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gdflp
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No bromine oxides are unstable at STP and they are not easily prepared. They cannot simply be prepared by the direct union of the two diatomic
elements, rather they can be prepared by the reaction of ozone with bromine at -50C, or by passing an arc through bromine and diatomic oxygen at low
temperatures and pressures. And on what basis is oxygen a stronger oxidant than chlorine?
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Upsilon
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Quote: Originally posted by gdflp  | | No bromine oxides are unstable at STP and they are not easily prepared. They cannot simply be prepared by the direct union of the two diatomic
elements, rather they can be prepared by the reaction of ozone with bromine at -50C, or by passing an arc through bromine and diatomic oxygen at low
temperatures and pressures. And on what basis is oxygen a stronger oxidant than chlorine? |
Ok, looks like I'll take the H2O2 route then.
Also I may be relating information incorrectly but I was going based on electronegativity and the fact that oxygen oxidizes chlorine and never the
other way around. Similar to how fluorine is a stronger oxidizer than oxygen, and thus oxidizes oxygen.
Chlorine's higher electron affinity may make it a better oxidant in general though, I just don't really know.
[Edited on 3-11-2015 by Upsilon]
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Upsilon
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Hold on - couldn't I just use concentrated sulfuric acid on the KBr? HBr is formed to be immediately oxidized by more sulfuric acid into bromine with
sulfur dioxide as a byproduct.
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j_sum1
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You could. But IME, yield is low without an oxidant. Even with oxidising iodide you don't get a clean product.
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Upsilon
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| Quote: Originally posted by j_sum1 | | You could. But IME, yield is low without an oxidant. Even with oxidising iodide you don't get a clean product. |
Maybe using nitric acid instead? Oxidation by H2O2 requires a proton source meaning using HBr as an intermediate. It would be nice to have it
generated and oxidized in situ.
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MrHomeScientist
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Despite the multitude of bromine preparation threads here, this topic keeps resurfacing. I might as well copy and paste a previous post I made:
| Quote: | Anyways every time this is brought up again, I feel compelled to share the electrolysis method since very few people seem to know about it. I found
the procedure on woelen's site, and made a 2-part video of it here: http://www.youtube.com/watch?v=NKjyM2AkZSY
It's very simple - uses only OTC pool chemicals (except potassium dichromate), and an electrolysis setup with graphite electrodes. The oxidizer,
sodium bromate, is created in-situ by the electrolysis. It also greatly cuts down on vapors. The downside is that it takes quite a while to
electrolyze - 26 hours in my case. Lots of fun though.
...
Here's woelen's webpage on the method that I based my video off of: http://woelen.homescience.net/science/chem/exps/OTC_bromine/...
I followed his procedure pretty much exactly, and got a very similar yield.
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Fascinating element. I love seeing preparations of it. A minimum of site searching will save you a lot of trouble!
[Edited on 11-5-2015 by MrHomeScientist]
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Upsilon
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Today I started up the manganese acetate experiment again, this time with GAA. It is bubbling much more rapidly this time around, no surprise there.
I'll let it run overnight and come back with results tomorrow.
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Agari
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Since the OP of this thread is collecting elements, and so am I, is there any way to obtain fluorine from a fluoride compound?
Edit:Bromine can be obtained by reacting HCl with TCCA and passing the resulting chlorine gas into a saturated solution of water and a bromine salt
(Available at pool stores),then attaching the reaction flask to a condenser for a simple distillation,cooling it using ice water. The receiving flask
should be placed in an ice bath to prevent bromine vapor release. Once water starts coming over, the distillation is over. Remove the receiving
flask,place the contents in a separatory funnel, and add 98% sulfuric acid to remove water. Cap,shake,and vent the sep funnel. After shaking,allow the
funnel to remain undisturbed to form 2 layers,the top being the layer of sulfuric acid and water,the bottom being mostly bromine,drain the bottom
layer into a suitable container, and now you have bromine!
The overall reaction is as follows when using TCCA tablets for chlorine generation.
