Sciencemadness Discussion Board

The Short Questions Thread (4)

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DutchChemistryBox - 10-12-2014 at 12:36

Hello,

This weekend I'm going to make methysalicylate with the procedure from the book of Arthur Vogel.

It says:
28 g. of salicylic acid and 64 g. (81 ml.) of absolute methyl alcohol, and add 8 ml. of concentrated sulphuric acid.

It should yield about 25g, after the reaction the methanol will be boiled of.

After a rough calculation I see that about 25% of the mixture will consist of sulfuric acid. Isn't this wayyyy to much? I'm afraid it will hydrolyse when I'm going to wash it with water?

How is such amount of sulfuric acid possible? Can somebody explain me why I'm wrong?

xfusion44 - 10-12-2014 at 15:34

@greenlight
@forgottenpasword

Thanks! :) Unfortunately I'm very busy these days, so I'll try again, when I'll have more time :)

xfusion44 - 10-12-2014 at 15:48

Quote: Originally posted by DutchChemistryBox  
Hello,

This weekend I'm going to make methysalicylate with the procedure from the book of Arthur Vogel.

It says:
28 g. of salicylic acid and 64 g. (81 ml.) of absolute methyl alcohol, and add 8 ml. of concentrated sulphuric acid.

It should yield about 25g, after the reaction the methanol will be boiled of.

After a rough calculation I see that about 25% of the mixture will consist of sulfuric acid. Isn't this wayyyy to much? I'm afraid it will hydrolyse when I'm going to wash it with water?

How is such amount of sulfuric acid possible? Can somebody explain me why I'm wrong?


Before you boil off methanol, there should be about 13.8% by weight sulfuric acid, but I don't know this reaction, so I don't know if H2SO4 reacts with something, or is there just to bind water to itself? However, if it doesn't react, after boiling, there should be 34.5% by weight H2SO4.

DutchChemistryBox - 11-12-2014 at 01:31

Quote: Originally posted by xfusion44  
Quote: Originally posted by DutchChemistryBox  
Hello,

This weekend I'm going to make methysalicylate with the procedure from the book of Arthur Vogel.

It says:
28 g. of salicylic acid and 64 g. (81 ml.) of absolute methyl alcohol, and add 8 ml. of concentrated sulphuric acid.

It should yield about 25g, after the reaction the methanol will be boiled of.

After a rough calculation I see that about 25% of the mixture will consist of sulfuric acid. Isn't this wayyyy to much? I'm afraid it will hydrolyse when I'm going to wash it with water?

How is such amount of sulfuric acid possible? Can somebody explain me why I'm wrong?


Before you boil off methanol, there should be about 13.8% by weight sulfuric acid, but I don't know this reaction, so I don't know if H2SO4 reacts with something, or is there just to bind water to itself? However, if it doesn't react, after boiling, there should be 34.5% by weight H2SO4.


It is just a fisher esterfication, the sulfuric acid acts as a catalyst. It will also force the equilibrium to the right due to the hygroscopic caracter of the acid.

Till that part, it won't be rocket science.

I'm trying to understand why such a big amount of sulfuric acid won't hydrolyse my product.

[Edited on 11-12-2014 by DutchChemistryBox]

Amos - 11-12-2014 at 06:27

DutchChemistryBox, you need all that sulfuric acid in order to sequester water produced as a result of the esterification, which will produce a molecule of water for each molecule of methyl salicylate esterified. The percentage of sulfuric acid will decrease as the amount of water increases, though why percentages matter so much I'm not sure. It won't hydrolyze as long as it's in the methanol.

Now, in regards to your fear of hydrolysis occuring, it has been my experience that if you decide to boil off the methanol from your reaction mixture IT WILL hydrolyse your product, and then more bad stuff happens. Trying this yielded a lot of foul-smelling black goo, probably due to the hydrolysis of the ester followed by an unfavorable reaction of the salicylic acid with so much now-concentrated sulfuric acid. In order to prevent this, I would recommend adding calcium carbonate to your reaction mixture when you believe it is sufficiently reacted. This will effectively remove the dangerous sulfuric acid, and the calcium sulfate formed should absorb water present in the mixture to form the dihydrate, which should take care of the water as well. Since it's insoluble, you can filter it off and squeeze any methanol/methyl salicylate mixture through the filter paper, and then simply boil off the methanol. There'll be some mechanical losses on such a small scale, though.



[Edited on 12-11-2014 by No Tears Only Dreams Now]

DutchChemistryBox - 11-12-2014 at 07:02

Quote: Originally posted by No Tears Only Dreams Now  
DutchChemistryBox, you need all that sulfuric acid in order to sequester water produced as a result of the esterification, which will produce a molecule of water for each molecule of methyl salicylate esterified. The percentage of sulfuric acid will decrease as the amount of water increases, though why percentages matter so much I'm not sure. It won't hydrolyze as long as it's in the methanol.

Now, in regards to your fear of hydrolysis occuring, it has been my experience that if you decide to boil off the methanol from your reaction mixture IT WILL hydrolyse your product, and then more bad stuff happens. Trying this yielded a lot of foul-smelling black goo, probably due to the hydrolysis of the ester followed by an unfavorable reaction of the salicylic acid with so much now-concentrated sulfuric acid. In order to prevent this, I would recommend adding calcium carbonate to your reaction mixture when you believe it is sufficiently reacted. This will effectively remove the dangerous sulfuric acid, and the calcium sulfate formed should absorb water present in the mixture to form the dihydrate, which should take care of the water as well. Since it's insoluble, you can filter it off and squeeze any methanol/methyl salicylate mixture through the filter paper, and then simply boil off the methanol. There'll be some mechanical losses on such a small scale, though.



[Edited on 12-11-2014 by No Tears Only Dreams Now]


Thank you for the input!

