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Author: Subject: The short questions thread (2)
ketel-one
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[*] posted on 30-9-2009 at 23:06


Thanks, I've found it NaOH + alkyl halide, in retrospect that should have been really easy to realize but alright I got it, thanks.
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dann2
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[*] posted on 1-10-2009 at 06:00


Just wondering (as one does)
What is the reaction product between Titanium Tetrachloride and Acetone?


http://cgi.ebay.co.uk/THE-REACTION-PRODUCT-OF-ACETONE-TITANI...

Dann2
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[*] posted on 1-10-2009 at 10:41


Quote: Originally posted by not_important  
HDPE is in effect an alkane, mostly CH2 groups with a small amount of branching giving tertiary hydrogens, and a tiny amount of unsaturation. Anything immiscible with petroleum ether and that doesn't readily chemically attack the same is not going to do much to HDPE. ...


Thanks not_important, I will be trying the hot hydrocarbon when I get the chance to.
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[*] posted on 1-10-2009 at 10:43


Quote: Originally posted by densest  
A (hopefully) easy question: how to destroy nitrate (NO3-) ions in an aqua regia solution after dissolving metals presumed to contain precious ones. The protocol which I have read says to boil the liquid down until it's syrupy, add concentrated HCl, repeat twice more. This seems to me both wasteful and toxic. What I propose to do is to add urea (CO(NH2)2) and heat, assuming that the reaction 6 HNO3 + 5 CO(NH2)2 -> 5 CO2 +8 N2 + 13 H2O is possible. I haven't found anything explicit either way: what I do find is that urea is used to remove excess NO2 from nitric acid solutions.... is the analogy likely? Thanks for any pointers!


Urea will destroy nitrogen oxides from HNO3, but with aq. HNO3, you get urea nitrate. Aqua regia and urea forms something like near quantitative amount of nitrous oxide.
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[*] posted on 1-10-2009 at 19:28
Reduction Stannic to Stannous


I am looking to recycle Sn(II) used for a reduction. I was wondering if I could reduce Sn(IV) to Sn(II) in an acidic solution by means of SO2.

I have looked at the reduction potential tables and it seems that SO2 is a strong enough reducing agent; I wonder If I am missing something which may make this fail in practice.

I do know that Ferrous salts have been used for this reduction, but I would like to achieve this reduction without adding any metal salts that would precipitate under basic conditions (I plan to isolate the tin from solution by basifying it to collect the hydrated oxide).

Thanks




"Titanium tetrachloride…You sly temptress." --Walter Bishop
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[*] posted on 1-10-2009 at 23:21


Quote: Originally posted by dann2  
Just wondering (as one does)
What is the reaction product between Titanium Tetrachloride and Acetone?


http://cgi.ebay.co.uk/THE-REACTION-PRODUCT-OF-ACETONE-TITANI...

Dann2

Titanium tetrachloride is a relatively strong acid so its reaction with enolisable ketones should give normal acid catalysed self condensation products (in amounts and ratios depending on temperature and other conditions). The acid catalysed self condensation products of acetone are numerous, but the most famous ones are mesityl oxide, phorone and mesitylene.




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ketel-one
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[*] posted on 2-10-2009 at 21:13


Does anyone know any common high BP boiling solvents? What I'm looking at (decarboxylation of amino acids to corresponding amines) asks for cyclohexanol (BP 160C) but I'm wondering if something common has similar or higher boiling point.
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[*] posted on 2-10-2009 at 21:19


Quote: Originally posted by ketel-one  
Does anyone know any common high BP boiling solvents? What I'm looking at (decarboxylation of amino acids to corresponding amines) asks for cyclohexanol (BP 160C) but I'm wondering if something common has similar or higher boiling point.


Xylene is probably the most common solvent with the highest BP.




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[*] posted on 4-10-2009 at 21:30


It's alright, I've found that solvent is not totally necessary, strongly heating amino acid with burner makes it fume a lot of white CO2 (fumes don't smell bad or anything). If you heat it too strongly it will start to turn brownish and melt and will easily ignite even if it never directly touches flame. I think the main disadvantage over doing it with solvent is that it might get heated unevenly, some parts burning and some parts undecarboxylated.

A new question: This doesn't work for acetic anhydride, but can nicotinic anhydride be made by nicotinic acid + conc. H2SO4? I'm basing this on the fact that nicotinic acid is a solid while glacial acetic acid is a liquid, only the former I would believe is capable of dissolving sulfuric acid without dehydrating.
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[*] posted on 5-10-2009 at 05:48


Quote: Originally posted by smuv  
I am looking to recycle Sn(II) used for a reduction. I was wondering if I could reduce Sn(IV) to Sn(II) in an acidic solution by means of SO2.

I have looked at the reduction potential tables and it seems that SO2 is a strong enough reducing agent; I wonder If I am missing something which may make this fail in practice.

I do know that Ferrous salts have been used for this reduction, but I would like to achieve this reduction without adding any metal salts that would precipitate under basic conditions (I plan to isolate the tin from solution by basifying it to collect the hydrated oxide).
Thanks


Crystallize your tin chloride and then dissolve in warm concentrated hydrochloric acid. Add tin metal in small pieces and keep warm until dissolved. Leave a single piece of tin at the bottom to keep the stannous ions "fresh".




