Sciencemadness Discussion Board - The Short Questions Thread (4)

The Short Questions Thread (4)

Pages: 1 .. 4 5 6 7 8 .. 26

Brain&Force - 18-5-2014 at 21:43

People who join metal together are welders.

In the case of burning hydrocarbons, the blue part of the flame is an emission band. I've seen this myself with a diffraction grating. It's not blackbody radiation because you need temperatures of tens of thousands of degrees Kelvin. And there wouldn't be bands.

In the case of wood fires the bluest part of the flame is the hottest because it's more exposed to oxygen and therefore burns clean and hot. DraconicAcid's suggestion of copper perchlorate works because a large amount of oxidizer (solid perchlorate) is available and the flame burns clean, thus permitting coloration.

alexleyenda - 19-5-2014 at 09:16

You're right about the blue hydrocarbon flame, as soon as I read what you wrote it flashed in my head, i'm so stupid, it takes a temperature of around 6500 K to get a maximum blackbody radiations around blue wavelenghts while a propane flame's maximum is around 2000k... I'm so ashamed, I really contradicted myself :p Anyways, they say intelligent people are those who made every mistake, i'm on my way for that :p

About the fact that the hottest part is the bluest, then, I guess it is because the energy needed to excite the electrons to the level from which they emit blue (in the case of hydrocarbons) is high, so the flame needs to be really hot (burn well, be oxidized completely), if not it is only the blackbody radiation that is seen. Right?

[Edited on 19-5-2014 by alexleyenda]

arkoma - 21-5-2014 at 20:34

Can I dry acetone with NaOH? Happen to not have MgSO4 at the moment.

Mailinmypocket - 21-5-2014 at 20:54

No. It will cause an aldol condensation with acetone (http://www.chem.ucalgary.ca/courses/351/Carey5th/Ch18/ch18-3...) you could always use anhydrous copper sulfate though!

Haber - 22-5-2014 at 12:27

I recently made some 5-bromovanillin.
It was an in situ bromine method, which was something like this:
To a mixture of vanillin, potassium bromide, water and methanol there is added sulphuric acid and hydrogen peroxide, and stirred for 45 minutes, Temperature were kept at 18-20C under the procedure.
The mixture was added to cold water and filtered and then washed with cold water.

The crude product obtained had a a quite strong yellow color. mp: 162-164C
It was not recrystallized.

Now i few days later the product has changed color to something more green, some parts have even turned black.

My first thought was that it has oxidated to the carboxylic acid since it was not stored air-tight, but I have also read somewhere that these substituted aromatic aldehydes dont oxidize very easily. Not sure what to think.

I should also note that the sulphuric acid used was a dirty (purple color) drain cleaner.

Anyone have some thoughts on this?

Help would be greatly appreciated, thanks!

[Edited on 22-5-2014 by Haber]

Bezaleel - 24-5-2014 at 15:20

A while ago I put a spatula of microcrystalline nickelhydroxide in some 50ml water and added three spatulas of ammoniumthiocyanate to it. After some swirling, the solution became green and a faint smell of ammonia was noticed. I let it stand for a week, swirling the flask and heating gently now and then. Only a small portion of the nickelhydroxide dissolved.

I then decided to filter the undissolved hydroxide off and let the clear green solution evaporate on a petri dish. When it dried, a black substance with a hint of yellow/brown formed. This seems to be the normal colour for nickelthiocyanate and its dihydrate, according to atomistry.

The question I am left with is where the excess ammoniumthiosulfate has gone. Are there any double salts known?

[Edited on 24-5-2014 by Bezaleel]

solo - 24-5-2014 at 16:42

Quote: Originally posted by Cheddite Cheese

Quote: Originally posted by solo

Note- March 5th edition page 914 section 14-3

edition of what?

Also, what would prevent the alkene from being halogenated?

Can you be more specific about conditions, substrate, etc.?

March's Advanced Organic Chemistry: Reactions, Mechanisms, and Structure [Hardcover] 5th edition

"Also, what would prevent the alkene from being halogenated?" ......thats my question

.....any aldehyde with the coditions stated i.e. an aldehyde with no alpha carbon hydrogen , with a double bond adjacent to the alpha carbon....solo

TheAlchemistPirate - 24-5-2014 at 19:58

Has anyone here synthesized ruby? I have been looking around on the internet and am interested in making synthetic rubies, and have some questions. I know that synthetic ruby is made by reacting chromium oxide with aluminum oxide at 2000+ degrees Celsius, but I have seen on this forum where someone who did the reaction used a platinum-lined crucible to contain the reactants. I don't see why this would be necessary, since as far as I know the reagents aren't acidic and don't need platinum as a catalyst. If a platinum crucible is required I will probably forget the idea since platinum is so ridiculously expensive.
I also cant seem to find out how long the reactants need to be heated, I did see an article on the industrial process which apparently takes around 20 minutes...? The process involved using oxygen and hydrogen to heat the mixture and dripping the melted product onto a seed crystal, of course I cant do this and was thinking of another method which involves dipping a rod with a seed crystal into a crucible filled with the mixture. The crystals would then form on the rod.
I figured that I could use the latter method by making a furnace with firebrick and a graphite crucible, putting a graphite rod into the melted mixture and getting the rubies. I would very much hope this would be the case as it would be very awesome to be able to make your own gemstones.
Unfortunately there isn't very much talk about homemade gemstone synthesis on the internet and all I can find is industrial processes, I hope you guys can help.

