Sciencemadness Discussion Board

The Short Questions Thread (4)

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j_sum1 - 31-7-2017 at 21:16

I have been cleaning up some old stock solutions in my school lab and have come across some stannous chloride solutions that are quite discoloured. I am used to SnCl2 being an off-white but this is something different.

The first solution was actually one I prepared myself about a year ago by dissolving some tin pellets in HCl. It has been stored in a PE dropper bottle for over that time. It is now a bright yellow colour.

The second is in a glass reagent bottle dating back at least ten years. It seems someone inexpertly attempted to dissolve SnCl2 in neutral distilled water. There is a thick white sludge on the bottom of the bottle. Stirring it up reveals some pale yellow discoloration. There are also some greenish lumps in the very bottom of the sludge which I assume originate from the bottle not having been properly cleaned.

Anyway, I am curious as to what the yellow colour is. My intention is to purify and recycle if possible. there is quite a bit there.

Actually, advice on purification would be welcome too. I am predicting that it will be a bit of a dog to work with.

[Edited on 1-8-2017 by j_sum1]

Sulaiman - 31-7-2017 at 23:25

Could the yellow precipitate be insoluble lead stannate, Pb2SnO4, due to lead (plus O2 from air) contamination ?
https://en.wikipedia.org/wiki/Lead-tin-yellow

j_sum1 - 1-8-2017 at 02:51

Quote: Originally posted by Sulaiman  
Could the yellow precipitate be insoluble lead stannate, Pb2SnO4, due to lead (plus O2 from air) contamination ?
https://en.wikipedia.org/wiki/Lead-tin-yellow

Lead contamination in the reagents used? Maybe. I'll check the source of the tin metal tomorrow. I have since mixed the two samples together but I can test for lead -- I think. Ill have to work out a way of detecting trace lead and differentiating it from the tin.

It is likely that the sludge came from the one of the (very old) bottles of SnCl2 we have on the shelf. One is the dihydrate and the other anhydrous. One might have an assay on it but the other is pre-metric and poorly labelled.

It's nice to have a little puzzle.

I will see if I can test the solution for lead.

j_sum1 - 1-8-2017 at 20:03

Update.

All three sources of tin -- the granules I used, the SnCl2 and the SnCl2.2H2O te technical grade. Which means that lead contamination is reasonably likely.

The 1 year old solution prepared from metal dissolved in HCl was a yellow solution with no precipitate. It was quite strongly coloured. I can't figure out what this is yet. I am picking that Sn is in the +4 state as a stannate as a result of available oxygen but then it should precipitate if lead is present. It certainly had sufficient time to settle out and so I am not thinking of fine particulate material.


I combined the two mixtures, partially boiled it down to get rid of water and left it to settle. What remains is a cream-yellow fine powder and a slightly cloudy yellowish liquor. The liquid is highly acidic. It tests positive for chloride with AgNO3 (hardly surprising) and negative for sulfates with Ba(NO3)2.

Testing with potassium iodide causes a deeper yellowing but no precipitate. This could easily be lead iodide at low concentrations. I am not seeing any red/orange precipitate of tin(II) iodide.

Neutralising with NaOH gives interesting results. Cloudiness begins to appear at around pH 4. By the time it becomes alkaline the cloudiness disappears and is replaced by a grey-green flocculant material. This resembles the greenish lumps that appeared in the sludge.
On filtering it separates into a yellow precipitate (Pb2SnO4? PBO?) and a small amount of grey material that settles last. The filtrate does not appear to contain any lead when tested with a carbonate solution.

I would be tempted to think the grey material is metallic except that it was the last to settle in the filter paper. It looks pretty dark. It is also extremely fine. Tin and lead convert from 2+ to elemental at almost the same reduction potential and so if it is metallic it could be both. I am not picking what would have done the reduction though. Perhaps I am missing something obvious.

Returning to the original liquor, I have carefully adjusted the pH to 6 using NaOH. A white precipitate is slowly settling out. At least I think it is white. It is a bit hard to tell in the presence of universal indicator. I intend to test the liquid portion with carbonate. If any lead remains then it should be visible.


I think the conclusion is that Pb is likely in my original reagents and that I have a bit to learn about distinguishing between tin and lead.

Ammonium thiocyanate test colors

experimenter_ - 2-8-2017 at 05:06

It is known that a qualitative test for iron ions (Fe+3) is the blood red color they form with a thiocyanate ( link )

Do you know any other metal that forms the same blood red color when reacted with a thiocyanate? (e.g. heavy metals, precious metals?)


I see this blood red color when adding thiocyanate in an Au chloride solution (Au was of .999 purity). Does this mean that the Au is contaminated with iron?

(the acids used were tested and found to be free of iron)

Cryolite. - 2-8-2017 at 12:56

I recently purchased a Corning PC220 hot plate stirrer off eBay. The stirring functions normally, but when I turn the heating on it makes a loudish buzzing noise when starting up. This noise subsides as the unit warms up. Is this anything I should worry about? I did remove the ceramic top before powering it on to clean out some accumulated dust.

Question and Answer

agent_entropy - 3-8-2017 at 03:59

I recently received 5 kg of cobalt (II) sulfate heptahydrate as a gift. Any ideas for interesting things to do with it? I'm not too keen on messing around with heavy metal salts (cumulative toxicity makes me nervous), but this seems like too large a quantity to ignore if there is anything worth doing with it.

Edit: Cryolite, I wouldn't worry too much about the buzzing unless parts other the ones that are supposed to are heating up. The oscillating magnetic field induced by the alternating current can cause slight vibrations in the conductors/resistors that result in audible frequencies.

[Edited on 3-8-2017 by agent_entropy]

[Edited on 3-8-2017 by agent_entropy]

j_sum1 - 3-8-2017 at 04:49

Cobalt chemistry is fascinating -- lots of complexes and interesting colours and a couple of oxidation states. But 5kg is a lot to play with. More than you would need. If it was me I would keep 500g and barter the rest.

