Sciencemadness Discussion Board - The short questions thread (2)

The short questions thread (2)

Pages: 1 2 3 .. 8

Nicodem - 10-4-2009 at 00:13

The old thread has become too long so please continue here. Those who never received a proper reply in the old thread, please feel free to repost the question.

sakshaug007 - 10-4-2009 at 18:39

Hello everyone,

I was wondering if anyone knows where to get teflon wool? I want to use it for packing my vigreux column for distillation of corrosive compounds.

Thanks a lot.

smuv - 10-4-2009 at 18:54

Vigreux columns are not meant to be packed. Additionally, i think it would be difficult to pack a vigreux column without breaking it.

A hempel column is meant to be packed, and there is little reason pack it with teflon wool as raschig rings, lessing rings, copper wool, stainless steel wool and even glass wool are cheeper, more readily available options. There is nothing you can distill without eating through the glass column that at least one of these packing materials cannot stand up to.

sakshaug007 - 10-4-2009 at 19:53

Quote: Originally posted by smuv

Vigreux columns are not meant to be packed. Additionally, i think it would be difficult to pack a vigreux column without breaking it.

A hempel column is meant to be packed, and there is little reason pack it with teflon wool as raschig rings, lessing rings, copper wool, stainless steel wool and even glass wool are cheeper, more readily available options. There is nothing you can distill without eating through the glass column that at least one of these packing materials cannot stand up to.

I can easily pack my column and have with steel wool before, it gives that much more surface area. I'm interested in Teflon wool because I want the most chemically inert substance for packing i.e. acid distillations, but I suppose I could use glass wool in this case.

Thanks for the reply

not_important - 10-4-2009 at 21:33

We are talking about the same here, yes? Vigreux columns are indeed not meant to be packed, doing so greatly increases their holdup, increases the pressure drop across them (they work well for vacuum distillations), and may actually decrease increase their HETP value.

For packed columns the better the thermal conductivity of the packing, the better the column works. Metal packing is preferable over glass or plastic, and beads over wool. Fibrous packings (wools and the like) tend to increase the holdup over other types, lower density Teflon's porosity also increases holdup; I think that Teflon wool would be a poor choice for any situation except those where a great deal of material will be distilled, such as solvent reclaiming.

With any decently constructed Vigreux there's not too much room for packing, and getting wool type packings into the interior would be quite difficult. On most of my Vigreuxs there are almost no clear paths through the length, if any at all; several are designed where opposing horns are of different lengths so there is no central path clear. And because on most the horns are tilted slightly downwards you'd never get a wool type packing back out.

If you bought a Vigreux that is easy to pack, you were ripped off badly. It's just plain shoddily constructed, there should not be anywhere that much open space in it.

sakshaug007 - 10-4-2009 at 22:26

Quote: Originally posted by not_important

On most of my Vigreuxs there are almost no clear paths through the length, if any at all; several are designed where opposing horns are of different lengths so there is no central path clear. And because on most the horns are tilted slightly downwards you'd never get a wool type packing back out.

If you bought a Vigreux that is easy to pack, you were ripped off badly. It's just plain shoddily constructed, there should not be anywhere that much open space in it.

Actually you know what I might have been ripped off, I purchased this vigreux column off of ebay and it does not have the obstructed horn pathways your referring to. On the outside it looks like the one you have pictured but the interior of the tube is completely open and only the walls are roughened (bumpy). Maybe I should have researched these columns more before making an impulse purchase.

Thanks

[Edited on 11-4-2009 by sakshaug007]

Electrolysis of LiCl in DMSO

Strepta - 11-4-2009 at 04:43

@ sakshaug007: I was very interested in your results with LiCl in acetone which you mentioned in the first short question thread. As far as solubility of LiCl in acetone, you stated that you dissolved .35 g LiCl in 100cc acetone. Did you mean 3.5g? CRC tables I have state a solubility of 4.1 g LiCl in 100 ml acetone at 25C.

I have used LiCl in DMSO (5g in 60 cc--result is a viscous solution) and electrolyzed with a graphite anode and aluminum cathode. Distance between electrodes was ~ 1cm, voltage 15 V and current .25A. A black amorphous deposit was formed on the Al cathode which was very reactive with water. Almost no visible activity (no visible bubbles) at the anode. There is some bubbling at the cathode.

If I use graphite for the cathode, it begins to decompose, fouling the solution.

Panache - 11-4-2009 at 17:49

Quote: Originally posted by sakshaug007

Hello everyone,

I was wondering if anyone knows where to get teflon wool? I want to use it for packing my vigreux column for distillation of corrosive compounds.

Thanks a lot.

NI is correct on every detail, in fact incredibly succinct, i don't suffer from economy of detail so let me blather on a bit here.
For solvent reclaiming teflon wool is an excellent packing material, especially given you simply make it yourself. Important though, as not important stated (ha! what a sentence), that it is high or medium density not the porous low density type.

Get any old scrap solid teflon block and on your drill press get a 10-20 mm drill bit and drill out holes slowly, the wool shaves off as one continuous strand, fill a bag with this and then shake up the bag to intertwine it all. Takes about 15min from start to finish (assuming you have some scrap teflon, a drill press, 15mm bit and a plastic bag)

sparkgap - 11-4-2009 at 19:34

My questions in the previous thread seem to have been missed:

Quote:

1. Can diphenyl carbonate and diphenyl oxalate undergo Fries rearrangement?

2. Would the conditions for hydrogenolysis of a benzyl group reduce double and triple bonds too?

sparky (~_~)

Lambda-Eyde - 11-4-2009 at 19:45

What's the easiest way of cleaning commercial glass wool meant for house isolation? I.e. removing the yellow color/contaminants. Acetone, sulfuric acid?