Zyklon-A - 24-5-2014 at 20:05

I've thought about it, but never got around to it. How are you going to make the furnace? What fuel will you use.
I would use electricity as a heat source, rather than any sort of furnace.
Do you want to make them just to make them (basically to show off a homemade ruby), or do you have a practical application for it? For me it would certainly be the former hehe. Sapphire would be nice to... I'm guessing they contain copper?

TheAlchemistPirate - 24-5-2014 at 20:16

I would make a coal furnace and load it with charcoal, and use a graphite crucible as said above. I would also connect a steel pipe to the walls of the furnace and insert a vacuum cleaner on reverse, I think that would be enough. I would preferably use an electric furnace(obviously) but I don't have $700. I would initially make them to show off that I can make a homemade ruby, then I would experiment with making various shapes and sizes. I also think that it would be cast-able into various shapes and maybe even into lab equipment, imagine blast-resistant and acid-resistant labware for energetics syntheses?! I believe sapphire can be synthesized with nearly the same method as it also contains mostly corundum.

[Edited on 25-5-2014 by TheAlchemistPirate]

Zyklon-A - 24-5-2014 at 20:58

2000°C with a charcoal furnace? I highly doubt that's even possible. Even with pure oxygen it would take a huge furnace, not even sure if it would work at all.
Just make an electric furnace, wont cost even close to $700.

TheAlchemistPirate - 24-5-2014 at 21:32

I will probably look into making an electric furnace. I still haven't found anything involving using a platinum crucible, it might've been something completely different lol. Another thing I noticed is the essential requirement of very pure reagents, or the crystal will be opaque. I plan on spinning a graphite rod skimming the mixture and making large indestructible crystal clusters at first, then I will work on making uniform ones to possibly sell. I will post my results and possibly start a thread on this if I get it to work.

Haber - 25-5-2014 at 03:55

Quote: Originally posted by Haber

text...

All of the product have now turned into a black dusty powder. Vanillin smell almost completely gone.
Even if it was 5-bromovanillic acid it shouldn't be black.
I'm very curious what has happened. Nobody who have any guess?

Zyklon-A - 26-5-2014 at 16:29

When gasses that exist at equilibrium are produced, are they made already at equilibrium?
For instance, in this reaction: 2 NO + O₂ → 2 NO₂, is the resulting nitrogen dioxide already at equilibrium with dinitrogen tetroxide? Or does it have to reach it's equilibrium after it's formation.
Also, at lower temperatures, would this reaction happen: NO + O₂ → N₂O₄? Or does it make nitrogen dioxide, and then the slowly form dinitrogen tetroxide as it would at whatever the temperature is: 2 NO₂ ↔ N₂O₄.

[Edited on 27-5-2014 by Zyklonb]

The Volatile Chemist - 26-5-2014 at 18:05

Quote: Originally posted by TheAlchemistPirate

I will probably look into making an electric furnace. I still haven't found anything involving using a platinum crucible, it might've been something completely different lol. Another thing I noticed is the essential requirement of very pure reagents, or the crystal will be opaque. I plan on spinning a graphite rod skimming the mixture and making large indestructible crystal clusters at first, then I will work on making uniform ones to possibly sell. I will post my results and possibly start a thread on this if I get it to work.

The crucible material probably has to do with the corrosiveness of molten oxides. I'm not entirely sure, but the periodic table of videos' video on aluminum talks on some odd properties of aluminum oxide.

bismuthate - 29-5-2014 at 14:42

So I was just wondering: why is it OK for citizens to dump copper sulfate root killer down their toilets yet we can't dump copper sulfate waste from the lab? The same question goes with conc. sulphuric acid.

Hegi - 29-5-2014 at 22:11

Quote: Originally posted by bismuthate

So I was just wondering: why is it OK for citizens to dump copper sulfate root killer down their toilets yet we can't dump copper sulfate waste from the lab? The same question goes with conc. sulphuric acid.

You can´t? ´Cause we can ..

Also interested in this question.

arkoma - 30-5-2014 at 13:26

anyone know how to deposit copper on glass? Be a bad ass retro mirror.

DraconicAcid - 30-5-2014 at 13:38

Quote: Originally posted by arkoma

anyone know how to deposit copper on glass? Be a bad ass retro mirror.

"Chemical Curiosities" states...

30 mL of 1 mol/L copper(II) acetate solution is added to a well-cleaned 500 mL round bottom flask (without a ground glass joint), followed by 5 mL of 80% hydrazine hydrate solution. The blue solution immediately turns a brownish green. The copper mirror is formed when the flask is rotated slowly while being heated gently in the gas flame.

It does not suggest alternative reducing agents.

bismuthate - 30-5-2014 at 13:45

http://www.realscience.breckschool.org/upper/fruen/files/res... check out this paper. It talks a good bit about reducing agents for copper mirrors.

