woelen

CH3COOH instead of CH3OOH....ah, now the reaction makes sense....

You'll get CH3COONa (sodium acetate), Fe2O3 and most likely also quite some Cl2 (so you MUST do this outside).

The reaction, however, is not clean at all. You'll get a really terrible mess of all kinds of compounds. Fe(3+) forms complexes with acetate, you'll get chlorine in the mix (acid + bleach gives chlorine). Don't expect this mix to be really useful.

If you really want nice Fe2O3, then I would suggest you to take iron wool and immerse this in water and add some bleach to that (without the acetic acid). If you leave that overnight, then most of the iron wool is oxidized.

Getting concentrated acid from vinegar can be done by distillation, but it is not that easy. Getiing pure acetic acid will be hard, because the boiling point of acetic acid (118 C) and water are quite close to each other. But by boiling away half the liquid of household cleaning vinegar you can double its concentration without too much trouble.

Another source for concentrated acetic acid may be photographers (acid stop bath, sold as 80% acetic acid or as 28% acetic acid). If you play with acetic acid, so not simply regard this as a 'stronger vinegar'. This stuff must be handled with respect and in fact is quite corrosive and has a very pungent vapor.

I'm sorry if this is obvious.. but can you displace ANY hydroxide with HCl? If so, I can finally do a test I have been wanting to do.

Only the basic ones and amphoteric ones (the hydroxides of metals in lower oxidation states). Higher oxidation state hydroxides may be a problem (e.g. SnO(OH)2, stannic acid) some are even impossible (e.g. SiO(OH)2, silicic acid).

Beware, however, that some basic hydroxides may be remarkably inert and although in theory they dissolve, in practice they only dissolve with very great difficulty.

NaOCl + CH3OOH + Fe + H2O -> KABOOM

Aluminum is a HELLUVA lot more reactive than lead is. As a result, the particles are so fine that they immediately react with oxygen and you form aluminum oxide instead of aluminum metal. Lead metal is non-reactive enough where the lead metal that's formed won't oxidize immediately.

Given this explanation, you could try in vacuum or in an inert atmosphere, where there is no oxygen.

I think it is even stronger. Even in an inert environment, I expect that aluminium oxide is formed. The tartrate ion contains quite some oxygen and because of the strong electropositive nature of aluminium I expect that the oxygen will bind to the aluminium, as soon as the tartrate decomposes. A similar difference can befound between HgC2O4 and MgC2O4 (mercury and magnesium oxalate). On heating, the Hg-salt gives CO2 and metallic Hg, the Mg-salt (and also the Ca-salt) gives CO, CO2 and MgO or CaO, even in an inert atmosphere.

If that works it would likely result in really, disturbingly fine powder.

No, it does not work. You'll end up with aluminium oxide.

but my teacher give that chemistry menu to me...

the solid ammonium iron(II) sulphate first dissolve in dil.sulphuric acid ' date=' and then water...

how come??[/quote']

The solubility cannot be the reason of this. I have this stuff and it dissolves in water quite well and easily.

 

The reason he wants you to dissolve it in dil. H2SO4 first may be that it is somewhat prone to aerial oxidation. The sensitivity depends on pH. The ferrous ion, Fe(2+) is sensitive to aerial oxidation, but only at high pH. The salt ferrous ammonium sulfate is somewhat acidic on its own, because of the ammonium, so it hardly is oxidized and I personally think it is not necessary to dissolve it in dilute acid. With plain ferrous sulfate, this effect is much stronger and yes, if you do not want much ferric ion in your solution, then you have to dissolve that in dilute acid.

You are wrong. It does dissolve in water quite well.

2NaNO3(l) -> 2Na(l) + N2(g) + 3O2(g)

 

For doing that I would advise you be very careful and be sure you have the right equiment of good qualty to do this. I also advise you to use NaNO3 instead of NaHCO3 because NaHCO3 can give you H2O and this can be very dangerous! Sodium reacts very violently with water' date=' it can be explosive!

 

When doing electrolysis of molten NaNO3 you must seperate the Sodium completely from the oxygen because it will oxydize and am pretty sure you want Na(s) and not Na2O(s). Be sure to get the MSDS of Sodium before you do this, so you know how to handle and store you Sodium so that no explosion occurs!!!!!(MSDS= Material Safety Data Sheets) [/quote']

Sorry, but what you are saying here is wrong. You cannot obtain sodium from molden sodium nitrate. Besides that, you will get nitrogen oxides at the anode, according to the following reactions:

 

NO3(-) ---> NO + O2 + e

2NO3(-) ---> NO2 + O2 + 2e

 

The Na, formed at the cathode, reacts immediately with the strongly oxidizing nitrate ions, giving sodium oxide and nitrogen, you may also get some sodium nitrite.

Keep in mind that nitrate ion is a strong oxidizer and hence any sodium formed will be destroyed immediately. Nitrate already reacts violently with many other metals powders (used in protechnics), so it certainly will react with the much more reactive sodium metal.

