Sciencemadness Discussion Board - The short questions thread (2)

I tried "DE98070", but my search was unsuccessful. I'm thinking that there is a special format for entering old patents. Unfortunately, I dont know the year in which the English patent was published.

You can use DEPATISnet if you want the old German patents and don't have the year. As to the British patent, no idea how to get it without the year. Maybe it's referenced somewhere else.

hissingnoise - 12-7-2009 at 11:57

Quote: Originally posted by itchyfruit

Is cobalt nitrate supposed to be purple ?

It's a deep red colour on wikipedia. . .

querjek - 15-7-2009 at 10:01

The expected melting point of a compound is 189deg.C. When I tested a sample of mine, it melted around 204deg.C.

I remember something from general chemistry stating that impurities would lower the melting point of a compound (though honestly, that doesn't make sense to me if the impurities themselves have higher melting points).

Is the previous claim about impurities always true, or could my sample still be ok (by "ok", I mean, "at least could contain some of the compound")?

12AX7 - 15-7-2009 at 11:12

Different amount of moisture maybe?

Can you titrate the amount of anion or cation in the material?

Tim

merrlin - 15-7-2009 at 13:25

Quote: Originally posted by querjek

The expected melting point of a compound is 189deg.C. When I tested a sample of mine, it melted around 204deg.C.

I remember something from general chemistry stating that impurities would lower the melting point of a compound (though honestly, that doesn't make sense to me if the impurities themselves have higher melting points).

Is the previous claim about impurities always true, or could my sample still be ok (by "ok", I mean, "at least could contain some of the compound")?

The addition of impurities may lower or raise the melting point of a substance. The addition of chromium to gold can raise the melting point of gold, whereas the addition of about 3% silicon to gold will lower the melting point to about 363° C. Search for "eutectic" on Wikipedia and then check out the the reference link at the bottom of the page:

http://ocw.mit.edu/NR/rdonlyres/Materials-Science-and-Engine...

manimal - 15-7-2009 at 16:42

Since concentrated ammonia solution is so useful but rather scarce, I was thinking that a promising way to prepare it would be to heat intimately mixed ammonium sulfate and calcium hydroxide in a metal can and pipe the fumes into water.

I suppose that to judge the concentration, I can measure the volume increase. I am going for 100% increase in volume, because I intend to heat the CaOH2 and (NH4)2SO4 to a high temperature, and I am taking into account the 2 moles of water released by this reaction (I am heating it that hot to encourage the reaction to go to completion and to liberate the h2o from the CaSO4, which likes to hold onto it's water).

entropy51 - 15-7-2009 at 16:55

Manimal, maybe better to measure the density of your NH3 soln. in H2O. The density of strong ammonia (28%) is about 0.9 gm/mL. Weigh 10 mL to get an accurate density.

Be sure to cool the water you're absorbing the NH3 in, since absorbtion is exorthermic and NH3 is much more soluble in cold H2O.

UnintentionalChaos - 15-7-2009 at 18:06

Quote: Originally posted by entropy51

Manimal, maybe better to measure the density of your NH3 soln. in H2O. The density of strong ammonia (28%) is about 0.9 gm/mL. Weigh 10 mL to get an accurate density.

An appropriate volume depends largely on the capabilities of the scale being used. Hopefully, there are some volumetric flasks around as well.

crazyboy - 18-7-2009 at 14:18

What is a reasonable price for a rotovap with and without glassware? I am looking for one and the best I have seen is about $1200

. A little out of my price range but it is in excellent condition w/ glassware.

Any ideas where to get one? Anyone want to get rid of theirs?

