Sciencemadness Discussion Board - The Short Questions Thread (4)

The Short Questions Thread (4)

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DraconicAcid - 26-1-2014 at 13:41

Quote: Originally posted by HeYBrO

Just to clarify, would adding HCl to potassium hydrogen phthalate produce the free acid and KCl?

Yes.

Flash powder + sulfuric acid--> bang?

Zyklon-A - 26-1-2014 at 15:13

I made some KCLO₃ and Al flash powder, and put less than half a gram outside, then I added a drop of 98% H₂SO₄, I waited about 10 seconds, and then it popped. It was really loud, it reminded me of NI₃. The flash powder did not burn at all though, it was scattered over the ground. Is this a well known fact? Has anyone else done this before?

Edit: Maybe the ClO₂ that is formed is exploding?

[Edited on 26-1-2014 by Zyklonb]

Metacelsus - 26-1-2014 at 15:20

Chlorates and sulfuric acid explode, especially in the presence of fuels. (The chloric acid formed is VERY reactive.)

Zyklon-A - 26-1-2014 at 16:02

I just took about .25 grams of the same flash, added the drop of sulfuric acid, and proceeded to hit it with a hammer, before it popped by itself. The boom was massive! My ears are still ringing, it actually wasn't that loud for me, because I was so close, but my mom, in the house said it was really loud. I also felt the shock wave. I am quite impressed.

alexleyenda - 26-1-2014 at 17:33

I've got an Ethanol/Isopropyl alcohol mix, is it possible to isolate ethanol? Their BP is only 3-4 °C different so i'm not sure distillation would be effective. BTW, I only have a graham condenser.

DraconicAcid - 26-1-2014 at 17:58

Quote: Originally posted by alexleyenda

I've got an Ethanol/Isopropyl alcohol mix, is it possible to isolate ethanol? Their BP is only 3-4 °C different so i'm not sure distillation would be effective. BTW, I only have a graham condenser.

Even if they had a twenty degree difference in boiling point, these two liquids are likely to form an ideal solution, and will not separate very well with distillation.

alexleyenda - 26-1-2014 at 19:19

I see, they probably made it on purpose so it is not possible to distill it. I will use the solution as a solvant then, correct me if i'm wrong but it should show approximately the same properties as pure ethanol.

Question:

Zyklon-A - 27-1-2014 at 20:09

Why is strontium so cheap? strontium. All the other alkali metals/alkaline earth metals are much more expensive.(except Mg)
Lithium=$150/100g
Sodium=$89/140g
Potassium=$75/5g
Rubidium=$140/1g
Cesium=$140/1g

Beryllium=$15/1g
Magnesium=$3.50/25g
Calcium=$23/50g
Strontium=$85/5lbs
Barium=$140/50g
prices from Gallium Source.com
You have to pay $60 for shipping though

[Edited on 28-1-2014 by Zyklonb]

thebean - 27-1-2014 at 22:26

How can I tell when an organic compound has decomposed in solution? Today I was trying to recrystallize acetaminophen when the solution began to turn a yellowy/tan color. I've seen this when recrystallizing other organic compounds on occasion, so I began wondering if that was a sign that I was losing yield.

Question:

Zyklon-A - 29-1-2014 at 09:47

I have a sample of a metal wire, I don't know what it is.
It came with my little brothers chemistry set, and for some reason it wasn't labeled.
It's very soft and it felt like solder, but I think it's a pure element. At first I though it was lead or tin, but when I tried to melt it, I found out otherwise, I could melt it, but the spoon (I melted it in a spoon) was red hot- a lot hotter than necessary to melt lead or tin.
Any quantitative tests I can do?

Edit: Wow I'm dumb, it's zinc, I found two labels in a corner of the pack, one was for iron, the other zinc,I knew my sample Zn immediately, because, it wasn't magnetic, and I can't melt Iron at all.

[Edited on 30-1-2014 by Zyklonb]

thebean - 31-1-2014 at 09:42

I'm confused by my theoretical yield calculations. I found a piece of literature on the production of diphenylamine via a phosphoric acid catalyzed reaction between phenol and aniline. I decided that I would carry the reaction out but on a different scale. After my calculations the molar ratio between aniline and diphenylamine appeared to be 1:2. I was going to used .1 moles of aniline and thus if everything went perfectly, .2 moles of diphenylamine. Then I looked at the masses of the aniline and diphenylamine. 9.31g of aniline to yield 38.45g of diphenylamine. The ratio of aniline to phenol is 1:1 so I would be using 9.41g of phenol. This appears to violate the law of conservation of matter, so I was wondering if I was wrong and if possible where I may have gone wrong.