Brain&Force - 30-5-2014 at 14:57

I've read that ascorbic acid works well for mirror-making as well.

Zwitterion copper-plated his HTC HD2: http://www.sciencemadness.org/talk/viewthread.php?tid=14644&... (the image is missing)

alexleyenda - 31-5-2014 at 00:45

I read a lot about KClO3 cells and I have two questions:

1- Most people say the max voltage should be 5.5V, however some say 6V is fine too... Would I be fine with a 6V power supply or would i end up with perchlorate?

2- I have trouble understanding current density, I have not found clear explanations to it. What does the current density changes in term of efficiency in reactions? I know for some reactions it has to be really high, while sometime it has to be low.. why? /what is the best current density in this case (I think I saw it was 0.2A/cm²?)/ at what point does it starts to really damage the MMO electrodes/ How do I calculate the current density? (I guess Current/ area of the smaller electrode ?)

Thanks !

jock88 - 31-5-2014 at 04:11

http://oxidizing.typhoonguitars.com/

Current density at the anode is:
current into cell divided by the area of anode in use (in solution).

x amps per cm squared

Keep cathode area approx. (very approx) the same area to half the area of anode area.
Arranged in a sensible manner. It is good to have a cathode each side of a big flat anode.
Distance between anode and cathode not important as far as making product is concerned but you can vary the distance to suit the power supply voltage that you are STUCK with.

Stop obsessing over the voltage.

alexleyenda - 2-6-2014 at 23:03

A very random question I wanted to ask for a long time: When I cook instant noodle, as soon as I dump them in the boiling water it starts to boil 3 times stronger, and if I push the block of noodles to the bottom, it boils even 3 times stronger again. Why the heck does it do that? I guess more area for the bubbles to form like for boiling chips is part of the answer ?

jock88 - 3-6-2014 at 02:53

I think noodles are the new boiling chips.

MrHomeScientist - 3-6-2014 at 05:21

Yes, like boiling chips the noodles (like just about anything else) will increase the rate of bubbling by providing lots of nucleation sites. Try the same thing with salt instead of noodles and see if you notice the same effect.

alexleyenda - 3-6-2014 at 07:14

Alright, I know I noticed similar effect with other things, but the effect with these noodles is extreme, far stronger than almost anything else, probably because of their huge surface area.

Suggest an OTC-ish olefine!

Pumukli - 4-6-2014 at 09:19

Hello All,

Could someone suggest an OTC-ish olefine, which is:
- preferably has the R-CH=CH2 or R-CH=CH-R' structure,
- preferably has its boiling point above 60 Celsius,
- and is not an aromatic one? (I mean one without an aromatic ring.)

By "OTC-ish" I mean: one can get it as an OTC product or can prepare it with a very simple one-step synthesis from an OTC product.

I can think of eugenole and styrene but these are aromatic. Cyclohexene on the other hand is not OTC-ish afaik.

Good brain-storming!

[Edited on 4-6-2014 by Pumukli]

EdMeese - 4-6-2014 at 10:13

High-oleic vegetable oil is the first that comes to mind.

Pumukli - 4-6-2014 at 11:04

Ah, yeah, unsaturated oils are OTC and boil above 60 C it is sure.

But I'd prefer something more "well-defined", not an oily mixture. And definitely not something poly-olefinic.

But thanks anyway, your suggestion is appropriate for the original (lax) requirements.

EdMeese - 4-6-2014 at 13:06

Pinene then. Picky picky :-)

Dehydration of menthol should be pretty straighforward.
Maltol isn't "really" aromatic.
If nonconjugated dienes will do you've got a bunch of monoterpenes from flavors: limonene, gerionol, etc.
Allyl hexanoate is pineapple flavor.

Fumaric acid is in some supplement shops, maybe homebrewing.

So, get to it.

Pumukli - 6-6-2014 at 07:12

Search is over it seems.
Fumaric acid looks promising.
And is OTC here as well.
And the dehydrated menthol suggestion was also a good one!
Thanks!

alexleyenda - 6-6-2014 at 10:09

What would the greenish deposit on copper wire in a chlorate cell (not used as electrode in the liquid, only in the air) be? My guess is copper chloride, am I right?

PS: I know it is not a good idea to put copper in there, I was juste making a small scale test before I get my MMO electrodes for the real thing

[Edited on 6-6-2014 by alexleyenda]

Zyklon-A - 6-6-2014 at 13:26

alexleyenda , probably. It will be quite impure, most likely it will be a mixture of copper chloride, copper hydroxide hydrates and lots of other crap.
I once used a copper wire to suspend an MMO anode above the solution. It soon was covered in a green coating as well. Lots of oxygen, minute amounts of chlorine and maybe chlorine oxides are evolved. Also acidified chloride aerosols from bubbling, so it could be many things. The wire soon was so deteriorated, that it just snapped off.
Due to the high amperage I was using, the wire was over 100°C. Hot copper reacts pretty fast with oxygen and chlorine.

alexleyenda - 6-6-2014 at 14:40

Hmmm you are right about the hydroxide, I forgot about that one ! And yeah it did the same thing to me, the wire broke :p That's why I'm testing before I receive my real electrodes !