DJF90 - 18-7-2009 at 15:20

I would say about £500 second hand. Varies by make and model, and there are some bargains out there (under £300 I have seen once), as well as some overpriced equipment. Obviously this probably depends on location too.

manimal - 19-7-2009 at 14:43

I want to extract paracetamol from pain medication. (I intend to invistigate it's uses as a polymerization inhibitor). This medicine contains polyethylene glycol, however, which is soluble in water and ethanol. Will the PEG recrystallize w/ the paracetamol after undertaking a water or ethanol extraction?

ammonium isocyanate - 19-7-2009 at 18:34

I'm not sure about recrystalization from water or ethanol (in part it would depend upon the concentrations of the two chemicals), but I know that PEG is insoluble in diethyl ether and straight chain hydrocarbons. Paracetamol is slightly soluble in diethyl ether, so if you have some lying around, it's worth a try. Another method would be to dissolve both in ethanol, evaporate/boil it away, and wash it with cold water (PEG is much more soluble in water than paracetamol).

DJF90 - 19-7-2009 at 19:46

Perhaps you could check a paper or something... Perhaps a paper like the one included below

Attachment: Paracetamol synthesis and extraction.pdf (144kB)
This file has been downloaded 4610 times

Attachment: Paracetamol solubility.pdf (98kB)
This file has been downloaded 34458 times

ammonium isocyanate - 19-7-2009 at 20:42

Can carbon dioxide bubbled through an hot aqueous solution of alkali long chain carboxylates (i.e. from biodiesel saponification) acidify the salts to the corresponding carboxylic acids and alkali carbonates?

I ask because this would be useful for recycling the alkali hydroxides (which can be made more easily from alkali carbonates than alkali chlorides) used in the splitting of triglycerides into free carboxylic acids.

UnintentionalChaos - 19-7-2009 at 20:49

Can carbon dioxide bubbled through an hot aqueous solution of alkali long chain carboxylates (i.e. from biodiesel saponification) acidify the salts to the corresponding carboxylic acids and alkali carbonates?

I ask because this would be useful for recycling the alkali hydroxides (which can be made more easily from alkali carbonates than alkali chlorides) used in the splitting of triglycerides into free carboxylic acids.

No. The pKa of the long chain fatty acids is lower than the pka1 of carbonic acid. If this were not the case, bars of soap would react with air to free fatty acids and sodium bicarbonate over time.

ammonium isocyanate - 19-7-2009 at 21:44

Ok thanks.

I couldn't find a source for the pKa of specific acids in peanut oil (the feedstock I'm using) other than the really big ones (oleic and linoleic). Do you know where I could find a reference? (Although now that I think about it yeah its pretty obvious that CO2 couldn't degrade alkali carboxylate salts.)

I suppose I could use nitric acid and decompose the resulting nitrate to an oxide and hydrolyze that to the hydroxide, but HNO3 is pretty precious stuff so it would probably only be worthwhile with lithium salts and even then I'm not so sure.

UnintentionalChaos - 19-7-2009 at 22:37

The question is why? NaOH and KOH are incredibly cheap feedstocks.

I suspect that all the long chain fatty acids have largely the same pKa values since they all have a terminal carboxylic acid at the end of a lot of methylene units. Whatever double bonds the acid may have are so far downchain as to have negligible effect on the acidity of that terminal group

Nicodem - 19-7-2009 at 22:59

What other fatty acids besides linoleic and oleic acids do you get from peanut oil anyway. There are also some palmitic acid and arachidic acid triglycerides in peanut oil along with some minor ones. For example, arachidic acid has the same pKa like all fatty acids (4.78 in water), but the basicity of fatty acid salts depends extremely on many factors, because the anions aggregate to higher structures depending on concentration, counterion, temperature, etc. So it makes no sense to rely on these pKa numbers. Just imagine that the apparent pKa of fatty acids ordered in monolayers or bilayers can reach up to 8 or more! It suffices to think of fatty acids in their molecular form as acids of similar strength as their shorter aliphatic carboxylic acids (pKa ~ 4.7), though they are rarely or ever in molecular form when in aqueous solutions.

I assume the easiest way to recycle alkali hydroxides from your waste (NaCl or KCl) is to again make them like they are often made from chlorides in the first place, via electrolysis in a separated electrolytic cell (see Kirk-Othmer or other sources for references). You get the aqueous solution of NaOH (or KOH) which takes a lot of energy to dry. You can also recycle HCl (and some energy as heat) as well if you lead the Cl2 and H2 in a burner. Of course, unless you have an industrial production this would make no sense to get involved in such a infrastructural endeavour. NaOH and HCl are just too cheap and you will never beat the economy of the industry with anything you build for your small scale work.

not_important - 19-7-2009 at 23:34

Yes, under moderately high pressure. Ten to twenty bar of CO2 should convert the sodium salts mostly to the free acids and NaHCO3. You need to maintain the pressure until the carboxylic acids and the NaHCO3 have been separated.

