Sciencemadness Discussion Board

The Short Questions Thread (4)

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DraconicAcid - 16-5-2016 at 14:22

Quote: Originally posted by glymes  
Why is it that there are a lot of organic nitrates (nitrotoluene, nitrobenzene etc.) but virtually no chlorates of the same ilk? This interests me.


Nitrobenzene is not an organic nitrate; phenyl nitrate would have an oxygen between the carbon and the oxygen.

Perchlorylbenzene does exist (http://pubs.acs.org/doi/abs/10.1021/i360060a011 ), but it takes courage to work with such materials.

Prep: https://books.google.ca/books?id=vD-qCgAAQBAJ&pg=PA624&a...

Who the hell wants to work with perchloryl fluoride in order to make the compound?

woelen - 16-5-2016 at 14:40

Nitrotoluene and nitrobenzene are not nitrates, they are nitro-compounds.

Nitrates have structure R-O-NO2.
Nitro's have structure R-NO2
Nitrites have structure R-O-NO (isomers of nitro's)

Organic nitrates are not that stable. There are a few (e.g. glycerine trinitrate) but these are sensitive explosive compounds.
Organic chlorates simply are too unstable to exist. Chlorate only is stable in ionized form, as ClO3(-). The covalent HClO3 does not exist, chloric acid only exists in aqueous solution up to appr. 40% in which it is fully ionized to H3O(+) and ClO3(-). Organic chlorates would be covalent and hence they are so unstable that in practice they do not exist at room temperature.
The ionization of chlorate stabilizes the ions by means of resonance. The covalent chlorates do not have this advantage.

A similar thing exists for perchlorates. Ionic perchlorates are very stable, covalent perchlorates are unstable. E.g. anhydrous HClO4 does exist, but is extremely dangerous, as it can explode without apparent reason. Aqueous HClO4 on the other hand is quite inert and hardly shows oxidizing properties, just acid properties.
There are a few organic perchlorates though. Someone on sciencemadness has actually done some experiments with organic covalent perchlorates. IIRC this was user axt, who did experiments with ethylperchlorate. I myself have made quite a few organic perchlorates, but all of these are salts of amines, or complexes of metals with organic ligands, with the perchlorate ion as anion.

glymes - 16-5-2016 at 22:10

AXT had the ethyl perchlorate 'spoon shot' didn't he (where a drop of ethyl perchlorate blew a metal spoon out of his hand?) Ethyl perchlorate looks very nasty.Apparently, methyl perchlorate is more unstable, more explosive and around as toxic as dimethyl mercury.

TinSandwich - 18-5-2016 at 05:08

Why do alchemical manuscripts always show tons of different still head while we modern chemists only have one and a couple different condensers?

Photos: "1606" https://en.wikipedia.org/wiki/File:Alembics_from_Andreas_Lib... "2016" https://commons.wikimedia.org/wiki/File:Simple_distillation_...

Jstuyfzand - 19-5-2016 at 13:22

When So2 gets dissolved in water it forms the unstable sulfurous acid, what does it decompose into if it is unstable?

DraconicAcid - 19-5-2016 at 14:31

Quote: Originally posted by Jstuyfzand  
When So2 gets dissolved in water it forms the unstable sulfurous acid, what does it decompose into if it is unstable?

Sulphur dioxide and water.

Jstuyfzand - 20-5-2016 at 13:02

Thank you

mintreina - 23-5-2016 at 04:58

Question


Does an iron nail displace cobalt ions from CoSO4 solution?
Will it form FeSO4 and Co metal?

Metacelsus - 23-5-2016 at 08:44

Cobalt(ii)'s standard reduction potential is -0.28 V, and iron(ii)'s is -0.44 V.

Therefore, yes.

TinSandwich

Sulaiman - 23-5-2016 at 10:01

It looks to me that some are simple still heads, also I see equivalents to;

. anti-splash function
. fractionating function
. condensing/refluxing function

it appears that some were for specific processes such as mercury distillation,
just as some manufacturing stills look quite different from lab versions.

if anything we now have more complex arrangements;
flask capacity, shape, wall thickness, number and type of neck
column types
still adapter types
condenser types
take-off types
receiving flasks
rotovaps, pig adapters, stirrers, perkin triangles, pressure-equalised funnels ..............

I want ALL of them :D

and I think an alchemist of times gone bye would freak out if he was given a Quickfit catalogue to choose from :o


[Edited on 23-5-2016 by Sulaiman]

hissingnoise - 23-5-2016 at 10:06

Quote:
Quote:
When So2 gets dissolved in water it forms the unstable sulfurous acid, what does it decompose into if it is unstable?

Sulphur dioxide and water.

Sulphurous acid is unstable because air slowly oxidises it to H2SO4!


aga - 23-5-2016 at 12:44

Quote: Originally posted by Sulaiman  
I want ALL of them :D

What for ?

The Volatile Chemist - 24-5-2016 at 15:24

I thought at first he was wanting all of the styles of stillheads combined. that would be a sight to see...

Distilling - question about condenser

RogueRose - 24-5-2016 at 15:24

I'm curious how to tell whether all the gases that come through a condenser and out any "extending" hose has captured all available liquids/gases. Let's use ethanol as an example. The temp of the condenser is 70 degrees F and the temp of the gas coming out the end of the 12" hose at the end is 85-115 F.

The thing is that at times when distilling I'll hold the hose in a "U" shape, fill it with condensate, and watch to see if gas vapors pass/push through.

It seems that at times the gas pushes through and I see a mist on the collecting vessel but the temp is 100F where the evap point is ~170.

Is it possible that there is still something in gaseous form at those temps?


The Volatile Chemist - 24-5-2016 at 15:30

Well, when you boil water, it can condense in the air as a mist....but if the so-called mist is a gas still (punintended), then it's possible the gas simply hasn't liquified yet. Cooling down and condensing and coalescing into a liquid in a flask take time.

Seperation of sodium sulfates

Emlingur - 27-5-2016 at 08:24

I have a mixture of sodium sulfate and sodium bisulfate, the former of about 5% by weight. I need to get rid of those 5%, and I was wondering if there was a solvent that only one of these two is soluble in, or if I have to do a recrystallization, and if so, which solvent would be preferrable to do so.

