Sciencemadness Discussion Board

The short questions thread (2)

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densest - 4-2-2010 at 14:14

Perhaps this should be under "Acquisition"... if so, I'll redirect this.

Is P2O5 (phosphorous pentoxide, really P4O10) a suspect reagent in the USA? It's awfully useful for thoroughly removing H2O....

I haven't seen it on any lists, but that isn't necessarily good enough to go buy 500g of something.

Sedit - 14-2-2010 at 11:02

I wanted to further my Lithium complex thread but due to heavy snow I dont have access to the Ammonium nitrate or the money for it either so I made some Ammonium Chloride from the materials I had on hand.

I slightly over acidified the solution of Ammonium hydroxide and Muratic acid and I want to remove this excess HCl from the crystals else they will screw with my calculations for freeing the NH3 and waste my NaOH on top of that.

I have no acetone which would be my solvent of choice to wash the ammonium chloride so I was woundering about other things like EtOH or if a simple cold water wash would be enough to remove the excess acid.


I recrystalised it using plain old H2O but there is still a little excess acid I want to remove.

S.C. Wack - 14-2-2010 at 11:36

Quote: Originally posted by Sedit  
I recrystalised it using plain old H2O but there is still a little excess acid I want to remove.


Are you sure?

Nicodem - 14-2-2010 at 15:37

Sedit, what S.C. Wack is probably trying to tell you (my guess on why he asks) is that you can not judge the presence of residues of HCl in ammonium chloride using something like a litmus paper (just my guess of what you did) because ammonium chloride is an acid itself. You would get an acidic reaction even if there is no HCl there (which should not be if you dried it thoroughly).

Sedit - 14-2-2010 at 16:48

Didn't use litmus to tell much acid was left, when some of the solution that it was recrystalizing from dripped onto a concrete floor it fizzed... I doubt NH4Cl is acidic enough to produce this effect. I just figured a quick wash in Acetone would handle it but I have no acetone or atlest not enough to be of anyuse. So perhaps I'll just recrystalize it once more which should remove more then enough to be of no problem to me. I only need it for Ammonia generation but didn't want to waste more NaOH then needed neutralizing residual HCl.

Nicodem - 15-2-2010 at 01:20

Whether it be fizzling in the reaction with your concrete floor if it was pure saturated NH4Cl solution depends on what material it is on your floor. If the concrete is exposing calcite sand, then it may fizz somewhat. But that is not really scientific way of testing it, isn't it?

Is there a particular reason on why you can not properly dry it? Surely you can improvise a desiccator to dry it over anhydrous Na2CO3 or something? And yes, you could wash it with ethanol. It is nearly insoluble there.

entropy51 - 15-2-2010 at 08:48

Somewhat surprisingly the Merck Index (9th & 10th editions) says NH4Cl is soluble in MeOH and EtOH.

This seemed so odd that I tried a quick test tube experiment. There was no appreciable dissolving, even when warmed, in 95% EtOH (Everclear).

gsd - 15-2-2010 at 09:44


As per the wiki page of Ammonium Chloride, its Solubility in (ethyl)alcohol is 0.6 g/100 mL (19 °C).

gsd

jwarr - 15-2-2010 at 10:19

I managed to get a teflon thermometer adapter stuck in the ground glass joint of a flask by leaving it in there without lubricant for a long long time. I tried pulling it out with pliers to no avail. I'm hesitant to use a lot of force because I dont want to mess up threading. How can I get this to come out easily?

entropy51 - 15-2-2010 at 11:10

UTFSE

chemrox - 16-2-2010 at 10:58

Entropy posted a link to the thread under his UTFSE. I read through the thread and pulled out what would seem to apply to your situation. I often use gentle heat via hot water from the tap. Sometimes an overnight soak in the wash basin does the trick if the hot water doesn't. Since a thermometer is involved flames are not recommended. Careful use of a narrow heat gun judiciously applied might work but is risky. Forget MID's with a thermometer. Forget about beer or soft drinks (too much sugar and other crap) but carbonated water could get the result. From here on out, NEVER NEVER leave ground glass joints any longer than needed for the application and don't use them without lubrication or sleeves. Teflon grease is spendy but a tube will last a long time. Teflon sleeves are good for non-vacuum uses. The tradition that st joints can be used without grease or sleeves at atmospheric pressure has not been born out by my experience. It might have been true when all st joints were precision ground but this is sadly no longer the case.

DJF90 - 16-2-2010 at 12:39

Chemrox: I disagree here. All the glass I've used at uni has always been without lubrication. The only time I've had problems with stuck joints was back at school, and I happened to use grease that time! I think the trick lies in NOT twisting the joint; push to fit and leave it be.

entropy51 - 16-2-2010 at 13:32

Quote:
The tradition that st joints can be used without grease or sleeves at atmospheric pressure has not been born out by my experience.
I tend to agree with DJF. I only grease when using vacuum or strong bases and I've never had a joint to stick so tight that gentle warming wouldn't open it. I think the real trick is to take the joints apart when you're done. That said, I do grease glass stoppered bottles used for storage because they don't seal too well without grease.
Quote:
It might have been true when all st joints were precision ground but this is sadly no longer the case.
I can't argue with that because almost all my glass is at least 25 years old.

Cu2O to CuO

Hamilton - 20-2-2010 at 08:26

An easy one i think,

How to oxidize Cu2O to CuO in an aqueous mixture?

i tried KClO3, KNO3 and H2O2 and got very bad result, the mixture is still very brownish red. i was thinking about hypochlorite next, but I'm bored washing my Cu2O after each try. Any easy way with home chemical?

---
I generally use thermal expansion to solve any stuck glass joint, put the joint in the freezer and when it is cold try to separate it, if it still doesn't work get it out of the freezer and the outer glass will expand before the inner glass and they always (until now) separate.


[Edited on 20-2-2010 by Hamilton]

DJF90 - 20-2-2010 at 08:42

I'm not sure its going to oxidise efficiently as its insoluble in aqueous solution, just like CuO. Personally I would try heating it strongly in air ("roasting" it); this should be sufficient.

Magpie - 20-2-2010 at 09:17

Quote: Originally posted by chemrox  

It might have been true when all st joints were precision ground but this is sadly no longer the case.


I use a thin coat of grease on my tapered joints and twist them into place. Then I separate them for cleanup right after use. I've never used a Teflon sleeve. I've never had a frozen joint.

