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Author: Subject: The short questions thread (3)
Mixell
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[*] posted on 27-6-2011 at 07:45


What is the structure of copper (II) formate? Just as the oxalate (Cu2(HCOO)4)? Or is is Cu(HCOO)2?
Or a completely different structure?
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Nicodem
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[*] posted on 28-6-2011 at 11:07


Quote: Originally posted by Mixell  
What is the structure of copper (II) formate?

The structure of the dihydrate has been determined:
Acta Cryst. 1965, 19, 357-362. DOI:10.1107/S0365110X65003456
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barley81
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[*] posted on 28-6-2011 at 20:01


If I only want a vacuum pump for filtration (and maybe distillation) is it better to use an aspirator with a jet pump, the faucet, use a hand pump, or buy a 100$ vacuum pump from Harbor Freight? I am concerned that the local water supply does not have enough pressure to pull an acceptable vacuum on an aspirator.
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[*] posted on 29-6-2011 at 12:59


What is the simplest Styrene polymerization inhibitor that I can get over the counter?

Its oxygen that catalyses the reaction right? Would a small about of Ethanol or something of that sorts inhibit the polymerization? How about dissolving it in something like Xylene so that the concentration isn't very high?

I just want to attempt pyrolysis of PS and I want to be able to store it in order to perform various experiments with it.





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Jor
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[*] posted on 29-6-2011 at 13:20


I am not sure but isn't hydroquinone used for this purpose?
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[*] posted on 29-6-2011 at 13:50


Does anyone know where to buy 1,3-dichloro cyclobutane? Any other 1,3- derivitives would also be suitable.



I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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Magpie
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[*] posted on 29-6-2011 at 20:24


Quote: Originally posted by barley81  
If I only want a vacuum pump for filtration (and maybe distillation) is it better to use an aspirator with a jet pump, the faucet, use a hand pump, or buy a 100$ vacuum pump from Harbor Freight? I am concerned that the local water supply does not have enough pressure to pull an acceptable vacuum on an aspirator.


For 8 years and a lot of chemistry under the bridge I have only used a hydroaspirator. This has a lot of advantages if you have adequate water pressure. I don't know the minimum acceptable pressure, this would be equipment and task dependent, but I'm assuming 30 psig would do. Mine will draw down to a 50-60mmHg pressure, adequate for all filtrations and all the vacuum distillations to date. If the vacuum distillation procedure specifies 15-20mmHg pressure I have found that my 50-60mmHg will suffice. The product just comes off a little higher in temperature. So I lose a little more product due to decomposition.

So, unless you need a hard vacuum, or your aspirator is going to run for hours and the water consumption would be excessive, why buy a pump? If the pump is oil lubed you'll need a trap or have to change oil frequently.

I wouldn't buy a hand pump when an aspirator is just as cheap, unless water usage is a problem.

So you see, there's no direct answer. The choice very much depends on the individual's circumstances.




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smuv
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[*] posted on 29-6-2011 at 20:40


I use my fridge pump, before that I used a medical diaphragm vacuum pump. I think it is cheaper to go one of these routes than setting up an aspirator/recirculating system. On the other hand, if you just run it from your faucet, and don't care about the water usage, by all means, go for the aspirator.

The good thing about the aspirator is you don't have to worry about sucking evaporated solvent into the pump oil, but this is mostly preventable w/o a trap by not using extremely high vacuums during filtrations (you can use a bleeder valve to control vacuum level very easily).




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[*] posted on 29-6-2011 at 21:55


I use a plastic aspirator and can easily pull 25" @ 40PSI and 27" @ 70PSI

I also installed a 1-100PSI pressure gauge on the water line so that I can check the water main pressure and plan around its flux. Works fine for me.
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[*] posted on 30-6-2011 at 13:11


Thank you for the information. I'll buy an aspirator and use it with the garden hose. Would a water pressure of 30 PSI and a metal aspirator work? Or would I need to buy a jet pump or plastic aspirator?
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[*] posted on 4-7-2011 at 11:44


Does anyone know of a preparation of sulfur trioxide from chlorosulfonic acid? Industrially, chlorosulfonic acid is produced from HCl(g) and SO3.
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[*] posted on 5-7-2011 at 13:56


I've encountered someone on this site claiming that titanium dissolves better in 9% HCl then in 37%.
Is it true? Because I'm planning to dissolve some titanium flitter.
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[*] posted on 5-7-2011 at 14:03


Maybe do to passivization? Here is a PDF. I just scanned over it but I think the temperature of the acid is key.

http://www.archivesmse.org/vol28_6/2866.pdf




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[*] posted on 5-7-2011 at 14:41


Holy crap.
Who knew that 30g of titanium flitter and 200 ml of HCl heated to about 80C will continue reacting due to the heat of the reaction. the house smells of HCl. But at least I'm pleased.
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[*] posted on 5-7-2011 at 15:17


Quote: Originally posted by Mixell  
Holy crap.
Who knew that 30g of titanium flitter and 200 ml of HCl heated to about 80C will continue reacting due to the heat of the reaction. the house smells of HCl. But at least I'm pleased.


lol, careful. HCl is persistent. *cough* My throat hurts just thinking about it.




