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Author: Subject: Suitable test tubes for use with 48% HF
Mailinmypocket
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[*] posted on 21-1-2013 at 13:33


Here we go :)

http://www.firstaidproductsonline.com/calciumgluconate.aspx
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[*] posted on 21-1-2013 at 14:15


Quote: Originally posted by bfesser  

Out of personal curiosity, and possible benefit to others; I'm <a href="viewthread.php?tid=19098&page=23#pid272260">requesting the two articles</a> regarding the treatment of HF burns with <a href="http://en.wikipedia.org/wiki/Calcium_gluconate" target="_blank">calcium gluconate</a> <img src="../scipics/_wiki.png" />, which are cited in the Wikipedia article.

I managed to get the newer of the two references (posted in your request thread)...

<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: fixed broken image(s); link]

[Edited on 7/9/13 by bfesser]
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[*] posted on 22-1-2013 at 00:31


Quote: Originally posted by Mailinmypocket  
Here we go :)

http://www.firstaidproductsonline.com/calciumgluconate.aspx
Thanks for the link. I'll see if I can find a calgonate supplier in the EU, that probably saves a lot of money on shipping costs.

If a tube of 25 ml costs $29 (a little over EUR 20), then I am willing to invest that money on safety and then I'll not try something with CaCl2 or Ca(NO3)2. I do not think that the chloride or nitrate in these two chemicals do any harm as suggested before in this thread, but I can imagine that absorption of the calcium by the affected tissue is better with the gluconate. That probably is the secret of using gluconate instead of other more easy to obtain calcium salts.




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watson.fawkes
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[*] posted on 22-1-2013 at 05:50


Quote: Originally posted by woelen  
I do not think that the chloride or nitrate in these two chemicals do any harm as suggested before in this thread, but I can imagine that absorption of the calcium by the affected tissue is better with the gluconate.
My suspicious is exactly the opposite, that it's the counter-ion which is the problem. Gluconate is an oxidation product of glucose, more-or-less already in the metabolic chain. Chloride, for example, while common in the body, participates in the sodium-potassium pump, and (I'm surmising) a local excess of chloride has some kind of side effect.
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[*] posted on 27-1-2013 at 10:55


I now have the HF and I have the polypropylene plasticware and a small 100 ml high density polyethylene bottle. I did my first experiments with HF. Cold (appr. 10 C) HF at a concentration of 48% does not visibly fume in contact with air.

The polypropylene test tubes are not as nicely transparent as glass test tubes, but they are quite useful. In reality they are more transparent than one would conclude on the basis of pictures on eBay.

- Add a few drops of HF (48%) to a solution of ferric ammonium sulfate: The liquid becomes colorless at once. Initially it was brown, due to dissolved and partially hydrated ferric ions.
- Add a ml or so of 48% to some powdered Ti-metal: A very violent reaction occurs and a brown/green liquid is produced. Lots of hydrogen and fumes of HF are produced! If the same experiment is done with 37% HCl, then only a slow reaction occurs and it takes hours to dissolve most of the powder and a deep purple solution is obtained in that case.
- Add some solid MnO2 to 48% HF: No reaction occurs. Apparently, MnO2 is one of the oxides which is not affected by HF.
- Put a few drops of 48% HF in a glass test tube: The glass starts swelling and becomes like a milky gel. The glass heats up a little and the HF starts fuming.
- Add appr. 1 ml of conc. H2SO4 to 2 ml of 48% HF: The liquids mix and there is slight heating of the liquid. The resulting liquid is fuming quite strongly. Add a few crystals of solid KMnO4 to this liquid. The KMnO4 dissolves in the liquid and it does so more quickly than it dissolves in water. The color of the solution is green. The intensity of the color is MUCH lower than the intensity of the color of a similar solution of KMnO4 in water.
When the green solution is added to 100 ml of water, then the solutions obtains a deep purple color.