Sedit - 11-4-2009 at 22:17

Can someone tell me what the name of FeS2O8 is? And a possible synthesis if your feeling up to it.

not_important - 11-4-2009 at 22:38

Persulfate AKA peroxydisulfate, and I suspect that it is not synthesed but a solution of Na/K/NH4 peroxydisulfate and a ferrous compound used.

while we're at it,

1. Can diphenyl carbonate and diphenyl oxalate undergo Fries rearrangement?

Quote:

In all instances only esters can be used with stable acyl components that can withstand the harsh conditions of the Fries rearrangement. If the aromatic or the acyl component is heavily substituted then the chemical yield will drop due to steric constraints.

So I doubt the oxalte will work to give standard Fries products, but diphenyl carbonate might yield the phenyl ester of hydroxybenzoic acids.

2. Would the conditions for hydrogenolysis of a benzyl group reduce double and triple bonds too?

Mostly likely, but there are some special reduction tricks around, so I can't say that with 100% confidence. Maybe use trityl instead?

Aubrey - 12-4-2009 at 12:26

Does anyone know the best route or have a reference to 2-bromobenzamide? I imagine a halogenation of benzamide, but expect this may result in an o,p mix which may be difficult to separate.

Arrhenius - 12-4-2009 at 14:19

The o-bromo product is the "kinetic" product, whereas the p-bromo is "thermodynamic". Running the aryl bromination with Fe + Br2 at reflux in say ether, or maybe dioxane. I had a quick search of the literature, but this stuff is now so readily available (albeit not cheaply) that it's hard to find the old preparation. The other option I see is accessing the bromine via diazotization of 2-aminobenzamide.

[Edited on 12-4-2009 by Arrhenius]

DJF90 - 12-4-2009 at 15:18

Arrhenius: If the target material is the kinetic product then how will carrying out the reaction at reflux help you, especially in dioxane (bp ~100C) :S Any kinetic product is generally favoured by low temperatures (generally below 0C, and usually at -78C) and short reaction times.

One possible route from benzamide is sulfonation at the 4 position, then mono nitration, desulfonation, reduction to 2-aminobenzamide and then a sandmeyer reaction of the diazo salt with CuBr.

Edit: Synthesis of benzamide has been covered here, and is a nice write up by smuv.

http://www.sciencemadness.org/talk/viewthread.php?tid=4201&a...

[Edited on 12-4-2009 by DJF90]

sakshaug007 - 12-4-2009 at 23:10

Hello,

Does anyone know if the Kolbe electrolysis reaction can be used for producing alkyl halides? I ask because I recently stumbled upon an article addressing the synthesis of methyl iodide by electrolysis of aqueous potassium acetate in the presence of iodine or potassium iodide.

Perhaps I need clarification on the mechanism because I'm not exactly sure how free radicals can react with anions to form covalent bonds, wouldn't there be an extra electron to account for? i.e. CH3* + I- => CH3-I + e-?

Thanks a lot!

[Edited on 13-4-2009 by sakshaug007]

Nicodem - 13-4-2009 at 02:39

Quote: Originally posted by Arrhenius

The o-bromo product is the "kinetic" product, whereas the p-bromo is "thermodynamic".

The -CONR2 group has -M and -I electronic effects and as such it directs the electrophilic aromatic substitution on the meta positions. But anyway, electrophilic bromination of aromatic compounds is not reversible enough that an equilibration of the products toward the thermodynamic ones would be feasible under normal reaction conditions. This was already discussed in a thread in relation to bromination of naphthalene when the desired product is 2-bromonaphthalene instead of the kinetic product (1-bromo...).

@sakshaug007: Yes, in principle, if you have a proper oxidant like elemental halogens present during the Kolbe electrolysis you should obtain the appropriate oxidation products of the radicals (like R-X) as the major product instead their coupling products (R-R). Check the literature, surely there must be a few papers on the topic. However, your equation is on the wrong track. Radicals rarely react with nucleophiles in a simple manner (but they can get reduced by them, form an anion radical, or abstract other radicals from it). Yet, they do very easily react with oxidants, in this case with elemental halogens: R* + X2 => R-X + X*. Of course, if you have KI then I2 (actually I3-

forms at the same electrode as the decarboxylation of the RCOO* radicals is going on.

@sparkgap: Though, ArO-CO-OAlkyl esters in theory should not work simply in the acid catalysed Fries rearrangement due to carbocation fragmentation problems, I can't see why a diaryl carbonate does not or would not work. With ArO-CO-OAr fragmentation should not be an issue and one should obtain the corresponding 2,2'-dihydroxybenzophenone. Yet there is nothing in SciFinder or Beilstein. With oxalates intermediate fragmentation is even more of an issue, and I could also find no literature examples.
All I could find is a reference for a photochemical Fries rearrangement of a phenyl carbonate (ArO-CO-OEt, 5 examples): Journal of Organic Chemistry, 37 (1972) 3160-3163. But this is mechanisticaly different from the normal Fries rearrangement.

sparkgap - 13-4-2009 at 07:31

Quote:

Yet there is nothing in SciFinder or Beilstein.

Just as I suspected.

