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Arthur Dent
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[*] posted on 9-6-2011 at 03:24
Strontium and Barium


I recently purchased some interesting chemicals from a local pottery store including fairly pure Cobalt Carbonate and Manganese Carbonate, and I also got half a kilo of Strontium Carbonate. Yay! :D

But sadly, upon reading the MSDS, I learned that this particular grade of pottery SrCO3 may contain up to 2% Barium Carbonate contamination...

I've noticed that most Strontium and Barium salts have quite similar solubilities, so separating the two looks like a problematic task. I'd rather have a cleaner strontium compound that wouldn't have the toxic barium impurities in it.

Would there be some way of separating the Barium from the Strontium that would be feasible in a home lab, with reasonable yield?

Robert




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[*] posted on 9-6-2011 at 04:20


This might be a poor method, as you would just end up adding more chemicals and having to remove them and I can't imagine you getting a large yield, but you could make use of differences in solubility, i.e. Ba(NO3)2 is much less soluble than Sr(NO3)2
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[*] posted on 9-6-2011 at 04:45


According to the Wikipedia article for strontium sulfate, it dissolves appreciably well in alkali metal chloride solutions. So react your SrCO3 + BaCO3 with H2SO4. It will produce a precipitate of SrSO4 + BaSO4. Dissolve the precipitate in an excess of NaCl solution and filter the famously insoluble BaSO4. React the filtrate with Na2CO3 to reprecipitate SrCO3. Filter and dry. End of long-winded purification procedure.
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[*] posted on 9-6-2011 at 05:25


I would do it as follows:

Dissolve your mix of carbonates in dil. HCl and filter the solution from insoluble impurities. Next make the solution concentrated enough so that when you add excess NaOH solution, the barium will remain dissolved together with some strontium. You will get a precitipate of strontium hydroxide wich you can filter, wash and dissolve in acid to obtain a usable strontium salt.
Because the barium impurity is minor, you only lose a small amount of strontium this way (strontium hydroxide is somewhat soluble so don't use too much water, but also don't use very little, this will probably make filtering more difficult and is more likely to contaminate your solid with some barium).
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Arthur Dent
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[*] posted on 9-6-2011 at 10:36


Thank you guys,

These seem like promising suggestions to purify my Strontium. Thank you.

@LanthanumK: So when I add Sodium Carbonate to the saturated solution of SrSO4 and NaCl, the Strontium will precipitate as a carbonate, leaving a solution of dissolved NaCl?
I guess I should do this outside because quite a bit of SO2 will be evolved, right?

Is Sodium Carbonate preferable to the Bicarbonate in this reaction?

This sounds quite easy to do. I'll try it this weekend!

Robert


[Edited on 9-6-2011 by Arthur Dent]




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[*] posted on 9-6-2011 at 11:25


If you use bicarbonate you must boil the solution to drive off CO2 as the alkaline earth bicarbonates are somewhat soluble.

An alternative is to dissolve the carbonate in an excess of good grade hydrochloric acid, then __slowly__ add enough __dilute__ H2SO4 to precipitate roughly 3 to 4 times as much Ba is in the carbonate, using good stirring. BaSO4 is much less soluble the SrSO4, the conditions I gave result in the Ba preferentially precipitating out along with a bit of the Sr (equal to the excess of H2SO4, but don't be greed - accept a loss of several percent Sr as the sulfate).

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[*] posted on 9-6-2011 at 11:50


Quote: Originally posted by Arthur Dent  


So when I add Sodium Carbonate to the saturated solution of SrSO4 and NaCl, the Strontium will precipitate as a carbonate, leaving a solution of dissolved NaCl?
I guess I should do this outside because quite a bit of SO2 will be evolved, right?

[Edited on 9-6-2011 by Arthur Dent]


SO2? Where would the SO2 come from? Isn't the reaction as follows?:

Na2CO3 + SrSO4 => SrCO3 + Na2SO4

The strontium carbonate precipitates out to drive the reaction to completion.