C3N3O3Cl3 + 3HCl -> 3Cl2 + C3N3O3H3
Cl2 + 2NaBr= Br2 + 2NaCl
[Edited on 11-11-2015 by Agari]
Element Collection Status:
Elements Acquired: 21/91
Latest: Lead (Pb)
Quantity: 12g
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Upsilon
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| Quote: Originally posted by Agari | | Since the OP of this thread is collecting elements, and so am I, is there any way to obtain fluorine from a fluoride compound? |
No known chemical is capable of oxidizing fluoride to fluorine, so it must be done explicitly via electrolysis. Setting up the electrolysis is a whole
challenge in itself since it must be done with a molten fluoride (using an aqueous solution, the fluorine generated will react with water to form HF
which will contaminate your sample). Best compound you can use for this is potassium bifluoride, melting at only 238C.
[Edited on 11-11-2015 by Upsilon]
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j_sum1
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Short answer.
Nope.
There is one chemical method that involves antimony-fluoro-something-hazardous-and-unavailable.
For fluorine in an element collection, your options are pretty limited.
You can use a stand-in compound such as CaCl2 or teflon.
You can use a real fluorine sample in a nickel container. You won't be able to see it, but it will be there.
You can spend gazillions and get a small amount of visible gas in a specially treated single crystal quartz ampoule. It has been done but I doubt you
would ever find one available.
You could buy an element sample of fluorine diluted with helium in a glass ampoule. You won't be able to see it and you won't actually have any
unreacted after a few months but it will relieve you of some dollars.
You could get a hold of some antozonite which is a pretty rare form of radioactive fluorite containing minute quantities of elemental fluorine trapped
in the crystal imperfections. It is enough to make it smell when a piece is broken off.
Or, if you can find a method for producing it (electrolysis of a fluoride salt I guess), then you might be able to devise a way of storing it in a
teflon ampoule. Teflon is tricky though. Good luck.
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Upsilon
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| Quote: Originally posted by j_sum1 | | then you might be able to devise a way of storing it in a teflon ampoule. Teflon is tricky though. Good luck. |
I actually thought about different ways of storing it visibly. Something could be done with PFA, abusing its higher transparency compared to other
fluoropolymers, as well as its ability to be melt-formed. Melting some down and spraying it inside of a regular glass ampoule to have a thin layer of
(transparent) PFA protection may work to hold the gas. If you do end up attempting storage of fluorine, no matter what vessel you use, keep it under a
barium hydroxide solution. If any fluorine gas escapes, it will be instantly neutralized and will leave insoluble BaF2 behind that will act as an
indicator. If you see insoluble material at the bottom of the solution, you know you've got a leak.
[Edited on 11-11-2015 by Upsilon]
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Agari
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| Quote: Originally posted by j_sum1 | Short answer.
Nope.
There is one chemical method that involves antimony-fluoro-something-hazardous-and-unavailable.
For fluorine in an element collection, your options are pretty limited.
You can use a stand-in compound such as CaCl2 or teflon.
You can use a real fluorine sample in a nickel container. You won't be able to see it, but it will be there.
You can spend gazillions and get a small amount of visible gas in a specially treated single crystal quartz ampoule. It has been done but I doubt you
would ever find one available.
You could buy an element sample of fluorine diluted with helium in a glass ampoule. You won't be able to see it and you won't actually have any
unreacted after a few months but it will relieve you of some dollars.
You could get a hold of some antozonite which is a pretty rare form of radioactive fluorite containing minute quantities of elemental fluorine trapped
in the crystal imperfections. It is enough to make it smell when a piece is broken off.
Or, if you can find a method for producing it (electrolysis of a fluoride salt I guess), then you might be able to devise a way of storing it in a
teflon ampoule. Teflon is tricky though. Good luck. |
I think that by "Antimony-fluoro-something" you mean "Fluoroantimonic Acid",it is a superacid, meaning it is more powerful than 100% sulfuric acid.
Source:Wikipedia.