Thanks, n_i and Nicodem, back to the drawing board...

~~~~~~~~~~~~

Another quickie: am I correct in assuming that perfoming a Baeyer-Villiger oxidation on a ketone with remote (i.e. non-conjugated) double bonds will also epoxidize those double bonds?

sparky (~_~)

Nicodem - 13-4-2009 at 07:55

Quote: Originally posted by sparkgap

Another quickie: am I correct in assuming that perfoming a Baeyer-Villiger oxidation on a ketone with remote (i.e. non-conjugated) double bonds will also epoxidize those double bonds?

I guess it would really be difficult to avoid epoxidation with the common reagents like MCPBA or other such peracids, particularly since double bonds, nonconjugated to any carbonyl or other electron withdrawing group, can be pretty reactive toward peracids and I think Baeyer-Villiger oxidation generally requires conditions harsher than epoxidation. Maybe, you could avoid epoxidation by using a persulfate in aqueous solution (with acetonitrile as cosolvent). I think persulfates might be faster at Baeyer-Villiger oxidation than alkene epoxidation (just guessing based on the mechanism). But again, this is a problem calling for a thorough literature search (when time allows...).

[Edited on 13/4/2009 by Nicodem]

Formatik - 13-4-2009 at 14:37

Quote: Originally posted by Lambda-Eyde

What's the easiest way of cleaning commercial glass wool meant for house isolation? I.e. removing the yellow color/contaminants. Acetone, sulfuric acid?

You could try boiling a piece in an excess of Piranha's acid for some time. That will oxidize organic compounds. But there might also be some other compounds in them.

Sedit - 13-4-2009 at 20:23

Im assuming he means insulation correct? Heating this material will turn it compleatly white but leave it a little brittle so its a bit of a trade off it you ask me. It still worked well enough for a filter pack if thats what you had in mind.

N-alkylation of piperazine

chemrox - 13-4-2009 at 21:54

I looked up an N-alkylation of piperazine in Org Syn and am puzzled by the rather straight forward procedure and results. I have two questions: why is it desireable to get the piperazine in its dihydrochloride form and why does the reaction with R-Cl yield the desired R-piperazine instead of a mix of monoalkylated and dialkylated products?

http://www.orgsyn.org/orgsyn/pdfs/CV5P0088.pdf

one of my professors was a checker on this reax so I feel especially confident in it.

DJF90 - 13-4-2009 at 23:04

When you add equal molar quantities of the hexahydrate and the dihydrochloride to the solution then there is only enough HCl "in solution" to protonate half the nitrogens. This should cause one nitrogen in each piperazine ring to be protonated and thus non-nucleophilic, and hence alkylation only occurs once.

[Edited on 14-4-2009 by DJF90]

solo - 16-4-2009 at 18:09

Will dehalogenation of an aliphatic compound by catalytic hydrogenation , using Pd/C also work in removing a Flouride, in the similar way as in the removal of Iodine, Bromine and Chloride....i know the Flouride is very reactive, hence it might be a challenge.....solo

sakshaug007 - 17-4-2009 at 10:22

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

sparkgap - 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 (~_~)

solo - 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

manimal - 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.

Paddywhacker - 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.

sakshaug007 - 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

Sedit - 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]

DJF90 - 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.

Sedit - 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.

Bolt - 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.

DJF90 - 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.

Bolt - 20-4-2009 at 17:31

I doubt it works. I believe that "The Chemistry of Phenols" is around here somewhere.

Sedit - 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.

Formatik - 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.

DJF90 - 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.

Formatik - 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?

DJF90 - 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]

Formatik - 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.

chemrox - 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.

Panache - 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.

Paddywhacker - 27-4-2009 at 02:16

Quote: Originally posted by chemrox

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.

Nicodem - 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.

Sedit - 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.

Maja - 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.

querjek - 29-4-2009 at 16:51

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

not_important - 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...

Jor - 30-4-2009 at 15:38

I have some 'barium carbonate', but i highly doubt it is this compound right now.
I added a spatula to 3M HNO3 and even after heating to boiling, none seems to dissolve at all.

So this means that it is not barium carbonate right?

Maybe my supplier confused and gave me the sulfate.

oh, and I am now an 'International Hazard' :cool:

[Edited on 30-4-2009 by Jor]

Globey - 30-4-2009 at 17:16

OK, so it's not a chem question...more physics. Let's say one took one's nice digital HDV camcorder, and set an ungodly frame rate like 1,000,000 FPS or higher (consumer models easily do 10K FPS). One's camera is in the sky very high up, looking down on an area of land let's say several hundred miles in length and maybe 150miles in width (not really important). Anyway, recording is commenced, and a powerful, visible laser (nicely expanded/culminated) and focused, is shone 90 Deg perpendicular, on a 2D plane to the camera. With a slowed down play back, would one be able to literally witness the beam of light as it travels in space? Where you could actually watch the leading edge of the laser light (like a slug of light) as it progresses in space? How cool would that be.