-----

However this won't help much, simple base transesterfication tends to be wasteful of raw materials; the use of a large excess of alcohol with the attendant need for recovery is one example, the problem with FFA is another. Small scale batch mode operations make it even worse, "backyard biodiesel" is rather dirty and inefficient.

Reactive distillation is likely much more efficient, as FFA cause it no problem and the aqueous glycerol stream is salt free.

[Edited on 20-7-2009 by not_important]

ammonium isocyanate - 20-7-2009 at 09:19

Quote: Originally posted by UnintentionalChaos

The question is why? NaOH and KOH are incredibly cheap feedstocks.

True, but LiOH is not. I don't plan on producing biodiesel for use as a fuel, but instead I plan to seperate out various fatty acids as an experiment in itself, and for use as reagents in their own right. As such, I would be performing a saponification reaction instead of transesterification.

I don't have a vacuum distillation aparatus, and don't plan on buying one. Therefore, the easiest way for me to seperate out polyunsaturated fatty acids from other fatty acids is to prepare their lithium salts and then dissolve them in acetone ,lithium salts of polyunsaturated fatty acids are soluble in acetone, but others are not (I don't have the reference on hand, but I could dig it up if you want). Problem is, all the fatty acids must be converted to lithium salts, and thus alot of waste lithium would be produced in this process that I would prefer to recycle or use for other experiment (I suppose preparing lithium chlorates/perchlorates could be interesting).

My proposed method would be as follows:

-Split triglycerides into glycerin and sodium salts using NaOH
-Acidify the sodium salts using HCl
-Prepare lithium salts by adding Li2CO3
-Seperate out the polyunsaturated fats using acetone
-Acidify the lithium salts using HCl

This produces alot of waste NaCl and LiCl. If I could recycle LiOH, it would be reasonable cheap for me to use this, thus cutting out alot of steps and drastically reducing the amount of waste produced. Obviously I could regenerate LiOH from LiCl, but I don't have a suitable cell with a semi-permeable membrane and can't find one at a reasonable price (and also I really, really don't want to use the mercury amalgam method). Additionally, what would I do with all the Cl2 generated?

The only other way I can think of to regenerate LiOH without much of an aparatus would be to react LiCl with Na2CO3, precipitating the slightly soluble Li2CO3, then react that with Ca2NO3, precipitating CaCO3, and decompose the resulting LiNO3.

[Edited on 20-7-2009 by ammonium isocyanate]

Paddywhacker - 20-7-2009 at 13:56

You could use an anion-exchange resin. Regenerate in the OH- form with dilute NaOH. Then pass your Li salt solution through.

The trouble is that you will have to do it in small batches to avoid overloading the resin capacity, and you will end up with a dilute solution requiring a lot of CO2-free evaporation.

not_important - 20-7-2009 at 21:52

Prolonged boiling of fats with Li2CO3 in water will saponify the fats, as escape of CO2 drives the hydrolysis reaction towards completion.

Alternatively accept the NaCl waste formed, how big of a scale are you doing this on? Use acetone or MEK or even toluene to extract the free fatty acids after removing most of the water under reduced pressure. Evaporate the solvent, and do a preliminary separation by cooling the FFA to remove saturated acids, then mono-unsaturated ones. Repeat on each fraction to improve separation.

A trick sometimes used was to take the solids rich in higher melting acids, drop them in a Buchner supported by a flask, put the thing in an oven of just large box with a small heater, and slowly raise the temperature. The lower melting stuff 'sweats' out of the higher melting, although some of that does come with the sweat, and drips into the flask.