Does water aspirator flow rate affect max vacuum, or only the rate the vacuum is pulled?

FireLion3 - 6-8-2017 at 00:44

Googled, didn't see anyone else asking this. I understand that the maximum vacuum is related to the vapor pressure of the liquid at the given temperature, but does the flow rate of the liquid have an effect on the vacuum?

XeonTheMGPony - 6-8-2017 at 03:34

Yes it does, you need the pressure to push the flow to 5gpm for mine any ways it is the brass one sold from deschem.

Making esters from butyric acid

bobjgalindo - 9-8-2017 at 10:45

In the spirit of making odors by combining carboxylic acids (ie propanoic acid) and alcohols (ie ethanol) with a drop of sulfuric acid, is the following combination safe to mix:
Butyric acid (abt 10 drops)
Octanol (abot 10 drops)
Sulfuric acid

Texium (zts16) - 9-8-2017 at 10:51

Yes

Making benzyl alc from polystyrene

Ramium - 11-8-2017 at 01:29

I thought I might try producing benzyl alcohol from polystyrene just for fun.

To accomplish this, I was thinking:

Step 1. Depolymerize polystyrene to give styrene, this can be done by distilling the polystyrene. Use immediately in step 2 to prevent repolymerization

step 2. Reflux styrene with dilute sulfuric acid solution and sulphur (inhibits the polymerization of styrene). The styrene should undergo addition with water to produce the benzyl alcohol.

step 3 isolate the benzyl alcohol

Sound potentually feasible?


JJay - 11-8-2017 at 03:27

Hydration of styrene would typically result in 1-phenylethanol. You'd have to demethylate it somehow.

Ramium - 11-8-2017 at 13:09

Thank for the respones.

Oh yeah your right, forgot about markovnikov's rule. 1-phenylethanol would be the major product, and phenethyl alcohol the minor.

1-phenylethanol is a secondary alcohol and should be easily oxidized to give the corrosponding ketone, Acetophenone! which is very useful. This is even cooler then making benzyl alcohol:).

Polystyrene could be a very cheap source of acetophenone.

I don't know much about hydration reactions in practice. Is a reflux with water and a bit of sulphuric acid enough to hydrate an alkene?

myristicinaldehyde - 11-8-2017 at 15:34

So I made some isopropyl bromide (9.0g total). My next planned step is to try to see if I can alkylate salicylic acid with it.

I have 2 options for reaction conditions- anh. K2CO3 in butanone (never bothered getting acetone- but I might for this rxn) or I could try lithium isopropoxide in isopropanol, similar to sodium methoxide in methanol: the pKa's arent wildly different, and I have a lot of IPA. I could also try magnesium methoxide in methanol, lithium methoxide, or magnesium isopropoxide. I would rather not use methanol, as I am running low but I could always buy more.

Which is better? The pKa of salicylic acid's phenol is 13.6 so an alkoxide might be the better option. Or am I completely off-base here?

JJay - 12-8-2017 at 03:41

Quote: Originally posted by Ramium  
Thank for the respones.

Oh yeah your right, forgot about markovnikov's rule. 1-phenylethanol would be the major product, and phenethyl alcohol the minor.

1-phenylethanol is a secondary alcohol and should be easily oxidized to give the corrosponding ketone, Acetophenone! which is very useful. This is even cooler then making benzyl alcohol:).

Polystyrene could be a very cheap source of acetophenone.

I don't know much about hydration reactions in practice. Is a reflux with water and a bit of sulphuric acid enough to hydrate an alkene?


I've never hydrated an alkene, but I think typically they use high temperatures and pressures to make the reaction proceed at a reasonable rate. It may be possible to do it by simply putting dilute sulfuric acid and styrene in a test tube, putting it in an iron pipe, sealing both ends, and heating the pipe in a metal bath; I don't know. Magpie could do it: http://www.sciencemadness.org/talk/viewthread.php?tid=14652

Velzee - 14-8-2017 at 10:22

Anyone know what the energy of a Cl=O double bond is?

xfusion44 - 16-8-2017 at 10:35

Is nickel resistant to HCl and H2SO4? I'm trying to find out if what I've ordered really is nickel... There was no reaction between it and hydrochloric acid (31%) but it was slowly reacting when I added 30% H2O2 (water turned green and there was strong smell of chlorine). With H2SO4 there was something to see (nickel turned dark gray but H2SO4 color hasn't changed and reaction was very slow).

Thanks!

Zephyr - 16-8-2017 at 15:23

Quote: Originally posted by xfusion44  
Is nickel resistant to HCl and H2SO4? I'm trying to find out if what I've ordered really is nickel... There was no reaction between it and hydrochloric acid (31%) but it was slowly reacting when I added 30% H2O2 (water turned green and there was strong smell of chlorine). With H2SO4 there was something to see (nickel turned dark gray but H2SO4 color hasn't changed and reaction was very slow).

Thanks!


Yep, that's nickel. If you heat it to boiling in HCl it will react as well. Another way to identify it is using a magnet, it's one of the few ferromagnetic metals.. (Although I think it's weaker than iron?)

xfusion44 - 16-8-2017 at 16:38

Quote: Originally posted by Zephyr  
Quote: Originally posted by xfusion44  
Is nickel resistant to HCl and H2SO4? I'm trying to find out if what I've ordered really is nickel... There was no reaction between it and hydrochloric acid (31%) but it was slowly reacting when I added 30% H2O2 (water turned green and there was strong smell of chlorine). With H2SO4 there was something to see (nickel turned dark gray but H2SO4 color hasn't changed and reaction was very slow).

Thanks!


Yep, that's nickel. If you heat it to boiling in HCl it will react as well. Another way to identify it is using a magnet, it's one of the few ferromagnetic metals.. (Although I think it's weaker than iron?)