Paddywhacker - 30-4-2009 at 23:22

The effect can be seen in astronomy, where a the light front from a nova, for example, lights up a dust cloud that lies in space at an angle to the light wavefront. Observers here see an illuminated patch of cloud that appears to be moving faster than light, but it is only different parts of the light wavefront impacting the extended dust cloud.

appetsbud - 1-5-2009 at 11:12

I read: "Also, low concentrations (4 ppm) will anesthetize the nose, thus creating a potential for overexposure." on wiki, does this mean one will not be able to smell anything? or just the chemical in question? (NO2)

S.C. Wack - 2-5-2009 at 20:39

Quote: Originally posted by Jor

Maybe my supplier confused and gave me the sulfate.

I got sulfide in mine. Lots of H2S being carried by CO2 on acidification is not a good thing when you are expecting just CO2. The sulfate can be converted to the carbonate by boiling in baking soda solution for some hours.

Does anyone know where to get some Buchi condenser vacuum/water threaded connectors (GL-14) for cheap? Bought an older coil condenser, vapor duct, and receiver for $4 today. But the condenser didn't have the connectors and this bothers me. It would be a shame to rig something ugly to it.

What bothers me more is that I stopped bidding just last weekend at $40 (the winning bid) for one of the older Rotavap models, because it was older and worn, I already have a perfect + complete R110, and because it was missing the condenser and receiver.

kclo4 - 2-5-2009 at 22:11

Chloroplasts can reduce the blue DCPIP to a colorless compound. What is the reduction potential, or what else could chloroplasts reduce? I don't know why, but I am intrigued by the idea of stealing protons from chloroplasts.

Aubrey - 4-5-2009 at 08:17

I wish to dehydrate some Benzamide to Benzonitrile using Phosphorus Pentoxide. Does anyone have a reference to this reaction or know how I may go about it? Should i just heat the two until they form a melt for several hours?

Nicodem - 4-5-2009 at 09:07

There is a thread where the dehydration of benzamide to benzonitrile was discussed and some references posted. I take it that you did use the search engine and checked that out before asking, so here is what I found:

Chem. Listy 97 (2003) 1079-1082 give up to a 69% yield for this reaction (they get some of the trimeric product as well). I don't understand Czech, so you might want to use machine translation on that and, of course, also follow their references for this particular reaction since they made a literature review in the introduction section.

Then there is also a three sentence long experimental in patent US6617454:

Quote:

EXAMPLE 32

Regeneration of Benzonitrile From Benzamide

A two-necked flask was charged with of benzamide (0.0054 mol) and phosphorous pentoxide (1.2 g) in 10 ml of methanol. The mixture was stirred for 1 h at room temperature. After the completion of the reaction (followed by TLC), the mixture was concentrated on a rotovapor and benzonitrile was obtained quantitatively by vacuum distillation.

In my opinion the simplest method (for the amateurs) is described in Zhurnal Organicheskoi Khimii, 23 (1987) 1799 (there is also the English version of this journal, as far as I know). They use H2SO4 in acetonitrile to dehydrate benzamide (or other amides) while the acetonitrile gets hydrated.

Phenylpropanolamine from alanine and benzoic anhydride

turd - 4-5-2009 at 11:18

The thread on L-PAC got me thinking:
Wouldn't it be possible to obtain phenylpropanolamine by Dakin-West reaction between acetyl-alanine and benzoic anhydride, followed by reduction of the ketone and saponification?

J. Am. Chem. Soc., 1950, 72 (7), 3200-3203 says that reaction between benzoic anhydride and amino acids is slow and low yielding.
But in J. Org. Chem., 2003, 68 (7), 2623-2632 they claim a 73% yield with 3 eq. of benzoic anhydride in pyridine at 90°, but on a completely different substrate.

So?

User - 5-5-2009 at 11:20

I was wondering/wandering..

I could wel be that this is covered already.

I have nitric acid with a concentration of 53%.
Now i would like to bring up the concentration to azeotropic forming.
Well as far as iam concerned this can be achived by boiling it down.
The last time ive done this the fumes where acidic even before the 68% is reached, so this means loss of HNO3.

Could it be a nice idea to put my vigreux (60cm in length) on top so only the highest boiling fraction steams off?
Anyone any ideas on this, i thought it might be a neat trick.

regards, user

hissingnoise - 5-5-2009 at 11:47

Quote: Originally posted by User

Could it be a nice idea to put my vigreux (60cm in length) on top so only the highest boiling fraction steams off?
Anyone any ideas on this, i thought it might be a neat trick.

The lowest boiling fraction will boil off first. . .
It would be simpler to distill from H2SO4 and dilute afterwards.
Ooops!
Sorry User, I wasn't thinking---the solution will lose water until the azeotrope is reached. . .

[Edited on 5-5-2009 by hissingnoise]

User - 5-5-2009 at 12:13

Damn your right, man that was dumb.
The point is that a vacuum source is still a pain in the ass.

[Edited on 5-5-2009 by User]

hissingnoise - 5-5-2009 at 12:25

Let's see if I can get it right this time, User---your dilute HNO3 contains water and the azeotrope which boils at ~120.5, making water the lowest-boiling fraction. . .

User - 5-5-2009 at 13:20

I am not making myself real clear i guess, sorry 4 that.
The vapor coming from the boiling down mixture (which hasnt reached azeotrope) can be measured acidic with a piece of PH paper, so i figured HNO3 is lost.
To what extend is not really measurable by this method.

Has anyone destillated H2SO4/HNO3 under normal pressure with good results?

Would it be an idea to destillate HNO3 with sulfuric acid and afterwards boiling down the sufuric acid so that it is recovered.
And there would be no actual loss of chemicals.

hissingnoise - 5-5-2009 at 13:37

Your post was quite clear, User, it was I who got diverted. . .