I use a heavier coat of grease for vacuum distillations. Even with that I don't always get a good seal because I believe the joint is not precision ground. The faulty joint always has a cheap knock-off adapter or flask. When I replace these with a good brand like Kontes the joint seals.

bbartlog - 20-2-2010 at 18:16

>How to oxidize Cu2O to CuO in an aqueous mixture?

These aren't soluble in water. However in strong base they can dissolve, so I'd look for an oxidizer that is either basic or compatible with same. Sodium hypochlorite with added sodium hydroxide would probably work (then you'd have to neutralize to precipitate the CuO).
But roasting in air seems like a much easier plan.

Bolt - 3-3-2010 at 23:18

Is there a method of preparing an alkyl halide from an alkyl amine besides using NOCl/NOBr or the Sandmeyer rxn? Thanks for any answer!

[Edited on 4-3-2010 by Bolt]

Hamilton - 14-3-2010 at 12:04

here's another one,

can small particles of PVA like common white glue be chlorinated with HCl or (something similarly available) to produce PVC and acetic acid?
Juts like copper acetate turn to copper chloride and acetic acid when mixed with HCl.

For the CuO reply, in presence of NaOH, sodium hypochlorite do oxidize but not completly as it is decompose in O2 bubbles. The result was particules of Cu2O coated with CuO. As i found after dissolving in common vinegar. The roasting process work great.

Thanks

[Edited on 14-3-2010 by Hamilton]

jokull - 14-3-2010 at 16:34

Quote: Originally posted by Hamilton  
here's another one,

can small particles of PVA like common white glue be chlorinated with HCl or (something similarly available) to produce PVC and acetic acid?
Juts like copper acetate turn to copper chloride and acetic acid when mixed with HCl.


[Edited on 14-3-2010 by Hamilton]


Definitely NOT, with HCl you will hydrolyse the Polyvinylacetate (PVAc) thus obtaining Polyvinylalcohol (PVA).

[Edited on 15-3-2010 by jokull]

Hamilton - 16-3-2010 at 08:33

thx,

I tried to mix hcl and PVAc (white glue), i got a liquid more translucid than dilute PVAc. (same volume of PVA in HCl and in H2O) It's look partially soluble like PVOH should.(PVA, not to make me confuse will be write PVOH). Maybe there is some white pigments in white glue.

so i tried to write and equation. but can't figure it out very well.

(C<sub>4</sub>H<sub>6</sub>O<sub>2</sub>;)n + (HCl)m = (C<sub>2</sub>H<sub>4</sub>O)x + something else.

by smplicity (and because I don't know enough about possible resulting organic substance) i proceed with m=n

[C<sub>2</sub>H<sub>3</sub>(CH<sub>3</sub>COO)] +HCl = [C<sub>2</sub>H<sub>3</sub>(OH)] +[Cl + C<sub>2</sub>H<sub>2</sub>O]

The resulting substance doesn't smell Cl nor HCl nor CH3COOH. and C<sub>2</sub>H<sub>2</sub>OCl doesn't exist (well i think)

can you help me on the equation?

[Edited on 16-3-2010 by Hamilton]

i just return to my experiment and the solution of PVAc and HCl have become totaly translucid and very slightly yellow. On addition of water the PVAc (or whatever it is now) precipitate again. It look like it is soluble in very acid solution but not on dilute HCl.

[Edited on 16-3-2010 by Hamilton]

DJF90 - 16-3-2010 at 09:17

Bolt, use a pyryllium salt and some sodium/potassium halide. The pyrylium converts the amine into a good leaving group (pyridine) and a simple Sn2 ensues with the halide ion.

Nicodem - 18-3-2010 at 09:49

Quote: Originally posted by Hamilton  

so i tried to write and equation. but can't figure it out very well.

(C<sub>4</sub>H<sub>6</sub>O<sub>2</sub>;)n + (HCl)m = (C<sub>2</sub>H<sub>4</sub>O)x + something else.

by smplicity (and because I don't know enough about possible resulting organic substance) i proceed with m=n

[C<sub>2</sub>H<sub>3</sub>(CH<sub>3</sub>COO)] +HCl = [C<sub>2</sub>H<sub>3</sub>(OH)] +[Cl + C<sub>2</sub>H<sub>2</sub>O]

The resulting substance doesn't smell Cl nor HCl nor CH3COOH. and C<sub>2</sub>H<sub>2</sub>OCl doesn't exist (well i think)

can you help me on the equation?

You mean the hydrolysis equation? You wrote no water in the equations above, so it is not clear what you want. Perhaps you should first read something about ester hydrolysis, because from what you wrote above it looks like you should. Try putting "ester hydrolysis" keywords in a search engine or just open a basic level organic chemistry book (most will have a chapter on various hydrolyses). Besides "Cl" is chlorine radical and as such can not be the end product of any reaction (it combines either with itself giving Cl2 or with something else). Products are end results of reaction pathways and not of acrobatic writing with chemical symbols. Each bond that is broken must be accounted for by the formation of new ones.

Anyway, the hydrolysis equation would be:

[-CH<sub>2</sub>-CH(OCOCH<sub>3</sub>;)-]<sub>n</sub> + nH<sub>2</sub>O <=> [-CH<sub>2</sub>-CH(OH)-]<sub>n</sub> + CH<sub>3</sub>COOH

HCl is not a reagent in the reaction. It is there just to catalyse hydrolysis. Of course, with acid catalysed hydrolysis you can not fully hydrolyse polyvinyl acetate (unless you somehow remove acetic acid during the reaction). You should check a few examples of how polyvinyl acetate is hydrolysed. The patent literature is relatively abundant with such examples for it is a reaction used industrially. So try spending some time at Espacenet using the search engine.

manimal - 25-3-2010 at 21:44

I am looking into the stability of Cu(I) oxide. According to wiki, it can be oxidized by "moist air". Do you suppose simply drying a copper(i) oxide precipitate in air would cause significant oxidation?

bbartlog - 26-3-2010 at 07:23

Based on my experiences with CuCl, Cu(OH) and Cu2CO3 (or whatever passes for it) I'd say yes. All of these oxidize fairly quickly to Cu(II) when wet and exposed to air. Of course it still depends on what you would consider 'significant', the particle size, drying temp and so on; nonetheless I think that you would need to do something more than just pop it in the oven to dry if you want it to remain as Cu(I). Acetone and vacuum drying might help.

crazyboy - 27-3-2010 at 21:46

I recently got a buchner flask and funnel, when I put a piece of filter paper in the flask with some liquid and turn the vacuum on the filter paper rips several holes. I am wetting the filter paper prior to applying vacuum. I have tried using two pieces of filter paper at the same time, this works better but it still rips occasionally. What can I do? This is an electric pump not an aspirator so I can't really adjust the vacuum without a bleeder valve.