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[*] posted on 5-10-2009 at 13:16


Has anyone ever made sodium pyrosulfate from sodium bisulfate by heating? I am having trouble. I can't tell if the molten yellow liquid is pyrosulfate or still bisulfate. Also, there is much smoking.
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[*] posted on 5-10-2009 at 13:53


Manimal- check out the oleum from NaHSO4 thread.
The smoking is most probably the decomp of the sodium pyrosulphate to SO3 which immedietly reacts with water vapour to produce fumes of H2SO4.
To make the pyrosulphate slowly heat the bisulphate to +-10deg C 315 degrees, till it melts and starts to release water. when the water is removed (may take a few hours) then SO3 (lots of fumes) should form. This is the point you want to stop at to make the pyrosulphate.

[Edited on 5-10-2009 by Picric-A]
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manimal
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[*] posted on 5-10-2009 at 15:19


Yes, I suppose you're right. I didn't think it was H2SO4 mist, because I thought H2SO4 mist would be intensely irritating, whereas with this, it only had an odor that I would describe as like hot ashes or rock dust.
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[*] posted on 6-10-2009 at 07:12


I am interested in the use of sulfur dioxide gas to reduce chlorate ion contamination in electrolytically-prepared perchlorate. In the past, I have used (with success) Potassium Metabisulfite added to hot, acidic aqueous solutions, but the contact of the evolved SO2 gas with the solution is very brief, and the amount of potassium metabisulfite necessary to "clean" the perchlorate is higher than I would like.

SO2 gas has an inverse solubility, with higher SO2 solubilities at lower temperatures. This is in conflict with the salt of interest, potassium perchlorate, which has a lousy solubility except at higher temperatures.

Essentially, the challenge is to maximize contact of the SO2 gas and use the least amount of metabisulfite per unit of chlorate contamination, which is an estimated 0.5%.

I'm thinking perhaps the best way is to start cold, introduce the SO2, then slowly heat, which will dissolve the contaminated perchlorate, with hopefully enough SO2 remaining to reduce the contamination.

Any thoughts on obtaining the most economical use of the metabisulfite?
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Picric-A
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[*] posted on 6-10-2009 at 07:19


Quote: Originally posted by manimal  
Yes, I suppose you're right. I didn't think it was H2SO4 mist, because I thought H2SO4 mist would be intensely irritating, whereas with this, it only had an odor that I would describe as like hot ashes or rock dust.


Maybe becuase SO3 evoloutiuon had just started so it was a vry dilute acid mist. Either way i dont reccomend smelling it to detect H2SO4 presence. use indicator paper suspended above the molten mix. When all its red stop heating becuase you have the pyrosulphate
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watson.fawkes
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[*] posted on 6-10-2009 at 07:42


Quote: Originally posted by Swede  
Any thoughts on obtaining the most economical use of the metabisulfite?
I'd modify a fractional distillation column, using counter-current circulation of the gas (up) and the liquid (down). Really, this just means a pump and a way of introducing the gas. You could continue to use the metabisulfite, or alternately use a sulfur candle. The goal in intimate gas-liquid contact, reusing rather than venting the gas.
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[*] posted on 6-10-2009 at 23:17


I know this might be a stupid question, but is it possible to drink lab reagent grade ethanol that is diluted to 40%? This is no different to vodka yes? I already have ethanol however i dont want to be buying vodka if what i have is perfectly safe when diluted appropriately.
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[*] posted on 7-10-2009 at 10:36


Depends if it is denatured or not. Reagent grade ethanol means nothing.
Rectified spirit (~98% ethanol) can be drunk if diluted however it would be cheaper to buy vodka than use that.
Most 100% ethanol has been made by azeotropic distillation with benzene to remove water and as such the ethanol may contain traces of benzene with its accociated hazards.
Just buy vodka and dotn try to use laboratory chems :) safer and cheaper
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[*] posted on 7-10-2009 at 12:43


Thanks for your reply, it is definately not denatured and i believe it is 100% ethanol but the thought of trace amounts of benzene is enough to put me off.
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[*] posted on 16-10-2009 at 11:32


I can currently obtain a lot of silica gel, however, it is a powder used for chromotography. Can this also be used for used a dessicant in my dessicator, instead of the granules?
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[*] posted on 16-10-2009 at 12:30


Yes i use a powder however everytime you change the thing you are drying i reccomend you give the powder a bit of a stir up.
Granules are of course better but the powder works nonetheless
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[*] posted on 16-10-2009 at 20:07


Quickie: I tried checking in Google Scholar but can't seem to sort it out; has epichlorhydrin ever been used in a Friedel-Crafts alkylation?

sparky (~_~)




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[*] posted on 17-10-2009 at 08:11


The problem with epichlorohydrin is that it has two labile positions. The epoxide moiety will react under FC conditions to yield an aryl-substituted alcohol (the ratio of primary/secondary depends on the nucleophilicity of the aromatic you are reacting it with). Additionally epichlorohydrin is chloro-substituted which will as well lead to FC products. One could even make the argument that the alcohol formed from the ring opening of the epoxide would as well react under some conditions.

Overall you have 2 or 3 possible reactions, based on this, I think it would be very difficult to do any FC chemistry on this compound without observing a mixture of many products (including polymers).

[Edited on 10-17-2009 by smuv]




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[*] posted on 18-10-2009 at 07:35


Hmm, I forgot to take into account that alkyl groups activate aromatic rings towards FC alkylation. :( I was hoping 1,3-diphenyl-2-propanol was as easy as epichlorhydrin + benzene, but I guess I'll look for something else.

Thanks smuv.

sparky (~_~)

[Edited on 18-10-2009 by sparkgap]




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[*] posted on 18-10-2009 at 08:03


You could try using 2eq. Phenylmagnesium bromide? It should attack at the least hindered position (unless you add a strong lewis acid) of the epoxide, yielding a halohydrin, which would then react with the second equivalent to displace chlorine. Shouldnt be too hard?
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