Mailinmypocket - 7-6-2014 at 06:45

When an experiment calls for k50 montmorillonite clay, is there anything for which it can be substituted for?

sparkgap - 7-6-2014 at 06:49

Depends. For one, you didn't quite elaborate on the nature of the "experiment" you speak of. For all we know, this could be something for which kitty litter might suffice.

sparky (~_~)

Mailinmypocket - 7-6-2014 at 07:37

Thanks, I should have elaborated on that, it's used as a catalyst in the synthesis of 7-hydroxy 4-methylcoumarin as per experiment PG-2 in the following document:

http://www.dst.gov.in/green-chem.pdf

Gooferking Science - 8-6-2014 at 06:58

I have an oil burner ignition transformer that puts out 10kv at 23 mA. Is there an easy way to limit that current down to just a few mA? Could a resistor be used to limit the current?

arkoma - 8-6-2014 at 08:40

^^yes. Use Ohm's law. Remember that only several milliamps is FATAL if they flow through your heart.

Metacelsus - 8-6-2014 at 14:16

If you use a resistor, you will also drop voltage (the amount depends on your load). Make sure that you use a resistor that can take the power (up to 230 W).

An inductive ballast is also an option. This has the advantage of dissipating less power.

alexleyenda - 10-6-2014 at 12:05

In a chlorate cell, the hydroxide solution you bubbles Cl2 into, does it needs to be replaced from time to time? I guess it does, I just found strange that it was not said in any thread I read about it.

Also, something strange happened in my chlorate cell. It's been running for around 4 hours, I stopped it, but the anode kept bubbling weakly. Any Idea why?

Another also :p I tested the pH of my cell with universal stripes, it did not work, it seemed like it was destroyed by the sample of water. I tested with phenol red, same thing, it became colorless. I guess the chlorine oxidate the dye and destroys it? Btw, I then tested it with bromothymol blue and finaly got a result : the pH was between 1 and 6.

[Edited on 11-6-2014 by alexleyenda]

[Edited on 11-6-2014 by alexleyenda]

plastics - 11-6-2014 at 03:24

How much sulphur trioxide would I need to dissolve in 100ml of 98% sulphuric acid to obtain 65% oleum (w/w)?

I assume I initially need to add enough to convert the outstanding 2% water to obtain 100% sulphuric acid?

arkoma - 14-6-2014 at 07:04

Is my hydrometer/specific gravity affected by altitude?

Metacelsus - 14-6-2014 at 07:30

Yes, but not by much. The density of air will be less, which will cause you to measure a slightly lower specific gravity (as some of your sample volume is air with a hydrometer).

Gravitational effects do not matter, as you are measuring relative specific gravity with a hydrometer.

Zyklon-A - 14-6-2014 at 07:48

Quote: Originally posted by plastics

How much sulphur trioxide would I need to dissolve in 100ml of 98% sulphuric acid to obtain 65% oleum (w/w)?

I assume I initially need to add enough to convert the outstanding 2% water to obtain 100% sulphuric acid?

You probably should not be making pyrosulfuric acid if you can't answer this question.

arkoma - 14-6-2014 at 08:13

Quote: Originally posted by Cheddite Cheese

Yes, but not by much. The density of air will be less, which will cause you to measure a slightly lower specific gravity (as some of your sample volume is air with a hydrometer).

Gravitational effects do not matter, as you are measuring relative specific gravity with a hydrometer.

That was what I thought, but wanted some confirmation. Thanx cheddite.

Edit--@4000 feet above sea level here, twice as high as "home", Twentynine Palms

[Edited on 6-14-2014 by arkoma]

plastics - 15-6-2014 at 12:25

Quote: Originally posted by Zyklonb

Quote: Originally posted by plastics

You probably should not be making pyrosulfuric acid if you can't answer this question.

OK smart arse. If you can't answer the question just keep quiet and stop filling the forum up with pointless shite.

Just for your delectation here is a picture of 120g SO3 sitting in a flask next to the tube furnace that generated it, waiting to be turned into oleum. Seeing as you think I am so stupid, perhaps I should just add 100ml warm water and see what happens?

So do you know the answer or not?

bismuthate - 15-6-2014 at 12:32

NO NO NO. Do not ever just add warm water (or any water at that) to that much SO3. That will send sulfuric acid everywhere and you won't get out unscathed.

elementcollector1 - 15-6-2014 at 12:43

*Intensive sigh*
bismuthate, he was being sarcastic.
plastics, first calculate stoichiometrically how much SO₃ is needed for that last 2% of water, then calculate how much additional SO₃ you would need to add so that 65% of the H₂SO₄ is converted to oleum.
Funny thing is, I'm not sure of the chemical formula of oleum, so I can't help you there.

alexleyenda - 15-6-2014 at 15:13

Be careful though, you should listen to this at 1:30 :p https://www.youtube.com/watch?v=u6MfZbCvPCw

TheAlchemistPirate - 16-6-2014 at 20:14

Hello I have some questions about the "synthesis?" of blasting gelatin, I am planning to blow up some stumps with dynamite sticks but I'm not 100% sure about the procedure. First I was wondering what they mean by "low colloid nitrocellulose" I figure this is partially-nitrated nitrocellulose meant for being more easily dissolved by solvents but I cant find anywhere showing the ratios for synthesizing it. Secondly I was wondering about the sensitivity, I have seen some sources saying blasting gelatin is very shock-sensitive but others saying it requires high-level blasting caps??? But yeah I hope you guys can help me out.