Rather than converting all of the rough cut of unsaturated acids to the Li salt, dissolve the free acids in acetone and add LiOH in IPA slowly. The more saturated acids should precipitate as their Li salts, when precipitate stops forming you've a solution that's mostly the polyunsaturated acids with small amounts of their Li salts and some of the more saturated acids per the solubility of their Li salts - those most be low but non-zero.

The combination of preliminary removal of the more saturated acids by cooling, and then fractional precipitation of their Li salts, should noticeably reduce the consumption of LiOH over simply converting all the FFA to Li salts.

At that point the use of CO2 under pressure as a way of recovering the lithium becomes attractive. Treat the acetone solution with 20 bar CO2, bleed the solution off through a filter to leave behind the lithium as carbonate. Treatment of that with a suspension of Ca(OH)2 in water gives a solution of mostly LiOH; evaporate in stainless steel 'flask' and extract the LiOH out with IPA to get the solution used to ppt the saturated fatty acids as Li salts.

Main consumables are NaOH, HCl or H2SO4, Ca(OH)2, and CO2. Waste is NaCl or Na2SO4, CaCO3.

ammonium isocyanate - 21-7-2009 at 19:55

Thanks not_important for all the suggestions!

I only plan on using about a gallon of peanut oil, so it's not worth building an expensive aparatus. The method I plan on using is as follows:

1. Saponify the triglycerides with dirt-cheap NaOH.
2. Seperate out the glycerin and possibly purify it.
3. Acidify with HCl, H2SO4, or H3PO4 (haven't decided).
4. Remove H2O and extract the fatty acids with acetone.
5. Evaporate acetone.
6. Remove saturated fats by cooling in an ice bath.
7. Seperated out all but the most saturated fats in a dry ice/isopropanol bath.
8. Dissolve the remaining liquid fatty acids in acetone.
9. Add the LiOH in isopropanol until phenolphthalein indicator turns pink.
10. Remove the precipitate and react with iodomethane to produce the methyl ester.

Anyone have trouble buying this chemical?

Agent MadHatter - 22-7-2009 at 22:43

Has anyone ever had trouble buying phosphorus oxychloride before? Has someone had the DEA come knocking just because they bought it?

JohnWW - 22-7-2009 at 22:56

There are various possible alternatives to POCl3 for organic chlorinations e.g. to form acyl chlorides, namely: PCl5, PCl3, AsCl3, SiCl4, BCl3, SCl2, S2Cl2, although these may require different reaction conditions. If you need to make an acyl chloride or similar just as an intermediate, the corresponding bromides should also work.

manimal - 23-7-2009 at 13:54

I'm not sure about recrystalization from water or ethanol (in part it would depend upon the concentrations of the two chemicals), but I know that PEG is insoluble in diethyl ether and straight chain hydrocarbons. Paracetamol is slightly soluble in diethyl ether, so if you have some lying around, it's worth a try. Another method would be to dissolve both in ethanol, evaporate/boil it away, and wash it with cold water (PEG is much more soluble in water than paracetamol).

I carried out an ethanol extraction with good results. However, I used a brass rod to stir the solution around, and the solution turned red. Obviously a complex of some kind, caused by residual copper corrosion. Anyone familiar with that sort?

bfesser - 23-7-2009 at 14:09

I'm refluxing an aqueous mixture of sodium hydroxide and acetylsalicylic acid to prepare salicylic acid. Under these conditions, is there any possibility of forming residual benzene (concerned with safety, but my bet is no):
~1.55g ASA
~1.37g NaOH
~25 mL solution

Am I driving it too hard, i.e. to sodium phenoxide?

[Edited on 7/23/09 by bfesser]

[Edited on 7/23/09 by bfesser]

crazyboy - 23-7-2009 at 14:29

I suggest acid hydrolysis it is far more effective and cleaner. From my notes:

10g ASA are added to a 1L Erlenmeyer flask. 700ml hot water are added and the ASA dissolves with stirring and gentle heating. A small amount of concentrated hydrochloric acid is slowly added until a pH of 2 is reached. Several boiling stones are added to promote even boiling and the mixture is heated to a light boil for one hour. The solution is poured into a 600ml beaker and SA precipitates as the solution cools.