Thank you very much! Yes, it is attracted by a magnet. One more question: why the chlorine was produced when it reacted with HCl and H2O2? And at the same time the green color appeared - nickel chloride? But why chlorine and chloride at the same time? Or was it some other compound?

Geocachmaster - 16-8-2017 at 17:41

Acidic solutions of H2O2 are strong enough to oxidize chloride into chlorine. The green color can be attributed mostly to dissolved chlorine, although I'll bet there's some nickel chloride in there as well. The HCl and H2O2 react to make chlorine and HCl reacts with the Ni to make NiCl2 (Edit: although the latter reaction should be pretty slow at room temp)

[Edited on 8/17/2017 by Geocachmaster]

physics inclination - 18-8-2017 at 13:04

Would urea and a strong base react to make ammonia and/or methanol from the urea molecule? I wonder because bases are proton donors, and it seems you could split the urea molecule into 2 ammonia molecules and 1 methanol molecule if you "add" enough hydrogens.

Crowfjord - 18-8-2017 at 13:27

Quote: Originally posted by physics inclination  
Would urea and a strong base react to make ammonia and/or methanol from the urea molecule? I wonder because bases are proton donors, and it seems you could split the urea molecule into 2 ammonia molecules and 1 methanol molecule if you "add" enough hydrogens.


Bases are proton acceptors. Acids are proton donors. In Bronsted-Lowrey theory at least. Urea hydrolysis would give ammonia and carbonate in base, and ammonium salt and carbon dioxide in acid.

For methanol to form, electrons as well as protons would need to be added to that carbon. That is, a reduction would need to take place.

[Edited on 18-8-2017 by Crowfjord]

Question

TheNerdyFarmer - 3-9-2017 at 12:32

Hello.
Does anyone know of a good way to test for hydrofluoric acid?
The only method that I have seen is by treating calcium carbonate with it and it produces an insoluble precipitate. But I'm not sure how well this works. Would be a tremendous help if someone could mention a great way to test for the presence of it so I could further preform a titration to determine the concentration. Thanks!!

ninhydric1 - 3-9-2017 at 17:45

Magnesium also works. I'm not sure if this works immediately, but HF reacts with glass. Maybe that could be used as a test?

[Edited on 9-4-2017 by ninhydric1]

j_sum1 - 3-9-2017 at 18:01

Calcium fluorite is amongst the most insoluble things there is. Other calcium halides are soluble and so a calcium salt makes for a sensible way of differentiating fluoride from other halides.

As for any test of this type, you presumably have a limited number of possibilities of what your compound contains and your test differentiates between those possibilities. So, context helps in answering this question. What other acid might it be if it is not hydrofluoric?

Testing a drop on a glass slide for 24 hours would seem to be a good idea. Even low concentrations will etch the glass in time. There really is not much else that will attack glass and so this would be pretty definitive.

TheNerdyFarmer - 4-9-2017 at 09:14

Here was my procedure. I acquired some 1,1 difluoroethane (DFE) as electronics duster. I did some research on hydrofluorocarbons and found that when these burn they often produce hydrogen fluoride and carbonyl fluoride. When carbonyl fluoride comes in contact with water it hydrolyses to HF and CO2.
My main dilemma was the carbonyl fluoride. I was pretty scared to mess around with carbonyl fluoride considering that it is toxic at 2 ppm. But I decided to do a small scale reaction anyway. I filled a test tube about halfway full of 1,1 DFE. I then put about 1ml of water in the bottom of the test tube. I bent a wooden splint so it would fit down in the test tube with my hand out of the way. I lit it on fire and stuck it down in the tube. The DFE took longer to burn than expected. About 5-7 seconds. Once it was done burning I quickly turned on the fume hood and stoppered it with a butyl rubber stopper. I vigorously shook the tube. Then I opened the tube and squirted a little bit on the walls of the test tube to get the excess HF of the walls. I was left with an acidic clear liquid.
The only things that I could imagine being in the liquid would be Hydrofluoric acid and a small amount of Carbonic acid. Obviously I would expect any carbonic acid to decompose shortly. Thus I would be left with reasonably pure Hydroflouric acid.
But I am not 100% sure if it is or not. That is why I asked. Like I said, I am working small scale. This means that I only have <1ml of solution at dilute concentrations so I don't have a lot to work with. I may try the magnesium method and as a secondary test the glass slide.
It would be much appreciated if I could get some input on the reaction going on here and tell me if I'm doing something wrong.
I cannot find any information online referring to the mechanism or the chemical equation to how this works so I will have to check out the library at my college to see if I can find any information.

j_sum1 - 4-9-2017 at 14:42

Given that, I would smear a tiny portion on a microscope slide and leave it for a number of hours. i would attempt to protect it from evaporation -- perhaps a cover slip and an evacuated zip-seal plastic bag. Then examine the surface of the slide under a microscope and see if any etching had taken place.

Question: Solubility of multiple ions

Plunkett - 8-9-2017 at 11:34

I am working on making a safer version of a lead acid battery that uses sulfate salts instead of sulfuric acid for the electrolyte. The two salts I am considering are ammonium sulfate and magnesium sulfate, but neither are soluble enough on their own to put enough sulfate ions into solution for what I want to do. What limits how much of these salts can dissolve? Is it the magnesium/ammonium or the sulfate? What I am really asking is is the solubility of of ammonium sulfate and magnesium sulfate in the same solution the same as if they were in separate solutions?

I am away from my lab so I cannot test this experimentally.

[Edited on 8-9-2017 by Plunkett]

DraconicAcid - 8-9-2017 at 13:27

Quote: Originally posted by Plunkett  
What I am really asking is is the solubility of of ammonium sulfate and magnesium sulfate in the same solution the same as if they were in separate solutions?

I am away from my lab so I cannot test this experimentally.


No. There's two factors at play here- first, an ionic compound will be more soluble in a solution with high ionic strength than in pure water. Secondly (and more importantly), two compounds with common ions will be less soluble in the same solution than in pure water, or a solution without common ions.