Using a column should minimise HNO3 losses, but they should be small, anyway, and more or less unavoidable!
The H2SO4 route isn't actually necessary, after all.
When the azeotrope is reached, the temp. will go up.

Sedit - 8-5-2009 at 07:19

I have some old Manganese welding rods that I was going to toss out. Does anyone have any suggestions as to what use they could be put to and what the other contents of these maybe?

DJF90 - 8-5-2009 at 09:15

Thats something I've never heard of before! Are they for welding manganese? (Who builds anything out of manganese lol :S). Perhaps you could use them to make manganese salts? pretty boring to be honest but I cant think of much else manganese can be used for. As to other contents in the rods... There is likely to be some kind of flux to help keep the weld pool from being oxidised, and the manganese might be alloyed for better mechanical properties.

Sedit - 8-5-2009 at 11:00

I have no idea what there for they came with an old welder I got for junkin a while ago. I was clearing my garage the other day and found a ton of them. Its either toss them out for make something of it. I was considering MnOx but the fluxes are my main concern because if there is some contaminations I want to know what they are.

DJF90 - 8-5-2009 at 12:54

Could you not purify the metal like they do with copper using electrolysis? The impure manganese ionises to form Mn2+ (I expect) and then deposit on a small piece of pure manganese used as the other electrode (needs to be anode).

Intergalactic_Captain - 11-5-2009 at 09:31

Are there any shipping regulations (in the continental US) on ammonium iodide? I picked up a pound of Malinkrodt AR grade as the price was right, but it's way more than I need - I'd like to repack and resell it on ebay (or here, PM me if interested)... I'm guessing its either unregulated or ORM-D, but I can't find any solid info.

kclo4 - 11-5-2009 at 19:03

From what I know in chemistry, it seems like one could react Sodium Methyl Sulfate with Sodium phenolate to produce methoxybenzene, do you think this is true?
The producure on orgsyn to produce methylbenzene uses DMS and sodium phenolate, and claims sodium sulfate is produced in the end. I'm guessing sodium methyl sulfate is an intermediate in that reaction, and so I guess that gives me reason to believe, or at least gets my hopes up that it will work.

I have been meaning to try it, but I haven't got the time and figured I might as well ask.

EDIT: Well I'm a damn fool!
it says it works on the orgsyn document.
Sorry about that, but hey I guess I brought it to your attention that phenols can be methylated with the easily prepared sodium methyl sulfate!

[Edited on 12-5-2009 by kclo4]

Nicodem - 12-5-2009 at 00:18

Kclo4, please UTFSE. This has already been discussed and there are a few literature examples posted at this forum.

Panache - 12-5-2009 at 16:04

Two Questions

1) Cation exchange Resins
I have some acidic resin that i want to use to make some HI from KI. I have already performed a quick try and it worked fine, however a dilute KI solution was used. I don't have the solubility of KI at RT, but at 0C its 127.5g/100cc H20 and this increases to 210g at 100C.
Assuming linearity, this gives me a solubility around 150gKI/100c H20 at 25C. At this concentration one should be obtaining near azeotropic HI out of the bottom of the column (well ~54%HI).

That is the background.

My question is, would it be normal procedure to exchange the cation using such a saturated solution (and hence save time and material losses in distilling the HI to 57%) or am i going to run into problems doing this exchange and hence would be better off running at say 10% and doing a more tedious distillation? Any thoughts?

2) Citric acid solution (~10%) (further acidified with ~3%HCl) and 50% H202 reacts quite vigorously after a few minutes onset. Gas evolution occurs with heat evolution and the slight yellowness of the solution clears completely to water clear. What is going here? I assume the citric is simply forming the per-acid, which because of the condensed structure of the tri-carboxylic acid is inherently unstable and decomposes rapidly evolving oxygen. I tried to find some literature on forming per-citric acids but found none. (disclaimer-i din't look really hard). If my supposition is plausible am I reasonable in perhaps supposing then that the percitric acid could be a more active per-acid than the typical per-formic and per-acectic, (solutions of which can be stablised enough to make them shelf stable.) My question pertains mainly to per-acetic acids' excellent ability to disinfect. Solutions of said are commonly used in wash downs in the food industry.

[Edited on 13-5-2009 by Panache]

UnintentionalChaos - 12-5-2009 at 16:23

I'd suspect that the per-citric acid may form, but decomposes, ejecting O2 and CO2. I would suspect that the citric acid is completely consumed in the process with things like the very oxidation-prone acetone dicarboxylic acid being generated as intermediates.

Keep in mind that the HCl will generate HClO and Cl2 with the H2O2. Perhaps you should test it with dilute H2SO4 instead.

[Edited on 5-13-09 by UnintentionalChaos]

Panache - 12-5-2009 at 16:37

Quote: Originally posted by UnintentionalChaos

I'd suspect that the per-citric acid may form, but decomposes, ejecting O2 and CO2. I would suspect that the citric acid is completely consumed in the process with things like the very oxidation-prone acetone dicarboxylic acid being generated as intermediates.

Keep in mind that the HCl will generate HClO and Cl2 with the H2O2. Perhaps you should test it with dilute H2SO4 instead.

[Edited on 5-13-09 by UnintentionalChaos]

Apologies i should have mentioned the evolved gas was colourless and without odour.

UnintentionalChaos - 13-5-2009 at 09:41

Quote: Originally posted by Panache

Quote: Originally posted by UnintentionalChaos

Apologies i should have mentioned the evolved gas was colourless and without odour.