The fluffy mass is filtered and washed with 200ml ice cold water. Yield: 5.15g.

As for your original question, no it is very unlikely that you will produce benzene.

DJF90 - 23-7-2009 at 14:30

You're not going to get phenoxide from a salicylic acid derivative unless it decarboxylates. The concentration of your NaOH solution is less than 1M, so there is little danger of that! I would probably use a 2M solution myself, i.e. 2g NaOH/ 25mls.

Acid hydrolysis is an equilibrium reaction so you would have to distil out a component to shift the equilirbium and make it tend towards a quantitative yield. Personally in this case I would use base hydrolysis as neither component of the ester (Acetic acid and salicylic acid) are particularly volatile at the temperatures used.

[Edited on 23-7-2009 by DJF90]

bfesser - 23-7-2009 at 15:44

crazyboy:
Thanks for the advice, but I have no hydrochloric acid, only sulfuric. And I don't want to waste any of that preparing hydrochloric. Could I substitute the sulfuric in for the hydrochloric without appreciable side reactions?

DJF90:
That was the last of my NaOH, unfortunately. Thanks for the advice, though.

crazyboy - 23-7-2009 at 15:48

Quote: Originally posted by bfesser

crazyboy:
Thanks for the advice, but I have no hydrochloric acid, only sulfuric. And I don't want to waste any of that preparing hydrochloric. Could I substitute the sulfuric in for the hydrochloric without appreciable side reactions?

Never tried it but I don't see why not.

crazyboy - 23-7-2009 at 20:05

A few random questions:

1. How interchangeable are potassium and sodium salts? Assuming you adjust the amount to the right molarity can NaOH be substituted for KOH and K2CO3 be used in place of Na2CO3?

2. Is xylene a suitable substitution for toluene as an organic solvent?

[Edited on 24-7-2009 by crazyboy]

ammonium isocyanate - 24-7-2009 at 09:45

Crazyboy,

1. It all depends on the intended purpose of the salts. If, say, all you are trying to do is neutralize an acid and don't care about what salt is formed, then they are interchangable. However, sodium and potassium salts do exhibit different solubility characteristics, so if you are trying to perform a double-displacement reaction, it may be effected (although most would probably still work). Overall, they are pretty much the same for most purposes.

2. Usually. Xylene is usually a mix of isomers, which would cause problems as a reactant but probably not as a solvent. The main difference is that the bp of xylene is about 30*C higher, which may be good or bad depending on the application.

kclo4 - 25-7-2009 at 00:42

Also potassium hydroxide, and potassium carbonate are stronger bases then their corresponding sodium salts. That could be important in some situations.

User - 25-7-2009 at 05:41

I was wondering, is there any reaction between sodiumbicarbonate x hydrate and etOH.
I am asking this because i would like to neutralize/clean up an alcoholic solution.
Then it could be distilled so any residue would be eliminated.
Anyone any ideas about this.

manimal - 25-7-2009 at 11:52

No, barcarbonate will not react with ethanol. Nor is it soluble to any appreciable extent.

Regarding this, I found this method to work, but it required strong heating, was slow, and gave variable yields of ammonia with unknown amounts of water. I also tried pyrolysis of urea, which was a disaster. It expanded in volume and flowed into the recieving flask, and set up into a rock-hard chunk that made it necessary to discard my 'flask' (actually a metal can).

The best results I got were from heating 10% ammonia under reflux and piping the fumes into water. I would recommend this for the purpose of preparing concentrated ammonia.

querjek - 25-7-2009 at 13:02

I have a solution of 90% acetic acid.

The bps of acetic acid and water only vary by 18deg.C.

I don't necessarily need all of the water out, but it would be nice, and googling "acetic acid distillation" has only sorta helped.

What would be my best be to concentrate the acetic acid?