alexleyenda - 31-1-2014 at 10:58

Well yes, you are obviously wrong, around 20 g of matter can't be created out of nowhere, as you said, you violate the law of conservation of matter. Your aniline : diphenylamine 1:2 ratio makes no sense, you double your reagents out of nowhere. To make a simplified equation, (A being aniline, B phenol and AB the products)
you say : A + B ---> 2AB
while it should be : 2A + 2B --> 2AB
or in fact simply if we divide by 2 everywhere : A + B --> AB.
It is like if you say I will use one TV screen and make two TV out of it :p (yeah I know, the exemple sucks but you get the point). You did not give enough details about your calculations and did not give your reference, so it is not really possible to help you more than that, but you clearly made a major error somewhere. The only thing here that is near to a 1:2 ratio is molar mass of aniline : molar mass of diphenylamine, or the ratio amine group:benzene ring in the product. You may somehow have given the first one, while the second one is quite random :p

One possibility would be that they asked you to use a 2:1 ratio aniline : diphenylamine, in other words to use double the amount aniline so that the reaction goes faster and somehow you understood it the other way. It is hard to say without your source and your calculations details.

[Edited on 31-1-2014 by alexleyenda]

[Edited on 31-1-2014 by alexleyenda]

thebean - 31-1-2014 at 14:24

You're exactly right! It's a 2:1

Sanity check on sulfur analogs

Metacelsus - 31-1-2014 at 15:58

A carbonyl is more electrophilic than a thiocarbonyl, right?

Also, a thiolate is a better leaving group (less basic) than an alkoxide, right?

So, for nucleophilic acyl substitution, which would be a better substrate: an O-ethyl thioester or an S-ethyl thioester? Does the leaving group matter more than the carbonyl?

RC=S-OEt vs. RC=O-SEt

Does nucleophilic acyl substitution even make sense with thioesters?

[Edited on 1-2-2014 by Cheddite Cheese]

Topic for high school research project?

blargish - 2-2-2014 at 19:29

Hey guys,
I'm doing a chemistry research project at school and I'm having difficulty in finding a topic. My original idea was to synthesize hydrazine sulfate via the hofmann rearrangement and the hypochlorite - ketazine process, then compare purity with iodimetry. This got rejected because apparently we are not allowed to be in contact with any substances that may possibly be carcinogenic (hydrazine sulfate). I was wondering if anyone had any interesting ideas of an area I could explore; not too complicated, but not too simple either. Something that would require a nice experiment and give me good data to work with. I don't mean to be asking for spoon-feeding here, but any ideas would be appreciated.

alexleyenda - 4-2-2014 at 16:45

How can I clean bismuth/bismuth oxide from a stainless steel cauldron if it is even possible?

[Edited on 5-2-2014 by alexleyenda]

elementcollector1 - 4-2-2014 at 16:57

Quote: Originally posted by blargish

I was wondering if anyone had any interesting ideas of an area I could explore; not too complicated, but not too simple either. Something that would require a nice experiment and give me good data to work with. I don't mean to be asking for spoon-feeding here, but any ideas would be appreciated.

Quantitative electrolysis of silver nitrate to form metallic silver?
Titration of NaOH/ HCl?
Thermite, with % yields?
Spitballing here.

Brain&Force - 4-2-2014 at 18:52

Quote: Originally posted by blargish

Hey guys,
I'm doing a chemistry research project at school and I'm having difficulty in finding a topic. My original idea was to synthesize hydrazine sulfate via the hofmann rearrangement and the hypochlorite - ketazine process, then compare purity with iodimetry. This got rejected because apparently we are not allowed to be in contact with any substances that may possibly be carcinogenic (hydrazine sulfate). I was wondering if anyone had any interesting ideas of an area I could explore; not too complicated, but not too simple either. Something that would require a nice experiment and give me good data to work with. I don't mean to be asking for spoon-feeding here, but any ideas would be appreciated.

Tollen's reagent?
Coordination chemistry of Cr, Co, Ni, Cu?
Rare earth salts fluorescence/magnetism?
Effectiveness of transition metal catalysts in the decomposition of hydrogen peroxide?
pH-mediated fluorescence of pyridine?

Zyklon-A - 7-2-2014 at 07:06

I found this Liebig condenser for only $10.00, it is ground glass jointed and 350mm long. Does it look like it will be suitable for laboratory use?

Mailinmypocket - 7-2-2014 at 07:11

Wow, its actually the price for two of them! Doesn't mention the brand or anything though... Sometimes with super cheap glass I have found that the tapers do not make a good seal, then again...at that price there isn't much to lose!

Zyklon-A - 7-2-2014 at 09:38

Yeah, I saw it said ''Package Quantity 2'', but I wasn't sure what it meant. I think I might get it, that's the cheapest I've ever seen for a Liebig condenser.