What can I do?

1281371269 - 28-3-2010 at 05:58

Is it a good quality filter paper? You could use the acid resistant glass based ones which would be much stronger.

DJF90 - 28-3-2010 at 11:05

It sounds to me like you're trying to do vacuum filtration with a "high vacuum pump". If this is the case then you will probably need that air bleed, or to use an aspirator. Good quality filter paper (look for Whatman) will also help.

crazyboy - 28-3-2010 at 15:00

Yes I am using a high vacuum pump. And yes I am using high quality filter paper (it is Double Rings 9cm ashless 202) am not using coffee filters.

How would I establish and air bleed? Can it be made with parts from a hardware store or does it require a special part?

Panache - 31-3-2010 at 20:39

---->Yes I am using a high vacuum pump. And yes I am using high quality filter paper (it ---->is Double Rings 9cm ashless 202) am not using coffee filters.
---->How would I establish and air bleed? Can it be made with parts from a hardware -------->store or does it require a special part?


its as simple as a t-piece in the vacuum line running to the filter funnel with a short piece of hose that you clamp to the bench anywhere from open (ie no vacuum to your funnel) to closed (the situation you have currently), and an infinite no. of increments between. Of course a needle valve makes life easier.

[Edited on 1-4-2010 by Panache]

bbartlog - 10-4-2010 at 13:48

I notice after filling my lab with HCl fumes that some of the borosilicate glass surfaces are covered with a white haze. I see the same thing on the inside of some test tubes that have HCl solution in them and have been sitting closed for a few months.
What is this white haze? I don't think HCl can react directly with borosilicate glass, but something seems to be happening.

aonomus - 10-4-2010 at 13:59

Quote: Originally posted by bbartlog  
I notice after filling my lab with HCl fumes that some of the borosilicate glass surfaces are covered with a white haze. I see the same thing on the inside of some test tubes that have HCl solution in them and have been sitting closed for a few months.
What is this white haze? I don't think HCl can react directly with borosilicate glass, but something seems to be happening.


I store my HCl in a glass bottle which hazes up, and I keep that bottle inside a LDPE bag, which also gets a white haze/coating on it. I don't think its connected to the glass itself, but HCl tends to form that haze that can be wiped off with a damp cloth; the haze is also corrosive over long periods of time to tools and metals (which my pile of iron/steel stock has found out unfortunately).

Lambda-Eyde - 10-4-2010 at 16:06

Quote: Originally posted by crazyboy  
I recently got a buchner flask and funnel, when I put a piece of filter paper in the flask with some liquid and turn the vacuum on the filter paper rips several holes. I am wetting the filter paper prior to applying vacuum. I have tried using two pieces of filter paper at the same time, this works better but it still rips occasionally. What can I do? This is an electric pump not an aspirator so I can't really adjust the vacuum without a bleeder valve.

What can I do?


How large are the holes in your buchner funnel? Crappy chinese funnels tend to have fewer, larger holes (3-4 mm) compared to Coors porcelain funnels which have many smaller holes. The paper is easily ripped apart in a funnel with larger holes.

Ephoton - 17-4-2010 at 23:27

why is every one so down with chinese goods. seriously the glass I buy is chinese and ive used a lot of
glass.

I never had a problem with my supplier they are fast understanding and very very high quality and
I often do crazy stuff with my glass at high temps never a problem.

I will never go back to western glass now I have found my supplier its quicker and easier and a lot cheaper
to use them.

densest - 18-4-2010 at 11:08

@ephoton - if you're having good luck with Chinese glass, good for you. You may be dealing with someone other than the ubiquitous Bomex. I've found Bomex to be irregular in shape, uneven in thickness, and with obvious flaws where the glass is either impure or was not melted sufficiently. It seems to break more easily than my Pyrex or Kimax which would be easy to blame on the flaws. If you have a better vendor for inexpensive good glass, I'm all ears!

I do flameworking of borosilicate glass. The Chinese glass is much less expensive than US or European glass, but in general is "lumpy": parts of the glass melt at higher or lower temperatures and when they do it's more or less viscous than other parts. I blame that on insufficient mixing and melting time and perhaps variations in composition.

Unfortunately, at this end of the supply chain it's very hard to figure out exactly which factory something comes from and whether or not the manager demanded better quality that day.

H2SO4 - KClO3 limerick

The WiZard is In - 19-4-2010 at 10:52

There is a limerick 'bout a student who
mixes sulphuric acid and potassium chlorate
to his demise! Can anyone bring it to mind.

Ephoton - 29-4-2010 at 06:34

can nitrogen triiodide ammonium adduct be oxidised back to iodine with out detonation.

I am finding it hard to find information on this any help would be great.

of course it is under water and not dry :).


[Edited on 29-4-2010 by Ephoton]

not_important - 29-4-2010 at 07:26

Not directly, SFAIK. Alkali plus H2S will convert it to the alkali iodide. But the stuff is so darn sensitive, alpha particles can set it off, that I'd not convert a bunch at one shot. Perhaps just plain alkali, Na/KOH, would do the same.

manimal - 2-5-2010 at 19:50

Which is more correct in the verb usage, "sublime" or "sublimate"?

DJF90 - 3-5-2010 at 05:06

A material sublimes, the collected solid is the sublimate.

Paddywhacker - 3-5-2010 at 14:15

Or you can sublimate your passion for chemistry by indulging in the preparation of sublime meals.

Panache - 3-5-2010 at 17:34

Is there a trick to fitting thread parts together so that the orientation you require can be assured at the correct level of tightness. Inevitably my t-pieces always end up pointing at the wall or something, very frustrating.

jokull - 4-5-2010 at 06:07

Dear Panache,

if your are talking about Swagelok-like connections, a good result can be achieved if you insert the T against the tubing, then you can screw the nut while keeping the T in a desired position.

manimal - 7-5-2010 at 16:22

What is the approximate solubility of hydrogen chloride in ethanol?

entropy51 - 7-5-2010 at 17:06

41 gm per 100 mL. Get a Merck Index; pure gold.