PHILOU Zrealone - 17-6-2014 at 09:36

Quote: Originally posted by TheAlchemistPirate

Has anyone here synthesized ruby? I have been looking around on the internet and am interested in making synthetic rubies, and have some questions. I know that synthetic ruby is made by reacting chromium oxide with aluminum oxide at 2000+ degrees Celsius, but I have seen on this forum where someone who did the reaction used a platinum-lined crucible to contain the reactants. I don't see why this would be necessary, since as far as I know the reagents aren't acidic and don't need platinum as a catalyst. If a platinum crucible is required I will probably forget the idea since platinum is so ridiculously expensive.
I also cant seem to find out how long the reactants need to be heated, I did see an article on the industrial process which apparently takes around 20 minutes...? The process involved using oxygen and hydrogen to heat the mixture and dripping the melted product onto a seed crystal, of course I cant do this and was thinking of another method which involves dipping a rod with a seed crystal into a crucible filled with the mixture. The crystals would then form on the rod.
I figured that I could use the latter method by making a furnace with firebrick and a graphite crucible, putting a graphite rod into the melted mixture and getting the rubies. I would very much hope this would be the case as it would be very awesome to be able to make your own gemstones.
Unfortunately there isn't very much talk about homemade gemstone synthesis on the internet and all I can find is industrial processes, I hope you guys can help.

I did with Al2O3 powder, Cr2O3 powder into a refractive cement/clay reactor and with hobby acetylen/aceton/N2O welding blowtorch...as explained in a topic initiated by you under general chemistry with "Subject: aluminium oxide crystal" : it is not very practical
CoO or FeO/TiO2 is used for saphire (blue corundum)

[Edited on 17-6-2014 by PHILOU Zrealone]

PHILOU Zrealone - 17-6-2014 at 09:43

Quote: Originally posted by alexleyenda

A very random question I wanted to ask for a long time: When I cook instant noodle, as soon as I dump them in the boiling water it starts to boil 3 times stronger, and if I push the block of noodles to the bottom, it boils even 3 times stronger again. Why the heck does it do that? I guess more area for the bubbles to form like for boiling chips is part of the answer ?

For the distillation pumice stones are used (those are kind of light lava stones with a lot of air bubbles (stone foam).
But one can also use sand.
The famous menthos into a sparkling soda proceeds by the same effect...but with addition of solubilisation of the menthos suggar and expelling the dissolved gas.

PHILOU Zrealone - 17-6-2014 at 09:45

Quote: Originally posted by Pumukli

Hello All,

Could someone suggest an OTC-ish olefine, which is:
- preferably has the R-CH=CH2 or R-CH=CH-R' structure,
- preferably has its boiling point above 60 Celsius,
- and is not an aromatic one? (I mean one without an aromatic ring.)

By "OTC-ish" I mean: one can get it as an OTC product or can prepare it with a very simple one-step synthesis from an OTC product.

I can think of eugenole and styrene but these are aromatic. Cyclohexene on the other hand is not OTC-ish afaik.

Good brain-storming!

[Edited on 4-6-2014 by Pumukli]

Lycopene from tomatoes.
Potassium sorbate.
Accrylic acid.

PHILOU Zrealone - 17-6-2014 at 09:50

Quote: Originally posted by alexleyenda

What would the greenish deposit on copper wire in a chlorate cell (not used as electrode in the liquid, only in the air) be? My guess is copper chloride, am I right?

PS: I know it is not a good idea to put copper in there, I was juste making a small scale test before I get my MMO electrodes for the real thing

[Edited on 6-6-2014 by alexleyenda]

If it is soluble an emerald green CuCl2; if it is unsoluble and like tuquoise green Cu(OH)Cl (copper oxychloride).

papaya - 17-6-2014 at 12:08

Sorbitol

Where I live I can freely obtain sorbitol sugar substitute (crystalline) for experiments, however a quick google search reveals, that sorbitol hydrates also exist besides anhydrous form
https://www.sciencedirect.com/science/article/pii/0040603188...
https://www.sciencedirect.com/science/article/pii/S000862159...

so how can I get sure if my stuff is anhydrous, or if even not - how to get the anhydrous form from it? I ask, because I'm nearly sure they wouldn't put anhydrous form regardless what they're writing on the box, isn't it? Can it be dried by melting and keeping the melt at high T for a while?

[Edited on 17-6-2014 by papaya]

[Edited on 17-6-2014 by papaya]

alexleyenda - 17-6-2014 at 13:17

Quote: Originally posted by PHILOU Zrealone

Quote: Originally posted by alexleyenda

If it is soluble an emerald green CuCl2; if it is unsoluble and like tuquoise green Cu(OH)Cl (copper oxychloride).