[Edit] Ok, I just ordered it, and it should be here in ~5 days.

[Edited on 7-2-2014 by Zyklonb]

GoldGuy - 7-2-2014 at 13:49

Anybody know a methylamine synth (preferably from hexamine) in which NO di or trimethylamine is formed?

I saw a synth where the hexamine is reduced by iron/zinc/al or alltogether in the presence of HCl.... The synth didn't specify how much metal to use so basically my other question is:

Can hexamine be over reduced this type of synth?

[Edited on 7-2-2014 by GoldGuy]

jock88 - 10-2-2014 at 13:04

I wish to alloy some bismuth with Iron. But bismuth will not dissolve in Iron.
What could you recomment to add to the mix this is similar to bismuth (if it is even sensible to say that) so that I may get a few percent of bismuth to alloy with the Iron. I am thinking of Antimony?

forgottenpassword - 10-2-2014 at 14:48

Quote: Originally posted by GoldGuy

Anybody know a methylamine synth (preferably from hexamine) in which NO di or trimethylamine is formed?

I saw a synth where the hexamine is reduced by iron/zinc/al or alltogether in the presence of HCl.... The synth didn't specify how much metal to use so basically my other question is:

Can hexamine be over reduced this type of synth?

[Edited on 7-2-2014 by GoldGuy]

Reduction of nitromethane.

Boron Trioxide - 10-2-2014 at 16:44

I have a question about divided cell electrolysis, in most descriptions the cell has a semi-permeable membrane and salt bridge presumably to allow the flow of ions and therefore electric current.

However as a proper semi-permeable membrane is hard to find, would it not be possible to replace the membrane with a specifically chosen non-reactive metal to allow much greater current to flow, then adding a salt bridge for ion transfer? This would provide much extra current flow, which has always been a problem for me with only a salt-bridge.

So why would this not work?

Zyklon-A - 10-2-2014 at 16:46

It's not hard to make one.http://chem.wisc.edu/deptfiles/genchem/lab/labdocs/modules/e....

Metal will not work.
I don't know what you are trying to make, but I made this and it will likely work for whatever you are trying to make.

[Edited on 11-2-2014 by Zyklonb]

Boron Trioxide - 10-2-2014 at 17:28

Sodium hydroxide was one thing I was looking at making, but the main problem was with using soaked paper or similar device is that such a device only allows a small amount of current, but I will give the agar a try if I can find some.

Though won't the metal allow current just like the salt bridge, I know for sodium hydroxide you couldn't use aluminium but couldn't iron or zinc work?

But thanks for your links and help

Zyklon-A - 10-2-2014 at 17:42

Yes I ran into that exact problem, soaked paper only allows a small amount of current. It did work, but at about 3 amps it took about 2 days of constant running just to get to 14 Ph, almost nothing. I eventually gave up and just bought my NaOH.
I don't know why metal wont work, but I know it wont.

elementcollector1 - 10-2-2014 at 18:00

I experimented with alternative salt bridges such as gelatin and other such, to no effect. My advice would be to take a length of rubber tubing, and fill it with concentrated salt solution before stuffing it at both ends with cotton. This *might* allow for more current...

Oscilllator - 10-2-2014 at 20:53

When sodium thiosulfate reduces potassium permanganate in a basic solution, a green colour even stronger than the permanganate forms. Why is this?
I couldn't find any info on this online.

HeYBrO - 11-2-2014 at 01:26

Quote: Originally posted by Oscilllator

When sodium thiosulfate reduces potassium permanganate in a basic solution, a green colour even stronger than the permanganate forms. Why is this?

I would say it is because of the reduction to manganate, the same can be seen in the famous chemical chameleon reaction. equation for reduction :MnO4- + e- → MnO4-2
Other members can fill in any gaps as I'm not sure other wise for this one.

TheChemiKid - 11-2-2014 at 04:56

What is the easiest way to separate Ethylbenzene and Xylene?
I have hardware store "Xylene" which contains 70-90% mixed Xylenes and 10-30% Ethylbenzene.

testimento - 11-2-2014 at 10:55

Just a theory, but ethylbenzene melts at -95C where p-Xylene melts at 13C, so could it be crystallized out of the ethylbenzene?

Is this reaction real and balanced:

1 Ca(OH)2 + 1 CO(NH2)2 = 1 CaCO3 + 2 NH3

Theory: the calcium hydroxide should donate two hydrogens for ammonia and obtain the carbon monoxide and two oxygens to make calcium carbonate.

If not, are there any decent ways to make ammonia from urea except pyrolysis?

[Edited on 11-2-2014 by testimento]

TheChemiKid - 11-2-2014 at 12:54

A mix of xylenes melts at -47.4°C, though.