No sulfur dioxide is evolved, right?




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[*] posted on 9-6-2011 at 12:04


Quote: Originally posted by LanthanumK  
According to the Wikipedia article for strontium sulfate, it dissolves appreciably well in alkali metal chloride solutions. So react your SrCO3 + BaCO3 with H2SO4. It will produce a precipitate of SrSO4 + BaSO4. Dissolve the precipitate in an excess of NaCl solution and filter the famously insoluble BaSO4. React the filtrate with Na2CO3 to reprecipitate SrCO3. Filter and dry. End of long-winded purification procedure.


OK how do you get the chloride out?

I know there is a book on this, however, I will take the
inexpensive way out and consult a free catalogue.

S/P Mallinckrodt Laboratory Chemicals

Strontium carbonate AR
Barium 0.01%

Byda - Linke is 1 914 pages.


Linke-Cover.jpg - 161kB SrSO4-Solubility.jpg - 193kB


djh
-----
"The most interesting part of this spectrum are the emission bands at 624 nm,
636 nm, 648 nm, and 675 nm. These bands are a result of SrCl emission and
are not a part of any barium atomic or molecular emission system as previously
reported [1]. The strontium is most likely present as an impurity in the barium
nitrate. While emissions from impurity constituents are quite common in
pyrotechnic flares, they are usually a neglible part of the total part. The most
common impurity is sodium which seems to be present in almost all
compositions. However, even with its excellent emission properties, as an
impurity it seldom accounts for more than 1% of the total power. The molecular
emission from SrCl in the case of this flare accounts for almost 22% of the total
radiant power. It should also be pointed out that this SrCl emission is not an
isolated event but is observed in all compositions containing barium nitrate and a
chlorine donor."

View TR Citation | View Full Text pdf - 638 KB
Title: Spectral Distributions V. Visible Spectra of Standard and
Improved Green Flare Compositions.
Personal Author: Webster, III, Henry A
Corporate Author: Naval Weapons Support Center Crane In
Source Code: 409351
Page Count: 18 page(s)
AD Number: ADA089055
Report Date: 01 AUG 1980
http://tinyurl.com/3p3chyl

[1] B Jackson & et al.
Improved Green Signal Compositions
Proceedings of Second International Pyrotechnics Seminar 1970

See also —

Title: Visible Spectra of Standard Navy Colored Flares.
Personal Author: Webster,H A , III
Corporate Author: Naval Weapons Support Center Crane In.
Applied Sciences Dept
Page Count: 21 page(s)
AD Number: ADA140342
Report Date: 15 AUG 1983

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[*] posted on 9-6-2011 at 13:32


The NaCl is soluble, while the SrCO3 is insoluble. Washing the precipitate several times with water should remove the chloride.
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[*] posted on 9-6-2011 at 15:00


Ooh, okay. I see, sodium sulfate along with the sodium chloride will be the soluble byproducts of this reaction...

Good! I'll try it with a bit of my Strontium Carbonate. Sounds pretty straightforward. According to the solubility chart from Wizard though , I'll need a heck of a lot of NaCl to actually dissolve all the precipitate after treating my compound with the sulfuric acid... As for the Sodium Bicarbonate, i'll simply boil-up my solution before chucking it into the SrSO4 + NaCl solution.

LanthanumK, I'll drink a pint of ale in your honor if this works! :D :D

Robert




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[*] posted on 9-6-2011 at 15:15


Don't use me as an excuse to drink. :( If you have too much, just use it in chemistry experiments, such as coloring an alcohol flame red with strontium compounds.
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Arthur Dent
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[*] posted on 9-6-2011 at 15:39


LOL, never too much my friend, I drink in moderation but i do love my bitter ales and fine Islay Single Malts! :D

Speaking of which, I believe it's time for a bit of Laphroaig Quarter Cask. ;)

Robert




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[*] posted on 9-6-2011 at 15:57


Won't the barium sulfate also dissolve in the NaCl solution?