It still makes me wonder why Vogel does it that way.

In his procedure he washes his product with water after boiling of the methanol. Would it be possible to forget about whole the idea of boilling of and just start directly with washing with water and bicarbonate?

Amos - 11-12-2014 at 18:22

If you left the methanol around it would probably greatly increase the solubility of the methyl salicylate in the solution. I don't see how you could possibly boil off the methanol without charring and destroying the methyl salicylate, so the Vogel method really doesn't seem viable to me. But I haven't tried it, so I could be wrong.

gdflp - 11-12-2014 at 21:01

Washing the acidic ester with saturated NaHCO3 to remove excess acid, then washing with saturated NaCl to remove water would most likely eliminate most mechanical losses rather than adding calcium carbonate. The solution should then be able to boil without fear of decomposition.

Amos - 11-12-2014 at 21:36

Quote: Originally posted by gdflp  
Washing the acidic ester with saturated NaHCO3 to remove excess acid, then washing with saturated NaCl to remove water would most likely eliminate most mechanical losses rather than adding calcium carbonate. The solution should then be able to boil without fear of decomposition.


You can't merely wash the ester until you've isolated the ester from all of the methanol and acid it's stuck in. If you boil off the methanol without neutralizing the acid, your ester will decompose. If you wash the whole contents of the reaction vessel with sodium bicarbonate solution, you'll be stuck with an aqueous/methanolic solution of methyl salicylate that is rapidly hydrolyzing thanks to all of the water. At that point I don't think you could boil off the water and methanol without completely ruining the ester, as removing methanol while water is present will shift the equilibrium even further to the left.

Given that sodium sulfate is insoluble in methanol, you could maybe add solid sodium bicarbonate to avoid introducing water, but the neutralization still produces water. When you went to boil the methanol off, your now hydrated sodium sulfate would melt, which might cause hydrolysis of the ester at such a high temp; not sure.

gdflp - 11-12-2014 at 21:46

I have done this before, an organic layer of methanol and methyl salicylate will form while the sulfuric acid migrates to the aqueous layer and reacts. Also, water will not hydrolyze methyl salicylate on its own, it needs an acid catalyst for the hydrolysis to occur. After the ester has been washed, it can simply be distilled to separate the methanol and methyl salicylate.

Romain - 12-12-2014 at 15:33

Hi, is CaCl2 soluble in dichloromethane? I couldn't find any data on that online.

I want to remove water and ethanol from dichloromethane, but I can't distill the mixture with the CaCl2 (don't have the apparatus), so I thought I'd simply decant the CaCl2/ethanol/water and retrieve the dichloromethane. Now if the DCM is contaminated with CaCl2 that's not exactly ideal...

DraconicAcid - 12-12-2014 at 16:35

Quote: Originally posted by Romain  
Hi, is CaCl2 soluble in dichloromethane? I couldn't find any data on that online.

I want to remove water and ethanol from dichloromethane, but I can't distill the mixture with the CaCl2 (don't have the apparatus), so I thought I'd simply decant the CaCl2/ethanol/water and retrieve the dichloromethane. Now if the DCM is contaminated with CaCl2 that's not exactly ideal...


No, it is not soluble in dichloromethane.

Justin Blaise - 12-12-2014 at 21:55

I've followed the Vogel procedure for methyl benzoate, which also calls for distilling the methanol off at the end of the reaction if I remember correctly, and achieved the reported yields.

DutchChemistryBox - 13-12-2014 at 04:43

Quote: Originally posted by Justin Blaise  
I've followed the Vogel procedure for methyl benzoate, which also calls for distilling the methanol off at the end of the reaction if I remember correctly, and achieved the reported yields.


So it seems to work, I've been searching for people who followed the Vogel procedure. I've heard good stories about it.

I'm quite confused now. Almost everbody seems to think that it won't be good practise due to the risk of hydrolysing the product. But it does seem to work.

The question is why. Maybe there is not enough water present for hydrolysing (due to the hygroscopic caracter of Sulfuric acid)?
That it will only hydrolyse if there is more water or less acid present? That boiling of the methanol only is possible because of the high amount of acid?

xfusion44 - 13-12-2014 at 19:22

I have a few questions about chloroform...

I've made probably about 14ml of CHCl3, using 1l 67g/l NaClO and approx. 14ml of acetone.

First: balanced equation says, that there should be 1 mole of chloroform, when using 3moles NaClO and 1 mole acetone - in my case .9 and .3 moles of each, so I should get about 35g of chloroform. I didn' measure it yet, but it looks more like 15ml (approx. 20g), than 35, so how is that? I've also heard, that there should be about as much chloroform as the amount of acetone, that was used, is that true?

Also, how would I neutralise reaction mixture afterwards? Can I get anything useful from it? NaOH or sodium acetate maybe? Probably I can't just pour it down the drain, so there must be a way to neutralise it.

And the last question is about stabilising it. Would 1% bw of ethanol be enough?

Thanks and sorry for my english :)

xfusion44 - 13-12-2014 at 20:04

PS: also, what color should reaction mixture be? Last time I used excess acetone and after 3h mixtute was completely clear - like water, but now it's not clear at all and a little yellowish in color - did I use too little acetone this time?

Thanks!

DrMario - 14-12-2014 at 01:01

I am vaguely familiar with the haloform reaction, and how it can be used to make iodoform. But what would happen if one added HCl or other strong acid to acetone + iodine, instead of a strong base?

confused - 14-12-2014 at 01:33

im guessing that nothing would happen seeing as HCL and iodine will not react together, the iodine might dissolve a bit in the water from the HCL but that should be it

DrMario - 14-12-2014 at 03:05

I started with a dark brown solution, and after adding HCl I have now a clear solution with a light coloured precipitate (can't tell the colour exactly as it's in an amber bottle.