[Edited on 11-11-2015 by Agari]
[Edited on 11-11-2015 by Agari]
Element Collection Status:
Elements Acquired: 21/91
Latest: Lead (Pb)
Quantity: 12g
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MolecularWorld
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It's antimony pentafluoride.
https://en.wikipedia.org/wiki/Fluorine#Chemical
| Quote: | While preparing for a 1986 conference to celebrate the centennial of Moissan's achievement, Karl O. Christe reasoned that chemical fluorine generation
should be feasible since some metal fluoride anions have no stable neutral counterparts; their acidification potentially triggers oxidation instead.
He devised a method which evolves fluorine at high yield and atmospheric pressure:
2 KMnO4 + 2 KF + 10 HF + 3 H2O2 → 2 K2MnF6 + 8 H2O + 3 O2↑
2 K2MnF6 + 4 SbF5 → 4 KSbF6 + 2 MnF3 + F2↑
Christe later commented that the reactants "had been known for more than 100 years and even Moissan could have come up with this scheme." As late as
2008, some references still asserted that fluorine was too reactive for any chemical isolation.
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j_sum1
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That's the one. Actually, it is the 10HF that is scary there.
Not a synthesis for this little duck.
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Upsilon
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There's honestly no point in doing this anyway, electrolysis is by far the easier option to F2 gas.
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Agari
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We are getting a bit off-topic here. Anyways,I might not even get to obtaining fluorine gas and instead aim for another element in the meantime.
Element Collection Status:
Elements Acquired: 21/91
Latest: Lead (Pb)
Quantity: 12g
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MolecularWorld
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Easier still to buy it. But easy rarely factors into why and how amateur chemists do things.
The first reaction isn't that scary, it proceeds in solution, right? So you could use extremely dilute hydrofluoric acid.
But according to the relevant Wikipedia page, preparation of antimony pentafluoride requires either anhydrous hydrogen fluoride or fluorine gas, and
high temperatures.
Much scarier.
Though I agree that electrolysis is more practical, if one really must generate fluorine in quantity.
Personally, I'm satisfied with a nice fluorite to represent fluorine in my element collection.
| Quote: Originally posted by Agari | | Anyways, I might not even get to obtaining fluorine gas and instead aim for another element in the meantime. |
That is most wise.
Sorry, I thought the question was:
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Agari
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I think we have settled the matter of obtaining Fluorine gas, MolecularWorld, and should move on to whatever the OP is posting or other elements.
Element Collection Status:
Elements Acquired: 21/91
Latest: Lead (Pb)
Quantity: 12g
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MolecularWorld
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@Agari:
And you're entitled to think that.
Just as I'm entitled to post my thoughts and ideas when I think they'll add something to the discussion.
I suppose I'm a bit of a gadfly. It speaks to my subconscious that I registered on mischief night (something I just realized).
I can't see any posts in this thread on elements other than fluorine since you asked about it, so I can't quite understand how we
"should move on to whatever the OP is posting or other elements".
Whatever is the OP posting?
What other elements did you have in mind?
I will require this information in order to generate posts that meet your exacting standards.
[Edited on 11-11-2015 by MolecularWorld]
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woelen
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Both of Christe's reactions for making F2 without electrolysis require strictly anhydrous conditions. The first reaction hence requires anhydrous HF,
anhydrous H2O2 and a very strong drying agent to absorb the water, formed in the reaction.
The second reaction in turn also must be in strictly anhydrous conditions. Any water present will hydrolyse the SbF5 and will destroy the formed F2.
It is not without reason why it took so long (until 1986) before this reaction was carried out. Handling a combination of anhydrous HF, absolutely
dry/pure H2O2 and SbF5 only can be done in a very well equipped lab and I am quite sure that nearly every lab in the world is not capable of handling
these things safely.
I am not sure about buying F2-samples. In the past it seems that there have been fake samples with dried Cl2 and NO2 and dried air to make the color
less intense which were mixed into an ordinary glass ampoule to get something which is supposed to look like F2. I am not sure about this, but I can
certainly imagine this, because F2-samples are sold for prices of EUR 100 or so, while the gasses and ampoules needed for making the fakes only cost a
fraction of that. I once read on Theodore's website how a real F2 sample can be made, using special quartz ampoules which are dried meticulously. The
entire process, described on his site, must cost hundreds of euros/dollars for a single ampoule!
