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sparkgap
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[*] posted on 17-4-2009 at 10:34


@solo: Pd/C is definitely not powerful enough a catalyst to strip fluorine off from a fluoroaliphatic. It might reduce any multiple bonds present (e.g. this), but no fluorine removal can happen.

sparky (~_~)




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solo
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[*] posted on 17-4-2009 at 10:46


Quote: Originally posted by sparkgap  
@solo: Pd/C is definitely not powerful enough a catalyst to strip fluorine off from a fluoroaliphatic. It might reduce any multiple bonds present (e.g. this), but no fluorine removal can happen.

sparky (~_~)


......I was afraid of that.....thanks , solo




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manimal
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[*] posted on 18-4-2009 at 13:46


Can I use sodium carbonate or bicarbonate and ammonium sulfate to create concentrated ammonia solution? I supposed that the reaction would be slow, because due to the weak acidity of ammonium sulfate, only a small percentage of it would be deprotonated at a given time in order to acidify the carbonate.

I tried mixing bicarbonate and amm. sulfate and bubbles were indeed evolved, although slowly, and the odor of ammonia became noticeable after a few minutes. I am wondering what percentage of the ammonia will react with the CO2 to form ammonium carbonate, and what can be done to minimize excape of nh3 gas.




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[*] posted on 18-4-2009 at 17:30


Quote: Originally posted by manimal  
Can I use sodium carbonate or bicarbonate and ammonium sulfate to create concentrated ammonia solution? I supposed that the reaction would be slow, because due to the weak acidity of ammonium sulfate, only a small percentage of it would be deprotonated at a given time in order to acidify the carbonate.

I tried mixing bicarbonate and amm. sulfate and bubbles were indeed evolved, although slowly, and the odor of ammonia became noticeable after a few minutes. I am wondering what percentage of the ammonia will react with the CO2 to form ammonium carbonate, and what can be done to minimize excape of nh3 gas.


With heat, ammonium sulphate will evolve ammonia with almost any basic compound, including metallic oxides such as magnesium and zinc oxides. With carbonates you will also get CO2 evolved. This is the dry powders, mixed. Trap and dissolved the NH3 in water.




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[*] posted on 18-4-2009 at 18:41


Hello everyone,

I have recently become very interested in Kolbe electrolysis particularly concerning alkyl halide synthesis. I want to attempt a kolbe electrolysis of a solution of sodium acetate and potassium iodide to form methyl iodide and would like to ask your advice for optimizing efficiency of the electrochemical cell so that I get the most methyl iodide with the least amount of ethane and methanol byproducts. For instance, should the solution be hot or cold? Should I have it stirring or not? Should I have divided cells? Should I use a dilute or saturated solution of iodide and acetate (I assume saturated)? What voltage, current density etc.

I have read previous posts concerning reaction conditions for Kolbe electrolysis and it was stated that the highest yields are obtained using smooth platinum electrodes (which I of course don't have). I will instead use graphite electrodes which have shown to produce moderate yields of the kolbe products (i.e. ethane). I thought about platinized titanium electrodes but they are apparently detrimental to the reaction (not sure why). Does anyone know where I can get "cheap" (Ha ha!) platinum?

If anyone could provide references or textbook information regarding Kolbe electrolysis would be very helpful as well.

Thank you very much




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Sedit
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[*] posted on 20-4-2009 at 15:01


Does any one know of a way to cleave the methoxy down to benzene? If nothing else Phenol? I faintly remeber seeing SeO2 used for this reason but I cant find additional information on it.


[Edited on 20-4-2009 by Sedit]





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DJF90
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[*] posted on 20-4-2009 at 15:42


With HBr or HI then the ether is cleaved on the alkyl side, leaving you with phenol. You could then react this with zinc dust at elevated temperatures to form benzene.
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Sedit
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[*] posted on 20-4-2009 at 15:51


What exactly is elevated temperatures? Do I need a tube furnace or is reflux sufficiant? Sorry for being persistant Iv just always wanted to know how to cleave a methoxy and cant find any good data on it. Will this work with carbon substitutes on the ring also.




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[*] posted on 20-4-2009 at 17:02


The selenium dioxide method is mentioned in Love Drugs. I doubt that it works well (if at all). As for the demethylation, search for demethylation, and you will find lots of schemes.
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DJF90
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[*] posted on 20-4-2009 at 17:11


I'm not entirely sure to be honest, If I were going to try it I would make sure the phenol and zinc dust were mixed homogeneously (as much as possible) and then heat the distillation flask until benzene starts (and finishes) coming over.
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[*] posted on 20-4-2009 at 17:31


I doubt it works. I believe that "The Chemistry of Phenols" is around here somewhere.
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[*] posted on 20-4-2009 at 18:03


It wasnt in Love Drugs that I seen the SeO2 method it was a fleeting pictogram in a threed somewhere here at SM. I will have to look at Love Drugs to see if there is any additional information about the SeO2 process. Iv seen many schemes for the demethylation but they all proceed thru the phenol so if there is a way to do away with that It would be nice.