There was some potassium sulfate and traces of potassium iodide in the solution as well, so this isn't a well controlled experiment. I will have to repeat it.

EDIT: after pouring in a pyrex beaker, I see that the liquid is light yellow. The precipitate is brown.

[Edited on 14-12-2014 by DrMario]

greenlight - 14-12-2014 at 06:16

@ Xfusion,

I have always made chloroform similar to the way you specified with 12.5% Sodium hypochlorite solution which I buy in the pool section at the hardware store.
My ratios are 1000ml NaCl to 45.5ml Acetone but sometimes I scale it up and do 4 litres of NaCl at a time in a bucket if I need a larger amount.
I chill the Sodium hypochlorite to about 0-5 Degrees Celcius, then add the acetone, stir and keep it on ice for a couple hours afterwards before I seperate it in a sep funnel.
The yields aren't that amazing.
My solution usually goes from the original yellow-green colour of the NaCl to yellow to a milky white with a blob of chloroform on the bottom of the vessel.
I only make it when I need it now and rarely store it, but when I did i added a small amount of Ethanol to be safe.

xfusion44 - 14-12-2014 at 16:53

Quote: Originally posted by greenlight  
@ Xfusion,

I have always made chloroform similar to the way you specified with 12.5% Sodium hypochlorite solution which I buy in the pool section at the hardware store.
My ratios are 1000ml NaCl to 45.5ml Acetone but sometimes I scale it up and do 4 litres of NaCl at a time in a bucket if I need a larger amount.
I chill the Sodium hypochlorite to about 0-5 Degrees Celcius, then add the acetone, stir and keep it on ice for a couple hours afterwards before I seperate it in a sep funnel.
The yields aren't that amazing.
My solution usually goes from the original yellow-green colour of the NaCl to yellow to a milky white with a blob of chloroform on the bottom of the vessel.
I only make it when I need it now and rarely store it, but when I did i added a small amount of Ethanol to be safe.


Thanks! Today I did it again and noticed, that I've made mistake yesterday in calculations. I should use about 17.4ml, not 14ml, so maybe that's why it was still little yellowish in color.

How's about rxn mixture, if you know, maybe? Can I just pour it down the drain? Probably not?

Thanks ;)

gdflp - 14-12-2014 at 17:12

Once you've separated the chloroform, simply let the water portion of the reaction mixture sit for a few days. The haloform reaction produces 2 equivalents of sodium hydroxide which will slowly destroy the residual chloroform and convert it to a mixture of chloride and formate which can be poured down the drain.

xfusion44 - 14-12-2014 at 19:37

Quote: Originally posted by gdflp  
Once you've separated the chloroform, simply let the water portion of the reaction mixture sit for a few days. The haloform reaction produces 2 equivalents of sodium hydroxide which will slowly destroy the residual chloroform and convert it to a mixture of chloride and formate which can be poured down the drain.


Thanks! I've just saw that in one of yt videos on how to make chloroform :)

I Like Dots - 18-12-2014 at 14:37

Hey guys, today I was in a thrift store, and a shiny butter knife caught my eye.

The knife is silver plated, on copper on an unknown base metal/alloy, The purpose of this post is identifying the base metal/s. Im fairly sure its pewter, but just curious anyway. Im trying to design an experiment to identify the different elements.

Physical properties
Chemical properties





The metal in question


The knife, melted and broken in half


The plating layers visible on the knife


The Knife


The emblem on the knife

Metacelsus - 21-12-2014 at 20:07

I want to turn p-nitrotoluene into p-aminobenzoic acid. I plan to oxidize the methyl group (with permanganate) and reduce the nitro group (with Sn/HCl).

I've read a reference which calls for acetylation with acetic anhydride before the oxidation (http://pubs.acs.org/doi/abs/10.1021/ed033p71). I have some acetic anhydride, but would rather not use it for this, as it was not cheap to get. Is this acetylation strictly necessary?

Could I oxidize the methyl group first, and then reduce the nitro group?

(After more research, yes, I can. I guess I answered my own question.)

[Edited on 22-12-2014 by Cheddite Cheese]

AlphaDecay - 22-12-2014 at 11:19

I have some MnO2 from zinc-carbon batteries, but I won't use it since it is too impure, and I don't know what is the best way to dispose of it. Any ideias? I searched for information on MnO2 MSDS about disposal and found nothing.

Metacelsus - 22-12-2014 at 11:45

I would just put it in the trash. Manganese dioxide is insoluble, so it won't migrate out of a landfill.

bismuthate - 26-12-2014 at 13:13

So, how could I get a cerium containing solution from CeO2 (without H2SO4).
Also why does Na2MoO4 not react with HCl to give MoO3?

[Edited on 26-12-2014 by bismuthate]

Oscilllator - 26-12-2014 at 17:23

Quote: Originally posted by bismuthate  
So, how could I get a cerium containing solution from CeO2 (without H2SO4).
Also why does Na2MoO4 not react with HCl to give MoO3?

[Edited on 26-12-2014 by bismuthate]

I believe it does. In the past I have observed a white precipitate upon acidification of Na2MoO4

Argentum - 28-12-2014 at 13:29

Can I evaporate a solution of KOH with heating in a glass beaker? I know molten hydroxides can eat glass but just heating it up to evaporate water will eat through my beaker?

Also, do anybody knows how to purify slaked lime, Ca(OH)2?
Thanks!

DraconicAcid - 28-12-2014 at 13:58

Hot concentrated KOH solution will etch glass.

bismuthate - 28-12-2014 at 14:14

Add vinegar to the Ca(OH)2 and then add KOH to that and filter.