JohnWW - 7-5-2010 at 18:45

Quote: Originally posted by entropy51  
(cut) Get a Merck Index; pure gold.
I have uploaded the editions 12.3 and 11 of the Merck Index to my Rapidshare premium account, and have posted the links for downloading them as DJVU, PDF, and EXE files in appropriate threads in the References section. In fact, I have today had to reupload the 11th edition, the previous link as a DJVU file of about 4 years' standing for which had been broken this morning, thanks to the Copyright Gestapo (still operating on the 65th anniversary of VE Day, to remind us that they are still around!). At some stage soon, I will be uploading the 13th and 14th editions, as EXE files which require installing a setup program.

[Edited on 8-5-10 by JohnWW]

Hamilton - 11-5-2010 at 15:15

hello! a not so small question for you all

i was experimenting copper chemistry again the other day when i find myself "carbonating" CuCl2 with NaHCO3. I was thinking that the Cu will precipitate with the CO3 anion to form insoluble CuCO3. I didn't ask myself where the H would go in this reaction.

I was using a left over CuCl2 "I think" solution made by dissolving Cu2O in HCl in excess, until all the white CuCl get in solution. The solution also contain copper acetate and acetic acid for some reason. it is important to note that no nitrogen compound were present in the solution. The blue solution had the property of turning to a dark green one on heating and revert as the temperature drop.

A small sample (100 ml) was moved to another beaker and heat in the microwave. To the heated solution, a hot NaHCO3 solution was add slowly. Lots of bubbles come off. At a certain point the solution turn dark blue with a bright blue precipitate. I separate the blue solution from the precipitate to make some investigation later as i never saw this before. I neutralizes everything else with NaHCO3 until all was in a bright blue precipitate. Then got bored and stop...

The next day, I found that the dark blue precipitate had crystallized in small needle turquoise crystals with the mother solution still around them.

The interesting property "for me at least" of those crystals is that they don't look soluble at all but was made from a water soluble 'state'.

I tried adding more (a lot) water to dissolve but they stay out of solution. On adding NaOH, the turquoise crystals turn sky blue (the same blue as the originating solution) but they don't dissolve.

They dissolve in acid to give the respective copper salt, without giving any gas off. ( i tried with vinegar in excess they return to the turquoise color before getting dissolved)

I am clueless here about their composition. Some articles talk about a Na and copper hydroxide like compound but they don't describe it enough.

what do you think about those crystals?




12AX7 - 11-5-2010 at 15:39

Copper forms a carbonate complex.

I once tried electrolyzing copper metal in a sodium carbonate solution. Didn't work. Make the most amazing cobalt-but-not-actually-cobalt blue solution.

I don't know if this complex is stable, or if it's metastable in that it will precipitate a basic carbonate over time (you seem to have grown crystals of azurite or malachite, which is neat).

Neutralizing a green (concentrated) CuCl2 solution tends to produce a lot of pastel green oxyhydroxide, Cu(OH,Cl)2. When diluted, the chloride complex is decomposed, the solution becomes bluish, and the majority product is light blue Cu(OH)2.

With an excess of carbonate, there may be a lot of Cu(OH,CO3) as well.

The extra H on NaHCO3 goes off as H2O, while the extra CO2 goes off as CO2 gas (hence the foaming).

Tim

Hamilton - 11-5-2010 at 16:47

can i test them further? if it is malachite, dissolving in vinegar should release CO2 gas but it don't. For azurite, does it get dark blue in NaOH solution?

I will tried to make more crystals later!

thx for your fast answer

12AX7 - 11-5-2010 at 18:23

If you dissolve only a small amount, the CO2 may remain in solution. Can you get a pile of the stuff and drop acid on it (>10% HCl would be best)?

Tim

stygian - 11-5-2010 at 18:31

Can nitroalkenes (b-nitro alcohols) undergo a retro-aldol type reaction? Or shall I say, what conditions would be needed to coerce this into happening?

Hamilton - 12-5-2010 at 05:17

Quote: Originally posted by 12AX7  
If you dissolve only a small amount, the CO2 may remain in solution. Can you get a pile of the stuff and drop acid on it (>10% HCl would be best)?

Tim


I will try making more but right now i don't even have 1 gram.

then i will use 20 B degree HCL to try.

Pope - 13-5-2010 at 04:16

Can Phenylpropanolamine be produced by the reductive amination of L-Pac Via the below method except the catalyst (raney nickel) is substituted for Al/Hg?

Run down for people to lazy to open the link,
Quote:

Compound 1 5g [α]D25 -155° (containing 4.28 g pure (R)-(–)-1-hydroxy-1-phenyl-2-propanone, 0.03 mol·dm-3) was taken up in 15 mL of benzene. 11.4 mL (0.17 mol·dm-3) of 25% ammonia and 0.5 g Raney nickel were added and hydrogen was passed for 2.30 h at 50 psi. After completion of the reaction, the reaction mixture was filtered off. concentrated hydrochloric acid (7 mL) was added to adjust the pH to 5.5. The solution was concentrated under reduced pressure. The viscous mass was dissolved in absolute ethanol and dry ether was added to the solution. A colourless compound crystallised from absolute ethanol, mp 192–194°C (literature cited 190-194°C, [α]D25 0°, was obtained. Yield: 30%, 1.61 g.


http://www.erowid.org/archive/rhodium/chemistry/ppa.l-pac.ra...

Thanks for any help given.

manimal - 15-5-2010 at 15:44

Yes, but yields will likewise be low.

manimal - 15-5-2010 at 15:49

I have in mind a reaction that alkylates an m-nitrobenzyl-alkyl-ketone at the alpha-carbon with an alkyl chloride in a biphasic PTC environment. Is anyone cognizant of references that affect this or any analogous reactions?

[Edited on 15-5-2010 by manimal]

entropy51 - 15-5-2010 at 16:58

Quote: Originally posted by Pope  
Can Phenylpropanolamine be produced by the reductive amination of L-Pac Via the below method except the catalyst (raney nickel) is substituted for Al/Hg?
L-Pac? Is that some of that drug cook's terminology that isn't appropriate on a chemistry forum?

manimal - 15-5-2010 at 19:21

Quote: Originally posted by entropy51  
L-Pac? Is that some of that drug cook's terminology that isn't appropriate on a chemistry forum?


I don't think so. It stands for l-phenylacetylcarbinol, which is legit nomenclature.

Nicodem - 18-5-2010 at 04:33

Quote: Originally posted by manimal  
I have in mind a reaction that alkylates an m-nitrobenzyl-alkyl-ketone at the alpha-carbon with an alkyl chloride in a biphasic PTC environment. Is anyone cognizant of references that affect this or any analogous reactions?