I had meant more along the lines that hypochlorite or chlorine would be a powerful intermediate oxidizer that might have screwed with your results. Perhaps straight citric acid and H2O2 acidified with H3PO4 or H2SO4 would behave differently.

Lambda-Eyde - 14-5-2009 at 15:28

First off, let me put something straight. I have absolutely no experience in practical "large-scale" chemistry, such as distillations and such.

So tonight I attempted my first distillation. I decided on distilling 50 ml of 30% technical grade hydrochloric acid.
As I suspected, HCl is a real bitch to distill.

But let's not blame the acid alone.
Knowing that aqueous HCl forms an azeotrope at 20,2% HCl I confered with ScienceGeek for some advice on the distilliation. He adviced me to start the distillation slowly, so that as little HCl gas as possible escaped from the solution, increasing the yield.
A happy first year student then set out on this adventure, plugged in the heating mantle, assembling the apparatus, greasing the joints, rigging cooling and covering the still head in aluminium foil.

It took me two hours to reach the boiling point of 30% HCl, 90*C. Don't laugh!

I see that I was waaaay too careful in the beginning, and I could have sped it up many times.
Yes, I know a distillation isn't done in 15 minutes, but 2 hours to reach the boiling point is just ridiculous.
During the distillation I observed lots of white smoke at the vacuum port which obviously was HCl gas escaping from solution, so the concentration in the distillation flask must have changed.
I have approx. 5 ml of nice and clear HCl solution in the receiver flask of unknown concentration, but about 20 minutes ago the distillation came to a halt. The temperature is now 98*C and I suspect that the concentration in the distillation flask has reached something around 20% (b.p. 108*C).
I don't dare to take this distillation any further as my thermometer only goes to 110*C.

But at least I got a few milliliters of nice reagent grade HCl!

So, do any kind souls feel like sharing any advice on such procedures? What did I do wrong, what could I do to speed it up without getting a horrible yield and so forth. Any advice is appreciated.

(Yes, I'm gonna buy another thermometer that goes beyond 110*C!)

[Edited on 14-5-2009 by Lambda-Eyde]

Jor - 15-5-2009 at 04:05

You should just heat the flask with a hotplate/mantle/burner.
When you heat concentrated hydrochloric acid, pure HCl will evolve until the concentration of the solution has reached about 20% (azeotrope). In the same matter, when you heat dilute 10% acid, you will drive off water until the concentration will be about 20%. Then the temperature of the thermometer will rise, and the azeotrope comes over.

I see no benefit in doing the distillation slowly. Just lead the evolved HCl in water for extra hydrochloric acid. You do this by attaching a tube at the vaccuum port and leading this into water.

woelen - 15-5-2009 at 06:46

Jor, what you are saying is dangerous. You should not simply attach the vacuum port to a tube and put this in water. As you already stated, with 30% HCl, you first will get fairly pure HCl gas when the solution is heated. When the pure gas is hitting the water, it will dissolve and then at once, the water flashes back into the apparatus. This is going fast, amazingly fast and the cold water will suck back into the distillation setup in a fraction of a second and the apparatus will crack

I once was surprised by suckback of water in an HCl-filled system and this is going fast, soooo fast! I was really shocked to see this happening.

What you can do is take an inverted funnel, which just is a few mm below the water surface, with the water in a beaker, which is just somewhat wider than the funnel. When the HCl is absorbed, then the water is sucked into the funnel, but this causes the water surface to lower and then the funnel looses contact with the water surface and no suckback occurs into the distillation setup.

Another option is to dilute the acid somewhat before distilling (take 2 parts of acid and 1 part of water) and then distill this acid. The thermometer must go well beyond 120 C if you want to distill the water/HCl azeotrope.

Lambda-Eyde - 15-5-2009 at 09:54

Thanks for your answers!

Quote: Originally posted by woelen

When the pure gas is hitting the water, it will dissolve and then at once, the water flashes back into the apparatus. This is going fast, amazingly fast and the cold water will suck back into the distillation setup in a fraction of a second and the apparatus will crack

I would of course have used a trap in the line. The water would anyways only be sucked into the cold end of the distillation setup, wouldn't it?

Yes, I thought about diluting the HCl to 20,2%, but in the end I decided against it for some reason. Azeotropes are cumbersome!

Wikipedia says (yes, I know. I'm saving up for a CRC) the boiling point for a 20% solution is 108*C, but at 10% it drops to 103*C. No data is provided of the boiling points between those two, but you say that it will go beyond 120*C ? Sigh... I hate limitations.

[Edited on 15-5-2009 by Lambda-Eyde]

[Edited on 16-5-2009 by Lambda-Eyde]

querjek - 15-5-2009 at 19:22

Can anybody give me a quick rundown as to what a spectrophotometer does? I bought one (cheaply) today, thinking, "oh, yeah, this'll be really useful", but now can't remember whether it measures concentration or (roughly) identifies samples.

Paddywhacker - 15-5-2009 at 22:36

Quote: Originally posted by querjek

Can anybody give me a quick rundown as to what a spectrophotometer does? I bought one (cheaply) today, thinking, "oh, yeah, this'll be really useful", but now can't remember whether it measures concentration or (roughly) identifies samples.

See http://en.wikipedia.org/wiki/Spectrophotometer

In order to make any use of it you will need the ability to weigh accurately, and to measure volumes of liquid accurately.

You get it at a good price? Going to resell? What make and model?

Panache - 16-5-2009 at 00:56