DraconicAcid - 11-2-2014 at 13:16

Quote: Originally posted by testimento

Just a theory, but ethylbenzene melts at -95C where p-Xylene melts at 13C, so could it be crystallized out of the ethylbenzene?

I really doubt it. You've got a four-component system there- ethylbenzene, o-xylene, m-xylene, and p-xylene. They are sufficiently similar that I'd use any two of them as a textbook example of an ideal solution. Even at temperatures well below 13 oC, p-xylene would be very soluble in ethylbenzene. If you cool it in dry ice, you may see some xylenes crystallizing out, and the ethylbenzene would stay in the liquid phase. It doesn't seem like a great separation method, though.

Why do you want to separate them?

vanillin

testimento - 11-2-2014 at 14:35

Could vanillin be synthetized how reasonably?

bfesser - 11-2-2014 at 14:40

Yes. Try Google. This [was] off-topic. Please don't ask for spoon-feeding.

[Edited on 11.2.14 by bfesser]

TheChemiKid - 11-2-2014 at 14:47

Can Phenylacetic Acid be synthesized using the same method as Benzoic Acid from Toluene or Phthalic Acid from o-Xylene (KMnO₄ and reflux). I expect it can, but I want to be sure before using my Ethylbenzene.

bfesser - 11-2-2014 at 15:35

No, KMnO<sub>4</sub> & reflux yields benzoic acid.

Attachment: Chem263_Oct12_notes_2010.pdf (475kB)
This file has been downloaded 1333 times

Crowfjord - 11-2-2014 at 15:36

Nope. The benzylic carbon would be oxidized, as it is much more reactive than the terminal methyl. Depending on how much oxidant is used (or how strong it is), one would expect to get phenylmethylcarbinol, acetophenone, or benzoic acid.

[Edit] Aw, darn it bfesser, you beat me to it! ;P

[Edited on 11-2-2014 by Crowfjord]

Zyklon-A - 12-2-2014 at 16:00

I search but couldn't find the solubility of cadmium sulfate in sulfuric acid. I plan on purifying some cadmium (from Ni-Cd batteries), by dissolving in H₂SO₄.
Does anyone know anything about it?
I plan on using an excess of sulfuric acid because it is contaminated with Cd(OH)2, but I don't know what the proportions are.
Cd + H₂SO₄ → CdSO₄ + H₂.
Cd(OH)2+ H₂SO₄→2H₂O+ CdSO₄.

[Edited on 13-2-2014 by Zyklonb]

[Edited on 13-2-2014 by Zyklonb]

DraconicAcid - 12-2-2014 at 16:34

If you don't use too much of an excess, you only have to worry about its solubility in water.

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: removed broken BBCode quote]

[Edited on 15.2.14 by bfesser]

Zyklon-A - 12-2-2014 at 17:53

I dissolved in slight excess sulfuric acid, soon the reaction got very hot, and turned green. I immediately knew there was nickel(II) sulfate forming as well. Is there a way to separate the two? Solubility of both sulfates are very close. Single displacement wont work except with cobalt, to displace nickel. I don't have cobalt.
Any other ways? I'm not trying to encourage spoonfeeding, I just don't know what to do.

[Edited on 13-2-2014 by Zyklonb]

DraconicAcid - 13-2-2014 at 21:16

Quote: Originally posted by Zyklonb

I dissolved in slight excess sulfuric acid, soon the reaction got very hot, and turned green. I immediately knew there was nickel(II) sulfate forming as well. Is there a way to separate the two? Solubility of both sulfates are very close. Single displacement wont work except with cobalt, to displace nickel. I don't have cobalt.
Any other ways? I'm not trying to encourage spoonfeeding, I just don't know what to do.

[Edited on 13-2-2014 by Zyklonb]

Cadmium carbonate has a much smaller Ksp than nickel carbonate. Maybe that would work.

testimento - 14-2-2014 at 16:38

Does anyone have data what is the electrical conductivity of water with 5-25% NaCl concentration per mm2?

bfesser - 14-2-2014 at 16:49

testimento (and others), please don't use this thread as a substitute for Google or the CRC Handbook. Also, what the heck do you mean by "concentration per mm2?"

Brain&Force - 14-2-2014 at 16:57

I think he means conduction per millimeter squared of solution (and that would probably millimeter cubed). Or would it just be millimeters (distance from one charge to the other)?

testimento - 14-2-2014 at 17:01

I presented it in mm2 because I was thinking of pipe, let's say 100mm2 cross sectional area, where brine would be, and through this should electricity be conducted. At least the current can be reported mm2 cross sectional area in metals.

And yes, I attempted to search for it for at least half an hour but managed to find no data which I could use as a reference, so I thought to ask.