A saturated solution of MX (generic ionic compound) will have the equilibrium MX(s) <=> M+(aq) + X-(aq). If you already have some X- in the solution from the presence of another compound (say, NaX or KX), the equilibrium will shift to the left (less MX dissolves).

gluon47 - 9-9-2017 at 21:42

Just out of curiosity, would copper(ii) form complexes with amphetamine or methamphetamine? It would be interesting if such complexes existed

Totally theoretical, I have no interest in drugs.

Oxalate or Carbonate

veerenyadav - 11-9-2017 at 03:05

I have a solution with in which iron is leached using oxalic acid. What will precipitate if I add NaHCO3 in the solution ?

it will be Sodium Oxalate or Iron Carbonate or both ?

ninhydric1 - 11-9-2017 at 20:01

Depends on how saturated your solution is:
Sodium oxalate has a solubility of 3.7 g per 100 mL of water at 20 degrees Celsius, so it depends on how much iron oxalate is in solution. If you are able to keep all the sodium oxalate in solution, iron carbonate will be almost all of the precipitate that forms from the addition.

Red P

fabtasticwill - 17-9-2017 at 16:05

Is it possible to get pure red phosphorus from red phosphorus fire retardant? Its added to some kind of resin. I wasnt sure if there would be a way to dissolve the resin leaving behind the red phosphorus?

Zephyr - 17-9-2017 at 22:55

Quote: Originally posted by fabtasticwill  
Is it possible to get pure red phosphorus from red phosphorus fire retardant? Its added to some kind of resin. I wasnt sure if there would be a way to dissolve the resin leaving behind the red phosphorus?


I'm not sure about how to actually isolate the phosphorus, you'll need to check what the other components are to get an idea of how to separate them, try checking an msds?

But I think it's really interesting that they ate using phosphorus to put out fires considering it's history... it seems the mechanism is the creation of phosphoric acid which is able to subdue plastic fires.

fabtasticwill - 18-9-2017 at 12:31

It says that there is a layer of polymer protective on the surface, though others seem to be in some sort of resin. Would it be simple to dissolve the polymer without damaging the red phosphorus?

symboom - 18-9-2017 at 15:15


Fire retardant doesn't attempt to put out wildfires or even necessarily halt flames in their advance. Consisting primarily of ammonium phosphate — fertilizer, basically — fire retardant is formulated to slow down the combustion of trees, brush and grass. The idea is to give firefighters time to mount a ground attack. Excerp wikipedia
There is no red phosphorous it is a dye
Aluminum and ammonium phosphate in a exothermic reaction
Will produce phosphorous but it will be white phosphorous.
Just use fertilizer it is a more pure source

fabtasticwill - 19-9-2017 at 12:10

My goal is actually to get white phosphorus. So if I burn it in an inert atmosphere so the phosphorus doesnt burn it will produce white phosphorus?

Panache - 21-9-2017 at 00:27

Quote: Originally posted by gluon47  
Just out of curiosity, would copper(ii) form complexes with amphetamine or methamphetamine? It would be interesting if such complexes existed

Totally theoretical, I have no interest in drugs.


Should do, at a guess, I imagine however that the freebase coordinated to the copper would still be susceptible to reaction with atmospheric CO2 as all such amines are, I would doubt you could isolate the complexes as solids though but really I actually wouldn't really know some stranger things form stable complexes, I think silver forms a volatile complex with pyridine if I recall.

edit - the qualifier 'I have no interest in drugs'...bit tedious, bit irrelevant, no one cares anyway and if they do, they are unlikely to believe your caveat anyway.

[Edited on 21-9-2017 by Panache]

JJay - 21-9-2017 at 06:31

What does fluorine smell like?

ninhydric1 - 21-9-2017 at 07:29

Probably similar to chlorine as in it attacks the nose.

xfusion44 - 1-10-2017 at 05:26

Would bubbling chlorine through a solution of KNO3 yield NO3 gas which would react with water to make HNO3?

Metacelsus - 1-10-2017 at 05:34

Quote: Originally posted by xfusion44  
Would bubbling chlorine through a solution of KNO3 yield NO3 gas which would react with water to make HNO3?


No.

Σldritch - 1-10-2017 at 05:36

Quote: Originally posted by xfusion44  
Would bubbling chlorine through a solution of KNO3 yield NO3 gas which would react with water to make HNO3?


No.

First NO3 is a radical and extremely unstable. But i guess you mean N2O5, which is also unstable but much more stable than NO3. It will react with water however so it will not be formed.

You would get a solution of K+, H3O+, NO3-, Cl- and ClO-.

xfusion44 - 1-10-2017 at 06:06

Oh well, just looking for a cheap and easy way to make HNO3

Thank you both for your answers.

I once tried NurdRage's method of making it and it works but the acid obtained this way is of rather low concentration and a lot of NO2 is wasted.

Pulverulescent - 1-10-2017 at 06:19

Quote:
Oh well, just looking for a cheap and easy way to make HNO3

Don't give up on chlorine, just yet ─ the dry gas will displace NO3 from AgNO3 as the anhydride, N2O5?

ninhydric1 - 6-10-2017 at 19:44

Which sulfamate salts are insoluble in water?

Stopcock size

Geocachmaster - 14-10-2017 at 19:11

Are the sizes of glass stockcocks standardized? I need one that has a 2mm bore, ~30mm long, ~13mm widest, ~10mm narrowest. This eBay listing has dimensions that are close, but not exactly the same (.4mm off). Is this just error of measurement or the wrong size?

Some pictures of what needs the stopcock:

IMG_0070.JPG - 1.2MB IMG_0071.JPG - 1.2MB

As always, thanks for any help :)

SWIM - 17-10-2017 at 09:03

Does PTFE adsorb much water?