How about preparing the nitrate solution and adding a percentage of H2SO4 or K2SO4/KHSO4. Equilibrium should precipitate the barium preferentially if its slow enough. Filter, recrystallize, etc.

I would also like to make some Sr nitrate to get some pretty red flames!




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[*] posted on 9-6-2011 at 16:02


I have heard that barium sulfate is insoluble in just about anything. I just Googled it (the solubility of BaSO4 in NaCl) and didn't see any results regarding solubility. The reaction BaCl2 + Na2SO4 --> BaSO4 + 2 NaCl is irreversible.
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[*] posted on 9-6-2011 at 16:32


Quote: Originally posted by LanthanumK  
I have heard that barium sulfate is insoluble in just about anything. I just Googled it (the solubility of BaSO4 in NaCl) and didn't see any results regarding solubility. The reaction BaCl2 + Na2SO4 --> BaSO4 + 2 NaCl is irreversible.


You get a 10% discount for two answers within 6-hours.

I included the BaSO4 - H2SO4 because I thought it curious.



BaSO4-Solubility.jpg - 150kB BaSO4-H2SO4-Solubility.jpg - 134kB

I would go with the lanthanum nitrate.


Speaking of hard to separate - remember the element
named for an occupant of The Infernal Regions
imprisoned with Ixion, Sisyphus, Tityus, the giant
whose form is so immense that as he lies stretched over
nine acres, while a vulture preys on his liver, which as
fast as it is devoured grows again, &c.?

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[*] posted on 9-6-2011 at 19:09


You enjoy being witty, don't you Wizard? Your posts hardly ever seem to be straight-forward haha.
Are you referencing tantalum and niobium?
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[*] posted on 10-6-2011 at 04:58


Whilst all this is very interesting, a simple recrystallisation from Sr/Ba compounds with similar solubilities (the chlorides should do it) should keep your minor constituent in the ‘mother liquor’. So, dissolve in HCl, filter, crystallise until almost no liquid is left. Liquid will contain the barium, crystals will be relatively free of Ba. Repeat crystallisation for even better purity. Then you can go and isolate the barium if needed…

Flame testing should be able to confirm this...

[Edited on 10-6-2011 by blogfast25]
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[*] posted on 10-6-2011 at 05:53


@blogfast25: That easy? Okay, I have to admit that this sounds much much simpler. So the Barium Chloride won't crystallize at all?

As for the barium impurity, i'll dispose of it safely, since I have an unused, sealed 500g jar of barium hydroxide (bought it years ago for pyro applications but never found any use for it so far).

Any tips on crystallization? i'm asking because my attempts at crystallization have always been very long in time, took several days to get nice CuSO4 crystals and tin chloride took me something like a month to start crystallizing...

So I have 32% HCl and i'll dissolve the Sr carbonate very carefully in it until it's all neutralized and it won't dissolve anymore. I'll filter that solution and then heat it gently? Or should I let it evaporate at room temp? Will putting a beaker of the saturated solution in my dessicator with a bed of Calcium Chloride at the bottom work faster?

Right now, my pyrex dessicator contains a beaker with Manganese Chloride solution with Calcium Chloride at the bottom, under vacuum and it still hasn't started crystallizing yet.

Robert






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[*] posted on 10-6-2011 at 05:59


Regarding manganese chloride; I dried a nice pink MnCl2 solution with a hair drier set to the highest setting. It needs to stay in a sealed container. An hour of sitting in air absorbed enough water to dissolve itself.

Blogfast25: Barium chloride has a solubility of 31.2g/100mL at 0C. Strontium chloride has a solubility of 106g/100mL at 0C for the hexahydrate, which will be formed by crystallization from solution. From this data it appears that the crystals will have a higher concentration of barium than the liquid. Do you have any reason why the Sr will preferentially crystallize?




hibernating...
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[*] posted on 10-6-2011 at 06:21


Quote: Originally posted by sternman318  
You enjoy being witty, don't you Wizard? Your posts hardly ever seem to be straight-forward haha.
Are you referencing tantalum and niobium?