These experiments were just some random experiments to get a feeling of the properties of HF and to assess the strength of the solution, but it was quite interesting already. More will follow and maybe this will lead to new web pages on my website.

I must say that I do feel somewhat uncomfortable when working with 48% HF. I feel much more comfortable when working with e.g. conc. H2SO4, 65% HNO3 or stuff like PCl5 or SOCl2. Everytime when I see a little wet spot somewhere I am wondering if that is a drop of HF or just some water.




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[*] posted on 27-1-2013 at 12:10


Nice to see you received the HF in good shape and have done some interesting experiments with it already!- it is quite fun stuff, treated with respect of course :)

I will take out my notes from experiments I did with it not long ago and share them here, they are similar experiments though. Just familiarizing myself with the substance.

[Edited on 27-1-2013 by Mailinmypocket]
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[*] posted on 27-1-2013 at 15:40


Woelen:

It's not called paranoid when you've got something to be paranoid about! ;-)

It'd be interesting to repeat some of these experiments with, say 50 % NH4HF2...




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[*] posted on 29-1-2013 at 13:47


As you suggested, I did some experiments with NH4HF2 and also with KHF2.

I prepared a (nearly) saturated solution of KHF2 and added some titanium metal to this. Initially no visible reaction occurs. After a few minutes VERY tiny bubbles of gas are produced, slowly. The reaction is so slow that only careful observation reveals that a reaction occurs.

I did the same experiment with a (nearly) saturated solution of NH4HF2. Quite some NH4HF2 dissolves in water. I would think 40% or so, but I did not measure things, it just is estimation by volume of added solid. I added so much NH4HF2 that some of it remained solid at the bottom of the test tube. Now, after addition of some titanium powder, there was a somewhat faster reaction as with KHF2, but still it was quite slow. As soon as some HCl (30%) was added to the solution of NH4HF2 the reaction was much faster and there was a steady production of hydrogen gas. A strong solution of NH4HF2 is not nearly as active towards titanium metal as a solution of similar strength of HF.

I also did the experiment with ferric ions. I added some fairly concentrated solution of FeCl3 (which is deep yellow) to a saturated solution of KHF2 and at once the solution turned opalescent white. Again, the color of the ferric complex (this time with chloride instead of simply hydrolysed ferric ions) at once disappears. The opalescence tells that some very fine precipitate is formed in low quantities. Apparently a colorless fluorocomplex is formed, which with potassium ions is sparingly soluble. I can imagine that the white opalescence is due to formation of solid K3[FeF6]. Iron(III) forms a complex FeF6(3-) with fluoride ions and it might be that the potassium salt of this is somewhat less soluble.




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[*] posted on 29-1-2013 at 14:01


The low reactivity of HF/HF2 slats towards the Ti metal must be the consequence of the low Ka of the acid.

I wonder if there's a (NH4)3FeF6, that could be a good route to anhydrous FeF3...




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[*] posted on 10-2-2013 at 11:52
Some further experiments


Up to now, I am somewhat disappointed about what I can do with 48% HF. I tried the following things, but I did not really get really interesting results:

- Add very fine white sand (mainly SiO2, the sand is off-white) to some 48% HF. The sand very slowly dissolves. No gas bubbles are produced, the liquid does not become entirely clear. Most likely that is due to some contaminants in the SiO2.
Add a lot of water to the HF with dissolved sand. The liquid become opalescent, apparently some SiO2 or H2SiO3 is formed again.

- Add some high-purity V2O5 to 48% HF: The solid slowly dissolves. After a few minutes a pale yellow and clear solution is obtained.
Add 10 times its volume of water: The liquid remains clear and very pale yellow. I expected the formation of orange hydrous V2O5, but apparently the fluoro-oxo species of vanadium, which is formed from V2O5 and HF is stable even at low concentration.
Add some very fine zinc powder to this liquid: The zinc powder hardly reacts, there is very slow production of hydrogen gas.
Add appr. 10% of total volume of conc. HCl and swirl: Now there is a somewhat faster reaction. The zinc dissolves, giving hydrogen gas and the liquid turns blue, due to reduction of vanadium(V) to vanadium(IV).