Praxichys - 27-5-2016 at 08:50

Make a solution of your contaminated bisulfate and add excess sulfuric acid. From this solution you will be able to crystallize sodium bisulfate of a higher purity. There will always be a tiny equilibrium of sodium sulfate that will coprecipitate, although a large excess of acid will make this negligible.

Why do you need pure sodium bisulfate? When dissolved it acts as 1eq of sodium sulfate mixed with 1eq of sulfuric acid. Unless you need it for very precise measurements like a titration (and this is a bad standard to use, by the way) you can always just calculate the excess you need to add to complete a reaction and not bother purifying it.

Emlingur - 27-5-2016 at 09:13

Quote: Originally posted by Praxichys  


Why do you need pure sodium bisulfate?


The sodium bisulfate will be dehydrated and then cracked to generate sulfur trioxide, although maybe the sodium sulfate will not have any effect on the process, but I'm not sure.

RogueRose - 27-5-2016 at 14:28

Is it possible to have a condenser that is too big (not like for distilling 200 gal w/ a 1L run) but something that may serve a 30L setup used in a 4-5L run?

Calculating moles of Acetic Acid in vinegar solution

RogueRose - 27-5-2016 at 19:39

This gets confusing to me when trying to figure out how much of something is in a solution with water.

The vinegar is 20% "200 grain" and 3785ml.
Acetic Acid 60g/mole & density of 1.049g/ml

3785ml @ 20% = 757ml
757ml @ 1.049g/ml = 794g
794g Acetic Acid @ 60g/mole = 13.234 moles

The math is correct but I'm not sure if my method is correct. Is this an accurate way of determining how many moles per gallon? I didn't weigh anything but only took the printed info for my numbers.

DraconicAcid - 27-5-2016 at 19:42

Your math is correct, but most of the vinegar I've seen is 5% v/v, which works out to almost exactly 1 mol/L.

Vacuum Gauge Noobness

Deathunter88 - 28-5-2016 at 00:17

Hello everyone, I have decided to buy a vacuum gauge so I can finally get into vacuum distillation. However, there are so many types of vacuum and so many units that I get hopelessly lost when trying to pick one. Attached is a photo of one that I plan to get, would this be good for most vacuum distillations?

Screen Shot 2016-05-28 at 4.16.22 PM.png - 690kB

Texium - 28-5-2016 at 06:33

Well, that looks like a decent gauge, and it does cover roughly the range that you need (-.1 MPa is pretty much full vacuum). However, from what I've seen MPa is not a unit you see used that often in chemistry papers or on this site, so you might find yourself converting a lot. Getting one that covers the same range but in mmHg (torr) would be easier as far as that goes as it is used more widely. Should read 0 to -760.

Ba(ClO3)2 - 31-5-2016 at 00:49

Would boiling a solution of sodium hydroxide with antimony trioxide result in a solution of sodium meta-antimonite ?

Sb2O3 + 2 NaOH ==> 2 NaSbO2 + H2O

If an excess of antimony trioxide were used then presumably a fairly pure solution of sodium meta-antimonite could be obtained?

The excess antimony trioxide would be easy to remove (filtration).

Come to think of it, is sodium meta-antimonite even stable in solution?

Sulaiman - 31-5-2016 at 05:10

Quote: Originally posted by Deathunter88  
Hello everyone, I have decided to buy a vacuum gauge so I can finally get into vacuum distillation. However, there are so many types of vacuum and so many units that I get hopelessly lost when trying to pick one. Attached is a photo of one that I plan to get, would this be good for most vacuum distillations?


I have one almost identical, the scale does not matter to me
because pressure is reported in different units by different reporters,
all you have to do is convert for recommended values then
keep notes in whatever units your gauge works.

This type of gauge is good for a rough estimate of vacuum,
as you get towards -0.1Mpa it is useless as a measure of absolute pressure,
they measure the difference between the pipe pressure and ambient atmospheric pressure, inaccurately,
but I consider it useful (especially for vacuum filtering) and good value.
I have not got a reliable low pressure (1 pa area) gauge yet ..... cost more than my pump !

Ramium - 8-6-2016 at 01:09

Can sulphur be used as a stabilizer for acrolein?
If not, are there any other effective stabilizers besides hydroquinone?

Glassware cleaning question

j_sum1 - 8-6-2016 at 14:39

I have managed to get some fine sulfur precipitate on the inside of a PE funnel. (Don't ask.) I would like a simple efficient way of removing it -- obviously there are parts inaccessible to mechanical scrubbing.
My first thought is to sluice a few times with warm xylene but there might be better options.

diddi - 8-6-2016 at 16:37

just hung it on a string in the flood waters j_sum :D

Orenousername - 8-6-2016 at 16:51

Quote: Originally posted by j_sum1  
I have managed to get some fine sulfur precipitate on the inside of a PE funnel. (Don't ask.) I would like a simple efficient way of removing it -- obviously there are parts inaccessible to mechanical scrubbing.
My first thought is to sluice a few times with warm xylene but there might be better options.


Hot H2SO4 will remove it but that might not be as attractive as xylenes. Also, the PE may melt. Xylenes is probably still your best bet.

Sulaiman - 8-6-2016 at 17:19

It depends upon what solvents that you have and their cost,
e.g. CS2 is a good solvent but expensive, DCM is a poor but cheap solvent,
I have read woelen recommend hot toluene on another forum.

j_sum1 - 8-6-2016 at 17:28

Yeah, I could use CS2. But the stuff is so volatile I am likely to have some evaporate and still leave a bit of residue behind.
Xylene I believe is better than toluene. But maybe I could react it with something?...