T and N? Yup. Gives me a chance to pull Bullfinch off the shelf.

Here is something to keep you(s) busy. Sorry to say I cannot
at the moment find my answerer sheet on my PC.


The WiZ's Elements Trivia Quiz.

Greenish yellow.
A metal not found alone.
The acid forming element.
The magnet element.
Three elements know to Tubal-Cain, seven generations from Adam.
The light bearing element.
From flint.
First sons of the earth.
Associated with Saturn.
Found in Egyptian tombs.
Heavy stone
Headache element.
First obtained from charcoal.
Brimstone.
Gold colored element.
Named for a town in Scotland.
The Greek element.
The heavy element.
From Cyprus
The stone element.
Charcoal.
The male element.
Colorful element.
First obtained from urine.

Separating radium from barium.

Metallic Radium
P. Curie and A. Debierne.
Comptes. rend., 1910, 151, 523-525.
In:— The Journal of the Society of Chemical Industry.
29 [18] 1111-1112. September 30, 1910.

By making use of the method described by Gruntz (this J., 1903, 800 ; 1905, 278
; 1906, 80) for the preparation of barium, the authors have succeeded in isolating
metallic radium. A liquid radium amalgam, unstable in the air, was prepared by
the electrolysis of a solution of 0.106 grm. of pure radium chloride, using mercury
cathode (10 grms.) dried, transferred to an iron boat and heated in a quartz tube,
in an atmosphere of pure hydrogen, the pressure of latter being kept above the
vapour pressure of mercury at the temperature of the iron boat. Most of the
mercury had distilled over at 270o C., and the temperature was then raised
gradually to 700o C., when the whole of the mercury appeared to have been
expelled, and the radium commenced to volatilize ; the radium vapour vigorously
attacked the quartz tube. The residue in the boat consisted of a bright white
metal, melting at 700o C., which adhered strongly to the iron. The metal is rapidly
attacked by the air, becoming black, probably owing to the formation of a nitride.
It blackens paper and decomposes water, passing for the most part into solution ;
the dark-coloured residue dissolves completely on addition of a few drops of
hydrochloric acid. The radio-active properties of the metal appear to be normal.
As radium is much more volatile than barium, the authors propose to purify the
metal by sublimation in vacuo.



You would think Nobel Prize Winners would know better....


Physiological Action of the Radiation from Radium.
H. Becquerel and P. Curie.
Comptes. Rend. 132 [22], 1289-1291.
In - The Journal of the Society of Chemical Industry, 8 [20], 845. August, 1901.

The action of the radiation from radium on the skin, announced by Walkoff and by
Giesel, has been confirmed by M. and Mme. Curie, and by Becquerel. Preparations
of radium, carried next to the arm or in the waistcoat pocket for periods of two to
six hours, gave rise to inflammation increasing very gradually, but lasting many
days, and leaving after treatment and recovery, little permanent scars. The intensity
of the physiological action depends on the activity of the radium and the duration
of its application. When handling radium and its compounds the finger tips become
hard and painful ; the pain often remains long after the inflammation has disappeared.
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[*] posted on 10-6-2011 at 06:40


Quote: Originally posted by Arthur Dent  
@blogfast25: That easy? Okay, I have to admit that this sounds much much simpler. So the Barium Chloride won't crystallize at all?

Any tips on crystallization


Quote: Originally posted by LanthanumK  
Blogfast25: Barium chloride has a solubility of 31.2g/100mL at 0C. Strontium chloride has a solubility of 106g/100mL at 0C for the hexahydrate, which will be formed by crystallization from solution. From this data it appears that the crystals will have a higher concentration of barium than the liquid. Do you have any reason why the Sr will preferentially crystallize?


No, the BaCl2 won’t crystallise AT ALL in this case. Because it’s there in such low amounts its solubility limit will NOT be exceeded. This is the principle of ANY recrystallisation, provided the contaminant is present in small amounts, has good solubility and you don’t boil your solution dry. From BaCl2’s solubility data you can even calculate how much liquor you should have at the end for all the barium to stay in solution.