I want to render an auger-type powder dispensing funnel rigorously dry, but the PTFE auger has a phenolic looking handle that isn't obviously removable so I would hesitate to heat it at all.

Is the PTFE already dry if it looks dry, or is there some procedure for drying it that doesn't involve the ovens or torches used to dry glass?

This is a pretty expensive piece of glassware for me, and I don't want to ruin it by doing something stupid, but I would really like to be sure of its dryness.

Recrystalizing of CuSO4 - odd results when adding Na2CO3 - brown liquid

RogueRose - 17-10-2017 at 11:41

I wanted to recrystalize some CuSO4 I had. After about 60% or so of the crystals had formed (by weight of original crystals), looked at the remaining liquid and it had a slight green tint as well as a black tinge to it. The process took place at room temp over 5-6 months.

I thought I would try to add some Na2CO3 to it to make copper carbonate but the result was very different than doing the same with a more pure CuSO4 sample I added carbonate to.

The pure sample ended up with a nice green precipitate with a light blue liquid above the settled precipitate.

The remaining liquid from the recrystalization resulted in a brown precipitate that turned almost gelatenous. The same Na2CO3 was used for both samples.


Now it is very possible that the recrystalization had some FeSO4 (1-3% possibly) mixed with the CuSO4 (which is why I wanted to purify). Would the CuSO4 crystalize leaving the FeSO4 in solution if there was this much of a difference in concentration? It seems that the brown precipitate may be iron carbonate though it seems that there has to be a lot of iron carbonate in the solution for it to turn so brown.

From pure CuSO4
CuCO3.jpg - 89kB CuCO3_2.jpg - 204kB


Remaining solution from recrystalization
recryst.jpg - 109kB


Above solution with Na2CO3 added
brown_1.jpg - 237kB brown_2.jpg - 74kB brown_3.jpg - 101kB

DraconicAcid - 17-10-2017 at 11:47

I don't think you'll form iron carbonate- as soon as the solution becomes slightly basic, iron(II) ions will react with oxygen to give iron(III) hydroxides.

Root Beer Flavoring

JJay - 18-10-2017 at 23:55

When I was very young I remember making root beer with sassafras oil. Obviously, no one is doing that any more. Does anyone know offhand what manufacturers use as root beer flavoring these days?

wg48 - 19-10-2017 at 01:06

Aldi (a European discount supermarket) has sold what it said on bottle was an American root beer. The beer was flavoured with what to me is germolene (an antiseptic ointment from my childhood) which contained methyl salicylate. Its weird that a beer would smell like that.

Crowfjord - 19-10-2017 at 07:08

The main root beer flavoring these days is methyl salicylate/oil of wintergreen. The original root beer recipe used equal amounts of sassafras and wintergreen, along with other roots and herbs like sarsaparilla, licorice, dandelion root, cherry bark, birch bark, and some others. Hansen's uses sassafras in their root beer, but it's hard to taste it over the wintergreen and vanilla.

S.C. Wack - 19-10-2017 at 11:23

Shirley there is more than (if any) methyl salicylate in A&W. Order sassafras oil and you'll get methyl salicylate, perhaps with eugenol. Arctander mentioned other esters and stuff like p-dimethoxybenzene WTF.

Who remembers typical root beer with the safrole, over 50 years ago? Because the frosty mugs of A&W of 40 years ago (and A&W's food) could hardly have been better. AFAIK. Gee if only I had a drop of safrole I could just add it to root beer...

[Edited on 19-10-2017 by S.C. Wack]

The Volatile Chemist - 19-10-2017 at 19:38

Hmm. I've seen DIY rootbeer kits, though it seems the flavoring they sell is likely mostly vanilla, etc. Is it just me that never knew Root beer had oil of wintergreen/a minty taste? I've never heard/noticed this till you-all just brought it up...

You could always grow a sassafras tree. Also, Sioux City Sarsaparilla is a manufactured root beer that I used to have at the local bean-soup historical festival (bean soup cooked in a giant iron cauldron over an open fire with lots of salt and meat is so good...), apparently they thought it was close enough to the 'original' that it was worth selling at a historical festival. You could try that. I think Cracker Barrel sells then in their stores, too, if you have one nearby.

[Edited on 10-20-2017 by The Volatile Chemist]

JJay - 19-10-2017 at 19:48

Birch bark contains methyl salicylate, so it seems reasonable to suppose that birch beer would contain it. But root beer has a distinctive smell and taste that does not remind me of wintergreen at all... *shrug*

Sassafras bark and leaves have a fragrant odor but I don't think they contain any safrole. I think the powdered leaves are used in some Creole dishes. (Oh and I'm well aware of how I could obtain some safrole if I wanted some but to me it just isn't worth it.)

Melgar - 19-10-2017 at 19:51

Lots of people seemed to be interested in sassafras root bark on eBay, who were also interested in my listing for palladium chloride... I don't imagine that stuff is artificial.

I think that on perhaps two occasions I may have inadvertently synthesized acetone peroxide. Once was when I noticed a precipitate when mixing used piranha solution into a temporary waste container that almost certainly contained acetone. This is something I've been very careful to avoid since then. The other time, I was attempting to dehydrate ethanol to ether, and added H2O2 to try to neutralize one of the more stubborn denaturing agents that was giving me problems. I saw a white precipitate there too, but only briefly, and it dissolved. In both cases, I immediately added copper sulfate, which decomposed the H2O2, and presumably would have decomposed any organic peroxides as well.

Neither of these was recent, but I just wanted to make sure for future reference, that transition metal salts would most likely be able to safely neutralize suspected organic peroxides, correct? After all, I'm not actually sure there was anything there to neutralize in either case, but I do want to make sure that if I DID have something dangerous, what I did earlier would actually work

[Edited on 10/20/17 by Melgar]

ninhydric1 - 20-10-2017 at 21:01

Anyone have a good source of ninhydrin (I gave up trying to obtain the materials to synthesize it)? 1-5 grams will suffice, as long as the price isn't above $20 USD.

