kilowatt
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Beryllium Separation
The complex nature of beryllium chemistry is causing me some confusion on this matter. I would like to isolate beryllium oxide (or other compounds)
from beryllium copper scrap (generally containing 0.3% to 2% beryllium by weight). My idea is to initially dissolve the entire scrap in HCl + H2O2
solution which acts quickly on copper. After boiling the solution to saturation or to dryness, the trick of separating the two salts arises.
One possible option is to add sodium hydroxide to the chloride bath. This should precipitate Cu(OH)2, but the beryllium should form the complex ionic
compound Na2Be(OH)4. I believe this is at least somewhat soluble and should remain in solution. Upon fluorination of the resultant solution,
beryllium fluoride and sodium fluoride should be separable with crystallization (I can't find any solubility figures for BeF2 to compare with NaF,
though). The vast overabundance of NaF compared with BaF2 could make this difficult.
I do not believe aqueous beryllium chloride may be heated to an anhydrous state from the tetrahydrate, and that beryllium oxide or various hydroxides
would instead tend to form with the liberation of HCl. This could be another possible route, as heating the dry Cu/Be chloride mix may leave behind
the insoluble oxide or hydroxide of beryllium. Depending on decomposition yields, another approach could also be used (possibly afterward) by heating
the mixture further distilling out BeCl2 which boils at only 520°C.
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not_important
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regarding BeF2 | Quote: | BeF2 is hygroscopic and excessively soluble in water, it does not dissolve in anhydrous HF and is only slightly souble in absolute alcohol.
...
It can not be made by treating the hydroxide with hydrofluoric acid as the salt hydrolyses on evaporation. |
Looking at its chemistry, up to the hydroxide step should be OK. Some copper may dissolve in the alkaline solution, as you'll want it fairly alkaline
to ensure all the beryllium is extracted into solution. After that I'd consider adding ammonium sulfate and strong aqueous ammonia to precipitate
Be(OH)2 again while keeping copper in solution. Repeat the solution in NaOH and precipitation to remove traces of copper co-precipitated or adhering
to the Be(OH)2. Give it some time to precipitate, there's two forms and the first formed is more soluble.
Various solvent extract methods are commonly used for isolating beryllium, but many require solvents that may be difficult or expensive to obtain.
The beryllium halides must be prepared in a dry fashion, their aqueous solutions hydrolyse on evaporation.
--------------
a bit more
2 N NaOH will dissolve 2/3 moles Be(OH)2 per liter. On long standing or boiling, dense crystalline Be(OH)2 precipitates.
Beryllium ion is soluble in 10% (NH4)2CO3 with a pH range of 8.5 - 9 I believe that copper will precipitate under those conditions. Boiling the
solution will precipitate the beryllium as a basic carbonate.
Beryllium has a noticeable tendency to absorb onto glass which starts at a pH of about 4.5 and increases sharply above pH 6. Given its toxicity this
should be kept in mind, for processes not involving heat polythene containers might be better but some absorption still happens.
Analytical references mention that ignited beryllium compounds tend to dust badly, again given its toxicity it might be better to avoid doing so
unless you are very careful.
[Edited on 7-2-2008 by not_important]
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kilowatt
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Interesting. That seems like some very helpful info, also thanks on the note about the absorption onto glass. Does it actually diffuse into the
glass or can it be cleaned out somehow? I do not wish to permanently contaminate any of my good glassware especially with such toxic compounds. Any
of the high temperature processes will be carried out in a small inconel retort. I will have the means to cast one before proceeding with this
(technically I could probably cast one right now If I was at home, but it would be really pushing the limits of my large furnace).
I'm a little unclear on some of the stuff in your first paragraph here. Adding ammonium sulfate to Na2Be(OH)4 solution will precipitate beryllium
hydroxide by forming sodium sulfate or bisulfate and liberating ammonia gas? What is the purpose of excess ammonia here? Does a high pH need to be
maintained?