A cursory SF search gives these related examples:

Bioorganic & Medicinal Chemistry, 2009, 17, 2963-2974 describes an alpha-C-alkylation of m-nitrobenzyl p-methoxyphenyl ketone with methyl bromoacetate (NaH/DMF);

WO2005070420 describes alpha-C-methylation of 4-fluoro-3-nitrobenzyl cyclohexyl ketone with methyl iodide (NaH/THF, 2h, rt);

Chemical & Pharmaceutical Bulletin, 1969, 17, 605-610 describes alpha-C-propargylation of 3-chloro-5-nitrobenzyl methyl ketone with propargyl bromide (NaOEt/EtOH);

These are only examples on substrates containing the 3-nitrobenzyl moiety. Otherwise, there are plenty of examples on other benzyl alkyl/aryl ketones. Some use KOtBu or NaH, others use NaOH under PTC conditions (quats or crown ethers; see Tetrahedron Letters, 1971, 1351-1352 for a method). If the electrophile is a fairly reactive primary alkyl bromide or iodide, the reaction should give reasonably good yields. With alkyl chlorides under PTC conditions, I doubt you would get anything useful, unless of course it is a benzyl or allyl chloride. Rather use an alkyl bromide/iodide/tosylate/etc. Likewise secondary alkyl halides are unlikely to work, and if they do the yields will suffer greatly.

Edit: In case you want to alkylate with beta-chloroethylamines in a biphasic system, consider that here the actual alkylating species (aziridinium ions) are cations and thus partition into the aqueous phase (and succumb to hydrolysis) rather than stay in the organic phase to react with the substrate. (just guessing about the nature of the electrophile you intent to use)

[Edited on 18/5/2010 by Nicodem]

497 - 20-5-2010 at 16:41

Nevermind, found the answer..

[Edited on 21-5-2010 by 497]

Pope - 20-5-2010 at 19:37

If possible could someone link me to a viable source, I've only found one place where they said you could substitute it for Al/Hg but there was no hard facts.

Panache - 25-5-2010 at 04:21

Is there an similar property for fluid flow in pipes cf/ starting friction and sliding friction for solid surfaces moving against one another?

watson.fawkes - 25-5-2010 at 06:38

Quote: Originally posted by Panache  
Is there an similar property for fluid flow in pipes cf/ starting friction and sliding friction for solid surfaces moving against one another?
For ordinary Newtonian fluids, no. The equivalent terms would be "standing viscosity" and "moving viscosity", but there's only a single viscosity. The Newtonian fluid model, though, doesn't nearly model all real fluids. One non-Newtonian model uses shear-dependent viscosity, which models thixotropic and shear-thinning materials.

Panache - 25-5-2010 at 07:44

Quote: Originally posted by watson.fawkes  
Quote: Originally posted by Panache  
Is there an similar property for fluid flow in pipes cf/ starting friction and sliding friction for solid surfaces moving against one another?
For ordinary Newtonian fluids, no. The equivalent terms would be "standing viscosity" and "moving viscosity", but there's only a single viscosity. The Newtonian fluid model, though, doesn't nearly model all real fluids. One non-Newtonian model uses shear-dependent viscosity, which models thixotropic and shear-thinning materials.


apologies then i guess this gets far more complex quickly however i'll ask anyway. The circumstance i'm interested in is a recirculating pump on startup running nothing other than water or water glycol. Does the pump initially do more work until the circuit is flowing or does the work being done by the pump simply gradient upwards leveling out once the pressure in the system is obtained. Are equations available to predict the pressure gradient in the system from startup until i guess you could call it steadystate.

solo - 25-5-2010 at 11:45

.....have been pondering on the Cleisen Rearrangement of c6h5-ch2-ch2-(0-ch3) , following the rearrangement it would make phenylacetone.......it's a question not a fact and what would the ether be called?......also can the rearrangement function as outline or does it only with an allyl group.......solo

Sedit - 25-5-2010 at 12:44

(2-methoxyethyl)benzene I know of no common names off the top of my head sorry but that may help you in finding the slang for it.

UnintentionalChaos - 25-5-2010 at 13:09

Quote: Originally posted by solo  
.....have been pondering on the Cleisen Rearrangement of c6h5-ch2-ch2-(0-ch3) , following the rearrangement it would make phenylacetone.......it's a question not a fact and what would the ether be called?......also can the rearrangement function as outline or does it only with an allyl group.......solo


I think you mean Claisen rearrangement. You need an allyl vinyl ether, which you do not have. Part of the aromatic ring can stand in for one of the double bonds, but you don't have that set up properly either.

watson.fawkes - 25-5-2010 at 17:03

Quote: Originally posted by Panache  
The circumstance i'm interested in is a recirculating pump on startup running nothing other than water or water glycol. Does the pump initially do more work until the circuit is flowing or does the work being done by the pump simply gradient upwards leveling out once the pressure in the system is obtained. Are equations available to predict the pressure gradient in the system from startup until i guess you could call it steadystate.
Both the fluids you mention are perfectly good Newtonian fluids, at least to an approximation better than you could (easily) measure at home. So to your first question, there's no extra work at startup due to viscosity effects.

Non-steady-state pressure gradients transmit at the speed of sound in the fluid, which is quite fast, about 1.5 km/s. The sonic transmission time across the diameter of your lab-scale apparatus is only a millisecond at most, which is nothing to worry about.

querjek - 26-5-2010 at 08:17

I've been googling all morning but can't seem to get anywhere with this: does anybody know of a non-electrolytic method of copper methoxide production?

bbartlog - 29-5-2010 at 17:38

Some text in patent 5006508 says:

'...A typical reaction may be that between lithium metal and anhydrous methanol to produce a methanolic solution of lithium methoxide. This lithium methoxide solution is filtered to remove any precipitate that may have formed and the filtered solution is added to a solution of anhydrous cupric chloride in anhydrous methanol. The lithium methoxide and cupric chloride react to form copper (II) methoxide which precipitates and lithium chloride which is soluble in methanol. The copper methoxide is collected by filtration and washed with several methanol washes...'

There is also this, of which I can however only see the abstract:
Kendric J. Nelson, Ilia A. Guzei, Gregory S. Lund, Robert W. McGaff, Copper(II) methoxide: direct solventothermal synthesis and X-ray crystal structure, Polyhedron, Volume 21, Issue 20, 1 September 2002, Pages 2017-2020, ISSN 0277-5387, DOI: 10.1016/S0277-5387(02)01132-4.
(http://www.sciencedirect.com/science/article/B6TH8-46MK5SK-2...)

