Ammonia:

http://www.sciencemadness.org/talk/viewthread.php?tid=23491
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[*] posted on 22-11-2014 at 02:24


Quote: Originally posted by greenlight  
I think you need a catalyst to aid the reaction.
Here's a link to a patent using Zinc oxide as a catalyst:

http://www.google.com/patents/US3040095

Also, here's an old thread that has a synthesis of Acetamide from Ethyl acetate and Ammonia:

http://www.sciencemadness.org/talk/viewthread.php?tid=23491


Thanks :)

But I'm still worried, as in the first link, they also mention heat and very high pressure - is that really necessary (high pressure)?

Thanks!




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[*] posted on 22-11-2014 at 03:15


On looking at the patent a second time it does appear to be a bit more complicated than the backyard lab unless you have a heating source to heat a vessel to 250 Degrees Celcius and a vessel or stainless steel "bomb" capable of handling 1000 psi pressure as it says here:

"The acetonitrile, water and zinc acetate are mixed in ice a closed vessel and heated to a temperature of 250 C., preferably 240 C. As the heating proceeds, the pressure rises. During the initial stages of the reaction, the pressure may reach as high as 1,000 p.s.i.g., preferably about 700 p.s.i.g. Thereafter, as the reaction proceeds, the pressure falls due to the formation of acetamide which has a lower vapor pressure than acetonitrile and hence as acetamide is formed, the pressure drops. The pressure at the conclusion of the reaction is usually from 40 to 50 p.s.i.g. "

I think the best way to go would be to obtain some Ethyl acetate and use the method posted in the thread link by Polesch which uses 9% Ammonia as it looks like less of a headache. The Ethyl acetate can be synthesized and the 9% Ammonia should be easy to get.

I have read that ethyl acetate is the primary solvent in non-acetone nail-polish remover but a distillation would be necessary.
Otherwise, here is a two links to a method of making Ethyl acetate from concentrated Sulphuric acid, ethanol and glacial Acetic acid:

https://www.erowid.org/archive/rhodium/chemistry/ethyl.aceta...

https://sites.google.com/site/mutludemirel/organic-chemistry...

There is also a video on youtube here:

http://www.youtube.com/watch?v=cFxZ0NircIk

Hope that helps.




[Edited on 22-11-2014 by greenlight]
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[*] posted on 22-11-2014 at 04:32


Why are there hardly any brown solutions?



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