Your carbonate precipitation looks like another good option.
| Quote: |
Analytical references mention that ignited beryllium compounds tend to dust badly |
While a respirator is obviously a good idea while working with beryllium bearing powders, none of the compounds I have mentioned ought to be remotely
flammable, though apparently the halides can form shock sensitive explosives in halocarbon solvents. How this is possible I'm not sure, but I read it
in an MSDS. If I go on further to extract the metal by reduction of a halide likely in a gaseous state (I know of no other good way, really),
obviously that can be a rather hazardous process which I think I will hold off on until fully prepared.
| Quote: |
Various solvent extract methods are commonly used for isolating beryllium, but many require solvents that may be difficult or expensive to obtain.
|
I would be interested in knowing more in this regard, it may be possible to obtain or synthesize some such solvents, but I realize that would be
adding more steps to an already complicated process.
| Quote: |
The beryllium halides must be prepared in a dry fashion, their aqueous solutions hydrolyse on evaporation. |
Yes; it is what they hydrolyze into that I am interested in. If they hydrolyze to an insoluble hydroxide or oxide, then I should be able to separate
them pretty easily from the soluble copper chloride mass after heating.
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not_important
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| Quote: | Originally posted by kilowatt
Interesting. That seems like some very helpful info, also thanks on the note about the absorption onto glass. Does it actually diffuse into the
glass or can it be cleaned out somehow? |
The reference does not say, but given Be chemistry I'd assume surface absorption. Citric or oxalic acid and dilute HCl (0,1 N) together is likely to
remove it.
| Quote: | ...
Adding ammonium sulfate to Na2Be(OH)4 solution will precipitate beryllium hydroxide by forming sodium sulfate or bisulfate and liberating ammonia
gas? What is the purpose of excess ammonia here? Does a high pH need to be maintained? |
Neutralise the NaOH, excess ammonia to keep any copper contaminating the Be extract as a complex in solution. If there is little copper carry-over
then the additional ammonia would not be needed, the ammonia released by netralising the NaOH will be sufficient. Ammonium sulfates solutions run a
pH in the range of 5 to 6, target an ending pH of 6.5 to 7.5, maybe 8.5 if there's a lot of copper and NH3 is added.
Don't use a pH meter, the Be will absorb onto the probe.
| Quote: |
| Quote: |
Analytical references mention that ignited beryllium compounds tend to dust badly |
While a respirator is obviously a good idea while working with beryllium bearing powders, none of the compounds I have mentioned ought to be remotely
flammable, though apparently the halides can form shock sensitive explosives in halocarbon solvents. |
This was mentioned in analytic methods where a beryllium compound was precipitate and heated (ignited) to the oxide, phosphate, or other stable
compound. From what I read they tend to form a fine, loose dust that flies around at the slightest provocation.
That information leads me to suggest that strong heating of Be compounds should be avoided if possible. Wet separation methods are preferable over dry
thermal ones.
| Quote: |
... If I go on further to extract the metal by reduction of a halide likely in a gaseous state (I know of no other good way, really), obviously that
can be a rather hazardous process which I think I will hold off on until fully prepared. |
A reference states that solid Be is obtained by electrolysis of a mixture of the fluorides of Na, Be, and Ba at 1350 C. There may be newer methods
than that.
Note that a slightly acid solution of Be(II) treated with F-, the Ba(II), will precipitate BaBeF4. This could be done by added NaF to the acid Be,
then adding BaCl2 or Ba(NO3) solution. Excess Ba will precipitate as BaF2, which may not be harmful if this is being used to prepare the electrolyte.
The mix should be held a bit below boiling for 10 to 15 minutes for better filtering properties. BaBeF4 is anhydrous, compact, and dense, making for
better handling. It can be redissolved in a mixture of HNO3 and H3BO3 (or B(OH)3 if you prefer).
| Quote: |
Yes; it is what they hydrolyze into that I am interested in. If they hydrolyze to an insoluble hydroxide or oxide, then I should be able to separate
them pretty easily from the soluble copper chloride mass after heating. |
I'm not sure you'll get a clean separation in this case. BeO and Be(OH)2 are soluble in solutions of beryllium salts, so an incomplete decomposition
will result in much more Be going into solution than the amount of undecomposed halide would suggest. Another issue might be that the environment of
CuCl2 might retard the decomposition of the hydrated beryllium chloride, with it's boiling point some could vapourise away from the CuCl2 even before
that reached its melting point.
Attachment: Extraction of Be & Al with di(2-ethylhexyl)phosphate.pdf (182kB)
This file has been downloaded 469 times
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12AX7
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Don't forget that copper is soluble in base. If beryllium is acidic enough, it may suffice to precipitate it by neutralizing the solution (just
before Cu(OH)2 precipitates).
Do you know how acidic and basic Be(OH)2 is?