Pope - 1-6-2010 at 04:54

Hmmm... I continued researching into it more and I don't know how it evaded me, but I would like to verify if that L-Pac is reduced to either Phenylpropanolamin or Ephedrine by the same way P2P is reduced to Amphetamine or Methamphetamine e.g. reductive alkylation/amintion. If so I feel like quite the moron and would like to apologise for the inconvenience.

This is not drug related but it was the only way to lay mention of the reduction processors.

Sedit - 5-6-2010 at 17:54

Pope you need to understand the mechanics of why it would form Ephedrine or PPA. Look into why normal Al/Hg reductive amination work and the mechanics behind it. PS: you will more then likely not find the mechanics information on Rhodiums archives. Look more for the reductive amination of ketones instead of limiting yourself to a certine substate in general. Also look into imine as well on wiki to get an idea of whats truely getting reduced.


==========================================
QUESTION:


Does anyone have the solublity data of the Vanillin Bisulfate adduct it either H2O or EtOH?

Im attempting to extract it from an impure OTC Vanilla extract which contains mostly H2O and some EtOH so I figured add sodium Metabisulfate till saturation and quickly filter then add EtOH to precipitate the adduct but I want to make sure I got what I desire.

[Edited on 6-6-2010 by Sedit]

UnintentionalChaos - 5-6-2010 at 18:27


QUESTION:


Does anyone have the solublity data of the Vanillin Bisulfate adduct it either H2O or EtOH?

Im attempting to extract it from an impure OTC Vanilla extract which contains mostly H2O and some EtOH so I figured add sodium Metabisulfate till saturation and quickly filter then add EtOH to precipitate the adduct but I want to make sure I got what I desire.

[Edited on 6-6-2010 by Sedit][/rquote]

Buy a cheap "white" decorating vanilla. Boil down with something to help exclude air until the fumes aren't flammable, or watch until the temp hits 100C. I topped the beaker with tinfoil and a pinhole to help with this. Then chill the mixture, as vanillin has a pretty steep solubility curve. You should get white to off white (off white more likely) needles forming- let it cool undisturbed and slowly or you get mushy fine crystals that trap the solvent- molten vanillin may also oil out as it has a low melting point out but should freeze.

I have NMR verification that (at least as of 4 years ago) Wilton's clear decorator vanilla was all methyl vanillin with no ethyl vanillin or other flavors (at least of the product as isolated above and recrystallized from boiling water).

Sedit - 6-6-2010 at 09:57

Thanks UC. I did a test tube run by adding "dirty" (ill explain in abit) Vanilla extract to Metabisulfate until no more would dissolve and filtered it. Upon adding more EtOH as expected the clear solution took a cloudy appreance and a fine white precipitate was noted but not isolated due to the small scale.

By Dirty I mean all kind of crap like glycols and sugers which I obviously don't want but its what I had on hand when i first decided to try it for the hell of it.

Im doing a small test as well along the lines of what you just stated and am evaporating the entire contents down to a syrup(more then likely) mainly to remove the EtOH and H2O then I was considering adding a .5:1:1 H2O/EtOH/Et2O mixture to dissolve the syrup and slowly allow to evaporate and see if any crystals form. If they indeed do I could adjust the ratios of the crystalizing solvents to get more bang for my buck. Indeed however I will search for the brand you made mention of because I want pure vanillin when im done just to say I did it.

PS: What about NH3 adduct, that should accomplish the same goal simular to how BnO is isolated forming the trimer and then a simple acid hydrolysis would recover my Vanillin correct. May be one of the simpler methods if proves true... No other way to find out:D Ill be right back , lol

[Edited on 6-6-2010 by Sedit]

Sedit - 6-6-2010 at 13:34

Nope not even the slightest bit of precipitate using (aq) NH3 although the R-OH contents in the mix may possibly be solvating the adduct but I don't know. Ill get some cleaner material tommorow and give your metho a shot UC.

I must ask. Whats the need to exclude air from the mixture needed for? Am I doing this just to remove the EtOH and allow the Vanilin to precipitate due to its low solubility in H2O or is there something im missing?

Lambda-Eyde - 8-6-2010 at 15:49

I made 500 ml of 0,1M (NH<sub>4</sub>;)<sub>2</sub>Fe(SO<sub>4</sub>;)<sub>2</sub> solution today by dissolving 19,61 g of the dry powder to 500 ml with deionized water in a volumetric flask. It is very old, however it had the characteristic blue color of Fe(II) it should have before I dissolved it. The solution, however, was more like a murky yellow. Now I'm seeing a brown precipitate, but not much.
Is this due to atmospheric oxidation or must it have been something in the water that caused it? Mohr's salt is said to be stable against air oxidation because of the acidic ammonium ions present in the solution.

Anyhow, would it be okay if I filtered off the fine precipitate and standardized the solution against 0,02M KMnO<sub>4</sub>? The permanganate is standardized against sodium oxalate, a primary standard.

mewrox99 - 10-6-2010 at 17:13

Why are potassium salts less soluble than sodium salts when the anion is oxidizing

JohnWW - 10-6-2010 at 18:08

Probably because of the larger size of the K+ cation, which must make formation of crystal structures on precipitation or crystallization less enthalpy-reducing than with Na+ cations and the same anions. This would be in spite of the larger size of the K+ cation resulting it being less strongly hydrated in solution (which explains why K salts, e.g. KMnO4, are frequently anhydrous in solid crystal form, while the corresponding crystalline Na salts, including the permanganate, are hydrated, or at least have more water molecules than the K salts.).