Argentum - 28-12-2014 at 14:31

Thanks Draconic and bismuthate for your help!

bismuthate - 28-12-2014 at 16:56

Why are solutions of Cr chloride and sulfate different colors when chloride and sulfate ions are colorless and they both have the hexaaquachromium ion?
Or does CrCl3 form a chloroaquachromium ion? And if so how come it doesn't go purple upon dilution?

[Edited on 29-12-2014 by bismuthate]

DraconicAcid - 28-12-2014 at 18:02

Chromium(III) is very slow to exchange ligands. If you were to dissolve chromium in hydrochloric acid, you could conceivably get a triaquotrichlorochromium(III) complex, a tetraaquodichlorochromium(III) complex (in cis and trans varieties, which will also be different colours), pentaaquochlorochromium(III) ion, or hexaaquochromium(III). Diluting the solution will not quickly convert the chloro complexes into the hexaaquochromium(III) ion- ligand exchange can take weeks, depending on the solution conditions.

Backyard Chemist - 28-12-2014 at 22:59

Recently bought a bag of Calcium ammonium nitrate and was wondering how to separate the ammonium nitrate. Repeated filtration? A simple method please. And btw i'm not planning on using it as an explosive, going to use it to make some potassium nitrate.

DraconicAcid - 29-12-2014 at 00:53

Quote: Originally posted by Backyard Chemist  
Recently bought a bag of Calcium ammonium nitrate and was wondering how to separate the ammonium nitrate. Repeated filtration? A simple method please. And btw i'm not planning on using it as an explosive, going to use it to make some potassium nitrate.


If you add potassium carbonate to that, the calcium should precipitate out as calcium carbonate. Filter that out, and heat the solution to drive off carbon dioxide and ammonia, leaving potassium nitrate.

Backyard Chemist - 29-12-2014 at 01:27

All right thanks!
I have a few questions if you don't mind:
1. If I was just to isolate the ammonium nitrate I could just heat the calcium ammonium nitrate solution and filter it right?

2. wouldn't a saturated solution of potassium chloride and ammonium nitrate produce potassium nitrate?

Thanks in advance:)

gdflp - 29-12-2014 at 08:22

1. Yes, though it wont be very efficient since there is approximately a 5:1 mol ratio of calcium to ammonium nitrates. A more efficient way would be to add ammonium sulfate just until calcium sulfate stopped precipitating, then filter the solution. This will leave you with approximately 6 equivalents of ammonium nitrate in solution.

2. Yes it will due to solubility differences if chilled to low enough temperatures. This will result in some chloride and ammonium contamination though, so recrystallizing it would be useful to increase the purity. In addition, potassium nitrate is still rather soluble at low temperatures, so a lot of product will be lost in the solution. Adding potassium hydroxide, or potassium carbonate and heating, will result in a much better, purer yield of potassium nitrate.

Backyard Chemist - 29-12-2014 at 17:40

How long would it have to cool for in order to let the AN crystals to appear? How much can you expect from this process?

Zephyr - 29-12-2014 at 18:50

All of you questions have been answered previously and it will save both you and me time if you first search for the answers yourself.
But while I'm here, I might as well help you out ;)
Firstly, when crystallizing something, the colder the solution gets the more of your product and impurities will crystallize out (Usually).
In the case of your ammonium nitrate, at 20 degrees Celsius every 10ml of solution will contain 15 grams of your product, meaning you will lose 15g of ammonium nitrate for every 10 ml of solution you crystallize. This means that in order to maximize your yield, you will need to boil off most of the water before crystallizing. And because of your (hopefully) low amounts of impurities, you final product should be reasonably pure.
Secondly, you need to find the best possible yield. Because the molecular weight of ammonium nitrate is 80 and the weight of calcium ammonium nitrate is 584, the yield will be about 13 percent of total ammonium nitrate based on calcium ammonium nitrate. Minus whatever you don't crystallize. So I'd you start with 1kg of CAN you can expect about 110 g total ammonium nitrate.

[Edited on 12-30-2014 by Pinkhippo11]

Backyard Chemist - 29-12-2014 at 19:32

Thanks mate!

Jylliana - 6-1-2015 at 05:02

Does Bismuth(aq) react with copper(s)?
Because my redox table says it should, but my eyes say it doesn't.

I'm confused.
I tried to make elemental bismuth out of bismuth subcarbonate.

EDIT: I think my Bismuth is readily dissolving in my HCl-solution where it's in.

!@#$ How am I gonna fix that.. Neutralizing wil just yield bismuth oxide again >.<

[Edited on 6-1-2015 by Jylliana]

Pasrules - 6-1-2015 at 05:21

Quote: Originally posted by Jylliana  
Does Bismuth(aq) react with copper(s)?
Because my redox table says it should, but my eyes say it doesn't.

I'm confused.
I tried to make elemental bismuth out of bismuth subcarbonate.


This reactivity series I found would say no. However it also says platinum is more reactive than gold which I had to argue at my chemistry teacher about because that's what school syllabuses teach. Most of these tables I find vary text book to text book and arn't all done in the same conditions.

Next questions which is more reactive gold or platinum...? Because I've seen platinum take a far longer amount of time in Aqua Regia.

gdflp - 6-1-2015 at 09:44

Based on the reactivity series, platinum is less reactive than gold. Though I don't believe that the speed of dissolution in acids is a good measure of it.

@Jylliana
Copper will plate onto bismuth metal, not the other way around. You will need a more reactive metal such as Fe, Zn, Mg, etc. This is a useful reference.

CuReUS - 6-1-2015 at 23:50

I have a few questions :) :

1.what is the best chemical test to confirm an aromatic aldehyde(benzaldehyde) from an aliphatic aldehyde other than burning it and seeing if the flame is sooty ?

2.why cant the H of an aldehyde group(CHO) be pulled out,why only the alpha H is removed ?