[Edited on 15-2-2014 by testimento]

Omnipresent75 - 15-2-2014 at 13:24

I saw a demonstration online, where potassium persulfate was made by following procedure:

45,6 G of ammonium persulpfate dissolved in 100 ml of water

30 G of potassium chloride dissolved in 100 ml of water

Mix the solutions, cool it in fridge, filter crystals out and rinse them with water

Said to yield 54 G K2S2O8, and a leftover solution of Ammonium Chloride.

The said vid on youtube shows said reaction, but .....

So i scoured the WWW for more info on this proces, and found none. Which is why i hope someone here will enlighten me, as i´m out of options.

Is it a Hoax?

Thanks!

papaya - 15-2-2014 at 13:45

Hello friends, I want to know is there a way to synthesize some Al(NO3)3 starting from Al metal and HNO3 ? I want to have the final salt in as pure as possible solution(not contaminated with other ions, which will happen if I obtain it by some sort of double decomposition reaction), but as I know Al doesn't react with nitric acid. Ideas?

elementcollector1 - 15-2-2014 at 13:56

I know about the passivation from concentrated nitric acid, but what about dilute? Copper metal seems to follow much the same rules, and it reacts vigorously when in *dilute* nitric acid.

papaya - 15-2-2014 at 14:06

I'm afraid even dilute acid won't work.
https://en.wikipedia.org/wiki/Al%28NO3%293
I can't figure out any clean route myself. Maybe if Al(OH)3 is obtained first, it could be dissolved in HNO3? But hydroxide is a gel as far as I know and will entrap all kind of ions (if for example obtained by AlCl3 + NaOH route) inside.

Question

Etaoin Shrdlu - 15-2-2014 at 14:42

methyl anthranilate + undecanal --> Schiff base

Is there any simple chemical method to check for the presence of leftover reactants? Short of a GC analysis I'm stumped. (Yes, I know they'll be there.)

Brain&Force - 15-2-2014 at 16:09

papaya, dissolve it in HCl first, then add an excess of nitric acid to the aluminum chloride. The acidic solution of nitrate and chloride will convert to nitrosyl chloride, and eventually there will be no more chloride in solution, just nitrate. Do it in a fume hood for safety.
Or why not make aluminum carbonate?

alexleyenda - 15-2-2014 at 17:07

A little bit out of subject, but I have to tell it, it is too funny: The first time I read your message ( I am very tired tonight and english is not my first language so I make connections slowly, and I didn't read the previous posts before ), I thought you told someone to dissolve a papaya in HCl and add HNO3 to the produced aluminium chloride etc... and I was really strongly thinking what the hell, meditating in front of my screen, trying to understand XD I had never seen that someone is called papaya before. I laughed a lot when my brain decided to go back to work :p

papaya - 16-2-2014 at 00:13

Quote: Originally posted by Brain&Force

papaya, dissolve it in HCl first, then add an excess of nitric acid to the aluminum chloride. The acidic solution of nitrate and chloride will convert to nitrosyl chloride, and eventually there will be no more chloride in solution, just nitrate. Do it in a fume hood for safety.
Or why not make aluminum carbonate?

AlCl3 + nitric acid looks interesting, if you are absolutely sure that it will work. Do I need anhydrous AlCl3 and conc. HNO3 or aqueous solutions will do? Is it that easy to get rid of all chloride or it'll be still impure? Some details (or references) please.
Also I thought Al carbonate cannot be made from water solutions since it'll tend to hydrolyze , wiki is not clear about that
https://en.wikipedia.org/wiki/Aluminium_carbonate
but if it works (Na2CO3 + Al salt) then I would like to go this way.

Brain&Force - 16-2-2014 at 01:06

https://www.youtube.com/watch?v=u4Ha1SJrazY
Just keep the solution acidic.

Nickdul - 18-2-2014 at 11:03

Does sodium/potassium metabisulfite degrade over time (say, 2-3 years), releasing SO₂ and leaving behind the sulfite? I have a bank of K₂S₂O₅ that has gotten quite lumpy and hard to dissolve in water.
Also, S₂O₅ + H₂O -> 2HSO₃ in solutions. Does any other anion, other than the bisulfite, exist in aqueous solution? I'll be buying new metabisulfite, but I wonder if it's also worthwhile to get what is sold as bisulfite (how can it be a dry powder?).

[Edited on 18-2-2014 by Nickdul]

bismuthate - 18-2-2014 at 15:10

I tried the briggs-rauscher oscillating iodine reaction but I replaced the H2SO4 with a stoichiemetric amount of NaHSO4. The mixture did nothing why is this?