Tim
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not_important
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Be(OH)2 is much more solubile in alkali than is Cu(OH)2. Cu(OH)2 in water is 3 x 10^-5 mole/liter, in 4.2 N KOH 1,8 x 10^-3 m/l; Be(OH)2 is 8 x 10^-5
in pure water, and described as "readily soluble in excess base" which generally implies a m/l or higher.
The solubility question is made difficult by the solubility of BeO and Be(OH)2 in solutions of Be salts, up to 4 moles of BeO dissolving per mole of
other Be compound. Aging of solutions can result in the precipitation of BeO/Be(OH)2, as can boiling. The aging also affects the precipitates, if
they are to be dissolved they should not be allowed to stand too long. In general the oxide and hydroxide and their solutions in alkali or carbonate
should be quickly processed further; solutions of salts such as the chloride can be allowed to stand without problem.
When washing Be(OH)2 or BeO precipitates a little ammonium acetate should be added to the wash water, in order to prevent the precipitate going
colloidal.
Another possible first separation is to bring the Cu+Be solution to near neutrality with NaOH, finish taking it slightly alkaline with aqueous
ammonia, collecting the mixed hydroxides without washing, add warm (25 to 30 C) saturated NaHCO3 solution to give 3 or 4 moles of NaHCO3 per mole Be,
more if needed to completely cover the ppt. The bring it to a boil for a short time, causing the Cu(OH)2 to convert to CuO and dissolving the
Be(OH)2. Filter hot, rinse the precipitate with dilute ammonium acetate adding the rinse to the filtrate, add enough saturated ammonium sulfate
solution to convert all the NaHCO3 to the sulfate and boiling for 20 minutes to precipitate the beryllium. Dissolve the precipitate in HCl or H2SO4,
use just a little more acid than needed to bring it into solution, add ammonia to a pH of 10 or higher to precipitate Be(OH)2 while keeping Cu in
solution.
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kilowatt
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| Quote: |
A reference states that solid Be is obtained by electrolysis of a mixture of the fluorides of Na, Be, and Ba at 1350 C. There may be newer methods
than that. |
1350°C is well above the boiling point of BeF2 at standard pressure. I'm sure the other salts would keep it in solution but the electrolysis would
still be problematic because fluorine will bond instantly with any electrode, especially at that temperature, and the toxic electrolyte would be
extremely volatile, fuming with beryllium salt. I read that beryllium can be produced by reducing the fluoride with magnesium metal at again around
1350°C, probably either under pressure or mixed with something else such as magnesium fluoride. Molten magnesium fluoride would sink to the bottom
while beryllium would float to the top. It could probably be done in a sealed vessel, made of what I'm not sure (probably graphite or a high melting
metal lined with ITC-213), by slowly heating it with the materials inside and monitoring pressure. Later today I should calculate the enthalpy of
this reaction.
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chemoleo
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According to Jander's and Blasius' analytical chemistry book - you should have the following method:
Dissolve all the Cu as indicated using H2O2 and HCl. The book states that Be is stable with oxidising acids, so you may even get Be powder, if you are
lucky, but which I doubt.
Assume you have a mixture of CuCl2 and BeCl2. Precipitate both with Ammonia, and keep adding MORE ammonia. Only the Cu(OH)2 will dissolve forming the
tetraaminhydroxide solution, while Be(OH)2 does *not* form beryllates (similar to Al). Therefore all your leftover hydroxide ((which is only soluble
at 0.0002 g in 100 ml H2O!) can be washed with more NH3 solution until white, and you should be left with pure Be(OH)2. Beautifully simple!
Alternatively, add to the CuCl2/BeCl2 solution KI. CuI2 precipitates, while BeI2 is soluble.
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kilowatt
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| Quote: |
while Be(OH)2 does *not* form beryllates |
What do you call Na2Be(OH)4 which forms in alkali solution then?
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chemoleo
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I don't understand.
Be(OH)2 does not form soluble beryllates with ammonia, i.e. ammonium beryllate does not exist, beryllates only form with Na/KOH etc.
So whilst Cu(OH)2 goes back into solution in the presence of NH3, Be(OH)2 doesn't, thus you are left with clean Be-hydroxide.
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not_important
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It doesn't form beryllates or complexes with aqueous ammonia, but it does with alkali hydroxides. It also forms a soluble complex with NaHCO3, which
is why I suggested that method. This is a complex and not a precipitation based on relative acid/base strengths.