3.why does ethanol on treatment with conc H2SO4 give an oily layer(ethyl hydrogen sulphate) at one time and a white precipitate at another ?

Jylliana - 13-1-2015 at 04:12

I always believed that Gallium was a pretty non-toxic metal. Nothing special about it, except it's low melting point.

Why is it that it requires HAZMAT shipping from everywhere I've looked?
I know it forms an alloy with a lot of stuff, but just shipping it in a vial or bottle shouldn't be a problem? Or is there something about this metal that I overlook?



[Edited on 13-1-2015 by Jylliana]

Brain&Force - 13-1-2015 at 08:27

No, it's not toxic, it'll just dissolve stuff too easily.

DraconicAcid - 13-1-2015 at 17:13

Quote: Originally posted by CuReUS  

2.why cant the H of an aldehyde group(CHO) be pulled out,why only the alpha H is removed ?


If you remove the alpha hydrogen, you get an anion that is stabilized by resonance. If you pull off the aldehyde carbon, you don't get any resonance.

Jylliana - 16-1-2015 at 06:36

I found a very old dessicator in my lab. I wouldn't be surprised if it hasn't been touched for over 20 years.
I want to use it, but the two halves are VERY stuck together. I think the silicon grease has dried out or something.
I've tried pulling very hard, prying with a screwdriver, asking it politely, but nothing works. I am afraid the thing may shatter if I use more force or heat.
Do you have any ideas?
I did manage to lift the stopper on top, so there is no vacuum.

gdflp - 16-1-2015 at 06:41

Try WD40, spray some around the joint and let it sit for a few minutes. Repeat this a few times and then let it sit for an hour. The goal is to try to get some WD40 to leak into the joint and loosen it. This is the only method I use for stuck glassware, but some other methods which I haven't tried are gently tapping the joint with a rubber mallet while trying to pull it apart, and using a mixture of acetone and acetic acid to replace the WD40. Vacuum desiccators generally have very thick glass, so I would imagine that trying to heat up and expand one side of the joint won't work, but you could try.

Metacelsus - 16-1-2015 at 07:17

I had a problem with a stuck vacuum dessicator once (at a lab at a local university), and it was solved by use of a heat gun. The nice thing is that it's a flat joint, so it's easy to heat up one side (the top) only.

Jylliana - 16-1-2015 at 08:35

Noted. I'll try it. Thanks, both :)

Zombie - 16-1-2015 at 19:28

Following Cheddite's lead...
Heat source on the outer join, and cold (Dry Ice) held on the inner join will separate them almost instantly unless there is a chemical bond.

I work with tapers, and press fit parts regularly, and nothing works better than heat/cold applied to opposite sides.

The other end of this spectrum is tapping. A rubber mallet will not work (well). Actually a hard plastic hammer is much better, and you do NOT hit the join. You hit the bottom edge of the flask/beaker/whatever. The idea is to create a shock wave that distorts the taper and the join will expel what ever is in it.

I've used shock cord/det cord to remove boat props, when heat/cold failed.

If it breaks after all of this... Good riddance!

The Volatile Chemist - 18-1-2015 at 18:01

I added a dilute solution of sodium bicarb. to a concentrated solution of cobalt chloride. A purple-ish precipitate formed after some time, but no carbon dioxide was formed. Furthermore, although I used stiochemically accurate quantities of reagents, the solution is still the color of cobalt chloride. I've found in the past I've had some trouble figuring out carbonates vs. hydroxides, and which are formed, some help would be nice.

Jylliana - 19-1-2015 at 03:05

My dessicator is unstuck!
Applying heat to the top half did the trick. I didn't have WD40 on hand so I tried the heat first. Thanks! :D

Brain&Force - 19-1-2015 at 11:24

Quote: Originally posted by The Volatile Chemist  
I added a dilute solution of sodium bicarb. to a concentrated solution of cobalt chloride. A purple-ish precipitate formed after some time, but no carbon dioxide was formed. Furthermore, although I used stiochemically accurate quantities of reagents, the solution is still the color of cobalt chloride. I've found in the past I've had some trouble figuring out carbonates vs. hydroxides, and which are formed, some help would be nice.


Don't add stoichiometric quantities when attempting precipitation.

All solids have a solubility product which defines the equilibrium between two dissolved species, in this case cobalt and carbonate. To drive the equilibrium to the right you need to add more of the precipitating agent. The excess carbonate drops more cobalt out of solution. This is known as the common-ion effect.

http://en.wikipedia.org/wiki/Solubility_equilibrium

The Volatile Chemist - 19-1-2015 at 14:43

Thanks.
And I'm not fond of your 'facebook' link. I'd seen it before, but without speakers........

Brain&Force - 19-1-2015 at 15:25

Just note that this represents what I think of Facebook...

The Volatile Chemist - 20-1-2015 at 09:09

Quote: Originally posted by Brain&Force  
Just note that this represents what I think of Facebook...

It represents what I think of Facebook, too, but still...
Thanks for your help, I precipitated all of my CoCO3.

SimpleChemist-238 - 22-1-2015 at 18:54

What is the safest and most legal way to get dispose of small amounts of uranium water such as .5g of UO2? Hypothetical.

j_sum1 - 22-1-2015 at 21:20

Quote:
most legal

Legal is not a sliding scale. It is either legal or illegal. And it would depend on your jurisdiction.
I doubt anyone is going to come knocking on your door though (Unless you have an active branch of the Nuclear Police armed with geiger counters patrolling your area.)
I think the question you need is, "What is the most environmentally responsible means for disposing of waste uranium solutions and precipitates?" For that, I would like to know the answer too. (Equally hypothetical.)
My very small quantities of solutions and precipitate wastes are (1) precipitated to stable safer forms as much as possible, (2) concentrated via evaporation and (3) mixed with cement to make concrete blocks for solid waste disposal. That said, the highest toxicity I play with is Pb. I am not sure that this procedure is suited for anything really nasty or for anything radioactive.