It might be copper oxychloride, however it's been a while now, I don't have it anymore I can't test its solubility. Thank you for the answer though.

papaya - 18-6-2014 at 04:40

So what about sorbitol?

alexleyenda - 18-6-2014 at 07:18

I tried to search about it and no one really seems to give any precise information about it :/

papaya - 18-6-2014 at 11:29

Well, I conducted an experiment, where I melted 5g of 'sorbitol' with unknown water content in a wide dish and let it stay at 130-150°C (at more than 250C it starts to fume - the polyol vaporizes too) for more than one hour. After cooling my 0.1g accuracy scales didn't detect any mass lose, so this means there's no water OR it can't be dehydrated just by heating. Better scales are not available, the problem is that if there's undetectable <0.1g loss, then it accounts as 2% water by weight, which is more than that in molar fractions. What do you think?

alexleyenda - 20-6-2014 at 13:56

Why is conc sulfuric acid slightly yellow most of the time? I have not found a credible source for an explanation. Some talk about sulfur, others about organic material contamination ?

Is there a way to remove the yellow? I know that when it is black you can oxidise it with peroxide, but it doesnt work on the yellow.

Edit: it became colorless when it cooled down, I guess this has to do with SOx gas then??

[Edited on 21-6-2014 by alexleyenda]

Hydrogen Chloride Fumes

Metacelsus - 22-6-2014 at 12:28

I was dehydrating manganese(ii) chloride, when it started releasing HCl fumes. I think the manganese chloride (from alkaline battery MnO₂) might have had traces of iron(iii) in it, but the fumes seemed too much to be solely from that.

Does manganese (ii) chloride hydrate release HCl upon heating (Wikipedia says no, but something must have been causing those fumes!).

Unknown purple spots

alexleyenda - 23-6-2014 at 21:20

I made salicylic acid from basic hydrolysis (NaOH) of methyl salicylate ans it worked well, but after the neutralisation with sulfuric acid I noticed the solution took a purple color and left a very strong purple color on the filter paper. Any Idea why and what it could be?

DraconicAcid - 23-6-2014 at 21:21

Quote: Originally posted by alexleyenda

I made salicylic acid from basic hydrolysis (NaOH) of methyl salicylate ans it worked well, but after the neutralisation with sulfuric acid I noticed the solution took a purple color and left a very strong purple color on the filter paper. Any Idea why and what it could be?

Ferric ion in the acid reacts with phenols (including salicylic acid) to give highly coloured complexes.

elementcollector1 - 23-6-2014 at 21:43

Quote: Originally posted by Cheddite Cheese

I was dehydrating manganese(ii) chloride, when it started releasing HCl fumes. I think the manganese chloride (from alkaline battery MnO₂) might have had traces of iron(iii) in it, but the fumes seemed too much to be solely from that.

Does manganese (ii) chloride hydrate release HCl upon heating (Wikipedia says no, but something must have been causing those fumes!).

Somewhat curiously, some MnO₂ I prepared once also gave off HCl (it had been prepared from battery sludge via hydrochloric acid). Related phenomenon?

alexleyenda - 23-6-2014 at 22:02

Quote: Originally posted by DraconicAcid

Quote: Originally posted by alexleyenda

Ferric ion in the acid reacts with phenols (including salicylic acid) to give highly coloured complexes.

So my acid would have ferric ions contamination? Is that common (because you seem to have directly targeted the acid as the source)?

[Edited on 24-6-2014 by alexleyenda]

DraconicAcid - 23-6-2014 at 23:33

Quote: Originally posted by alexleyenda

So my acid would have ferric ions contamination? Is that common (because you seem to have directly targeted the acid as the source)?

I don't know how common it is for sulphuric acid to be contaminated with iron, but hydrochloric commonly is. But purple with salicylic acid is definitely iron(III).

alexleyenda - 23-6-2014 at 23:48

Alright thanks for your very useful answer, as always. I'll make a few tests tomorrow to investigate the contamination.

alexleyenda - 25-6-2014 at 20:27

A question again, I seem to be over-using this thread but I've just always got hard to answer or random questions all the time^^

So, for some reasons i've got bismuth metal over silver metal, is there any way I can separate them to recover only the silver? From my researches I found that both almost only get dissolved in nitric acid and their nitrate salt's solubility is similar so... Is there any way I can dissolve the bismuth to keep only the silver? (or dissolve the silver to get rid of the bismuth. My final goal is to get pure silver nitrate)

[Edited on 26-6-2014 by alexleyenda]

bismuthate - 26-6-2014 at 08:23

Theoretically copper chloride should dissolve it slowly. But if I were you I would dissolve it all in nitric acid then displace the silver with copper.

alexleyenda - 26-6-2014 at 08:46

Right, bismuth's redox potential is right under copper...Then to get rid of the copper I could dissolve what's left in HCL+H2O2... Then dissolve the silver again in HNO3?