Sulaiman - 8-6-2016 at 18:09

Would Piranha solution attack the sulphur ?

j_sum1 - 8-6-2016 at 18:11

I guess it would but I am really not equipped for that.

gdflp - 8-6-2016 at 18:12

I'm not sure how available nitrates are in Aus currently, but one method which I find quite effective to remove organic tars(I've heard it works for sulfur too, haven't tried it) from condensers and flasks is to simply prepare some nitric acid with the contaminated glassware. The hot nitric acid oxidizes anything organic(or sulfur hopefully) to either gaseous(CO<sub>2</sub> and H<sub>2</sub> for organics) or liquid forms. With sulfur you would form either sulfur dioxide or sulfuric acid, neither one is a big deal since they're being used in the preparation anyway so the resulting acid wouldn't be significantly contaminated. Worst case, it needs to be redistilled, which isn't a bad idea anyway to remove NO<sub>x</sub> The result is clean glassware and clean nitric acid, a win-win in my book. This also works wonders for stained stir bars, one nitric acid prep and they're good as new. You could probably put the addition funnel in place of a column, and it would be subject to all of the acid vapors which would hopefully oxidize the sulfur.

j_sum1 - 8-6-2016 at 18:35

Thanks gdflp. That's a pretty cool idea. I think the ptfe stopcock will be up to that. I will have to give it a shot. It sounds a bit more appealing than sploshing around hot xylene.

j_sum1 - 8-6-2016 at 18:42

Quote: Originally posted by diddi  
just hung it on a string in the flood waters j_sum :D

Ha. Missed your comment diddi.
No flood waters up here. I am on the side of a hill 700m above sea level. Floods are a bit rare around here.

(Glad I don't have a multi-zillion mansion on the sand dunes overlooking the ocean though.)

gdflp - 8-6-2016 at 23:07

Quote: Originally posted by j_sum1  

(Glad I don't have a multi-zillion mansion on the sand dunes overlooking the ocean though.)

You mean you don't enjoy rebuilding your ridiculously expensive home every few years because it constantly gets hit with storms(on your dime I might add because you can't get insurance)? I thought everyone did:P

[Edited on 6-9-2016 by gdflp]

question on ethyl formate synth

Ramium - 12-6-2016 at 23:39

Today I started making a bit of ethyl formate via Fischer esterfication.
Apparently no catalyst is needed for this particular ester.

I refluxed 7ml of >90% formic acid and 6ml of ethanol for 11 hours.

The procedure I'm following says to reflux the mixture for 24 hours, but this seems like overkill to me. I can definitely smell the ethyl formate.

Will my yield be just as good if I stop now at 11 hours?




xfusion44 - 21-6-2016 at 14:42

Let's say we have a gas cylinder that can hold unlimited pressure and doesn't expand (not even a little bit). That cylinder is filled all the way to the top (no space left in it), with liquid gas (CO2 for example). Then we close the valve on the cylinder and start heating it. My question is: would the gas always stay in liquid state, no matter how hot the cylinder becomes? The heat, of course, wants to vaporise the gas, but there's no space and cylinder cannot expand, so how could the pressure build up? Also, if the cylinder couldn't hold unlimited pressure, would its failure still be as violent, if there would be no gas above the liquid phase gas?

Metacelsus - 21-6-2016 at 15:47

Quote: Originally posted by xfusion44  
would the gas always stay in liquid state, no matter how hot the cylinder becomes?


No. If the critical temperature is exceeded, it will be a supercritical fluid, not a liquid.

j_sum1 - 21-6-2016 at 19:50

Quote: Originally posted by gdflp  
I'm not sure how available nitrates are in Aus currently, but one method which I find quite effective to remove organic tars(I've heard it works for sulfur too, haven't tried it) from condensers and flasks is to simply prepare some nitric acid with the contaminated glassware. The hot nitric acid oxidizes anything organic(or sulfur hopefully) to either gaseous(CO<sub>2</sub> and H<sub>2</sub> for organics) or liquid forms. With sulfur you would form either sulfur dioxide or sulfuric acid, neither one is a big deal since they're being used in the preparation anyway so the resulting acid wouldn't be significantly contaminated. Worst case, it needs to be redistilled, which isn't a bad idea anyway to remove NO<sub>x</sub> The result is clean glassware and clean nitric acid, a win-win in my book. This also works wonders for stained stir bars, one nitric acid prep and they're good as new. You could probably put the addition funnel in place of a column, and it would be subject to all of the acid vapors which would hopefully oxidize the sulfur.

This is working beautifully. I am going to have to remember this trick for future grungy glassware.
Thanks.

Orenousername - 22-6-2016 at 01:00

Quote: Originally posted by Sulaiman  
Would Piranha solution attack the sulphur ?


In my experience, yes.

Orenousername - 22-6-2016 at 22:13

Can sodium nitrate be used in place of potassium permanganate to oxidize o-xylene to phthalic acid?

PHILOU Zrealone - 23-6-2016 at 04:35

Quote: Originally posted by Orenousername  
Can sodium nitrate be used in place of potassium permanganate to oxidize o-xylene to phthalic acid?

Not practically.
Maybe by adding HCl and H2O2 you would get a kind of Aqua Regia that may eventually oxydise (halogenate or nitrosate or both) the side chain depending on the Temperature...but reaction may also happen into the ring...

Afterwards a base treatment would hydrolysate the side chain nitroso or halide...

xfusion44 - 23-6-2016 at 11:10

Quote: Originally posted by Metacelsus  
Quote: Originally posted by xfusion44  
would the gas always stay in liquid state, no matter how hot the cylinder becomes?


No. If the critical temperature is exceeded, it will be a supercritical fluid, not a liquid.


Thanks!

The Volatile Chemist - 23-6-2016 at 11:39

What's the most basic, relatively safe way to make a black-powder-esque pulver from ammonium nitrate without sulfur, and how does one go about making a bang with it? I just need something small for independence day, to tear apart a small cheese block(speaking literally, sharp cheddar cheese)...
Definitely not a kewl, but one needs some fun once and a while. Thanks for any help.

ficolas - 23-6-2016 at 15:35

What is the composition of marble? I want to get calcium acetate and calcium chloride, and I thought about using marble, but I cant seem to find a place where it says what marble is composed of, since in some places it says its mainly just CaCO3 (in wich case a recrystalisation should be enough to purify?), and in other places I found it is CaCO3, CaO (no problem so far), SiO2 (wont react with the acid I think, so that leaves me with something to filter, no problem either), Na2CO3 (this is where it would get bad, however where I saw it it says it contains small amounts only, but I dont know how small) and other minor carbonates.