Nicodem - 14-6-2010 at 07:33

Quote: Originally posted by mewrox99  
Why are potassium salts less soluble than sodium salts when the anion is oxidizing

The oxidizing nature of the anion has nothing to do with the solubility (I assume you mean aqueous?). Solubility is a property that depends on lots of things. In the case of sodium vs. potassium salts it is mostly about the difference in hydration of the sodium cation vs. hydration of the potassium ion and the energy required to break/form the crystalline phase. Little can be said about the crystal structure of Na/K salts and its energies as it depends also on the anion (there are some useful estimates based on the difference in ion sizes, but I forgot the name of the theory). However, the sodium cation is more acidic than the potassium cation (due to the smaller size of Na(+) at the equal charge). Therefore the sodium cation solvates more strongly (more energetically) with a basic solvent like water. Still, this does not mean that every sodium salt will be more soluble than the corresponding potassium salt.

mewrox99 - 14-6-2010 at 12:28

Has anyone done any 'pyro' reactions with Calcium Hypochlorite

Seeing how it reacts with metallic powders, charcoal, sugar etc

Lambda-Eyde - 14-6-2010 at 16:19

I'm not sure if this is the right place to ask, but here goes anyways: Can anyone recommend an introductory book to chromatography? I'm especially interested in thin layer chromatography, not so much in modern methods requiring expensive instruments. The reason for this is that none of my basic chemistry books (two high school level, one university level) even mention TLC, and I would like to follow my organic reactions using TLC. I understand the basic principles, but I would like to learn it more in-depth and learn how to select various solvent systems, interpret results and solve practical problems. Also, it would be nice to have a physical book instead of having to resort to various online guides all the time. I wouldn't mind learning about GC, HPLC and column chromatography either, but that's not my primary interest.

The keywords are: Basic, nice price, focused on TLC and with a practical approach so I can relate it to real experiments.


I'm also looking for a book on basic practical electric "engineering", especially high voltage systems. I would love to build a Tesla coil someday, but I would also like to do so while understanding what I'm doing and not getting killed. I feel that I also need some literature on electricity that is more oriented around practical subjects, but not excluding the mathematics needed. The only knowledge of electricity I have is from a basic physics course (Chapters "Electricity" and "Semiconductor technology"), and I'm having a hard time relating this to real projects.

Again, the keywords are: Basic, nice price, practically oriented and preferably something related to HV projects.


Also, I'm going to shamelessly bump my question posted earlier:
Quote: Originally posted by Lambda-Eyde  
I made 500 ml of 0,1M (NH<sub>4</sub>;)<sub>2</sub>Fe(SO<sub>4</sub>;)<sub>2</sub> solution today by dissolving 19,61 g of the dry powder to 500 ml with deionized water in a volumetric flask. It is very old, however it had the characteristic blue color of Fe(II) it should have before I dissolved it. The solution, however, was more like a murky yellow. Now I'm seeing a brown precipitate, but not much.
Is this due to atmospheric oxidation or must it have been something in the water that caused it? Mohr's salt is said to be stable against air oxidation because of the acidic ammonium ions present in the solution.

Anyhow, would it be okay if I filtered off the fine precipitate and standardized the solution against 0,02M KMnO<sub>4</sub>? The permanganate is standardized against sodium oxalate, a primary standard.

More has precipitated and the solution is now a slight piss-yellow. Is there even any Fe(II) left in solution? I find this quite confusing as this compound is supposed to be stable with respect to aerial oxidation. If I remember I can get around to take a picture of it tomorrow.


Any input is GREATLY appreciated! :)

solo - 15-6-2010 at 09:36

.... a two part Question....
1. how can one degrade or oxidize a secondary aliphatic amine to its schiff base or imine,
2. if by using a single isomer of the secondary aliphatic amine , will the degradation or oxidation of the compound recimize it once its reduced back to the secondary aliphatic amine?

.......an interesting project, would appreciate some input, i have tried the benzaldehyde reflux , and reduced it back to the amine with NaBH4 but have noted no change ...that is chirality did not change, also tried the HI boil to destabilize the chiral center but no recimization occurred....solo

Nicodem - 15-6-2010 at 10:59

Quote: Originally posted by Lambda-Eyde  
I'm not sure if this is the right place to ask, but here goes anyways: Can anyone recommend an introductory book to chromatography? I'm especially interested in thin layer chromatography, not so much in modern methods requiring expensive instruments.

At least one book fully dedicated to TLC is freely available in the grey zones of internet: Thin Layer Chromatography 1 (Jork et al, 1990). It is available in the DJVU format so you will need a viewer (there are two or more freeware ones available).
Quote: Originally posted by solo  
.... a two part Question....
1. how can one degrade or oxidize a secondary aliphatic amine to its schiff base or imine,
2. if by using a single isomer of the secondary aliphatic amine , will the degradation or oxidation of the compound recimize it once its reduced back to the secondary aliphatic amine?

.......an interesting project, would appreciate some input, i have tried the benzaldehyde reflux , and reduced it back to the amine with NaBH4 but have noted no change ...that is chirality did not change, also tried the HI boil to destabilize the chiral center but no recimization occurred....solo

1. Depends on the amine. If it is pretty robust toward electrophiles, the amine can be N-halogenated and transformed to the imine via elimination in the presence of a base (for example using DIPEA) or by heating. Otherwise you are left with oxidation to the corresponding ketone and reforming the imine with the proper amine.

2. Assuming that by single isomer, you mean enantiomer, it then depends on where the chiral centre is. If it is on the carbon attached directly to the amine nitrogen then yes, you get a racemic amine after reduction, provided of course you do the reduction in a symmetric reaction environment. If the chiral centre is elsewhere, then such a treatment forms two diastereoisomers in a ratio depending on the asymmetric induction from the existing chirality (usually a near to 1:1 ratio unless the existing chiral centre is close to the imine group).

Refluxing with benzaldehyde will do nothing to a secondary amine except for forming an equilibrium with the hemiaminal (and/or iminium salt if an acid is added and water removed via Dean-Stark, etc.). So I really can't see what intermediate have you been reducing with NaBH4 and how it would have given you back the starting material (reduction of the hemiaminal or iminium salt would have given you the corresponding N-benzylamine).
Refluxing in HI certainly does not racemize an enantiomer having the chiral center at the carbon attached to the amine nitrogen, except if there is a neighbouring group allowing racemization via enolization or some similar mechanism. But if that was the case such harsh conditions would have most probably decomposed the substrate - just heating up the hydrochloride salt would have already been enough for a racemization via enolization if this was viable (one such examples are alpha-aminoketones or "activated" alpha-aminoacids). If there is no such neighbouring group then treatment with acid does nothing. In some cases a nickel catalysts can be used for a one step racemization via dehydrogenation/hydrogenation mechanism but the amine must be robust enough not to succumb to side reactions.

DJF90 - 15-6-2010 at 12:39

Lambda-Eyde: I have ten or so books on chromatography of several varieties. I will upload them and post a link after finals are over (finish on saturday, expect them sometime early next week).