Question

Zyklon-A - 19-2-2014 at 06:54

Does chlorine attack rubber stoppers?
For my element collection I have decided to store some chlorine now. I'm putting it in a big test tube, so the color will be visible, but will it disintegrate the stopper over time?

confused - 19-2-2014 at 07:27

yes, it'll cause the rubber to become brittle and eventually disintegrate the stopper, you might want to ampoule it, i recommend practicing a few times before storing the sample

Zyklon-A - 19-2-2014 at 08:20

Well I can't ampule it, but I have a roll of Teflon tape, So I'll rap the stopper with a few layers of that.

thebean - 19-2-2014 at 08:35

Honestly for something like chlorine the only way I would say you can safely store and be completely comfortable is ampouling. You can make an ampoule using a glass pipette. I know NurdRage has a video on it but I can't link to it right now.

Zyklon-A - 19-2-2014 at 08:44

I know I've seen the video, I need all my glass pipets right now, getting an ampule isn't the problem either, sealing it is, my torch recently broke. :mad:

I think the Teflon will work, I have stored chlorine before for about 3 months in a vial, as mentioned here.

jock88 - 19-2-2014 at 12:11

at Nickdul

alexleyenda - 19-2-2014 at 13:20

Quote: Originally posted by bismuthate

I tried the briggs-rauscher oscillating iodine reaction but I replaced the H2SO4 with a stoichiemetric amount of NaHSO4. The mixture did nothing why is this?

This reaction includes many equilibrium reactions that must all take place for the reaction to succeed. Quote from wikipedia "IO3− + 2H2O2 + CH2(COOH)2 + H+ → ICH(COOH)2 + 2O2 + 3H2O"

NaHSO4 being a neutral salt, it does not provide the H3O+ ions needed for the step cited above.

[Edited on 19-2-2014 by alexleyenda]

DraconicAcid - 19-2-2014 at 13:55

Quote: Originally posted by alexleyenda

Quote: Originally posted by bismuthate

I tried the briggs-rauscher oscillating iodine reaction but I replaced the H2SO4 with a stoichiemetric amount of NaHSO4. The mixture did nothing why is this?

Incorrect. NaHSO4 is hardly a neutral salt, being more acidic (Ka = 1 x 10^-2) than phosphoric acid (K1 = 7 x 10^-3). However, because it's not a strong acid, the solution will not be as acidic as a solution of sulphuric acid, which means that reactions that have hydronium ion in their rate law will go at a much different rate, thus mucking up the expected oscillation.

alexleyenda - 19-2-2014 at 16:00

My bad, it is sodium sulfate that is a neutral salt, I confused both. At least I had part of the answer

By the way I salute your knowledge Draconic, you always have the righ precise and well explained answer to everything

elementcollector1 - 20-2-2014 at 21:24

Not really an information-seeking question, more an opinion-seeking one.
I have a 1 troy oz. bar of .999 silver. I want to cut off a bit and put it in an ampoule. However, doing so would make me unable to sell the bar as anything other than bulk silver (which I guess is not so bad, but collector's value can be a powerful thing). Should I use this bar as my source for silver, or should I obtain it from another source and use that?

Bar is Pan American, and was bought about 5-6 years ago.

Nickdul - 21-2-2014 at 06:19

My opinion would be to keep it as an investment. If it is for an element collection, 1g bars are going for a decent price and (again, my opinion) are very cute

The same goes for 1g gold bars.

Zyklon-A - 21-2-2014 at 07:24

Keep it, I regret selling all of my silver.

Filter nitric acid

Zyklon-A - 26-2-2014 at 20:20

I made some dilute nitric acid according to this reaction:
Ba(NO₃) + H₂SO₄ → 2HNO₃ + BaSO₄ ↓.
(26.1 g Ba(NO₃) + 9.9 g H₂SO₄ + 7.0 g H₂O → 23.3 g BaSO₄ ↓ + 12.6g HNO₃ + 7 g H₂O.)
12.6g HNO₃ + 7 g H₂O = 19.6 g 66.3% HNO₃ = .2 mol HNO₃.
So my question is, How can I filter off the barium sulfate?
I have disposable coffee filters, a synthetic filter, and a filter made out of copper mesh.
Obviously, the copper will react, but would a disposable coffee filter work? Or will it react, I know con. sulfuric acid reacts with paper. There is slight excess sulfuric acid, because I wanted to make sure that all the nitrate would react, and form insoluble sulfate. Nitric acid + sulfuric acid + cellulose → potential nitrocellulose? Would that happen with 66% nitric acid and ~2% sulfuric acid?

[Edited on 27-2-2014 by Zyklonb]

S.C. Wack - 26-2-2014 at 20:29

Sand. Glass wool.

DraconicAcid - 26-2-2014 at 20:37

Or let it settle, and remove it with a pipette.

Zyklon-A - 27-2-2014 at 06:40

Ok, thanks. I don't have any glass wool, non do I have pure sand. I'll use a pipette.