PHILOU Zrealone - 23-6-2016 at 17:06

Quote: Originally posted by ficolas  
What is the composition of marble? I want to get calcium acetate and calcium chloride, and I thought about using marble, but I cant seem to find a place where it says what marble is composed of, since in some places it says its mainly just CaCO3 (in wich case a recrystalisation should be enough to purify?), and in other places I found it is CaCO3, CaO (no problem so far), SiO2 (wont react with the acid I think, so that leaves me with something to filter, no problem either), Na2CO3 (this is where it would get bad, however where I saw it it says it contains small amounts only, but I dont know how small) and other minor carbonates.

Reduce the marble into fine powder and wash with hot demi-water then filtrate...the filtrate will contain the Na2CO3, if ever present!

For other minor carbonates...most are unsoluble and hard to separate...MgCO3 (white), CuCO3 (green), FeCO3 (?), Fe2(CO3)3 (orange-red)...

Colour is an indication of pollution by other cations...so I advice you to use white chalk stick used in schools to write on the black board...then at least you have CaCO3 with only traces of white unsoluble carbonates...

PHILOU Zrealone - 23-6-2016 at 17:13

Quote: Originally posted by The Volatile Chemist  
What's the most basic, relatively safe way to make a black-powder-esque pulver from ammonium nitrate without sulfur, and how does one go about making a bang with it? I just need something small for independence day, to tear apart a small cheese block(speaking literally, sharp cheddar cheese)...
Definitely not a kewl, but one needs some fun once and a while. Thanks for any help.

NH4NO3/C to make a banger...that will be hard!
Maybe add a little very fine Al powder, but your NH4NO3 needs to be very dry (it is hygroscopic) to get a kind of flash.

Use C black of fume, it is very light and ultrafine for a better homogeneity of the mix.

Then do an ignition test in the open.

Then do a confinement test into a cardboard roll, with a fuse and cardboard stopers...or simpler into the famous triangle banger with a fuse out of a corner (triangle made from a ribbon of paper).

[Edited on 24-6-2016 by PHILOU Zrealone]

Morgan - 25-6-2016 at 19:03

I was wondering what salt has the highest melting point? Just looking up a few I could think of, calcium fluoride was the highest at 1418 C/2584 F. Or is there a combination of two salts that might work together that have some high melting point?

DraconicAcid - 25-6-2016 at 20:05

Quote: Originally posted by Morgan  
I was wondering what salt has the highest melting point? Just looking up a few I could think of, calcium fluoride was the highest at 1418 C/2584 F. Or is there a combination of two salts that might work together that have some high melting point?

Mixtures of salts will have lower melting points.

Do you consider oxides to be salts?

Morgan - 26-6-2016 at 08:28

It was just a passing thought after seeing a clip of NaCl in water which melts at 1474F/801C and wondering what other salts would do.
https://www.youtube.com/watch?v=PDRWQUUUCF0#t=2m


Metacelsus - 26-6-2016 at 15:31

SrF2 melts a bit higher, at 1477 C. Oxides and nitrides have much higher melting points.

xfusion44 - 26-6-2016 at 18:51

What would be the product of combining iodine chloride and ethene? Would this reaction require any catalysts or special conditions?

Thanks in advance

The Volatile Chemist - 26-6-2016 at 19:27

Quote:
Quote: Originally posted by PHILOU Zrealone  
Quote: Originally posted by The Volatile Chemist  
What's the most basic, relatively safe way to make a black-powder-esque pulver from ammonium nitrate without sulfur, and how does one go about making a bang with it? I just need something small for independence day, to tear apart a small cheese block(speaking literally, sharp cheddar cheese)...
Definitely not a kewl, but one needs some fun once and a while. Thanks for any help.

NH4NO3/C to make a banger...that will be hard!
Maybe add a little very fine Al powder, but your NH4NO3 needs to be very dry (it is hygroscopic) to get a kind of flash.

Use C black of fume, it is very light and ultrafine for a better homogeneity of the mix.

Then do an ignition test in the open.

Then do a confinement test into a cardboard roll, with a fuse and cardboard stopers...or simpler into the famous triangle banger with a fuse out of a corner (triangle made from a ribbon of paper).

[Edited on 24-6-2016 by PHILOU Zrealone]

Thanks! I'll try that. I don't have any Al, but I'll see what I can do. Hopefully ground up prills of AN will do, if I use the stuff quickly. I wasted a lot of time today trying to get some AN/Sucrose crap to light, wouldn't work. Was terribly humid out though. Is it really that hard to get AN mixed with a random fuel to at least burn?

PHILOU Zrealone - 27-6-2016 at 07:12

Quote:
Quote: Originally posted by The Volatile Chemist  
Quote: Originally posted by PHILOU Zrealone  
Quote: Originally posted by The Volatile Chemist  
What's the most basic, relatively safe way to make a black-powder-esque pulver from ammonium nitrate without sulfur, and how does one go about making a bang with it? I just need something small for independence day, to tear apart a small cheese block(speaking literally, sharp cheddar cheese)...
Definitely not a kewl, but one needs some fun once and a while. Thanks for any help.

NH4NO3/C to make a banger...that will be hard!
Maybe add a little very fine Al powder, but your NH4NO3 needs to be very dry (it is hygroscopic) to get a kind of flash.

Use C black of fume, it is very light and ultrafine for a better homogeneity of the mix.

Then do an ignition test in the open.

Then do a confinement test into a cardboard roll, with a fuse and cardboard stopers...or simpler into the famous triangle banger with a fuse out of a corner (triangle made from a ribbon of paper).

[Edited on 24-6-2016 by PHILOU Zrealone]

Thanks! I'll try that. I don't have any Al, but I'll see what I can do. Hopefully ground up prills of AN will do, if I use the stuff quickly. I wasted a lot of time today trying to get some AN/Sucrose crap to light, wouldn't work. Was terribly humid out though. Is it really that hard to get AN mixed with a random fuel to at least burn?

Al powder you can do from Al foil set into a ball form and scratched at emeri/sand paper...hard Al tool works also.
Yes AN is very hard to ignite especially with succrose that releases water when heated (polylol thus a lot of HO groups).
Also with AN/fuel the more the inimate the mix, the better it will start to burn and sustain...correct weight ratio is of course needed --> stoechiometry

[Edited on 27-6-2016 by PHILOU Zrealone]

PHILOU Zrealone - 27-6-2016 at 07:19

Quote: Originally posted by xfusion44  
What would be the product of combining iodine chloride and ethene? Would this reaction require any catalysts or special conditions?