Thanks!

ammonium isocyanate - 25-7-2009 at 13:09

There are a number of ways you could do it. One would be to react the solution with a base and then acidify the salt with sulfuric or phosphoric acid. This is the method I use, but only because I start with vinegar that is only 4% acetic acid by weight. For you, I would recommend the use of a dessicant such as MgSO4 or CaCl2.

manimal - 25-7-2009 at 14:40

There are a number of ways you could do it. One would be to react the solution with a base and then acidify the salt with sulfuric or phosphoric acid. This is the method I use, but only because I start with vinegar that is only 4% acetic acid by weight. For you, I would recommend the use of a dessicant such as MgSO4 or CaCl2.

How do you do this? Do you add the unhydrous sodium acetate to H2SO4, and then distill?

ammonium isocyanate - 25-7-2009 at 15:33

How do you do this? Do you add the unhydrous sodium acetate to H2SO4, and then distill?

You got it. In fact, if you use sulfuric acid you don't need totally anhydrous acetate as the sulfuric acid acts as a dessicant.

entropy51 - 25-7-2009 at 16:43

Quote: Originally posted by querjek

I have a solution of 90% acetic acid.

The bps of acetic acid and water only vary by 18deg.C.

I don't necessarily need all of the water out, but it would be nice, and googling "acetic acid distillation" has only sorta helped.

What would be my best be to concentrate the acetic acid?

Thanks!

This is an old problem, and there are threads on it that you should consult.

MgSO4 and CaCl2 won't remove the last of the water from 90% acetic acid. It's very difficult to fractionally distill off the water. Since you're already almost to glacial acid converting to NaOAc and distilling with H2SO4 is fairly wasteful. And the NaOAc does need to be anhydrous, I wish people wouldn't give advice when they're clueless. It doesn't make them look smart to say something when it's obvious that they are just blowing smoke out their arse. Distilling the NaOAc trihydrate with H2SO4 will give approximately the 90% acid you already have.

One way is to add enough EtOAc to form an azeotrope with the water in your 90% acid and distill with a good fractionating column. First the azeotrope will distill at about 70 C, then the EtOAc will come over at about 77 C, and finally glacial AcOH will distill at about 118 C.

But consult the existing threads. Their has been a lot of discussion about this.

In one of those threads, SC Wack posted a promising looking patent in which wet AcOH is frozen in vacuum and the H2O evaporated off leaving crystals of AcOH. It looked workable, but it didn't work for me. I think it should under the right conditions.

[Edited on 26-7-2009 by entropy51]

ammonium isocyanate - 26-7-2009 at 09:55

You could use P4O10 as a dessicant, but that is hard to come by (at least for me).

I have a question of my own. I had a solution of CuSO4 and tried to dispose of it with NaHCO3 and Na2CO3. A volumous greenish precipitate formed. However, a deep blue solution remained, presumably containing cupric cations. I can't think of any reason ammonia or any other chemical would be present that would complex with Cu++ ions. Does HCO3- form any complex with copper? I can't really dispose of this solution until I know it contains no copper ions. Any ideas for precipitating them?

turd - 27-7-2009 at 11:29

Is it possible to obtain 95% on a Reimer-Tiemann? I was always under the impression that yields of the Reimer-Tiemann are mediocre, yet SciFinder claims that some Chinese claim that they obtained 95% on a modified Reimer-Tiemann (with tertiary amine) on p-methoxyphenol:
Liaoning Shifan Daxue Xuebao, Ziran Kexueban (2007), 30(3), 328-329. (Journal written in Chinese)

Here: http://pubs.acs.org/doi/suppl/10.1021/ol800507m another group of Chinese claims 91% for Et3N, CH3CN, paraformaldehyde.
The base must be a tertiary amine, since other amines would condense with the aldehyde/formaldehyde and give water, right?
And triethanolamine (OTC for me) cannot be used, since the OH groups are too acidic, isn't it?

entropy51 - 27-7-2009 at 16:55

Is it possible to obtain 95% on a Reimer-Tiemann?

Is Chinese borosilicate glass the same as Pyrex?

not_important - 28-7-2009 at 05:46

..I had a solution of CuSO4 and tried to dispose of it with NaHCO3 and Na2CO3. A volumous greenish precipitate formed. However, a deep blue solution remained, presumably containing cupric cations. I can't think of any reason ammonia or any other chemical would be present that would complex with Cu++ ions. Does HCO3- form any complex with copper? I can't really dispose of this solution until I know it contains no copper ions. Any ideas for precipitating them?