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[*] posted on 20-4-2009 at 19:18


What is the oxidation potential of H2SO5? S2O8(2-) in acidic solution is 1.96V (Wiberg, Holleman) and USP6032682 gives 2V for H2S2O8. I suspect it is more than H2S2O8, since H2SO5 instantly liberates I2 from aq. KI, but H2S2O8 only slowly, though I can't find the value of H2SO5.
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DJF90
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[*] posted on 20-4-2009 at 19:29


Peroxodisulfate is a very strong oxidiser, but it is also quite slow. I assume this to be a kinetic effect, as adding a suitable catalyst ( I think we used a silver (I) salt (nitrate?) in the lab ) produces a much more rapid oxidising action. I believe it is a stronger oxidiser than the monopersulfate, and I don't think judging the oxidising capability on the speed at which the substance can liberate iodine from KI is a reliable method.
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[*] posted on 20-4-2009 at 20:24


Quote: Originally posted by DJF90  
Peroxodisulfate is a very strong oxidiser, but it is also quite slow. I assume this to be a kinetic effect, as adding a suitable catalyst ( I think we used a silver (I) salt (nitrate?) in the lab ) produces a much more rapid oxidising action.


True, AgNO3 will catalyse at least some persulfate oxidations.

Quote:
I believe it is a stronger oxidiser than the monopersulfate, and I don't think judging the oxidising capability on the speed at which the substance can liberate iodine from KI is a reliable method.


Do you have anything to corroborate the speculation that that might be the case?
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DJF90
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[*] posted on 20-4-2009 at 20:51


From what I remember, there isnt alot down that end of the table, and I only remember a few redox couples with a higher potential (although admittedly I dont remember seeing peroxymonosulfate at all).

The table has the S2O8(2-)/ 2 SO4(2-) couple at +2.010V. Wikipedia lists the standard electrode potential for HSO5(-) / HSO4(-) as +1.44 V; I know its wikipedia but its better than nothing eh?


[Edited on 21-4-2009 by DJF90]
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[*] posted on 20-4-2009 at 22:21


K2S2O8 was also given as 2.01 in this journal. But as to HSO5(-), I've found the following a paper concerning the KHSO5 triple salt, putting the value up higher:

Oxone (2KHSO5.KHSO4.K2SO4) is a triple salt of potassium and may be regarded as a mono-substituted derivative of hydrogen peroxide. However, it has an oxidation potential greater than that of hydrogen peroxide (E HSO5-/HSO4- = 1.82 eV compared to E H2O2/H2O = 1.77 eV). (link)

And then something in a paper picked up by search engine about : Addition of successive aliquots of hydrogen peroxide at the last step increases the oxidation potential of sulfuric acid to 1.81 V by forming monopersulfuric acid, H2SO5 (Caro’s acid) [19]. (link), putting the lower value of H2SO5 at 1.81.
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[*] posted on 21-4-2009 at 21:08


Has anyone here made the citrate of a solid amine? I noted the three acid positions and am asuming appropriate solvents are found and the acid is added in 1:3 molar concentration. I'm struggling with the solvents and would appreciate access to a compendium of solute/solvents.



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Panache
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[*] posted on 27-4-2009 at 00:34


Quote: Originally posted by chemrox  
Has anyone here made the citrate of a solid amine? I noted the three acid positions and am asuming appropriate solvents are found and the acid is added in 1:3 molar concentration. I'm struggling with the solvents and would appreciate access to a compendium of solute/solvents.


i think you would find the formed salt very very unstable if attempts were made to isolate it from solvent, perhaps you could do it at cryo temperatures, may hang around for a few milliseconds then.
When you say solid amine i assume you mean 'solid at STP' amine, not say methyl amine at -150C, which would be solid.




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[*] posted on 27-4-2009 at 02:16


Quote: Originally posted by chemrox  
Has anyone here made the citrate of a solid amine? I noted the three acid positions and am asuming appropriate solvents are found and the acid is added in 1:3 molar concentration. I'm struggling with the solvents and would appreciate access to a compendium of solute/solvents.


You mean like stearamine? The amine should dissolve in an aqueous solution of the acid. I would expect an increase in viscosity of the salt aqueous solution.

The high MW amines are only weakly basic, so expect only the 1:1 ratio salt to be easily formed.
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[*] posted on 27-4-2009 at 11:11


Quote: Originally posted by Sedit  
It wasnt in Love Drugs that I seen the SeO2 method it was a fleeting pictogram in a threed somewhere here at SM.

You probably mean this.
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[*] posted on 27-4-2009 at 13:31


Thats not the one but helpful indeed. Sadly I did know know the potential hazards associated with SeO2 so I think that method is out the window.




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[*] posted on 29-4-2009 at 13:59


I have some 70% acetic acid + 30% water solution and would like to convert this to glacial acetic acid. Distillation with 95% sulfuric acid will achieve this ? Or I need to go via acetates route + H2SO4 ? Thank you.
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[*] posted on 29-4-2009 at 16:51


Does anybody have a range around which an ethanol flame burns? How about isopropanol?



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[*] posted on 29-4-2009 at 22:50


Quote:
Does anybody have a range around which an ethanol flame burns?


Do you mean flame temperature or flammabilitry limits? Get this PDF for answers for both measurements.

http://www.vrac.iastate.edu/ethos/files/ethos2005/pdf/stokes...

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