Thanks in advance

This is a beginner question for the beginning section...

Unlike conventional halogenation with Cl2 that may sometimes require a catalyst to help making Cl(+) (like for benzene or unactivated alcenes (electron withdrawing groups rich alcenes)
AlCl3 + Cl2 --> AlCl4(-) + Cl(+)

For ethene normally a simple mix with Cl2 will result in addition.
CH2=CH2 + Cl2 --> Cl-CH2-CH2-Cl
This can happen by homolytic clivage of Cl2 into 2 Cl° or in heterolytic clivage into Cl(-) and Cl(+)...

This goes even better with I-Cl since it is naturally splitting into I(+) and Cl(-)

The Volatile Chemist - 27-6-2016 at 09:50

Quote:
Quote: Originally posted by PHILOU Zrealone  
Quote: Originally posted by The Volatile Chemist  
Quote: Originally posted by PHILOU Zrealone  
Quote: Originally posted by The Volatile Chemist  
What's the most basic, relatively safe way to make a black-powder-esque pulver from ammonium nitrate without sulfur, and how does one go about making a bang with it? I just need something small for independence day, to tear apart a small cheese block(speaking literally, sharp cheddar cheese)...
Definitely not a kewl, but one needs some fun once and a while. Thanks for any help.

NH4NO3/C to make a banger...that will be hard!
Maybe add a little very fine Al powder, but your NH4NO3 needs to be very dry (it is hygroscopic) to get a kind of flash.

Use C black of fume, it is very light and ultrafine for a better homogeneity of the mix.

Then do an ignition test in the open.

Then do a confinement test into a cardboard roll, with a fuse and cardboard stopers...or simpler into the famous triangle banger with a fuse out of a corner (triangle made from a ribbon of paper).

[Edited on 24-6-2016 by PHILOU Zrealone]

Thanks! I'll try that. I don't have any Al, but I'll see what I can do. Hopefully ground up prills of AN will do, if I use the stuff quickly. I wasted a lot of time today trying to get some AN/Sucrose crap to light, wouldn't work. Was terribly humid out though. Is it really that hard to get AN mixed with a random fuel to at least burn?

Al powder you can do from Al foil set into a ball form and scratched at emeri/sand paper...hard Al tool works also.
Yes AN is very hard to ignite especially with succrose that releases water when heated (polylol thus a lot of HO groups).
Also with AN/fuel the more the inimate the mix, the better it will start to burn and sustain...correct weight ratio is of course needed --> stoechiometry

[Edited on 27-6-2016 by PHILOU Zrealone]

Sorry, I don't know how the quotes in this got messed up. But I'll try the Al/file idea, I have a heatsink I could use. Thanks again for your assistance; I always assumed that sucrose would be a good fuel, but I guess since glucose used to be called carbon hydrate, sugars in general make terrible fuels. So hydroscopic oxidizers and things that burn with lots of water and not much thermal output don't work too well together.

Antimony(III) and (V)

nezza - 27-6-2016 at 23:10

Hi. I have an acid solution containing antimony(III) and (V). Is there some way of reducing antimony(V) to (III) and no further or a way of separating the two species.

xfusion44 - 28-6-2016 at 06:47

Quote: Originally posted by PHILOU Zrealone  
Quote: Originally posted by xfusion44  
What would be the product of combining iodine chloride and ethene? Would this reaction require any catalysts or special conditions?

Thanks in advance

This is a beginner question for the beginning section...

Unlike conventional halogenation with Cl2 that may sometimes require a catalyst to help making Cl(+) (like for benzene or unactivated alcenes (electron withdrawing groups rich alcenes)
AlCl3 + Cl2 --> AlCl4(-) + Cl(+)

For ethene normally a simple mix with Cl2 will result in addition.
CH2=CH2 + Cl2 --> Cl-CH2-CH2-Cl
This can happen by homolytic clivage of Cl2 into 2 Cl° or in heterolytic clivage into Cl(-) and Cl(+)...

This goes even better with I-Cl since it is naturally splitting into I(+) and Cl(-)


Thanks! So, this would result in 1,1 dichloroethane or 1,2 chloroiodoethane?

PHILOU Zrealone - 28-6-2016 at 08:26

Quote: Originally posted by xfusion44  
Quote: Originally posted by PHILOU Zrealone  
Quote: Originally posted by xfusion44  
What would be the product of combining iodine chloride and ethene? Would this reaction require any catalysts or special conditions?

Thanks in advance

This is a beginner question for the beginning section...

Unlike conventional halogenation with Cl2 that may sometimes require a catalyst to help making Cl(+) (like for benzene or unactivated alcenes (electron withdrawing groups rich alcenes)
AlCl3 + Cl2 --> AlCl4(-) + Cl(+)

For ethene normally a simple mix with Cl2 will result in addition.
CH2=CH2 + Cl2 --> Cl-CH2-CH2-Cl
This can happen by homolytic clivage of Cl2 into 2 Cl° or in heterolytic clivage into Cl(-) and Cl(+)...

This goes even better with I-Cl since it is naturally splitting into I(+) and Cl(-)


Thanks! So, this would result in 1,1 dichloroethane or 1,2 chloroiodoethane?

Almost correct!
1,2-dichlor(o)ethane and 1,2-chloroiod(o)ethane.

DraconicAcid - 28-6-2016 at 09:33

Quote: Originally posted by PHILOU Zrealone  

Almost correct!
1,2-dichlor(o)ethane and 1,2-chloroiod(o)ethane.


1,2-dichloroethane and 1-chloro-2-iodoethane.

PHILOU Zrealone - 28-6-2016 at 14:38

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by PHILOU Zrealone  

Almost correct!
1,2-dichlor(o)ethane and 1,2-chloroiod(o)ethane.


1,2-dichloroethane and 1-chloro-2-iodoethane.

Yes, thank you for the correction.
That's what happens when you do two things at the same time :D

Velzee - 28-6-2016 at 15:23

I understand that it could be highly dangerous, but how can one produce and collect dioxins from plastic? Or is it too foolish to even attempt such a thing?