Mailinmypocket - 27-2-2014 at 07:28

Try fiberglass insulation maybe?

Zyklon-A - 27-2-2014 at 07:58

Yeah, I thought about that, fiberglass isn't made out of pure silicon dioxide, more like SiO₂, Al₂O₃ B₂O₃, MgO ect.

Nickdul - 27-2-2014 at 08:26

For all I know, pure cellulose filter paper will withstand nitric acid, even with sulfuric acid present. At the amount of sulfuric acid you specify, I doubt that a nitration reaction would occur. If the precipitate if very fine, it might clog the filter quickly and slow the filtration down to a crawl. A large funnel with a decent wad of cotton may suffice for a first filtration. Then you could filter the remainder with filter paper.

Zyklon-A - 27-2-2014 at 10:24

That's a good idea. Or I'll pour off the top liquid though filter paper, and avoid the precipitate altogether.

Töilet Plünger - 27-2-2014 at 18:20

I have household ammonia with surfactants. What exactly are the surfactants, and how will they affect complexation and precipitation reactions?

ZIGZIGLAR - 27-2-2014 at 19:25

Howdy - I want to obtain a sample of the air (pollutants - smoke to be specific) to forward on for analysis. What is the best way to capture and transport a sample? Cheers

mnick12 - 27-2-2014 at 20:03

What brand of ammonia do you have? You can just look up the msds to find out, but IIRC its probably some benzalkonoium chloride or other quaternary ammonium salt. As to how they might affect your reactions, its hard to say. They are phase transfer catalysts and could screw up a reaction, or do nothing. Best bet is to buy some lab grade ammonia or go to a hardware store and try to find unscented pure ammonia.

DraconicAcid - 27-2-2014 at 20:45

Quote: Originally posted by ZIGZIGLAR

Howdy - I want to obtain a sample of the air (pollutants - smoke to be specific) to forward on for analysis. What is the best way to capture and transport a sample? Cheers

Fill a bottle with water. Dump it out where you want to collect your sample of air. Seal the bottle.

ZIGZIGLAR - 27-2-2014 at 21:10

Quote: Originally posted by DraconicAcid

Quote: Originally posted by ZIGZIGLAR

Howdy - I want to obtain a sample of the air (pollutants - smoke to be specific) to forward on for analysis. What is the best way to capture and transport a sample? Cheers

Fill a bottle with water. Dump it out where you want to collect your sample of air. Seal the bottle.

So just empty the water out (upside down) and immediately seal once the water is evacuated?

DraconicAcid - 27-2-2014 at 21:18

Quote: Originally posted by ZIGZIGLAR

Quote: Originally posted by DraconicAcid

Quote: Originally posted by ZIGZIGLAR

Howdy - I want to obtain a sample of the air (pollutants - smoke to be specific) to forward on for analysis. What is the best way to capture and transport a sample? Cheers

Fill a bottle with water. Dump it out where you want to collect your sample of air. Seal the bottle.

So just empty the water out (upside down) and immediately seal once the water is evacuated?

Yep. As the water leaves, air rushes in to replace it. Seal it, and your sample is prevented from mixing with air in other locations.

Panache - 28-2-2014 at 05:12

What is the standard method for obtaining a consistent mixture of two pressurised gases, obviously both flows are regulated, then a mixing chamber of sorts, is then followed by a further regulator.
The question stems from my concern that back pressure in the mixing chamber affects both initial gas regulators and one has a system where if an adjustment is made on one regulator you end up spending ages tweaking the other.
My two gases are lpg( propane) and dried compressed air

Oscilllator - 28-2-2014 at 17:07

Quote: Originally posted by Panache

What is the standard method for obtaining a consistent mixture of two pressurised gases, obviously both flows are regulated, then a mixing chamber of sorts, is then followed by a further regulator.
The question stems from my concern that back pressure in the mixing chamber affects both initial gas regulators and one has a system where if an adjustment is made on one regulator you end up spending ages tweaking the other.
My two gases are lpg( propane) and dried compressed air

Looking up how an oxy acetylene works will answer your question, since thats exactly what your talking about. One thing that oxy-acetylenes dont have that you mentioned is a regulator after the mixing chamber. It goes directly from mixing chamber to nozzle.

Question:

Zyklon-A - 28-2-2014 at 18:11

In an alloy, is the melting point conserved?
For example, if you alloy 3 parts copper(MP=1084°C x 3 parts=3252,) 5 parts nickel (1455°C x5 parts = 7275,) and 1 part silver (= 961°C.) =11488 ÷ 9 parts = 1276.44°C. So would that be the melting point of that an alloy. I don't intend on making such an alloy, just wanted to know.
[EDIT, just realized, it would have to be relative kelvin scale right?]