AngelEyes - 23-10-2017 at 06:09

Lithium Chloride

I have a solution of LiCl and I want to get the salt out without having to boil down a lot of solution. I tried adding methanol to the solution (roughly the same volume as there was in the beaker already...approx 100ml) in the hopes it would crash some of the chloride out, but nothing.

Would another alcohol work instead?
Are there any common salts I could use to 'flush out' the LiCl?

This LiCl would be used for a coloured flame on Nov 5th so I would like to avoid anything that's going to interfere with that; ie sodium chloride.

Cheers


Angel.

The Volatile Chemist - 23-10-2017 at 11:27

Quote: Originally posted by AngelEyes  
Lithium Chloride

I have a solution of LiCl and I want to get the salt out without having to boil down a lot of solution. I tried adding methanol to the solution (roughly the same volume as there was in the beaker already...approx 100ml) in the hopes it would crash some of the chloride out, but nothing.

Would another alcohol work instead?
Are there any common salts I could use to 'flush out' the LiCl?

This LiCl would be used for a coloured flame on Nov 5th so I would like to avoid anything that's going to interfere with that; ie sodium chloride.

Cheers


Angel.


Haha hmm... I'm looking for a method to precipitate out my lithium acetate myself, which has defied all of my own attempts. I can tell you that iPrOH does not work. I will have to try EtOH or acetone sometime soon.

JJay - 23-10-2017 at 20:44

The attached document discusses procedures established by IUPAC in the 1960s for preparing sodium carbonate and sulfamic acid as primary standards. Sodium carbonate is not exactly ideal because it emits carbon dioxide when neutralized, but sodium carbonate is widely accessible and easy to purify. Sulfamic acid looks harder to purify but it is also accessible, and the purification specified here doesn't look too hard, but you rarely hear of anyone using sulfamic acid in titrations. I guess.... Is there anything I'm likely to have sitting around the lab that is a better acidimetric primary standard than sulfamic acid?

Attachment: 1803x0443.pdf (240kB)
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j_sum1 - 23-10-2017 at 21:06

Thanks for that JJay. It is a good reference to have.
Short answer, I think, is "no". I doubt you will have access to better primary standards than sodium carbonate or sulfamic acid. And, I think your assessment is correct -- that sulfamic acid is likely to be the better of the two.

AngelEyes - 24-10-2017 at 00:29

Quote: Originally posted by The Volatile Chemist  
Quote: Originally posted by AngelEyes  
Lithium Chloride

I have a solution of LiCl and I want to get the salt out without having to boil down a lot of solution. I tried adding methanol to the solution (roughly the same volume as there was in the beaker already...approx 100ml) in the hopes it would crash some of the chloride out, but nothing.

Would another alcohol work instead?
Are there any common salts I could use to 'flush out' the LiCl?

This LiCl would be used for a coloured flame on Nov 5th so I would like to avoid anything that's going to interfere with that; ie sodium chloride.

Cheers


Angel.


Haha hmm... I'm looking for a method to precipitate out my lithium acetate myself, which has defied all of my own attempts. I can tell you that iPrOH does not work. I will have to try EtOH or acetone sometime soon.


I should have checked LiCl solubility...it's surprisingly soluble in water (69g/100ml @ 0c !!), but also fairly soluble in methanol. So my adding that solvent was never likely to help.

I might give acetone a try; otherwise a long and slow evaporation will be the only option...

j_sum1 - 24-10-2017 at 01:08

If you wish to avoid much of the evaporation process, you could always precipitate out Li2CO3. Then react that carbonate with some concentrated HCl at a much smaller volume.

Sigmatropic - 24-10-2017 at 12:16

solutions of LiCl in water are hygroscopic, which makes alot of sense since the solid is deliquescent. Hence my futile attempt at preparing it failed. None the less, I'm interested in a way of preparing it from lithium carbonate/hydroxide.


[Edited on 24-10-2017 by Sigmatropic]

DraconicAcid - 24-10-2017 at 14:50

Quote: Originally posted by AngelEyes  

I might give acetone a try; otherwise a long and slow evaporation will be the only option...


I suspect that you're more likely to salt out the acetone than to precipitate the lithium chloride.

AngelEyes - 25-10-2017 at 07:52

Quote: Originally posted by j_sum1  
If you wish to avoid much of the evaporation process, you could always precipitate out Li2CO3. Then react that carbonate with some concentrated HCl at a much smaller volume.


I bought the Li as the carbonate because it was cheap. And I am reacting it with 32% HCl, which I think is about as strong as I can get OTC.
Problem is, that reaction will produce water (and CO2) as by products, so it instantly dilutes the freshly-formed LiCl.

I think I have no option but to evaporate on the kitchen hob (I have no hotplate) and think of something to explain it to the wife.

Cheers though


Angel.


VSEPR_VOID - 30-10-2017 at 09:24

Quote: Originally posted by Vargouille  
Quote: Originally posted by papaya  
Why it is prohibited to extinguish a man on fire with the help of fire extinguisher ?


It depends on the type of extinguisher that you use. CO2 extinguishers are highly compressed, so when they come out they're very cold, which can cause frostbite. Also, if the person is on the ground, the CO2 can suffocate them. Some ABC dry chemical extinguishers are safe because their contents are of negligible toxicity, but you should check their MSDS to make sure. Of course, the person will want to cover their face with their hands if you use a dry chemical extinguisher on them, but other than that it's fairly safe.

http://www.homeofpoi.com/lessons_all/teach/Fire-Extinguisher...


I would have their head on a silver tray if someone sprayed me with a chemical fire extinguisher well I was on fire. The chemical inside of the spray MUST be removed from burns by scrubbing. This process of scrubbing raw burns is unimaginably painful and can last a large amount of time. Please use the CO2 as a little frostbite is preferable to hours of painful scrubbing at the ER.