[Edited on 6/29/2016 by Velzee]

Metacelsus - 28-6-2016 at 16:51

Dioxins are formed when PVC undergoes incomplete combustion. You could probably make enough dioxins this way to cause toxicity, but you'd have very little hope of separating them from all the other pyrolysis products. I really shouldn't have to say this, but making dioxins is a bad idea. Don't do it.

Velzee - 28-6-2016 at 18:15

Quote: Originally posted by Metacelsus  
Dioxins are formed when PVC undergoes incomplete combustion. You could probably make enough dioxins this way to cause toxicity, but you'd have very little hope of separating them from all the other pyrolysis products. I really shouldn't have to say this, but making dioxins is a bad idea. Don't do it.


I don't want anyone thinking I'm plan to poison anyone, so let it be know that I'm researching various chemicals that cause toxicity and death in past occurrences, for preparation for my future forensic career, and for organic chemistry.

PHILOU Zrealone - 29-6-2016 at 02:32

Quote: Originally posted by Velzee  
Quote: Originally posted by Metacelsus  
Dioxins are formed when PVC undergoes incomplete combustion. You could probably make enough dioxins this way to cause toxicity, but you'd have very little hope of separating them from all the other pyrolysis products. I really shouldn't have to say this, but making dioxins is a bad idea. Don't do it.


I don't want anyone thinking I'm plan to poison anyone, so let it be know that I'm researching various chemicals that cause toxicity and death in past occurrences, for preparation for my future forensic career, and for organic chemistry.

You will probably get more succes at synthetizing it from good precursors...
I have thought about it a lot but not for polyhalodioxines, for polynitrodioxines (as potent dense explosive/energetic materials and polymers of those for propellant applications)...
--> ortho-halo-phenates, orthohalophenol, o-hydroquinone and heating.

The rest of the halogens may be present into the initial molecule (example pentachlorophenol -heat-> octachlorodioxine) or added afterwards by conventional halogenation techniques although orientation selectivity may be difficult and may require synthetic pathway prevision to introduce desired groups at specific places (-NO2 or -CO2H to turn into -NH2 by reduction or rearrangement and then replace by -X (X=I,Br or Cl) by diazotation)

enammonium-iminium rearrangement

careysub - 29-6-2016 at 06:47

As I proceed with my project of learning some organic chemistry terminology (does anyone ever learn more than "some"?) I find myself perplexed about the term "enammonium".

The term "iminium" is well described, and widely used; but I am unclear about the term "enammonium" which does not appear even once in March, 6th edition, for example.

Crowfjord - 29-6-2016 at 07:20

I don't think Ive ever come across the term "enammonium," but it is easy enough to assume that it refers to the positively charged version of an enamine, which would perhaps show up in the mechanism of enamine/imine tautomerism. Is this a term you came across and did not understand, or one which you think should exist, but have not seen?

woelen - 29-6-2016 at 07:24

An enamine is an organic amine, with the nitrogen atom directly bonded to a carbon atom which has a double bond to another carbon atom. An example of an enamine is the following:

CH2=CH-NH2

Another example is:

CH2=CH-N(CH3)2 (both hydrogens on the N substituted by methyl)


When an acid is added to an enamine, then an enammonium salt is produced. In my example:

CH2=CH-NH2 + H(+) --> CH2=CH-NH3(+)

----------------------------------

An imine is an organic compound with a C=N-group in it, e.g. H2C=N-H. Another example is H2C=N-CH3.

-----------------------------------

Enamines with a hydrogen atom connected to the nitrogen atom can rearrange into imines. This reaction is reversible in some cases. An example:

CH2=CH-NH2 <---> CH3-CH=NH

The hydrogen atom "wanders" from the nitrogen atom to the C-atom at the other end of the molecule and the double bond flips from C/C to C/N.

This is possible for any imine with an H atom on the N-atom. Another example:

C(CH3)2=CH-NH(CH3) <---> CH(CH3)2-CH=N-CH3

The acid salts of imines are called iminium salts.

careysub - 29-6-2016 at 08:05

Quote: Originally posted by Crowfjord  
I don't think Ive ever come across the term "enammonium," but it is easy enough to assume that it refers to the positively charged version of an enamine, which would perhaps show up in the mechanism of enamine/imine tautomerism. Is this a term you came across and did not understand, or one which you think should exist, but have not seen?


I came across the rearrangement of my post's title. But to understand what it is, I need to understand both parts of the rearrangement, but I only could find definitions for "iminium" (which I already understood). "Enammonium" is used (you can Google it), but not often, and is absent from any of the standard texts on organic chemistry I consulted (like March) and nowhere did I find anyone offering a definition. As a relatively neophyte student of the subject I am cautious about trying to assume things - it is too easy to assume incorrectly.

Woelen seems to clarify the issue, I will need to study it a bit more to make sure it is clear to me.

[Edited on 29-6-2016 by careysub]

PHILOU Zrealone - 30-6-2016 at 14:52

Like Woelen said!

The equilibrium en(e)amine <--> imine is known and widely explained in organic textbooks ; it is similar to the en(e)ol-keton equilibrium but with an amine instead of an alcohol group.
R-CH=CH-NH2 <==> R-CH2-CH=NH
R-CH=CH-OH <==> R-CH2-CH=O

"EN(E)" thus refers to the alcene related structure...

Since imine may form an iminium cation when exposed to a proton R-CH2-CH=NH2(+)
maybe they are proposing the "quadrature of the circle" by talking about en(e)aminium (or what should be the same en(e)ammonium) --> R-CH=CH-NH3(+)

A typical "en(e)-aminium" cation is "anilinium" or "benzenaminium" because the carbon holding the -NH2 (or -NH3(+)) (of anilin/aminobenzene) is alcenic by nature (sp2 conformation involved into the aromatic ring resonance which is against the iminium structure what would imply a rupture of aromaticity --> minor equilibrium form).