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[*] posted on 5-7-2011 at 15:38


Ok, the reaction is done.
But I got quite strange results:
The solution is quite concentrated and looks black, when a small amount is taken it looks very dark purple-brown.
But when I dilute it, the solution turns quite brown, resembling this shade: http://ontariohottubcovers.com/imgs/colors/color-cedar-brown... .
When I add hydrogen peroxide to it, it first turns slightly yellow (very transparent) and if I add more it turns blood red.
But what troubles me is it does not resemble the purple solution at all, as shown on Wikipedia: http://en.wikipedia.org/wiki/File:TiCl3.jpg .
What could it be (the titanium is 99.5% pure and the HCl is lab grade)?
Also, the solution turns blue on addition of concentrated HCl.

[Edited on 5-7-2011 by Mixell]
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[*] posted on 5-7-2011 at 20:38


Well, I figured it out, the brown color is due to Ti[IV] contamination.
And the blue color is due to the TiCl6 ion? Is it titanium in its III or IV oxidation state?
Is there any good use for this solution? Or I'll just leave in on my desktop and see how it turns clear?
Shame that it will go to waste.
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[*] posted on 12-7-2011 at 08:06


What does the squiggle that is bound the carbon represent?

http://en.wikipedia.org/wiki/Acyl_group
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[*] posted on 12-7-2011 at 08:14


That's just a shorthand to say that anything could be connected to the other end of the bond. Rather like an engineering drawing, if you were drawing a closeup of the valve on the end of a pipe, you might make a wavy line across the end of your drawing of the pipe to show that the pipe continued beyond the edge of the page but wasn't important to what you were trying to illustrate.

In organic chemistry it's usually used when your reactions are occuring at a specific group on a rather large structure, like cholesterol. Rather than redrawing the entire cholesterol structure at each step of the reaction, you just draw the piece that is reacting in each step, and the reader can assume that the rest of the molecule remains untouched.

In your specific case, only the part of the molecule you show is considered an "acyl group" - some sort of carbon chain (R) attached to a carbonyl (C=O) attached to - something. An acyl group is just a fragment. If the "something" is an OH group, your acyl group is part of a carboxylic acid, if an OR' group it is part of an ester, if an NHR' or NR'R" it is part of an amide, etc.

[Edited on 12-7-2011 by fledarmus]
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[*] posted on 12-7-2011 at 19:48


What is the best way to dissolve nickel powder?
Concentrated sulfuric and nitric acid don't seem to do much, it reacts vigorously at first, but then the reaction stops.

[Edited on 13-7-2011 by Mixell]
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[*] posted on 12-7-2011 at 20:21


Thank you fledarmus!

And Mixell, have you tried diluting it? I have no idea if this is the case, but I know that concentrated nitric acid doesnt dissolve copper because it forms a passivation layer. Diluting prevents the formation in that case.
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[*] posted on 13-7-2011 at 07:36


I tried diluting both of the acids, but nothing seem to work, may be I just didn't find that "magic concentration".

[Edited on 13-7-2011 by Mixell]
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[*] posted on 15-7-2011 at 16:32
OK to boil sodium chlorate liquor electrolytically?


...instead of boiling after filtration? In other words, is there anything wrong with simply bumping the voltage/current up at the end of a run to boil the liquor, then filter and proceed as usual (aside from a little extra anode erosion)? The solution is maintained at 80-85 degrees C anyway and acidified throughout. I don't see how 2 or 3 extra amps would hurt but don't want to screw up this batch.

Thanks
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[*] posted on 18-7-2011 at 18:55


I'd like to remove carbonyl adulterants from my denatured ethanol. Would mixing it with sodium hydroxide and then distilling be sufficient?
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[*] posted on 18-7-2011 at 19:00


I'd like to remove carbonyl adulterants from my denatured ethanol. Would mixing it with sodium hydroxide and then distilling be sufficient?
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