Does anyone have a neat trick for reducing iodate back to iodine. I know of the common methods (KI, H2S, Hydrazine, SO2, ascorbic) but these all fail the 'i don't have any at hand nor am inclined to make some' test. The iodate is a potassium/sodium with about 20% citric acid and a few of salts of PO4 mixed in for good measure.
Many many references have said it undergoes oxygen loss reverting to KI with heat but i have not found this happens to any significant extent. However i only raised it to around 600C for some minutes. I can heat it heavily up to 1200, this would also remove the citrate. Is this foolish?

Oh it also now has some chromium in it as it turned green when i was evaporating off the water (stainless dish).

Oh also does anyone know of illegal/diversionary uses for benzotriazole, a sales rep asked me because he had a dodgy looking customer asking for it and when he asked the customer how much he was after the customer said 'how much have you got! Ha!'

[Edited on 16-5-2009 by Panache]

Lambda-Eyde - 16-5-2009 at 01:37

querjek: http://www.chemguide.co.uk/analysismenu.html

UnintentionalChaos - 16-5-2009 at 10:52

Panache- if you have some sulfite/bisulfite/metabisulfite, you have a source of SO2 sitting around, waiting for a little acid. That'll get you back to iodide, but getting it out may be a bit tricky.

S.C. Wack - 16-5-2009 at 12:29

A way to avoid an iodine clock reaction is hot aq. oxalic acid. It works.

[Edited on 16-5-2009 by S.C. Wack]

woelen - 17-5-2009 at 12:11

Quote: Originally posted by Lambda-Eyde

Thanks for your answers!

I would of course have used a trap in the line. The water would anyways only be sucked into the cold end of the distillation setup, wouldn't it?

Yes, I thought about diluting the HCl to 20,2%, but in the end I decided against it for some reason. Azeotropes are cumbersome!

Wikipedia says (yes, I know. I'm saving up for a CRC) the boiling point for a 20% solution is 108*C, but at 10% it drops to 103*C. No data is provided of the boiling points between those two, but you say that it will go beyond 120*C ? Sigh... I hate limitations.

Sorry for my mistake, I accidently exchanged the BP of nitric acid azeotrope with that of HCl azeotrope.
But even then, if your thermometer is going up to 110 C then you indeed should not attempt this distillation, you will be on the edge with that and just a little overheating will cause destruction of your thermometer. What can be done though is just putting in a stopper in the hole, where you normally put the thermometer. Best of course is a glass stopper, ground joint, but if you don't have that, then you also can use a rubber stopper, which is covered in a few layers of white teflon tape (from hardware stores, used for fixing water taps and the like), before it is pushed in the hole. The teflon tape will give sufficient protection, certainly for the duration of the distillation. With HCl you can do that, it is not that corrosive. Even a cork stopper can be used, but again, this must be well covered with teflon tape, no cork may be exposed directly to the acid vapors.

If you have the 10% mix, then the boiling point may be 103 C, but don't expect to get a 10% mix boiling off. At 10% concentration, you get almost pure water, and I think that it certainly will contain less than 1% of HCl.

Lambda-Eyde - 17-5-2009 at 14:43

Thanks again for a helpful answer, woelen.

My distillation apparatus is all 24/40 ground glass, however, my thermometer adapter has a rubber fitting. I'm going to purchase a teflon adapter from UGT along with a vigreux column very soon (and of course a new thermometer).

I didn't get 103 C on the thermometer when I distilled, that number was also taken from the wikipedia article. I'm curious as to what the bp's between 10% (103 C) and 20% (108 C) are, more specifically if they exceed 110 C. But like you say, slight overheating could reduce my thermometer to pieces. By the way, the graduation of the thermometer ends at 110 C, but the inner tube doesn't come to an end until about 3 centimeters above the end of the graduation. I understand that I shouldn't go above 110 C just because of that, but it doesn't mean that I shouldn't go to 108 C?

Regarding the concentration of the HCl, I don't know what its concentration is right now, either in the still pot nor the receiver. And I have no way of determining it either, as I lack a burette and a suitable titrant (not to mention an indicator...).
The starting concentration was 30% HCl, and I'm pretty sure it has reached the azeotropic concentration. At least it has according to my logics.

And a slight digression (yes, even more questions!): I have a one liter solvent can containing 70-100% dichloromethane, 5-10% formic acid, and 1-5% "anionic surfactants". Obviously, I want the CH2Cl2 for lab purposes, and I'm thinking of distilling it. The boiling point of CH2Cl2 is 40 C, and for HCOOH it is 101 C. I suppose I won't have to do a fractional distillation, a simple distillation will suffice, right?

And those surfactants, how will I get rid of those? I don't know anything about their chemical composition or physical properties.

UnintentionalChaos - 17-5-2009 at 14:59

Lambda-Eyde- If you add a base, you will make formate salts and not need to worry about the volatility of formic acid. The anionic surfactacants probably have a relatively high molecular weight and will be virtually nonvolatile under the temperatures needed to distill DCM. You will just need to let the DCM sit over a dessicant after distillation since neutralizing the formic acid will generate a small amount of water.

[Edited on 5-17-09 by UnintentionalChaos]

Lambda-Eyde - 17-5-2009 at 15:38

And I suppose neither NaOH or CaCl2 will react with the DCM?

Thanks!

stateofhack - 18-5-2009 at 04:00

Hello,

Quick thing that i am not sure about, i need to buffer 1.5 mls of 70 % ethylamine sol with Glacial acetic acid, how much would it take?

edit: found it.

[Edited on 18-5-2009 by stateofhack]

manimal - 18-5-2009 at 18:30

Some solubility values for substances list two different values: one for 'anhydrous' substance, and one for a 'hydrated' substance.