From what I can find, the HCl/H2O azeotrope happens at a 20.2% (by mass) HCl concentration and the solution boils at 108 degrees Celcius.

Nice to read this, I mixed up some numbers (the 20) .

I wonder if it would be possible to nitrate glucose?

As I mentioned before, you can do this, but be very careful. Before you know it, you have a runaway reaction (YT also mentioned it)!

In fact, there are explosives, based on sugars, e.g. nitromannitol. Chained structures like cellulose can also be nitrated. In cellulose, three OH-groups are present in all of the six-carbon units, (C6H10O5)n. During the nitration reaction, all of these hydroxyl-groups can be nitrated, resulting in trinitrocellulose, (C6H7O5(NO2)3)n.

If you really want to nitrate something with your nitric acid, then take some cotton wool (mostly cellulose, also a chain of sugar-like molecules), pluck this apart, and immerse the pieces in a mix of 1 part of conc. HNO3 and 1.5 parts of conc. H2SO4 for 30 minutes and occasional swirling. For 2 gram of cotton wool take appr. 20 ml of nitrating mix. After the reaction dump all of it at once in a large bucket of water (e.g. 10 liters of tap water) (do this at once, otherwise you get local hotspots!). Rinse the cotton wool with water, then rinse with dilute sodium bicarbonate, until there is no more bubbling, then again rinse with water to remove the bicarbonate and let dry by plucking it apart and putting it at a dry place. After a day or two, it will be dry and then the fun can start. Light a small piece of it. Do not confine the stuff to a small space.

With the nitrating reaction, there is a small chance of runaway, but with cellulose the risk is much lower than with sugar or other organics. If you see a slight reddening of the liquid or observe brown fumes during the reaction, IMMEDIATELY dump the whole lot in a bucket full of water, otherwise you may have a fire, an explosion or at least a huge cloud of NO2! These things are signs of onset of a runaway reaction (I know this from multiple personal experiences ).

so then you Do get your sugars that you need' date=' but it`s contaminated with the catalyst?

 

I should imagine in that case then that HCl would be the better of the 2 to use as the vapor pressure in a boiling medium will drive it off eventualy then?

whereas the citric acid will not.[/quote']

I think you are right, but I'm not quite sure what happens when the liquid is boiled to dryness. An azeotropic mix of H2O/HCl boils at 120C or something like that. With the sugars added in the boiling point may even be higher, it might be that they start decomposing. I would give it a try at test tube scale first.

Acid acts as a catalyst on the breakdown of alfa-d-glucopyranosyl-(1->2)-beta-d-fructofuranoside . So, your chlorine from the HCl does not go into an organic molecule.

The same is true for breakdown of even longer polysaccharides and also starch and the like. This is the reason, why the stomach contains some HCl. Combine with certain enzymes the breakdown in fact is quite fast at body temperature. Without the enzyme, indeed you might need to heat quite some time.

Ok guys... FYI.. I just bought some NaOH ... HAHA.. Ok Round 2..... ding ding ding...

You did a wise thing, my boy!

Yes, of course that should be possible. The problem here is more of a practical and instrumental nature than of a chemical nature. Indeed, getting the SO2 from the burning sulphur into the HNO3 may be a pain.

A practical setup may be the following:

Take a large glass pot or beaker (1 liter or something like that). Take some burning sulphur and put this in the pot. Cover this with a glass plate or screw cap. The sulphur burns for some time and then stops burning, due to lack of oxygen. Then, let cool down and take out the sulphur, which is connected to a small wire. Do this quickly, so that you don't loose your SO2. Then pour in your selenium solution and swirl a few times. SO2 dissolves in water quite well. This setup can easily be done without all kinds of glassware and apparatus. Even better would be if the air in the beaker is enriched with some oxygen, before the burning sulphur is put in.

no probs' date=' I have a few mins remaining, I`ll have a quick go at this now

 

edit: Aha yes, definately NO2 liberated! I can see the gas in the tube against a white paper, the liquid has a slight yellowing (nothing major though), and there`s a white precipitate along with one or 2 undissolved prills of NaNO3.

but Much more of reaction this time [/quote']

YT, that looks good. In fact, aqua regia is not deep yellow, the pale yellow color is quite good. The white stuff, which does not dissolve, is NaCl. Plain salt is next to insoluble in concentrated HCl. Your results will even be better if you crush your prills of NaNO3, before you add them to the HCl. Now, your prills become covered by a layer of insoluble NaCl.

 

Keep in mind though, that aqua regia is not stable on storage. It decomposes irreversibly, giving off NO2, ClNO and Cl2. So, make it just before use and assure that it is warm (not boiling hot!!!) when used. Do NOT store it for more than a day or so. If you store it in a tightly closed container, then that may rupture due to pressure buildup!

So YT's trying to make some aqua regia is he? You know, there are better ways to remove your wedding ring. Aqua regia is a bit extreme don't you think?

Extreme? Mwah....