[Edited on 30-10-2017 by VSEPR_VOID]

Solvent for metal carbides

veerenyadav - 1-11-2017 at 19:07

Can anybody suggest a suitable solvent for metal carbides ( M3C, MC, etc) ?

Material with minimum difference in melting and boiling temperature

veerenyadav - 1-11-2017 at 19:09

Which material has minimum difference in melting and boing temperature ?

SWIM - 1-11-2017 at 19:12

The one that's at its triple point.:D

RogueRose - 1-11-2017 at 20:06

I tried setting up a chlorate cell and the problem I was worried about reared it's head. I've tried to choose anode and cathode which are accessible to me so I chose stainless (cathode) and PbO2 (annode).


The PbO2 plates are from old SLA batteries which I oxidized in strong bleach. Since the plates are smallish, I stacked 3 vertically and used a copper wire to connect them at their tabs. The wire is varnished transformer wire (new) and I removed the insulation at the point where it contacts the plates, so there should be very little copper exposed in the solution.

The SS is non-magnetic and it is pretty shiny. I thought it was 304 or 316 but I think it is not. I do have another piece which is also non-magnetic which I am trying now and I think it is more likely the 304 or 316.

I'm using 5v from a computer PSU but would like to try 12v but don't know if that will effect anything in the setup adversely.

I'm worried about the green/blue color in the solution and am wondering if this may be CuCl2 or a chrome or nickel salt that has come out of the cathode - or is it the copper dissolving from the anode?

I'm planning on doing many re-crystalizations of the end product but IDK if any of these contaminates will dissolve in the chlorate solution. Should I be worried about what is in the solution at this point? It ran for about 3 hours at maybe 15a (probably less).


cath1.jpg - 337kB cell1.jpg - 231kB cath2.jpg - 409kB

ninhydric1 - 6-11-2017 at 21:21

Would this be a viable alternative to silicone vacuum grease?:

https://www.lowes.com/pd/Danco-Grease/1092167

The manufacturer claims it to 95% silicone, resistant to temperatures from -40 to 204 degrees Celsius and an ambiguous "most harsh chemicals". Would this be viable during distillations of relatively benign chemicals (alcohols, ketones, DCM, and other common solvents)? Would this also be viable for the distillation of nitric acid among others?

Melgar - 9-11-2017 at 12:20

@RogueRose I'm convinced that you live in a junkyard. IIRC, lead-acid battery chargers are often able to output at both 6V and 12V, and are probably the best power source for this type of thing. Every one that I've seen has an ammeter on it too. If you know any diesel mechanics, they'd be guaranteed to own a decent one, and would probably let you borrow it if you did them a favor. Same goes for anyone that owns a bulldozer or front-end loader or a lot of that type of heavy machinery. Your stainless steel cathode is probably not a good idea. Stainless steel is protected with a chromium oxide layer, and the cathode reduces things. All that in a chloride-rich bath just seems like a bad idea. They don't make "cathodized aluminum" for a reason, after all. Try copper or graphite, maybe? That way you could tell by the change in color if any lead was dissolving and plating.

@ninhydrin1 Yes. You could even use Vasaline if temperatures didn't get too high. Silicone is used mainly because its viscosity doesn't change much with temperature.

[Edited on 11/9/17 by Melgar]

j_sum1 - 10-11-2017 at 18:25

Negative X

I watched a King of random offering recently https://www.youtube.com/watch?v=I7vXvLvowgI. (It appears that Grant Thompson now has a side-kick.)
This video has a nice thermal mix of zinc metal, ammonium nitrate and ammonium chloride that is quite unstable. And of course it has a suitably hyperbolic but nondescript name. The guy appears to understand little chemistry but he makes a reasonable stab at presenting some chemical equations so that is at least something. (I haven't bothered working out the stoichiometry of his mix.)

The question I have, is why the green? The flame colour seen is not what I would normally associate with zinc. Anyone have any insight?

ninhydric1 - 10-11-2017 at 19:56

According to Thompson's Illustrated Guide, the flame test for ammonium ions is a faint green but is often masked by other species. That might explain it. The ammonium nitrate decomposes while the ammonium chloride present burns in the flames produced to form a green colored flame.

Melgar - 11-11-2017 at 01:34

Chlorine can make a flame green. And nitrate could oxidize chloride to chlorine.

Bert - 11-11-2017 at 08:11

I can personally report that one of the few pyrotechnic star compositions using Zn as a fuel (Weingart's "granite stars", AKA Zinc spreader stars) burns with a green flame.

The mixture is Zn fueled and KNO3 oxidized, so I think we can leave ammonium spectra out of the consideration?

(Edit)
Of course, S will react with Zn as well... Anyone recall the good old Zn/S rocket fuel flame color? Have not seen that burn since the 1970s.


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[Edited on 11-11-2017 by Bert]

j_sum1 - 11-11-2017 at 14:33

That's cool. I have only associated blue/white with zinc flames. Which is to say therd is a whole lot more going on in the spectra than the elementary textbooks suggest.

Diatomaceous Earth washed with HCl - Using the liquid for FeCl2?

RogueRose - 12-11-2017 at 09:12

I used some HCl (hardware store Muriatic - which is very clear and clean it seems) to clean some Diatomaceous Earth and the result was very yellow liquid and I think it still has a good bit of HCl and is not completely FeCl2. So I was wondering if this can be used to make some FeCl2 - or is there a good chance that there is other metals dissolved in the liquid. The DE was a high quality garden variety that is much whiter than many DE's I've seen so I think it was fresh water DE as opposed to salt water DE which tends to be darker and have more metal contaminates.

I'm not sure what else to do with the liquid and if there is something else that it could be useful for, I'd appreciate any input on that. Thanks!