Zyklon-A - 22-1-2015 at 21:42

Depleted uranium is barely radioactive.
It is a heavy metal however and is toxic for that reason. I'd. Precipitate it as the most insoluble salt as possible, that is not too toxic. Then bring it to the same place I take broken mercury lightbulbs to, Home Depot in my case.

SimpleChemist-238 - 22-1-2015 at 22:15

Can you make it into a peroxide thats insoluble and flush it into a utility drain with lots of water? Its depleted and the most dangerous aspect is the heavy metal part.

[Edited on 23-1-2015 by SimpleChemist-238]

j_sum1 - 22-1-2015 at 22:37

Quote: Originally posted by Zyklon-A  
Depleted uranium is barely radioactive.
It is a heavy metal however and is toxic for that reason. I'd. Precipitate it as the most insoluble salt as possible, that is not too toxic. Then bring it to the same place I take broken mercury lightbulbs to, Home Depot in my case.

Agreed on the depleted U. However, if ores are your start material (I am thinking of a particular video in the bad chemistry videos thread) then you are going to have mixed heavy metal radioactive salts as waste. I really don't know what is an acceptable disposal method for that. I certainly don't have anyone or anywhere I could "take it to".

It is a potential area of interest. I would want to have my disposal procedures fully sorted out before I began even buying anything. I know that there are a few chemists on this site that have experience of this. I would like to know what a sensible solution is.

SimpleChemist-238 - 23-1-2015 at 07:35

My starting material is depleted UO2. Not ore.

The Volatile Chemist - 23-1-2015 at 13:20

Quote: Originally posted by SimpleChemist-238  
My starting material is depleted UO2. Not ore.
Starting to sound less hypothetical. :) But that's OK, it's just best to be honest on this site.

AlphaDecay - 25-1-2015 at 12:38

What is the best way to remove colloidal particles in 98% sulfuric acid? I'd assume that concentrated H2SO4 would react with ordinary filter paper, wouldn't it?
Thanks

Chemosynthesis - 25-1-2015 at 12:48

Yes.
Vacuum distillation might be fastest. Depending on your particle size, a frit might clog.

SimpleChemist-238 - 27-1-2015 at 10:40

has any one had problems with a very old sample of HCl and or HCl from Home Depot. In 3 experiments the HCl was not doing what I needed it to do. I think that it is because of how old the sample of acid is or that I got it from Home Depot.

Zombie - 27-1-2015 at 10:49

You may have had another issue. I use HCL (Home Depot/Ace Hardware) almost daily for cleaning dark water stains off of fiberglass boat hulls. I have several cases I bought almost 5 years ago, and several opened bottles that have sat in different locations for years.
I have never had a bottle "go bad".
Perhaps if the bottle were left wide open to the environment for a year or more... but never anything like an "old stock" issue.

Molecular Manipulations - 27-1-2015 at 11:00

Quote: Originally posted by SimpleChemist-238  
has any one had problems with a very old sample of HCl and or HCl from Home Depot. In 3 experiments the HCl was not doing what I needed it to do. I think that it is because of how old the sample of acid is or that I got it from Home Depot.

If you want an answer you should give more details, what experiments? What did it do?
What's the brand name? I've used Home Depot HCl (aq) and never had any problems.

SimpleChemist-238 - 27-1-2015 at 11:12

I tired to dissolve gallium metal in the HCl and H2O2. No reaction even over a few days. Was it just the experiment? I know it reacts very slowly but after 3 days I would expect some reaction.

Butane cooling

math - 4-2-2015 at 11:38

Question:

can I safely put a small butane canister (lighter refills) at -15°C for some hours then expect to take it out and pour liquid butane out of it by pushing the valve (as usually done when filling a lighter) ?

Thank you :)

Argentum - 4-2-2015 at 13:44

Question

Is there any way of recognizing sulphuric acid without using a pH-meter or pH paper?

It's from a liquid drain cleaner, it should be H2SO4 but I'd like to be sure.
No, the label doesn't says the content of the bottle.

Thanks!

gdflp - 4-2-2015 at 13:56

Quote: Originally posted by Argentum  
Question

Is there any way of recognizing sulphuric acid without using a pH-meter or pH paper?

It's from a liquid drain cleaner, it should be H2SO4 but I'd like to be sure.
No, the label doesn't says the content of the bottle.

Thanks!


Are you assuming that it's concentrated? If so, place a few drops on some paper, sugar, or other organic material and wait for a few seconds. If quickly turns black, then it is most likely sulfuric acid. This occurs because the sulfuric acid rips the organic material apart and dehydrates it, leaving carbon.

Metacelsus - 4-2-2015 at 18:14

If you mix it with an iodide salt, you should get iodine and sulfur dioxide/sulfur/hydrogen sulfide (reduction products of sulfuric acid).

Other strong acids can dehydrate organic matter, but gdflp's test is generally good.

Alternatively, add a soluble barium salt, and look for a precipitate. This will rule out some acids like nitric and hydrochloric (but not phosphoric).

[Edited on 5-2-2015 by Cheddite Cheese]

math - 5-2-2015 at 14:50

Question:

can I safely put a small butane canister (lighter refills) at -15°C for some hours then expect to take it out and pour liquid butane out of it by pushing the valve (as usually done when filling a lighter) ?

Thank you :)

Zombie - 5-2-2015 at 16:08

Short answer is no. Not really.
Butane evaporates so quickly in air that you really have to keep it contained.
I suppose you could open the can in a -15* C environment, and potentially transfer it into another frozen container but simple things like body temp would cause it to boil off. Then there is the vaporization in air... It's really not practical unless the goal is to switch sealed containers. Then the answer would be yes. As long as it is maintained.