Brain&Force - 26-6-2014 at 09:00

Bismuth chloride is soluble in acids, silver chloride is insoluble. You could also try leaving silver in solution by means od a diammine complex.

alexleyenda - 26-6-2014 at 15:02

Right, so my idea to get rid of the copper with HCl + H2O2 (to get a solution of CuCl leaving the silver intact) should work, I'll try it.

Another question, something I could not explain arrived... I was preparing my electrolyte solution to top up my chlorate cell. I used the "waste" water from the recristallisation of my last batch to avoid wasting the ions and the dichromate left it in, so it contained very little K2Cr2O7 that gave it a pale yellow color, and the KCl I added. I left it outside for an hour and a half on a hot day under the sun because I had something else to do. When I came back, the solution had become colorless, it lost its slight chromium yellow color. Why ? I suspect electromagnetic rays have something to do with it but... ?

After that, I started heating it to dissolve the KCl to the temperature of my cell, but it didn't dissolve completely. I added some of the solution that was already in my cell (containing a 3g/L K2CR2O7 and what was left over of KCl and KClO3 from the last batch) to dissolve what was left of the KCl... The solution in the cell was a moderate yellow, the solution in the hot erlenmeyer was colorless... at the second they mixed, it became a very dark green ! Now for this one I would bet on Chromium(III) oxide, though I have no idea why it would have happened.

In brief: Why did it become colorless, am I right to think the green was caused by Chromium(III) oxide and if I am, how did Chromium(III) oxide form?

Additional Info: In the KCl used, there are these additives in small quantity: Potassium bitartrate, adipic acid silicon dioxide, mineral oil and fumaric acid.

bismuthate - 26-6-2014 at 15:30

Could the colourless stage somehow be potassium chromite? I have no idea how it could have formed but it would explain the clear colour.

alexleyenda - 26-6-2014 at 15:53

Quote: Originally posted by bismuthate

Could the colourless stage somehow be potassium chromite? I have no idea how it could have formed but it would explain the clear colour.

Well at first does potassium chromite even exist?? When I search for it I get no result.

bismuthate - 26-6-2014 at 16:50

Well sodium chromite exists so I'll bet that potassium chromite does too.

Zyklon-A - 26-6-2014 at 17:34

Does anyone know what reaction Dr. Steve Liddle is talking about in The Favorite Reactions video, by Periodic Table of videos? (1 minute, 45 seconds in)
He talks about reacting P₄ with Na-K alloy and then adding another reactant (of who's name he doesn't disclose).
He says its an extremely dangerous procedure (as evidenced by the description above.)

PHILOU Zrealone - 1-7-2014 at 09:48

Quote: Originally posted by Zyklon-A

Does anyone know what reaction Dr. Steve Liddle is talking about in The Favorite Reactions video, by Periodic Table of videos? (1 minute, 45 seconds in)
He talks about reacting P₄ with Na-K alloy and then adding another reactant (of who's name he doesn't disclose).
He says its an extremely dangerous procedure (as evidenced by the description above.)

Maybe the following?
3 Na + P --> Na3P(s)
3 K + P --> K3P(s)
Na3P(s) + 3 H2O(l) --> 3 NaOH(s) + PH3(g)
K3P(s) + 3 H2O(l) --> 3 KOH(s) + PH3(g)
2 PH3(g) +3 O2(g) --> P2O3(s) + 3 H2O(l) --> 2 H3PO3 (l)
phosphine

papaya - 6-7-2014 at 14:52

Can I use sodium carbonate as a standard for acid titrations ? The indicator used will be methyl orange, but I didn't hear that Na2CO3 is ever used as a standard, why? It has the advantage that it can be prepared completely free of water, it's fairly soluble, more stable to atmospheric conditions than NaOH etc.. I prepared it from sodium bicarbonate by heating the baking soda powder to 600C in the course of 1 hour, the mass change was from initial 19g of NaHCO3 to 11.9g, which is exactly what the theory predicts it should be when completely turned into carbonate. is there some serious drawback it can't be used for acid titrations ?

Metacelsus - 6-7-2014 at 14:58

No, it's fine (better than NaOH, IMO).

(Assuming you're titrating a reasonably strong acid, that is.) Carbonate isn't as strong a base as hydroxide.

[Edited on 6-7-2014 by Cheddite Cheese]

papaya - 6-7-2014 at 15:41

Yes cheddite, I need it for strong acids and if it works then some other related questions: which indicator is better in that case methyl orange(I think) or phenolph
thalein ? also I cannot figure out what concentrations I have to prepare both the titrant and titrand? Assuming I know very approximately that for example my sulfuric acid is 5M then I can initially take equal volume of 5mol/L Na2CO3 or 2x volume of 2.5mol/L or, or...
, or I could dilute both sides beforehand.. but why?
I just don't know which is more accurate and how to adjust concentrations for best results (how the precision is dependent on initial concentrations?)

Metacelsus - 6-7-2014 at 15:53

Use a relatively dilute solution of sodium carbonate (solubility 2.02 M at 25 C, your solution should be below 1 M) so that you can more accurately dispense a given amount and not have to deal with tiny volumes.