TheNerdyFarmer - 15-11-2017 at 11:56

Hello everyone. I am considering purchasing a thermocouple. I found some relatively affordable ones on eBay. I found the thermocouple reader w/ 2 extra probes, an additional probe that claims to have a temperature range of -100 to 700 degrees Celsius, and a thermometer well for just over 25 bucks. What I am asking is if it is worth it. I am a broke teenager so I don't want to spend anymore money than I have to :P. For those of you that use these, do you prefer a standard glass thermometer or do you like your thermocouple better?

Sulaiman - 16-11-2017 at 09:22

Thermocouples with digital readout are easier to read than conventional thermometers but not as accurate unless calibrated



Can anyone tell me,

What is a quick reliable easy method to differentiate between hdpe and PTFE ?
(I have no fluorine compounds, but many common solvents)

[Edited on 16-11-2017 by Sulaiman]

OldNubbins - 16-11-2017 at 10:11

Based on their densities, I imagine PTFE would sink in water while HDPE floats.

Bert - 16-11-2017 at 10:12

Shave a small sliver of the plastic off, hold it in a lighter flame. HDPE will easily melt and then burns like parafin with a smell like a wax candle.

PTFE does not burn in air, it will just degrade and smells quite different. Don't sniff too deeply or often of overheated Teflon.



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[Edited on 16-11-2017 by Bert]

Bert - 16-11-2017 at 10:46

Sounds like UHMW plastic sheet, lots of uses. Essentially solvent proof, non stick- almost no adhesive will bond, but can be thermally welded. It's self lubricating, tough, can usually be tapped for threaded holes. Makes a good sliding base for saws or routers to be run on surfaces that metal could mar, my usual use.

Sulaiman - 16-11-2017 at 10:49

Thanks for the quick responses,
Today a workmate gave me a 1" 24" x24" slab of white plastic that he thought may be ptfe,
I just tested it and unfortunately it burns nicely with a blue flame, dripping screaming blobs of molten plastic :(

Almost certainly hdpe/uhmwpe, which is still useful.

EDIT: Sorry Bert, this post was before yours, but I tried to edit it and messed up so posted it again, by which time you posted the above ... .
I use hdpe quite a lot for stupid-high-voltage stuff, almost perfect - until it melts.



[Edited on 16-11-2017 by Sulaiman]

elementcollector1 - 16-11-2017 at 20:43

Quote: Originally posted by Sulaiman  
Thanks for the quick responses,
Today a workmate gave me a 1" 24" x24" slab of white plastic that he thought may be ptfe,
I just tested it and unfortunately it burns nicely with a blue flame, dripping screaming blobs of molten plastic :(

Almost certainly hdpe/uhmwpe, which is still useful.

EDIT: Sorry Bert, this post was before yours, but I tried to edit it and messed up so posted it again, by which time you posted the above ... .
I use hdpe quite a lot for stupid-high-voltage stuff, almost perfect - until it melts.



[Edited on 16-11-2017 by Sulaiman]


You do realize 1" of UHMWPE will stop most conventional firearms, right? That's an armor plate you've got right there, and an expensive one at that - I'd give an arm and a leg to have that much UHMWPE at one time.

gluon47 - 21-11-2017 at 19:02

Could benzyl nitrate be prepared via nitration of benzyl alcohol,
or is the aromatic ring likely to be nitrated to, forming a mixture of nitrobenzyl nitrates?

Benzyl nitrate doesn't appear to be very well documented, but from what I can find it seems to be a liquid that's fairly stable at room temperature.


ninhydric1 - 21-11-2017 at 21:47

Would a paint can hold up against plante1999's synthesis of TiCl4, as quoted below (assuming it actually works)? If not, what is a suitable container for this reaction (which will take place at around 400 degrees Celsius)? I will be using a mini butane burner I happen to have.

Quote: Originally posted by plante1999  
Simply melt pyrosulphate with the ore in presence of NaCl, it will make HCl + TiCl4

Sources: my own work

In fact I already made most inorganic titanium compounds and many organic ones, I think that I have some(a lot) knowledge in this field.

[Edited on 17-5-2012 by plante1999]

[Edited on 17-5-2012 by plante1999]


EDIT: Added clarifications.

[Edited on 11-22-2017 by ninhydric1]

Morgan - 23-11-2017 at 12:18

I was needing a rubber stopper so I went to my box of neoprene rubber stoppers with each size grouped in a Ziplock bag. But they have become oily and was wondering what's going on with that? What might the oily substance be?

JJay - 12-12-2017 at 15:37

What is the name of the theoretical sieve that reverses entropy by allowing only high-energy molecules to penetrate while elastically bouncing off low-energy molecules?

DraconicAcid - 12-12-2017 at 16:10

Maxwell's demon?

JJay - 12-12-2017 at 17:21

Quote: Originally posted by DraconicAcid  
Maxwell's demon?


I think that's it. I remember hearing it described as something like a bucket that leaks boiling water as the water inside freezes.

DraconicAcid - 12-12-2017 at 20:54

Larry Niven's "Unfinished Story #1" (or #2) involved the idea.

Σldritch - 14-12-2017 at 13:52

Can sodium pyrosulfate be used as a dehydrating agent similar to, say, anhydrous copper sulfate?

Cryolite. - 17-12-2017 at 01:49

Does anyone know of a good way to distinguish rhenium and tungsten? I bought some cycloid tube filaments on ebay advertised as rhenium, but I've been having doubts on their authenticity.

j_sum1 - 17-12-2017 at 03:01

Well, tungsten will react with H2O2 to form tungstic acid which is a pretty yellow-coloured solid. Tungsten is also quite dense with a specific gravity of 19.3. I don't know if either of these properties is sufficient to ditingush it from rhenium but that is where I would start.

Cryolite. - 17-12-2017 at 03:08

I've gotten that far: rhenium also reacts with hydrogen peroxide to form perrhenic acid, while tungsten forms a soluble complex in excess peroxide. Rhenium is also denser than tungsten, so that rules that one out.

I do notice that some of the strips have an iridescent tarnish on them. Is this a bad sign?

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