[Edited on 30-6-2016 by PHILOU Zrealone]

Volanschemia - 30-6-2016 at 19:12

Quote: Originally posted by gdflp  
I'm not sure how available nitrates are in Aus currently, but one method which I find quite effective to remove organic tars(I've heard it works for sulfur too, haven't tried it) from condensers and flasks is to simply prepare some nitric acid with the contaminated glassware. The hot nitric acid oxidizes anything organic(or sulfur hopefully) to either gaseous(CO<sub>2</sub> and H<sub>2</sub> for organics) or liquid forms. With sulfur you would form either sulfur dioxide or sulfuric acid, neither one is a big deal since they're being used in the preparation anyway so the resulting acid wouldn't be significantly contaminated. Worst case, it needs to be redistilled, which isn't a bad idea anyway to remove NO<sub>x</sub> The result is clean glassware and clean nitric acid, a win-win in my book. This also works wonders for stained stir bars, one nitric acid prep and they're good as new. You could probably put the addition funnel in place of a column, and it would be subject to all of the acid vapors which would hopefully oxidize the sulfur.


I just tried this out yesterday on an a round bottomed flask that got a splattering of gelled up/burned Phenyl Salicylate syrup.
I had previously tried out hot Xylenes, Acetone, Ethanol, HCl, the works. No dice.
The second the Nitric Acid vapour front hit the stuff though, it liquefied and dribbled down into the KNO3/HNO3/H2SO4 mixture, where it promptly disappeared in a puff of CO2.

In short, I can confirm that this cleaning method definitely works on the most stubborn of stains.

morsagh - 30-6-2016 at 22:12

Quote: Originally posted by nezza  
Hi. I have an acid solution containing antimony(III) and (V). Is there some way of reducing antimony(V) to (III) and no further or a way of separating the two species.

Ascorbic acid should work for reduction

(Edit: added context)

[Edited on 7-1-2016 by zts16]

Heating wire

Romix - 4-7-2016 at 15:57

Please recommend alloy to use as a resistive heating wire.

I tried fat Ni-Chrome wire, from socket. It melted on halfs.
In the place were touched wood.

If placed correctly, will it last?

[Edited on 5-7-2016 by Romix]

acetol from sugar

madscience - 5-7-2016 at 16:44

how would you synthesis acetol from sugar

PHILOU Zrealone - 6-7-2016 at 08:44

Quote: Originally posted by madscience  
how would you synthesis acetol from sugar

Depends what sugar...
If normal table sugar; then if by acetol, you mean hydroxyaceton (1-hydroxy-propanone) ...
you may get a chance to go there from:
1)sugar alcoholic fermentation --> ethanol -oxydation-> acetic acid
2)sugar acetic fermentation --> acetic acid
3)making aceton from acetic acid (Ca acetate or Ba acetate pyrolysis...CaCO3 or BaCO3 recoverable and recycled with acetic acid)
4) sugar fermentation direct to aceton
5) aceton monohalogenation (Cl2 or Br2 in HCl or HBr) beware of the lachrymatory product (a wargas)--> Under the Hood
6) hydrolysis in neutral media to avoid as much polycondensation into resinous polyphénols

Not chemistry for a beginner!
Much easier to start directly from aceton...

Texium - 10-7-2016 at 05:03

I have an emulsion of water and isobutyl bromide. What would be the best way to break it up? I was thinking of adding some salt or sulfuric acid, but I don't want to screw it up.

Side question: Does anyone know what isobutyl bromide smells like? I can't find any description of it, but my product smells butterscotch-like, kind of along the lines of MEK but heavier and sweeter.

Edit: Never mind the thing about the emulsion, it cleared itself up after resting for a few minutes.

[Edited on 7-10-2016 by zts16]

Refinery - 10-7-2016 at 10:51

Can 2% MEK in ethanol be polymerized or otherwise converted to a more stable form for ethanol purification with NaOH?

PHILOU Zrealone - 10-7-2016 at 14:49

Quote: Originally posted by Refinery  
Can 2% MEK in ethanol be polymerized or otherwise converted to a more stable form for ethanol purification with NaOH?

Might work!
Butanone (MEK) may be subject to aldolisation-crotonisation thus yielding higher molecular weight molecules, less volatile and remaining in the flask with the NaOH during the distillation of Ethanol.

Refinery - 10-7-2016 at 23:23

Only way to find it out is to test it, then.

Haloform Reaction Question.

dawsonsuen - 11-7-2016 at 07:37

Greetings,
I carried out a haloform reaction between a solution of 5% sodium hypochlorite and acetone earlier today and I only used a 5% excess of bleach. Unlike my previous run that has a very pale greenish yellow, the end solution of this run was completely colourless. Hence I was wondering if it would be useful if I mix all my chloroform into 1L of bleach just to ensure all acetone has reacted and to avoid having to deal with the annoying azeotrope between acetone and chloroform before distilling my chloroform.
Thanks in advance!

leuckart reaction and formate salts

madscience - 19-7-2016 at 02:15

Would the salt of formic acid work in the leuckart reaction like sodium formate and potassium formate, and would it reduce the imine if you react aldehyde or ketone with a amine with formate salt as the reducing agent and would still form the n formyl impurities the reason I ask im trying to make alanine

[Edited on 19-7-2016 by madscience]

CRUSTY - 19-7-2016 at 12:22

No. The use of ammonium formate is critical for the synthesis of an amine. Without the ammonium cation, it is impossible to form the hemiaminal which is reduced to the final amine product. By the way, could you elaborate on what you mean by "reduce impurities"? What sort of impurities are you referring to?

Test for guanidine nitrate

Ba(ClO3)2 - 19-7-2016 at 13:53

We made some (hopefully) guanidine nitrate by following the procedure in this thread http://www.sciencemadness.org/talk/viewthread.php?tid=12938

Things got a bit out of hand near the end, so we're not sure If we actually have guanidine nitrate. Here's what we've got:

IMG_3257[1].JPG - 2.3MB

Is there any test for guanidine nitrate that would distinguish it from ammonium nitrate and urea?

Question – Is it really a catalyst?

j_sum1 - 19-7-2016 at 16:22

I have been thinking about a couple of well-known reactions.

First, Fischer esterification in which the added H2SO4 is often stated to be a catalyst. I am questioning this. It seems to me that its role is to push the equilibrium to the right rather than to catalyse. Sulfuric acid provides for an acidic environment and also absorbs water: both of which favour the formation of products. Is there any catalysis occurring also?

Second, the lovely Chameleon reaction demonstration of the oxidation states of manganese using (typically) sucrose. NaOH is added, purportedly as a catalyst. But it seems a strange chemical for a catalyst in my limited experience. And I cannot really figure out what its role is. I'd love some enlightenment.