- Add some GeO2 to 48% HF: The solid does not seem to react, no formation of gas, the solid does not dissolve. On initial contact there was some hissing noise, but no visible violent reaction occurred.

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

More interesting experiments most likely require anhydrous HF, but I do not have the apparatus to make and contain that stuff. I do have some oleum (65% SO3) though, maybe this can be used to make mixes of anhydrous HF and H2SO4 if added in the right amount. I see no way to mix these chemicals safely. I once added a drop of oleum to water and on contact with the water I heard a loud crackling noise, which was really scary. I am afraid that the mixing of the SO3 and aqeuous HF will be so violent that it may cause an eruption of HF/oleum mix and plumes of HF-fumes, even when working on ml-scale. If any of you has suggestions on how to mix these chemicals safely, then I am willing to do some experiments with anhydrous HF/H2SO4 mixes.




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[*] posted on 10-2-2013 at 13:10


Interesting stuff, woelen.

Perhaps it all shouldn't come as too much of a surprise: aqueuos HF is a weak Bronsted acid after all.

With anhydrous HF, quite 'inert' oxides like TiO2 are reported to cause the HF to boil because of reaction heat.

But it takes nothing away from HF's dangerous nature.

Since as I'm about to do some stuff with NH4HF2 - what kind of test tubes are you currently using?




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[*] posted on 10-2-2013 at 13:51


Curious, I've seen video of HF with SiO2 boiling from reaction heat -- the reaction itself, of course, doesn't release gas.

Do you have any elemental silicon? I forget which etch steps HF is used for in silicon processing, but it should corrode the element regardless, giving some hydrogen to better indicate activity. (I do, however, recall KOH is used for anisotropic etching; on certain cuts, I forget which crystal plane -- it makes pyramidal pits.)

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[*] posted on 10-2-2013 at 14:27


@blogfast25: I use these test tubes, from seller the_science_geek:

http://www.ebay.nl/itm/Test-Tubes-Tube-Plastic-Polypropylene...

This same seller also has PP-beakers, a HDPE-bottle for storing HF and solutions of ionic fluorides and PE-pipettes for transferring solutions. I ordered all in a single order such that I only had moderate shipping costs.

@12AX7: I have seen a video about boiling HF as well, but that was (nearly) anhydrous HF in a glass test tube. Mine is 48% HF and that water makes quite a lot of difference.

The reaction with SiO2 can release a gas, being SiF4. But in my case the reaction is slow and then the SiF4 does not escape as gas but remains in solution as H2SiF6.

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

I'll try reaction with hydrous Ta2O5 and Nb2O5. I have not been able to dissolve one of these oxides in anything, maybe HF can do the job. Right now it's too late (almost midnight where I live), maybe tomorrow or Tuesday.

This evening I also received a faint whiff of HF. Its smell is very similar to that of conc. HCl, but a little bit less pungent. A person who does not know he is working with HF and has experience with HCl will immediately say that he is smelling fumes of concentrated hydrochloric acid.

[Edited on 10-2-13 by woelen]




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[*] posted on 10-2-2013 at 14:42


Ok, there are nice but unfortunately (and unsurprisingly) slightly opaque.



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[*] posted on 10-2-2013 at 15:04


A nice little synthesis with aqueous HF is the preparation of potassium tetrafluoroborate, KBF4. I've done this myself a long time ago.
Slowly add 1 mole of boric acid to 4,5 moles of 40% aqueous HF while swirling and cooling. When the addition is finished and the solution has cooled down, slowly, with swirling, dropwise add a dilute aqueous solution of 0,5 moles K2CO3 (1 mole of K+).
Cool the solution and filter through a paper filter. Wash the crystalline solid with cold water and leave it to dry.
The solubility in cold water is almost the same as that of KClO4.
KBF4 crystals have almost the same refractive index as water so the crystalline precipitate is nearly invisible in the solution.