It also explains why on salt lakes the top layer is often the most interesting one: that’s where the rarer and more valuable materials tend to gather because they crystallised out last.

Tips on crystallisation?

There is an appreciable difference in solubility of SrCl2 at 100 C and at 0 C (look up in ‘wiki solubility table’). Carefully boil in your filtered solution until you get cloudiness or visible crystals: these are SrCl2. Allow to cool, then ice (best overnight). A seeding crystal (of SrCl2) would be handy but may not be strictly speaking necessary. Scratching the bottom of your glass container also helps to get crystallisation going. You should get a first crop of quite pure SrCl2. Separate these from the supernatant liquor (which contains the barium) and wash them with a bit of clean, iced water.

The liquor now needs to be boiled in until there’s perhaps about 10 - 20 % of it left (or until you're happy with the amount of solid crystals obtained), then allow to cool and as above. It's easy to boil in too much: on cooling everything then solidifies. Add some water and redissolve at 100 C, cool and make sure there's liquid left.

You can repeat the procedure once or twice, each time you will discard a bit of saturated (@ 0 c) SrCl2 solution which also contains the barium. This can still be separated if needed with the more sophisticated techniques above.

As regards MnCl2, the only time I made a considerable amount of that I had no problems at all. I just boiled in the solution till almost dry, discarded the ‘mother liquor’ (you see? An ‘in situ’ recrystallisation!). The crystals weren’t ‘pretty’ but they were pink, had the right amount of crystal water and were fairly pure.



[Edited on 10-6-2011 by blogfast25]
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[*] posted on 10-6-2011 at 09:43
Library time


Sorry.

purification strontium carbonate barium

Yields 3 270 hits @ Google.com/books
I sucked these out of the first page, I will leave
the remaining 3 260 to others.

Please report back any useful findings.

Physiochemical basis for the purification of strontium nitrate
from barium by crystallization

http://tinyurl.com/6c32lms

Journal of applied chemistry of the USSR.: Volume 59, Pages 1981-2567

Consultants Bureau, New York - 1986 - Snippet view
Existing methods of separating strontium from barium either fail to
achieve the required degree of purity or are of low productivity [1,
2]. In the process of developing a method for the purification of
strontium from barium, ...
books.google.com

-----------
Strontium—Preparation of Pure Salts.—Barthe and Falieres find
that the following method permits the preparation of pure
strontium salts with comparative ease, and has the advantages
that no heat is necessary, and that the chemicals employed need
not be pure.

Dissolve natural strontium carbonate (or the sulphide, obtained by
reduction of the sulphate) in just sufficient dilute hydrochloric acid
(1 :5); it will be of advantage that some of the carbonate or
sulphide remains undissolved. The clear liquid, which contains
calcium, barium and strontium salts, besides small quantities of
iron, alumina and magnesia, is decanted from the deposit, and a
small excess of ammonia added, which precipitates iron and
alumina. To the filtrate is added an excess of sulphuric acid; the
precipitate, consisting of the sulphates of strontium, barium and
calcium, is washed repeatedly by decantation with water
containing 1 to 2 per cent, of sulphuric acid, and finally with
distilled water: this will remove all traces of magnesia and calcium
sulphate. The precipitate is next treated in the cold with an excess
of a solution of ammonium or potassium carbonate (1 : 10),
stirring frequently for two days, and finally washing with distilled
water by decantation. The mixture of carbonate and sulphate is
treated with diluted hydrochloric acid, which dissolves strontium
carbonate and traces of baryta. After allowing to stand for twenty
-four hours, the clear liquid is decanted and filtered through a
filter, previously washed with diluted hydrochloric acid. To the
perfectly clear filtrate is added 200 gm. of hydrochloric acid (1.17)
per litre, and then 2 to 3 gm. of precipitated strontium sulphate,
which may contain barium sulphate; stirring frequently for several
hours. The strongly acid liquid dissolves a little strontium sulphate
(about 0.25 per cent.), but in proportion as the strontium is
dissolved, the barium takes hold of the sulphuric acid, while an
equivalent quantity of strontium chloride is formed. The strontium
sulphate being in excess, will insure the final elimination of all the
barium. The filtrate is evaporated to dryness, the salt dissolved in
three times its weight of distilled water, allowed to stand for
twenty-four hours, filtered, evaporated to crystallization, and
dried. The crystals showed in the spectroscope only the lines of
strontium.—Chem. Zeitg., Rep., 1892, 68, from Bull. Soc. Chim.,
1892, 104.