The basic copper carbonates, you're not going to see true CuCO3, form soluble complexes with excess carbonate. Try boiling in a wide mouthed container for a quarter hour or so, this may convert all the copper to cupric oxide.

Alternatively let the solution evaporate to dryness.

12AX7 - 28-7-2009 at 09:12

Indeed, it's better to neutralize a copper solution with sodium hydroxide than carbonate for exactly this reason.

You do run the risk of ammonia impurity and cuprate formation, but those should be controllable with purity and pH.

Boiling, so CuO forms immediately, is always a good idea. Besides precipitating a stable form, it also settles faster and denser.

Tim

ammonium isocyanate - 28-7-2009 at 10:48

Hmmm... I guess I'll have to try boiling, as evaporating isn't an option given that I have ~600mL of the solution.

An interesting side-note, I added more carbonate to one batch than the other, and then tried nutralizing that one with vinegar. It is now a more turquoise/light blue color, whereas the other a deeper blue. Addition of alkali halides to the light blue solution (just in case Cu++ ions had been reduced to Cu+ ions) results in the evolution of copious quantaties of gas, predsumably CO2 as no odor was observed, it did not support flame, and did not really affect a pH test strip very much, nor was any smoke-like affect shown, as would be expected from hydrogen halides. No precipitate forms. No reaction is observed with the other solution.

Addition of almunium or magnesium poweder to the dark blue solution results in a single-displacement reaction, yielding what appears to be copper metal and leaving a purple solution behind. No reaction is observed with the light blue solution.

turd - 28-7-2009 at 12:35

Quote: Originally posted by entropy51

Is it possible to obtain 95% on a Reimer-Tiemann?

Is Chinese borosilicate glass the same as Pyrex?

Why do you post if your intent is not to help?

Scifinder states the reagents as 4-methoxyphenol, CHCl3, EtOH, NaOH and Et3N. My bet is on translation error or nonsense publication. I've seen too many of both lately.

ammonium isocyanate - 29-7-2009 at 13:21

What is the best way to dispose of chromate/dichromate compounds? I don't have any soluble lead (II) compounds, so I can't go the lead chromate route.

UnintentionalChaos - 29-7-2009 at 13:52

What is the best way to dispose of chromate/dichromate compounds? I don't have any soluble lead (II) compounds, so I can't go the lead chromate route.

Flood the reaction mixture with sulfites and acid (generate SO2 in-situ), reducing all hexavalent chromium to relatively harmless Cr(III). You may want to heat a bit just to make sure it goes to completion.

[Edited on 7-29-09 by UnintentionalChaos]

JohnWW - 29-7-2009 at 14:20

PbCrO4 (which occurs as the rare mineral crocoite) is, like PbSO4, only very slightly soluble, which would, if a soluble Pb salt were available, allow it to be used to precipitate Cr(VI) from solution. It, and ZnCrO4 which is more soluble and less toxic, is of use as an intense red paint pigment, especially in anti-corrosive roof paints and metal primers.

[Edited on 30-7-09 by JohnWW]

entropy51 - 29-7-2009 at 15:39

Quote: Originally posted by entropy51

Is it possible to obtain 95% on a Reimer-Tiemann?

Is Chinese borosilicate glass the same as Pyrex?

The answer to both questions is no, obviously. Your inference of my intent was wrong, meadow muffin.

UnintentionalChaos - 29-7-2009 at 16:23

Quote: Originally posted by JohnWW

PbCrO4 (which occurs as the rare mineral crocoite) is, like PbSO4, only very slightly soluble, and could be used to precipitate Cr(VI) from solution. It, and ZnCrO4 which is more soluble, is of use as an intense red paint pigment, especially in anti-corrosive roof paints and primers.

1) he explicitly stated that he can't use lead as a precipitant.
2) must you include a wikipedia-like rundown of useless trivia with every post?

not_important - 29-7-2009 at 19:04