Because the alloy is at the most 2% Be, I suggested precipitating the copper while keeping the Be in solution (the CuO should be washed at
treated/extracted with a little fresh NaHCO3 solution, to extract any entangled Be) The CuO formed should be fairly dense and large enough grained to
be easy to filter. Bringing the Cu into solution leaves a somewhat gelatinous precipitate of Be(OH)2, which is more difficult to filter, in a large
volume of solution. The solution of Cu(II) in aqueous ammonia also tends to dissolve filter paper, so you'll want to use fritted glass.
| Quote: | | Alternatively, add to the CuCl2/BeCl2 solution KI. CuI2 precipitates, while BeI2 is soluble. |
CuI2 is not stable in water, decomposing into CuI and I2. Given the small amount of be in (most) copper alloys, it also will take a lot of KI.
Have you check the full spec on the Cu-Be alloy? Some contain Co, Ni, and/or Pb in amounts about the same as the Be. The ammonia separation will
also pull Co and Ni into solution, Pb should remain as a precipitate with the bicarbonate method.
As for making the metal, all I can tell you i that was the method given for making massive beryllium, as opposed to a powder. Note that Be combines
with carbon at 1300 C.
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chemoleo
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Well if there's contamination with other metals, you can combine any of the above methods - a cheap precipitation step such as the ammonia one, or
conversely of the copper with bicarbonate, then treatment of the residual precipitate/solution with alternative anions, such as sulfate, halogens,
thiocyanates, carbonates, complexing dyes etc. as well as oxidating/reducing agents to get certain transition metals into a higher/lower state for
removal.
Given all the anions commonly available, I should think a full purification of Be should be far from impossible.
PS whether BeOH2 is gelatinous or not doesn't really matter, as the solubility in H2O is extremely low, one just has to clean with H2O/NH3 several
times, or even centrifuge, losses should be minimal. Good point about the filter paper.
[Edited on 8-2-2008 by chemoleo]
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kilowatt
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Yeah last night I looked through most of the BeCu alloys on MatWeb and remember nickel specifically being common. If any lead is present in the final
product, it can be separated by sulfates as beryllium sulfate is soluble.
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12AX7
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Well copper forms carbonate complexes too; neutralizing a copper solution with excess carbonate leaves a blue solution. Neutralizing with hydroxide
leaves a clear solution, but excess (pH > 13?) forms cuprates.
Do you have the pKa values for Be?
Tim
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not_important
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The beryllium separation uses bicarbonate (OK, hydrogen carbonate, or acid carbonate to be trad) not carbonate. Potassium carbonate solution is a
decent solvent for copper carrbonate, KHCO3 is not very good, and the corresponding sodium salts are less effecxtive than the potassium ones.
A small amount of copper likely will stay in solution, thus the sequence NaHCO3 to remove most Cu, concentrate the filtrate some, and precipitate the
Be with NH4(I) and NH3 to keep the carried over copper (and Ni and Co) in solution.
When working with beryllium its complexing with oxygen is at least as important as pK values. Concentrated solutions BeCl2 or BeSO4 will dissolve
more than equal molar amounts of the basic carbonate, with evolution of CO2. Diluting these solutions precipitates the hydroxide or basic salts.
Concentrated solutions of Be salts may also fail to form a precipitate when enough NaOH is added to completely react with the Be salt (so precipitate
Be(OH)2 from moderately dilute solution).
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12AX7
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Complexes (plural): I have Cu(2+) + HCO3(-) <--> CuHCO3+ pKf = -1.8, so it's not entirely negligible (stronger than any chloride pKf, but weaker
than carbonate(2-), listed nearby).
I have all the numbers for copper (or at least a lot of them), but I have no clue about what beryllium does.
Tim
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blogfast25
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| Quote: | Originally posted by 12AX7
I have all the numbers for copper (or at least a lot of them), but I have no clue about what beryllium does.
Tim |
Tim, what's the pKf for CuCl + Cl- <--> CuCl2 (-) ?
And what about CuCl4 (2-), do you have a pKf value for that too?
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12AX7
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Hmm, I don't have anything for Cu(I).
Cu(2+) + Cl- <---> CuCl+ pKf = -0.20
CuCl+ + Cl- <---> CuCl2 pKf = 0.46
CuCl2 + Cl- <---> CuCl3- pKf = 2.03
CuCl3- + Cl- <---> CuCl4(2-) pKf = 2.30
Higher chloro-species form with difficulty, only in strong chloride concentration. It's no wonder it takes several M of Cl- to turn a solid green.
Tim
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blogfast25
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Thanks, Tim!
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