[Edited on 1-3-2014 by Zyklonb]

elementcollector1 - 28-2-2014 at 18:47

Nope. Alloys, like eutectics, usually have lower MP's than their respective constituents, and I don't think it's an average.

Zyklon-A - 28-2-2014 at 18:53

Oh, that's interesting. So, is NaK an alloy then, I always thought it was a compound.

confused - 28-2-2014 at 19:05

yes, NaK is an alloy
http://en.wikipedia.org/wiki/NaK

Panache - 2-3-2014 at 14:31

Quote: Originally posted by Oscilllator

Quote: Originally posted by Panache

Thanks so much Osc, that's an excellent example for me to refer to!. Also in my post I had punctuated incorrectly and forgetten the '?' after '...is then followed by a further regulator.' So now I can forget about whether this was the issue (edit-just looked up the wiki page on oxy/act, and my problem has an answer, rather obvious I didn't have a needle valve after the lpg regulator, thnx again)

That said I however have a further question.
Question Prologue
I have long gone by the assumption that vessels and associated piping etc designed to withstand pressure of any reasonable degree (~100psi-ish) can withstand full vacuum (well like 75torr).
This assumption generally holds also for o-ring seals and flanged seals, but I tend to decide upon it on a seal by seal basis (some seals to perform require fish).

Actual Question
Is this assumption generally safe and valid?

(The reason I ask is I was curious as to what would happen if I place a full bottle of liquid CO2 into my ultra-low freezer at -85. Liquid CO2 has a density half that of the solid and I was concerned regarding the integrity of the cylinder. I have since decided there was no real reason for doing it but it made me realise I have never questioned this assumption)

[Edited on 2-3-2014 by Panache]

Panache - 2-3-2014 at 14:38

Quote: Originally posted by DraconicAcid

Quote: Originally posted by ZIGZIGLAR

Quote: Originally posted by DraconicAcid

Quote: Originally posted by ZIGZIGLAR

Howdy - I want to obtain a sample of the air (pollutants - smoke to be specific) to forward on for analysis. What is the best way to capture and transport a sample? Cheers

Fill a bottle with water. Dump it out where you want to collect your sample of air. Seal the bottle.

So just empty the water out (upside down) and immediately seal once the water is evacuated?

Yep. As the water leaves, air rushes in to replace it. Seal it, and your sample is prevented from mixing with air in other locations.

Won't the water rushing out scrub the smoke coming it. An empty large plastic syringe (vet supplies horse size one) I would have thought more appropriate), or pull a full vaccum on a 9kg (bbg gas size) cylinder and then open it if you require a larger sample). Whatever happens the walls of the vessel will absorb or adsorb some of the smoke. I think the standard method is to suck through very fine ashless or glass filter paper a known volume of the smoke filled air, trapping the particulates and then analysing those through various really tedious and careful analytical techniques.

Deleted video

Zyklon-A - 2-3-2014 at 16:48

Does anyone know what method was used in this video? Nurdrage made a video on isolating potassium from hydroxide and Mg. The video was deleted, from the comments, it seems like the reaction used a catalyst, and occurred, dissolved in some inert liquid. Sounds, a lot like Make Potassium (from versuchemie.de).

Brain&Force - 2-3-2014 at 17:18

The catalyst is a tertiary alcohol (he used amyl alcohol), which was dissolved in a high-boiling hydrocarbon (he used tetrahydronaphthalene). I believe it was the same method as Pok's.

Zyklon-A - 2-3-2014 at 17:31

Ok, does anyone know of a video pertaining to this method?

elementcollector1 - 2-3-2014 at 17:34

Unfortunately, his (NurdRage) was the only one I know of.
Odd that of all his videos, the potassium ones in particular were deleted. Too many k3wls?

Zyklon-A - 2-3-2014 at 17:41

Too bad I didn't get to see it, I was too damn late. :mad:

Töilet Plünger - 2-3-2014 at 19:05

NurdRage also took down his magnesium-silver nitrate flash powder video.

I miss him a lot.

Mesa - 3-3-2014 at 05:21

I have a bottle of 'gelcoat restorer' that is usually used on fibreglass boat's.

The only MSDS I could find online lists the following contents(but not concentrations):
Hydrochloric acid
polyethylene glycol mono-p-nonylphenyl ether
water

The side of the bottle says it contains 680g/L hydrochloric acid, and given the colour/viscosity, I'm assuming the polymer is a fairly low molecular weight.

I have 2 questions regarding this;

-If I were to distill, would the polymer potentially come across with the HCl/H2O?

-Given the much higher acid concentration/lower water content, is there any possibility of using this product as-is as an acid catalyst in simple esterfications?

Edit: at current concentration, the liquid doesn't fume at all and only has a faint odour.

[Edited on 3-3-2014 by Mesa]

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