http://tinyurl.com/683jae2

----
Written for the American Druggist and Pharmaceutical Record.
PREPARATION OF PURE
STRONTIUM SALTS.
By H B. Dunham, Pa G., M.D ,
Boston, Mass.

http://tinyurl.com/3f33hun


djh
----
Do you believe that the sciences
would ever had arisen and became
great if there had not beforehand
been magicians, alchemists,
astrologers and WiZards, who
thirsted and hungered after
abscondite and forbidden powers?

Friedrich Nietzsche
Die fröhliche Wissenschaft, IV, 1886



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[*] posted on 10-6-2011 at 11:51


Quote: Originally posted by mr.crow  
...

How about preparing the nitrate solution and adding a percentage of H2SO4 or K2SO4/KHSO4. Equilibrium should precipitate the barium preferentially if its slow enough. Filter, recrystallize, etc.

...


Basically what I had said a bit earlier - use H2SO4 to ppt out the Ba and a bit of the Sr. That's related to the method just given by T.W.I.I.; however ceramics grade alkaline earth compounds intended for glaze use generally have been refined to eliminate iron and other common colour forming metals so the 1st stage bit using an excess of the SrCO3 isn't needed. However, that is a good way to remove a number of metals, have and excess of or separately form and add to the original solution the carbonate, hydroxide, &ct so as to have trace metals drop out onto the ppt. Oxidising conditions, either through the addition of HNO3 with later adding of the excess carbonate &ct, or by adding H2O2, help as the common impurities Fe and Mn readily form low solubility oxides/hydroxides in mildly alkaaline solution.

Be aware that crystallisation doesn't always do what you hope for, as some salts co-crystalise. It doesn't work for removing FeSO4 from CuSO4 for instance, the oxidation+alkali trick does the job in that case. Any time the ionic radii are similar and the ionic charge is the same, you'll find that simple crystallisation may fail to separate the compounds.


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[*] posted on 18-6-2011 at 06:35


I tried the experiment suggested by blogfast25, and took 200g of the carbonate in a beaker and poured (very slowly) conc. HCl until it didn't fizz anymore. I added a very small quantity of dist. water just to clean off the sides of the beaker, so the solution was intensely saturated and turbid. Then I brought the solution to a boil on my hotplate.

I then filterd the very hot solution and put the still hot solution in my freezer. Well it didn't take long for my solution to turn into a solid block that had the consistency of ice cream. I then filtered off the supernatant liquid and I was left with a generous quantity of Strontium Chloride, and i'm sure its purity is now much higher.

Here's a beaker-full of still slightly wet crystals of Strontium Chloride:



The Strontium Chloride took the form of long needlelike/hairlike crystals and is pure white.

i'll drop in in my dessicator to dehydrate it further.

Robert




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[*] posted on 18-6-2011 at 06:52


I have found strontium chloride annoying to crystallize, and it presumably forms the hexahydrate. Fortunately it is not so hygroscopic a solution that crystals will not form. Merely leave the solution out to evaporate; covered to keep dust out of it. Knock the crystal gunk off the walls occasionally since they wick solution up the sides, fighting against gravity until they saturate whatever you covered the beaker with and make a mess/ waste product. Filter off the glassy-looking crystals once in a while, dry, and return the filtrate to continue evaporating. Any sort of attempt at rapid cooling for crystal formation lead to a fibrous mess as shown in the above post



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