Phenolphthalein is a good indicator.

papaya - 6-7-2014 at 16:02

Thank you, I thought phenolphthalein is worse in this case because it's color transition occurs under basic pH 8.3–10.0 according to wiki and since carbonate is not a strong base this may be further from neutralization point compared with methylorange. But I agree, I also like it.

Metacelsus - 6-7-2014 at 18:49

I use it because it's easy(ish) to get. I assume the change occurs at pH 9 and adjust my calculations accordingly.

papaya - 8-7-2014 at 03:33

Hmm, interestingly I found here that phenolphtalein is not suitable for acid standardizations with carbonate, since it'll change color already on half neutralization
http://www.monzir-pal.net/Lab%20Manuals/Practical%20Quantita...

Other question: if I use methyl orange (full neutralization) which solution must be in the burete and which in the beaker ?

Metacelsus - 8-7-2014 at 06:56

If you account for the bicarbonate/carbonate equilibrium, it is possible to use phenolphthalein.

The acid solution with indicator should be in the beaker, and the carbonate solution should be in the burette.

papaya - 8-7-2014 at 07:04

Just read the page I provided
"however, in precence of phenolphthalein indicator, the end point of the reaction between Na2CO3 and HCl appears when the reaction proceeds to the point of NaHCO3 formation which requires half the volume of HCl consumed in the previous example.

Na2C03 + HCl = NaCl + NaHC03

The molarity of HCl in the previous example can be calculated by substitution of a 1:1 mole ratio of HCl to Na2CO3."

Thus, when titrating carbonate with acid (in burette) phenolphtalein color will fade when all carbonate is turned to bicarbonate - that's half way to neutralization.

papaya - 9-7-2014 at 04:57

I performed titration of most concentrated H2SO4 I could find( diluted by 1:50 ) with Na2CO3 solution,indicator is methyl orange( actually I titrated Na2CO3 with H2SO4, since the last was in the burette) which gave molarity of 18.3M (correspondents to 97% H2SO4). However I'm quite sure my sulfuric is not more than 90% conc - density is below 1.8 if I remember correctly from last measurement. This may indicate that carbonate is somewhat inconvenient for titrations, however what else can be used ?(I mean when no primary standards are commercially available) What about borax, is it better than carbonate?
I think determination of the concentrations of technical products available to amateur chemists is an important matter, however I don't see threads discussing this - what are most viable and accurate ways of estimation of acid strenghts, what standarts do You use?

arkoma - 9-7-2014 at 17:04

I salted out, then distilled isopropyl alcohol. Can I dry it further with lithium metal? Or will I be getting a visit from the fire department?

Brain&Force - 9-7-2014 at 17:30

Lithium reacts to form lithium isopropoxide in isopropanol, so you'll just be contaminating it more. Not explosively AFAIK, but don't risk it.

Metacelsus - 10-7-2014 at 06:55

You can dry it further by distilling a mixture of it and calcium oxide (quicklime). (I know this works for ethanol, and it should work for isopropanol).

AlphaDecay - 15-7-2014 at 16:04

I want to convert Calcium carbonate into a soluble calcium salt, but without the use of any acid. Is that possilble? Any suggestions?

Metacelsus - 15-7-2014 at 18:04

It's possible, but not worth it in my opinion.

1) Heat it strongly to decompose it to CaO.
2) Dissolve in water (Ca(OH)₂ is weakly soluble).
3) Precipitate with a water-soluble salt whose cation forms an insoluble hydroxide, and then filter. You'll end up with a solution of calcium ions and the anion you want.

AlphaDecay - 15-7-2014 at 18:33

Well, I do not have the required equipment to calcinate it. But I can try with the Calcium hydroxide instead. Thanks for suggestions.

Texium - 18-7-2014 at 06:40

Alright, I have a quick question. Is it ok to put a round bottom flask directly onto a hot plate if it isn't a coil type one? I've seen it done in many pictures on here before, but I have also read before that it's not a good thing to do.

gdflp - 18-7-2014 at 06:48

A rbf sitting directly on a hot plate will not be as effective as using a heating mantle or some other method of heating which has more surface area than the round bottom flask, but there is no reason you can't. To attain faster heat transfer an air bath, made by trapping air between the hot plate and the flask with a cloth or aluminum foil or such, or an oil or water bath will heat the flasks up much quicker and give you a better idea of what temperature the flask is at.

Texium - 18-7-2014 at 07:00

Alright cool, thanks, I'll use the air bath then.

arkoma - 18-7-2014 at 07:18

Quote: Originally posted by zts16

Alright, I have a quick question. Is it ok to put a round bottom flask directly onto a hot plate if it isn't a coil type one? I've seen it done in many pictures on here before, but I have also read before that it's not a good thing to do.

I put my chinese one directly on the hotplate, and the bottom fell out. deschem warrantied it, but I'll not try putting it directly on the element again

Texium - 18-7-2014 at 07:30

Is your hotplate of the type that has the direct metal coil? If so, that's probably why.
Also, what about a Florence Flask? That could be heated directly on a hotplate with ceramic wire gauze, right?

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