(I bet blogfast25 knows. But we have seen less of him lately.)

AvBaeyer - 19-7-2016 at 19:44

j-sum1:

Sulfuric acid is a catalyst in the Fischer esterification. It is chemically unchanged in the reaction. With some effort, you could isolate all of the H2SO4 that you started with at the end of the reaction which by definition makes it a catalyst. Mechanistically, the H2SO4 reacts with the carboxylic acid function to form a transient acylium ion which reacts with the alcohol to form the ester via a follow-on elimination of water. No water is "absorbed." Any good O-chem book should illustrate this mechanism.

AvB

j_sum1 - 19-7-2016 at 21:16

Quote: Originally posted by AvBaeyer  
j-sum1:

Sulfuric acid is a catalyst in the Fischer esterification. It is chemically unchanged in the reaction. With some effort, you could isolate all of the H2SO4 that you started with at the end of the reaction which by definition makes it a catalyst. Mechanistically, the H2SO4 reacts with the carboxylic acid function to form a transient acylium ion which reacts with the alcohol to form the ester via a follow-on elimination of water. No water is "absorbed." Any good O-chem book should illustrate this mechanism.

AvB

Thanks AvB.
I am yet to read a good O-Chem book. Long story.

So what is happening with the NaOH and the Mn reduction?

stoichiometric_steve - 20-7-2016 at 03:40

Can i dissolve Palladium powder in HCl (>30%) / H2O2 (~30%) to make PdCl2?

copper-glycerol complex

Ramium - 21-7-2016 at 00:44

I made this complex https://upload.wikimedia.org/wikipedia/commons/0/00/Glycerol...

I added some glycerol to a suspension of copper(ii) hydroxide in dilute sodium hydroxide solution and ended up with a beautiful solution of the complex.

My question is, how can I isolate the complex from this solution?

Will adding ethanol make the complex precipitate?

which functional group would aminate first

madscience - 21-7-2016 at 12:30

which of this functional groups would aminate first a benzoyl or an acetyl group if you do a reductive amination on these groups

The Volatile Chemist - 22-7-2016 at 12:29

Quote: Originally posted by dawsonsuen  
Greetings,
I carried out a haloform reaction between a solution of 5% sodium hypochlorite and acetone earlier today and I only used a 5% excess of bleach. Unlike my previous run that has a very pale greenish yellow, the end solution of this run was completely colourless. Hence I was wondering if it would be useful if I mix all my chloroform into 1L of bleach just to ensure all acetone has reacted and to avoid having to deal with the annoying azeotrope between acetone and chloroform before distilling my chloroform.
Thanks in advance!

From what I recall, the reaction runs various colors based on bleach quality/source. Watch chemplayer's synthesis of chlorobutanol to see if the reaction ran more like that preparation and the chlorobutanol sideproduct was formed due to excess acetone.

Distillation setup - 500ml vs 1L - can they be interchanged with vessel size?

RogueRose - 23-7-2016 at 17:10

I see 500ml and 1L setups advertised and price may be double for the 1L. Are general distillation (no columns - just boiling flask, riser from flask, condenser & receiving vessel) setups capable of just changing the vessel size but still using the same condenser and riser (assuming same joint size)?

CharlieA - 23-7-2016 at 17:18

Yes

MonoAmmonium Phosphate vs DiAmmonium Phosphate as yeast nutrient/energizer

RogueRose - 24-7-2016 at 20:02

I have seen a lot of yeast nutrients/energizers that contain DAP and I have also seen some that have ammonium sulfate in place of DAP. From what I have read the reason for putting these in with the yeast is to provide nitrogen to the yeast. My question is whether MAP can be used instead of DAP (with adjustments for the fact MAP has 1/2 the N as DAP). Also is there any reason that DAP is better than ammonium sulfate?

PHILOU Zrealone - 25-7-2016 at 03:59

Quote: Originally posted by j_sum1  

Second, the lovely Chameleon reaction demonstration of the oxidation states of manganese using (typically) sucrose. NaOH is added, purportedly as a catalyst. But it seems a strange chemical for a catalyst in my limited experience. And I cannot really figure out what its role is. I'd love some enlightenment.

It is an oxydoredox reaction as such the results depends on the pH, some reactions occurs into acidic media (H(+)) and others into basic media (OH(-))...see oxydoredox tables.

The reducer is the sucrose ...aldehyd turned into carboxylic and polyols turned into keto groups and further cleaved into oxalic acid and carbonic acid.

The oxydiser is the purple permanganate anion (MnO4(-)) turning into green manganate anion (MnO4(2-)) and then further into pale pink manganous cation (Mn(2+)).
So oxydation state of Mn goes from +7 to +6 to +2.

The H(+) or OH(-) are recovered from the water solvent.

PHILOU Zrealone - 25-7-2016 at 04:01

Quote: Originally posted by Ramium  
I made this complex https://upload.wikimedia.org/wikipedia/commons/0/00/Glycerol...

I added some glycerol to a suspension of copper(ii) hydroxide in dilute sodium hydroxide solution and ended up with a beautiful solution of the complex.

My question is, how can I isolate the complex from this solution?

Will adding ethanol make the complex precipitate?

Probably not with ethanol that will be a competitive complexator to 1,2,3-propantriol...maybe try adding some diethylether.

PHILOU Zrealone - 25-7-2016 at 04:03

Quote: Originally posted by RogueRose  
I have seen a lot of yeast nutrients/energizers that contain DAP and I have also seen some that have ammonium sulfate in place of DAP. From what I have read the reason for putting these in with the yeast is to provide nitrogen to the yeast. My question is whether MAP can be used instead of DAP (with adjustments for the fact MAP has 1/2 the N as DAP). Also is there any reason that DAP is better than ammonium sulfate?

Also pH addaptation because the trimming effect on pH will not be the same from MAP vs DAP...and of course the osmotic pressure...

DraconicAcid - 25-7-2016 at 07:42

Quote: Originally posted by PHILOU Zrealone  

Probably not with ethanol that will be a competitive complexator to 1,2,3-propantriol...maybe try adding some diethylether.

Ethanol will be monodentate, and unlikely to compete with the chelating gylerol.

palladium alternate

madscience - 26-7-2016 at 09:12

whats a good alternative to palladium acetate
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