This should work the same way for K2SiF6 by dissolving SiO2 in aqueous HF and adding the stochiometric amount of K2CO3. This salt is also very sparingly soluble in water.

As an antidote for HF burns I have bought myself a box of twenty 10ml ampoules with 10% calcium gluconate solution.
Since the contents of these ampoules are sterile and the ampoules are hermetically sealed, they will keep for a very long time (many years, much longer than the expiration date printed on the box).
This is the ebay link:
Calciumgluconat 10% Injektionslösung
In case of skin contact with HF I will soak a cotton bandage with the solution and keep it in contact with the affected area.
I haven't been able to find the calcium gluconate gel, and the shelf life of the gel is quite short as I've read. How long can your gel be kept, woelen?




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[*] posted on 10-2-2013 at 23:58


@blogfast25: These tubes indeed are somewhat opaque, but in reality they are not as opaque as they seem to be on the eBay picture. Unfortunately this opacity prevents me from making nice pictures for my website. Especially the experiments with titanium metal give nice colors, but they look dull when photographed in these test tubes.

I also have seen clear test tubes, but these are made of polystyrene and I have done an internet search on this and from multiple sources I understand that polystyrene is attacked by HF fairly quickly.


@garage chemist: I'll try that synthesis of KBF4. This KBF4 then in turn can be used for making BF3? That would be a really interesting thing.

I do not have calcium gluconate gel. I have the pure chemical as a dry solid. This has unlimited shelf life. When needed I can mix this with some gel and rub it on the skin. Calcium gluconate can be purchased easily as a pure chemical and is delivered without any questions.

The link you provide is even more interesting. It is affordable (much more so than the gel), but they do not ship to the Netherlands :( . I am somewhat reluctant to order medical stuff from foreign places. In the Netherlands there is a lot of hassle about so-called "Online Apotheken" and many many parcels from online pharmacies are captured and the people ordering from them get a lot of trouble (big fines). We now have special teams operating in the Netherlands whose task is to find and if possible shut down online pharmacies and pursue their customers. It now is more safe to order all kinds of chemicals online than ready for use medicins.

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

I also read about the gas ClO3F (perchloryl fluoride). This gas can be made from HF and KClO4 in the presence of a water absorbing agent (e.g. SO3). This sounds interesting. ClO3F is sufficiently inert that it is not hydrolysed at once by water and apparently can be collected over water. Unfortunately P4O10 cannot be used as drying agent, because this forms very stable fluoro complexes and fixes the fluoride, such that it is not available for reaction anymore.

[Edited on 11-2-13 by woelen]




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[*] posted on 11-2-2013 at 05:32


I'm surprised no one has considered PVC yet. That's transparent and chemically quite resistant...



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[*] posted on 11-2-2013 at 06:40
HF and Si/SiO2


For my masters degree, I performed EPR measurements on defects in silicon. The defects were due to ion implantation in float zone silicon wafers. After the samples were cut to size, they were etched with 35% (iirc) HF solution to remove the silicon dioxide.

When I asked whether not the whole sample might thus be dissolved in the HF solution, the answer was a convinced 'no'. Silicon was claimed not to dissolve in HF solutions. Is this correct?
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[*] posted on 11-2-2013 at 06:45


Quote: Originally posted by woelen  
- Add a few drops of HF (48%) to a solution of ferric ammonium sulfate: The liquid becomes colorless at once. Initially it was brown, due to dissolved and partially hydrated ferric ions.

Did you try to crystallise the fluoro ferric complex? I wonder wherther the flouroferrates can be isolated as stable compounds.
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[*] posted on 11-2-2013 at 08:42


Yes. Na3FeF6 for instance is commercially available. Hexafluoroferrates are quite stable.

In complexometric (EDTA) titrations the fluoride ion is used as a masking agent for Fe(III).

[Edited on 11-2-2013 by blogfast25]

[Edited on 11-2-2013 by blogfast25]




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