No, I do not think it is that extreme. Of course you need to be careful and know what you are doing. I've been playing around with mixtures of HF, HNO3 and H2SO4 in my home lab and I'm still here to amuse you with my posts .

Right now I (on Usenet my name is "Wilco Oelen") am doing experiments with NaCN in aqua regia or concentrated HCl. See this thread on sci.chem:

http://groups.google.com/group/sci.chem/browse_frm/thread/df430f74f5f79588/a43514c184e6f3f8?lnk=raot#a43514c184e6f3f8

It is just a matter of knowing what you do and knowing of the risks. Look at the URL posted in reply #22. This shows some pictures of the evil stuff I made .

If I look at your observations, then it's the ammonium ion, which is spoiling the thing. I write this in a previous post, that this might be a problem. Well, if I look at your observations, it is.

Repeat the experiment by adding NaNO3 (I understood from one of your posts on making HNO3 that you have this) to concentrated HCl (30+% is required for this). If you do that, then I bet that you get a yellow liquid and some brown/orange gas above it. Heat carefully, but you definitely should not boil it. Within a few minutes of heating you should have a nice yellow liquid.

Beware of the gases. These are very noxious!

Your question is VERY unclear. In fact I have not the faintest idea what you want. Please be more specific. What yellow solid do you want to make? What procedure do you intend to use? Now you just mention a bunch of chems, but I can't smell what you are going to do with these .

Does the liquid turn yellow? Do you have a brown gas above the liquid? Or is the gas orange? Or just colorless?

Ok' date=' I still reserve the right to try it myself before I fully commit but I must admit right now its not looking good

 

`Scott[/quote']

You not only have the right to do so, even stronger, I invite you to do so .

...NaCO3, stuff like that at a pharmacy...

Which pharmacy??? This would be a most interesting compound with remarkable redox-properties .

...MgCl, ZnCl and much more...

Wow, you are a really smart boy if you can make these!

Please tell me how, I'm really interested .

Yet another problem is how to store them for more than a fraction of a second .

Keep in mind that KClO3 mixtures are much less stable than KNO3 mixtures. I can tell you from personal experience that a KClO3-mixture ignited spontaneously. Not fun at all!!! Also, KClO3/sulphur mixes may sensitize in due time. Due to slight aerial oxidation, the sulphur forms tiny amounts of acid, and this acid in turn forms unstable compounds, such as ClO2 with KClO3. This may lead to spontaneous ignition of compositions, when they are stored. Especially if you have so-called 'flowers of sulphur' this is a risk to be taken very seriously! KNO3 and KClO4 compositions do not suffer from this insidiously dangerous property. With tiny amounts of acid, these do not form compounds like ClO2, which may lead to spontaneous ignition. So, be careful with KClO3. In fact, KClO3 hardly is used anymore in serious fireworks, KClO4 is the preferred oxidizer or KNO3.

If you perform the experiment with ammonium dichromate only, then the products formed are not that toxic. However, with this reaction one must be careful as well, because of the fact that not all ammonium dichromate reacts. The green stuff (which is Cr2O3 and which is only slightly toxic on its own) usually is contaminated with remains of (NH4)2Cr2O7 and the latter is very toxic. Dichromate ion, Cr2O7(2-), is a carcinogen.

If mercury thiocyanate, Hg(SCN)2, is mixed in, then the reaction becomes much more dangerous. The serpents consist of a mercury compound of unknown and variable composition and no single formula can be given to describe that compound. Besides that, fumes of metallic mercury and volatile mercury compounds are released into the air. This makes the experiment absolutely unfit for demonstration in confined spaces without a good fume hood.

im the pyrotechnic artist

You should be careful with how you formulate things. I totally understand YT's feelings. Telling that you are the 'skilled chemist', or the 'pyrotechnic artist', while you ask for even the most basic recipe things is not the smartest thing to do. We don't like buzzing like this.

Let's put it in other words. From a 'pyrotechnic artist' I would expect beautiful recipes for nice compositions with colors and the like, hands-on experience and the ability to explain that to others, etc. etc. The 'pyrotechnic artist' knows his compositions and recipes and does not need to ask for the basics.

There is nothing wrong with asking for the basics. Even more, I encourage you to ask for help on this forum and people certainly are willing to help you, but from your side, be humble and don't buzz and accept that you still have a lot to learn and present yourself as such.

So, I write this, not to piss you off, but in the hope that you learn something....

Very nice rod. How did you get such a nicely plated layer? I have done experiments with plating in the past, but I always got blisters and bladdering metal. It did not adhere to the underlying surface. What is the metal underneath?

the SO2+H2O2 idea is definitely good, but i would imagine that it could get pretty scary, as peroxymonosulfuric acid could be produced if there is an excess of H2O2

Don't worry about that. Only at very high concentrations of acid (50+ % H2SO4) this may become a risk. That's why I told to start from 10% H2O2. This does not lead to H2SO5 and on the other hand gives a solution of reasonable concentration.