Conclusion: Samples of elemental F2, visible as colored gas, are beyond reach of home chemists, unless you are able/willing to spend several hundreds
of dollars on a sample.
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Upsilon
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I think I'm going to go for a copper sample today. I was considering a CuO thermite, but seeing videos of this reaction it apparently occurs
explosively so I may not do that. I might go for a CuSO4 electrolysis, depositing the copper on a small copper wire.
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j_sum1
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CuO thermite is fun. You should do it. But yes, it is a bit vigorous. The copper disappears in a fine brown cloud of smoke.
You might make it workable by mixing in something to absorb the excess energy -- perhaps some CaF2 or CaO.
You might also try something a bit less energetic than Al. I am going to try one with zinc powder some time soon.
Electrolysis is easy. Make some sulfuric acid while you are at it. Lead anode, copper wire cathode, saturated solution of CuSO4 low current density,
24 hours.
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Upsilon
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| Quote: Originally posted by j_sum1 | CuO thermite is fun. You should do it. But yes, it is a bit vigorous. The copper disappears in a fine brown cloud of smoke.
You might make it workable by mixing in something to absorb the excess energy -- perhaps some CaF2 or CaO.
You might also try something a bit less energetic than Al. I am going to try one with zinc powder some time soon.
Electrolysis is easy. Make some sulfuric acid while you are at it. Lead anode, copper wire cathode, saturated solution of CuSO4 low current density,
24 hours. |
Yeah it does look cool, too bad I live in a suburban area and won't really be able to set off anything explosive without having the cops called. I
even get nervous when setting off thermites because of the smoke they cause.
I guess now would be a good time for me to learn about electrolytic deposition behavior. What are the criteria for having the metal deposit from
solution to form a solid mass as opposed to a powder that falls to the bottom? Does the electrode have to be made of the metal you're trying to
deposit? Or can it be any metal? I'm pretty sure graphite electrodes won't allow you to deposit any metals on it.
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Agari
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| Quote: Originally posted by Upsilon | | Quote: Originally posted by j_sum1 | CuO thermite is fun. You should do it. But yes, it is a bit vigorous. The copper disappears in a fine brown cloud of smoke.
You might make it workable by mixing in something to absorb the excess energy -- perhaps some CaF2 or CaO.
You might also try something a bit less energetic than Al. I am going to try one with zinc powder some time soon.
Electrolysis is easy. Make some sulfuric acid while you are at it. Lead anode, copper wire cathode, saturated solution of CuSO4 low current density,
24 hours. |
Yeah it does look cool, too bad I live in a suburban area and won't really be able to set off anything explosive without having the cops called. I
even get nervous when setting off thermites because of the smoke they cause.
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I never tried a copper thermite reaction,only the one with iron oxide as an oxidizer and aluminum as fuel,that's how I got iron. What is the more
energetic version of the copper thermite reaction,what are the components?
[Edited on 11-11-2015 by Agari]
Element Collection Status:
Elements Acquired: 21/91
Latest: Lead (Pb)
Quantity: 12g
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j_sum1
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Good question. I know little. Except that there are a large number of variables to play with -- substrate, concentration, current density,
temperature, agitation, surfactants, impurities and the presence of other ions.
My copper electrode is built up on a copper wire and is quite a dense solid mass. It looks ugly since I did a bit of electrolysis with carbon anodes
and so it has fine carbon dispersed throughout. I get a small amount of Cu powder deposit but that is really minor. I don't get long dendrites but I
guess the copper is probably dendritic on a smaller scale.
If I was aiming for a particular form, I would read the literature and follow a recipe.
My intention at some stage is to electrorefine my ugly electrode and produce a cleaner and more pleasant-looking copper sample for the element
collection. So... if you come across a method that results in beautifully-formed crystals, I am interested.
As for the CuO thermite -- it's all over in a few seconds. It would easily be mistaken as a dust cloud -- perhaps sawdust unless you were up close.
In a windy day, there would be perhaps a 20 second opportunity to notice something suspicious. For a small thermite you should get away with it
fairly easily.
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