Edit: By "pushing the valve" the flow rate would be too slow, and the exposure to air would dissipate it. Soooo Back to no. For that method.

[Edited on 6-2-2015 by Zombie]

Loptr - 5-2-2015 at 19:56

Will hydrolysis of acetamidine eventually lead to its acid, acetic acid?

I am jumping around in my old textbook, and I am not sure if amidines behave differently. I have found reference to them being easily hydrolyzed, but just not to what extent.

math - 6-2-2015 at 07:31

Do alcohols get more expensive (mole-wise, seems more comparative than weight- or volume-wise) the longer the carbon chain gets?

So, is isopropanol cheaper than butanol, pentanol and cetyl alcohols?

Zombie - 6-2-2015 at 14:04

It all depends on the manufacturing process. Some are easier to produce than others.

Pricing is easy...

https://www.spectrumchemical.com/OA_HTML/index.jsp?minisite=...

http://www.hobbychemicalsupply.com/

[Edited on 6-2-2015 by Zombie]

Loptr - 6-2-2015 at 15:59

Quote: Originally posted by math  
Do alcohols get more expensive (mole-wise, seems more comparative than weight- or volume-wise) the longer the carbon chain gets?

So, is isopropanol cheaper than butanol, pentanol and cetyl alcohols?



Cetyl alcohol
http://www.soapgoods.com/Cetyl-Alcohol-palmityl--p-675.html

Loptr - 6-2-2015 at 17:56

Quote: Originally posted by Loptr  
Will hydrolysis of acetamidine eventually lead to its acid, acetic acid?

I am jumping around in my old textbook, and I am not sure if amidines behave differently. I have found reference to them being easily hydrolyzed, but just not to what extent.


To answer my own question, yes it will.

Zombie - 7-2-2015 at 02:07

Quote: Originally posted by Loptr  
Quote: Originally posted by Loptr  
Will hydrolysis of acetamidine eventually lead to its acid, acetic acid?

I am jumping around in my old textbook, and I am not sure if amidines behave differently. I have found reference to them being easily hydrolyzed, but just not to what extent.


To answer my own question, yes it will.



I see how you did that. :)

Argentum - 8-2-2015 at 16:16

Quote: Originally posted by gdflp  
Quote: Originally posted by Argentum  
Question

Is there any way of recognizing sulphuric acid without using a pH-meter or pH paper?

It's from a liquid drain cleaner, it should be H2SO4 but I'd like to be sure.
No, the label doesn't says the content of the bottle.

Thanks!


Are you assuming that it's concentrated? If so, place a few drops on some paper, sugar, or other organic material and wait for a few seconds. If quickly turns black, then it is most likely sulfuric acid. This occurs because the sulfuric acid rips the organic material apart and dehydrates it, leaving carbon.


Nothing happened. Maybe it's muriatic acid, it looks like cheap hardware store muriatic acid.

Another question. Is cornstarch (the one sold in supermarkets) relatively pure starch?

Thanks

Zombie - 8-2-2015 at 16:24

Not really.
They are all going to contain at least some Iron, Ash to prevent clumping, and some Sulfur Dioxide from processing.
This applies to Pharmaceutical grade. Grocery or Food Grade will have much higher PPM.
I would stick to Pharma grade if purity is a concern.

The Volatile Chemist - 9-2-2015 at 10:18

If I boil a solution of Urea to saturation, and ammonia is able to be smelled coming off the beaker, how much has turned to biuret? Like, most of it, or almost none of it? (Or in between, I'm looking for 'about's here).

Loptr - 9-2-2015 at 11:08

Quote: Originally posted by The Volatile Chemist  
If I boil a solution of Urea to saturation, and ammonia is able to be smelled coming off the beaker, how much has turned to biuret? Like, most of it, or almost none of it? (Or in between, I'm looking for 'about's here).


I would argue that if you are using fertilizer grade urea, there is already biuret in the mixture. ;)

* Note: I have read that biuret is a common impurity in OTC urea. Source: Somewhere in the google-sphere.

If you heat Urea above its melting point, and smell ammonia, then you have likely produced some biuret. Whether this is the only product, I can't say at this time.
http://en.wikipedia.org/wiki/Biuret#Preparation

[Edited on 9-2-2015 by Loptr]

Sulaiman - 10-2-2015 at 03:49

math;
I use a combination of liquid and gaseous butane as a dielectric for one of my high voltage probes (P6015) and the answer to your question is YES.

I don't pre-cool the butane canister or the vessel that it will go into
and I do this indoors though outdoors would be wiser.

Just invert the butane can and depress the nozzle with pliers
(to avoid freezing your fingers)
squirt the liquid butane into your vessel
initially some liquid butane will boil off, cooling the receiving vessel,
then the vessel will begin to fill with liquid butane.

It's true that ambient heat will slowly boil the liquid butane
but the vapour pressure is about two atmospheres at room temperature
so it doesn't take much to seal your vessel and prevent boiling.
e.g. my probe is an aluminium tube sealed with a rubber compression washer,
I last re-filled it over two years ago and there is no sign of loss of liquid butane (the bottom of the probe is transparent plastic)

I have not tried but I expect that a rubber bung would seal a test tube of liquid butane at room temperature.

Update, just went to my lab/shed (presently c5 celcius!) for a quick test
no observable boiling in a glass vessel open to the air,
so I poured the liquid butane into a small plastic bottle and screwed the cap on .. not enough pressure to bulge the bottle.
BEWARE: the vapour pressure increases rapidly with temperature
see this wikipedia page http://en.wikipedia.org/wiki/Butane_(data_page)
and look at the graph 'Vapor pressure of n-butane'

/just did a bit of googling ...
a 2 litre coke bottle containing some liquid butane would probably burst if immersed in boiling water, so don't do that !
but it does indicate that at room temperature it should be ok.