My question is this: How do you have an aqueous solution of an anhydrous compound? For example, wikipedia lists the solubility for calcium sulfate (anhydrous) as .0021g/100ml and calcium sulfate (dihydrate) as .24g/100ml.

[Edited on 19-5-2009 by manimal]

Maja - 19-5-2009 at 02:14

Does sodium sulfite(Na2SO3) forms bisulfite adduct with aldehyde and ketone groups ?

woelen - 19-5-2009 at 03:31

Quote: Originally posted by manimal

Some solubility values for substances list two different values: one for 'anhydrous' substance, and one for a 'hydrated' substance.

My question is this: How do you have an aqueous solution of an anhydrous compound? For example, wikipedia lists the solubility for calcium sulfate (anhydrous) as .0021g/100ml and calcium sulfate (dihydrate) as .24g/100ml.

[Edited on 19-5-2009 by manimal]

There indeed can be quite a strong difference between solubilities of anhydrous compounds and hydrated compounds. I myself made anhydrous NiSO4 and that hardly dissolves in water, while NiSO4.6H2O dissolves very well. Another example is CrCl3 vs. CrCl3.6H2O.
This has to do with other structure of the compounds. Many of these anhydrous compounds are not really salts, but form covalent (sometimes polymeric) units. Some of these anhydrous salts very slowly are hydrated and if that is the case, then indeed you can distinguish two different solubility figures for the two. E.g. nickel sulfate anhydrous only dissolves with hundreds of mg in 100 ml of water, while the hydrated salt dissolves with tens of grams in 100 ml of water. If you wait for several days, then the anhydrous nickel sulfate also dissolves to a much larger extent and this is because it slowly is hydrated and then the solubility of the hydrated salt comes into play again.

This distinguishment only is important for anhydrous salts which slowly are hydrated in water. When the anhydrous salt immediately is hydrated (such as CuCl2 or CuSO4) then of course one cannot speak of different solubilities, then only the hydrated compound is present in aqueous solution. Calcium sulfate also slowly hydrates. Think of gypsum sculptures, which are mixed with water and which take a day or so to harden.

smuv - 19-5-2009 at 16:59

Quote:

Does sodium sulfite(Na2SO3) forms bisulfite adduct with aldehyde and ketone groups ?

No. The reaction would be reversible, forming 1eq hydroxide for every eq of adduct. Just buy sodium bisulfite....or if you are really desperate try halfway neutralizing a solution of it.

Formula409 - 23-5-2009 at 20:03

At school we were tasked to analyse an unknown white powder. Nitrate ions were confirmed with the "brown ring" test, and an emerald green colour was observed with a flame test. My first thought was that it was Zinc Nitrate, so I added a solution of sodium carbonate which garnered a white precipitate of Zinc Carbonate (so I thought). The teacher, however, suggested that I may be wrong and that I should look at other metal ions which give a green colour in the flame. I tested for barium with the addition of a sulfate which yielded nothing - I don't know what else to look for. This is only high school so it would not be anything TOO exotic.

Formula409.

DJF90 - 23-5-2009 at 20:35

According to this site:

http://chemistry.about.com/library/weekly/aa110401a.htm

Emerald green is copper (not a halide). Which makes perfect sense if you tested positive for nitrate. However one thing doesnt add up. You said it was a white powder... copper nitrate would be blue/green. Unless of course its copper (I) nitrate... that would be colourless right? But how stable would that be in aqueous solution?

Either that or its Thallium nitrate

EDIT: It doesnt look as if copper (I) nitrate exists... well at least not as a compound that is usable in the lab. But surely your teachers wouldnt give you thallium salts to vaporise in the flame... I would expect anhydrous Copper (II) nitrate to be "white", but surely upon adding water the Cu (II) ion would hydrate to the blue aqua ion [Cu(OH2)6]2+ ?

[Edited on 24-5-2009 by DJF90]

[Edited on 24-5-2009 by DJF90]

Formula409 - 23-5-2009 at 20:51

Quote: Originally posted by DJF90

According to this site:

http://chemistry.about.com/library/weekly/aa110401a.htm

Emerald green is copper (not a halide). Which makes perfect sense if you tested positive for nitrate. However one thing doesnt add up. You said it was a white powder... copper nitrate would be blue/green. Unless of course its copper (I) nitrate... that would be colourless right? But how stable would that be in aqueous solution?

Exactly what I said too! Plus the addition of sodium carbonate would yield a coloured precipitate!

Formula409.

[Edited on 24-5-2009 by Formula409]

UnintentionalChaos - 23-5-2009 at 21:54

Quote: Originally posted by DJF90

I would expect anhydrous Copper (II) nitrate to be "white", but surely upon adding water the Cu (II) ion would hydrate to the blue aqua ion [Cu(OH2)6]2+ ?

Anhydrous copper nitrate is also blue, so it definetly isn't that, plus a blue aqueous solution is pretty noticable.

Sedit - 28-5-2009 at 18:36

Two quick questions one more of an urgent nature..

First one. I have alot of good cloths of mine that when we put in our new well they forced us to go 125 ft down when our last one was only 50. Well the water sucks back and is loaded with iron. Almost all of my good cloths are ruined with iron hardwater stainds and a dingy orange yellow color. I have tried everything from bleach to hard water rust remover to peroxide soaks and nothing touches it. I am not considering a dilute muratic acid soak but thats a last resort.
Any suggestions as to what I could use to turn my whites white again?

Another question thats bugs me a bit.

Earlyer I made pasta. After it was done I started eating it with garlic bread and ate a decent amount as did my children until someone else came in and tasted the meat and said it was bad... Damned if it wasn't to. I couldn't tell the difference over the taste of garlic and my grandfather ate two plates... It was cooked very well but I was woundering if I should worry for myself/kids and more important my grandfather.

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