Lambda-Eyde - 15-6-2010 at 13:12

Nicodem, thanks for the recommendation. I couldn't find that book in references, though. And there was only one for sale on Amazon at 300$.

DJF90: Thanks. I'm looking forward to that.

Am I the only one on the forum who prefers books in paper form? :P

DJF90 - 15-6-2010 at 23:24

No not at all, theres nothing better than a real book!

not_important - 16-6-2010 at 00:33

Quote: Originally posted by Lambda-Eyde  
...
Am I the only one on the forum who prefers books in paper form? :P


I generally prefer physical books, but more recently have increasingly been switching to digital for simple reference materials.

There's several reasons for this. The first is that I already have around 4 thousand books and journal issues of all sorts, moving gets to be a bit challenging if many flights of stairs are involved.

Then there's the problem of I'm using more references than fit on the desk/table, which electronic media solves. Plus I can copy out the currently useful bits of references, and combine them in a single location as part of the documentation of what I am working on.

Finally:
Quote:

And there was only one for sale on Amazon at 300$.

need I say more?

Lambda-Eyde - 17-6-2010 at 06:00

Quote: Originally posted by Lambda-Eyde  
I made 500 ml of 0,1M (NH<sub>4</sub>;)<sub>2</sub>Fe(SO<sub>4</sub>;)<sub>2</sub> solution today by dissolving 19,61 g of the dry powder to 500 ml with deionized water in a volumetric flask. It is very old, however it had the characteristic blue color of Fe(II) it should have before I dissolved it. The solution, however, was more like a murky yellow. Now I'm seeing a brown precipitate, but not much.
Is this due to atmospheric oxidation or must it have been something in the water that caused it? Mohr's salt is said to be stable against air oxidation because of the acidic ammonium ions present in the solution.

Anyhow, would it be okay if I filtered off the fine precipitate and standardized the solution against 0,02M KMnO<sub>4</sub>? The permanganate is standardized against sodium oxalate, a primary standard.


Here is a picture of the solution, nine days old:




A close-up on the precipitate:

querjek - 17-6-2010 at 14:00

Would a mixture of sodium bisulfite, sodium sulfite, and sodium dithionite suffice to make bisulfite adducts?

Panache - 18-6-2010 at 04:21

Quote: Originally posted by querjek  
Would a mixture of sodium bisulfite, sodium sulfite, and sodium dithionite suffice to make bisulfite adducts?


Bisulfite adducts were a common technique used to elucidate structure of organic molecules prior to H1 NMR. They used many adducts back then because this was the part of the regime which was the only way to determine structure.
Not all aldehydes and ketones form bisulphite adducts easily or simply or stably, some do and you can even use the adduct formation as a crude purification step, in other examples it forms difficulty or not at all.
So in answer to your question, without the ketone/aldehyde your intending to formt the bisulphite adduct of, no-one could reliably answer your question, what is your material?

querjek - 18-6-2010 at 05:38

The substrate is normal benzaldehyde. I've made its adduct before with sodium bisulfite, but I ran out, and found a product whose constituents are the three compounds I've listed above.

I know that benzaldehyde will produce bisulfite adducts, so that's not where my concern arises from. It's that, previously, I've added bisulfite solution to a solvent which I had forgotten to neutralize excess acid in, and that produced a nice lot of SO2. Although all of the compounds I'm asking about are basic salts, I figured I should ask before trying: I'm not sure what the specific basic conditions to form bisulfite adducts are.

Addon - 18-6-2010 at 06:22

first post

How long is concetrated NH3-Solution 25-30%) , conc. HNO3 (65%) and concentrated HCl (37%) durable? I store it at 10 degree celsius in original and dense containers (HCl and HNO3 in glass bottles, NH3 in plastik bottles), and i want to buy stuff for the next 15 years or so...

:=)

[Edited on 18-6-2010 by Addon]

the Z man - 18-6-2010 at 14:57

I'm trying to make Pb acetate using excess white table vinegar and 30% H2O2. How do I get it to crystallize? When I boil it down and cool it only becomes a brown sticky mass. Also is it normal that when I diluite it becomes turbid and a white powder precipitates?

12AX7 - 18-6-2010 at 23:24

Dilution with tap water precipitates lead carbonate, mostly. Use distilled water.

Brown is either an organic or iron impurity. Let it sit for a long time and see if any crystals form; if not, precipitate and start over.

Tim

Nicodem - 19-6-2010 at 06:26

Quote: Originally posted by Addon  
first post

How long is concetrated NH3-Solution 25-30%) , conc. HNO3 (65%) and concentrated HCl (37%) durable? I store it at 10 degree celsius in original and dense containers (HCl and HNO3 in glass bottles, NH3 in plastik bottles), and i want to buy stuff for the next 15 years or so...

For NH3(aq) and HCl(aq) it depends only on the stability of the containers they are stored in. There is pressure building up during the warmer times in bottles containing concentrated aq. solutions of HCl or NH3, so it is best to make sure the container is robust enough and/or these are stored in cold.
For conc. HNO3 stability it depends also on the temperature and exposure to light. Generally it can last for years with minimal decomposition to NO2 if stored in a glass container, cold and dark, but even if gets saturated with NO2, there are simple means to remove it (UTFSE).

Addon - 19-6-2010 at 10:11

thx a lot!

:=)

mewrox99 - 21-6-2010 at 03:53

How do you make Barium Permanganate from Ba(NO3)2 + KMnO4

12AX7 - 21-6-2010 at 08:17

Combine solutions and filter the Ba(MnO4)2 precipitate.

Tim

Copenhagen - 21-6-2010 at 08:53

Quote: Originally posted by DJF90  
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.


Good point - surface area will be a concern for this one. It makes me think of the benzene from sodium benzoate thread where a major improvement was made by adding steel wool to the reaction vessel to increase thermal conductivity. I'm going to try a set up where I use a pressure equalizing addition funnel to feed phenol pre-heated to 165 C slowly onto red hot zinc dust inside a 3-neck flask connected to a distillation column. I think it might help by increasing the amount of phenol in contact with the zinc dust and increase the amount of time it has to react with the zinc.

mewrox99 - 22-6-2010 at 15:15

Quote: Originally posted by 12AX7  
Combine solutions and filter the Ba(MnO4)2 precipitate.

Tim


Thanks. For some crazy reason I thought Ba permanganate was soluble

JohnWW - 22-6-2010 at 15:41

Salts of univalent oxy-anions are usually soluble, but it appears that Ba(MnO4)2 is an exception.
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