[Edited on 10-2-2015 by Sulaiman]

[Edited on 10-2-2015 by Sulaiman]

[Edited on 10-2-2015 by Sulaiman]

The Volatile Chemist - 11-2-2015 at 09:14

Quote: Originally posted by Loptr  
Quote: Originally posted by The Volatile Chemist  
If I boil a solution of Urea to saturation, and ammonia is able to be smelled coming off the beaker, how much has turned to biuret? Like, most of it, or almost none of it? (Or in between, I'm looking for 'about's here).


I would argue that if you are using fertilizer grade urea, there is already biuret in the mixture. ;)

* Note: I have read that biuret is a common impurity in OTC urea. Source: Somewhere in the google-sphere.

If you heat Urea above its melting point, and smell ammonia, then you have likely produced some biuret. Whether this is the only product, I can't say at this time.
http://en.wikipedia.org/wiki/Biuret#Preparation

[Edited on 9-2-2015 by Loptr]

Thanks. It [the urea] is from driveway salt, one brand sells pet-safe stuff which is supposedly all urea. Of course, it has dye in it, but I found the dye is very soluble in isopropanol, thus I have been able to remove about 75-80% of the dye from the 'salt', and have turned it into a more crystalline form.

Sulaiman - 14-2-2015 at 05:13

Please excuse these noob questions but

Is there any kind of machine that can reliably identify all known compounds in a mixture?
and determine absolute amounts of each compound in the mixture?


On a more realistic level, at least all common pharmaceutical compounds.

Or at least,
is there some machine that can detect the presence of unlisted compounds in medicinal mixtures.

Potentially this would be for a commercial testing facility for food supplements etc,
(and probably a service to aid hospitals etc. in diagnoses
and other customers to be determined)

P.S of course a noob like me would not be the operator,
some kind of pharmacists/chemists would need to be hired and trained by the vendor(s) of this magic machine or suite of machines.

OR
which are the leading manufacturers of similar equipment?

[Edited on 14-2-2015 by Sulaiman]

fluorescence - 14-2-2015 at 05:29

Well I think there is no machine that can just look at it and tell you what it's made of.
But there are other possibilities. From NMR, IR to Mass spectrometry.
I have never really worked with them but they can show you the presence of functional groups.
It takes some time to read the diagramms but you could compare the signals to known substances and
so find out what a pharmaceutical compound is made of.

I've seen this documentation on new drugs since it is theoretically allowed to make something that resembles
a drug listed compound and then change some functional groups till you get a new molecule that is not enlisted.
On TV they bought some of this stuff and gave it to a laboratory and they were able to determine the structure.

So I'd say you can find out most of the molecular structure of a given compound but it might depend on what
machine you are using.

But as said, I'm no expert on this either, I might be wrong here.

Metacelsus - 14-2-2015 at 06:29

GC/MS might be your best bet, but even it has limitations.

Chemosynthesis - 14-2-2015 at 06:56

The pharmaceutical industry uses literally every analytical instrument type manufactured. It just depends what you are doing specifically. GC/MS is arguably the leading analytical test for things such as confirmatory drug testing in criminal cases (dealing, manufacture, sports doping). You have to know what you are looking for, though.
Adjunct confirmatory testing is often done with an IR or similar. Those two are how a local crime lab prefers to do business. In the actual pharmaceutical industry, NMR and IR spectra of each individual synthetic intermediate is taken, among other tests, and can be requested. Additional testing is often done.
The final pharmaceuticals tend to have USP-NF certified HPLC testing done, which is probably still the most common quantitative QC test in the pharmaceutical industry.

Individual instruments can vary, sensors vary, etc. Someone trained on an Agilent can't necessarily use a Waters. I'm not sure what to expect with NMR now, as Agilent bought Varian a few years back, and now both are exiting the NMR market. If this is a legitimate work thing, you need to just ask the employees what they need, and then start talking to regulatory agencies or start calling competitors/service providers and just ask what they use and why they think it's better than some other guy.

Sulaiman - 14-2-2015 at 07:13

That's the answer that I expected but didn't want,
so there is no universal analyser,
it is/was a potential business opportunity in a South East Asian country,
via relatives (I certainly could not afford the start up costs)

Chemosynthesis - 14-2-2015 at 09:22

One complication to keep in mind with overseas work is the labor force. Some standards of testing and regional markets may be different both in terms of employees and equipment.

battoussai114 - 19-2-2015 at 06:37

Are there any IR sources used in IR Spectrometry other than heated inert materials?

j_sum1 - 21-2-2015 at 05:25

Question: Slightly confused about carbon zinc battery.

I have disassembled a few lantern batteries -- for electrodes and other useful things.
First step in purifying the MnO2 paste is to mix with water to form a slurry and then filter. The intention is to dissolve the ZnCl2 or NH4Cl that acts as an electrolyte.
The other day I ran a current through it and discovered to my surprise that a nice grey metal plated out on the cathode. I thought I had a solution of ZnCl2 but of course Zn cannot be electrolytically reduced in aqueous solution. So, what is it that plated out? Do some batteries contain tin compounds or bismuth or something else strange? What is it that I have got?

What is the white frost on glass from being near HCl(aq)?

quantumcorespacealchemyst - 22-2-2015 at 08:49

What is the white frost on glass from being near HCl(aq)? It is wipable, smells odd and forms even near sealed acid containers. It seems to certainly be from HCl gas. What it is, I don't know.

[Edited on 22-2-2015 by quantumcorespacealchemyst]
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