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unionised
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Quote: Originally posted by kmno4  | Accidentally found article, may be interesting for some:
Compositional analysis of copper–silica precipitation tubes
| Code: | http://www.chem.fsu.edu/steinbock/papers/pccp07.pdf |
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Great!
Now we can check to see if there are bands in the raman spectra of the gel that are not due to silica gel or to copper hydroxide.
As far as I can see, they didn't run the raman spectrum of silica gel- they used Na silicate as their reference.
I found a spectrum for silica gel here
http://www.ias.ac.in/matersci/bmsapr2011/299.pdf
Fig2- the line labeled x=0
It has sharp bands at about 600 and 900 /cm and a broad band near 800 /cm
Those may arise specifically as a result of the method of preparation but, in any event, they are missing from the spectra recorded for the copper
containing gel in the paper KMnO4 found.
It looks to em like further work is needed in this field.
Anyway, to answer Blogfast's question, I can't see how the silicate wouldn't give silica gel on adding acid. It might be colloidal (particularly at
first) but this stuff didn't ppt even when the solution was left to dry/ settle in a shallow dish overnight.
I will get some batteries for my balance and do a titration on it to see where that gets me.
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blogfast25
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| Quote: Originally posted by turd | | Quote: | | Presumably this is BaSO4 and visually speaking in a quantity that would be consistent with about 0.01 mole of BaSO4. This is strong evidence
against a basic copper sulphate being the blue precipitate. |
You can visually distinguish 0.010 and 0.0075 mol BaSO4? |
Pondering the results last night it occurred to me that the precipitate may indeed contain sulphate and that this would make it more consistent with
Brochantite (Cu4(SO4)OH6). So now I'm going to chase up this sulphate.
Update:
The blue precipitate (third photo in my previous post) was scraped off the filter, crushed up a little and put on a new filter (Buchner), then washed
again with several aliquots of boiling DIW and sucked dry, in order to avoid any false positives by removing the last traces of soluble sulphate.
The washed filter cake was recovered into a 250 ml beaker and some water and a few ml of 37 % HCl was added. It dissolved easily to a clear green
solution. To it, 0.01 mole of Ba(NO3)2, dissolved in 50 ml of DIW was added. A white precipitate, presumably BaSO4, formed. I’ll leave the BaSO4 to
sink completely overnight, the quantity of precipitate can then be roughly compared to the BaSO4 from the previously obtained filtrate. But already it
looks like much less than that, consistent with (for stoichiometric purposes only):
4 CuSO4 + 6 OH- === > Cu4(SO4)(OH)6(s) + 3 SO4(2-)
… which makes one expect to find 3 times more sulphate in the filtrate than in the product.
All of this does seem to point strongly to Brochantite being the precipitate and not some copper silicate of sorts:
1. Precipitate doesn’t disintegrate to CuO on boiling, as Cu(OH)2 does
2. Precipitate seems to contain bound sulphate anions
3. On drying at about 200 C, precipitate turns green, not black
[Edited on 21-9-2013 by blogfast25]
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Eddygp
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That's a good idea, blogfast. Curiously enough, I have been looking for brochantite and other similar [copper sulfate with or without additional
anions] minerals in Mindat.
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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blogfast25
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Firstly, I can confirm my ‘Na2SiO3.5H2O’ is likely just that and I think unionised’s product probably is too. It’s just harder to make the
hydrated silica gel appear than one might superficially think. To do so I made a much more concentrated solution (of unknown molarity) and then added
a few ml of it to about the same amount of about 20 % HCl. The precipitate appeared slowly and there still wasn’t much of it. It would be very
easy to miss using much more diluted silicate solution. Unionised should try precipitating silica gel with a much more concentrated solution of
his reagent.
Secondly, here are the BaSO4 precipitates obtained from the dissolved blue precipitate (left) and from the filtrate (right):
The quantities are (subjectively) consistent with one third of sulphate being bound up in the product and two thirds remaining in solution.
I’m now fairly convinced the product is a basic copper sulphate. The question remains, why does it form in these precise conditions? Given that the
hydroxide ions are supplied by the hydrolysis of the conjugated base of a weak acid, could hydrolysis of other conjugated bases of weak acids be used
to prepare basic copper salts?
[Edited on 22-9-2013 by blogfast25]
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turd
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Ok.. You seem to have gotten what I got in my first experiment (crystalline basic copper sulfate). If you have time to waste you could also try the
conditions that led to the amorphous silica gel with an amorphous copper salt dispersed in it. Though convincing analytics will by much more difficult
in this case, since the microporous silica gel will adsorb everything from the solution.
Note that for example lack of black color is not proof that there is no CuO - the particles in such a gel might be smaller than half the wavelength of
visible light...
| Quote: Originally posted by blogfast25 | | The question remains, why does it form in these precise conditions? Given that the hydroxide ions are supplied by the hydrolysis of the conjugated
base of a weak acid, could hydrolysis of other conjugated bases of weak acids be used to prepare basic copper salts? |
I don't understand... The conjugated base of a weak acid is a strong base - you will have a high pH (just check with an indicator paper) and
accordingly many hydroxyls in that solution... So yes, there are certainly other strong bases that will give these salts.
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kmno4
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Everything is nice but......
This "copper basic sulfate" looks a little strange to me.
On the pictures it is perfectly blue, whereas all samples of basic sulfates I prepared (long time ago, by boiling sol. CuSO4 with CuO) and pictures of
minerals I have seen recently are green.
There are many articles about copper basic sulfates, for example
Thermal analysis, X-ray diffraction and infrared spectroscopic study of synthetic brochantite (DOI: 10.1007/BF01913606) and
it is written there: Sample SO is pale-green in colour and changes to brownish-black when heated beyond 300 C
(sample dried in air at 30 C. It is designated "SO" in this paper).
Results/products from sodium silicate are likely more complex than it seems.
Слава Україні !
Героям слава !
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blogfast25
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| Quote: Originally posted by turd | | I don't understand... The conjugated base of a weak acid is a strong base - you will have a high pH (just check with an indicator paper) and
accordingly many hydroxyls in that solution... So yes, there are certainly other strong bases that will give these salts. |
The concentration of OH<sup>-</sup> ions, [OH<sup>-</sup>], in a solution of (for instance) 0.4 M sodium acetate is
much lower than in an equivalent solution of 0.4 M NaOH. Respectively for sodium acetate (0.4 M), [OH<sup>-</sup>] ≈
10<sup>-4.4</sup> (pH ≈ 9.6) and for sodium hydroxide (0.4 M) [OH<sup>-</sup>] ≈ 10<sup>-0.4</sup> (pH
≈ 13.6), so much higher in the case of the real strong alkali. So I'm wondering whether this difference in
[OH<sup>-</sup>] is what causes a basic salt to precipitate (in the case of sodium silicate), rather than the straight hydroxide.
[Edited on 22-9-2013 by blogfast25]
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blogfast25
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| Quote: Originally posted by kmno4 | Everything is nice but......
This "copper basic sulfate" looks a little strange to me.
On the pictures it is perfectly blue, whereas all samples of basic sulfates I prepared (long time ago, by boiling sol. CuSO4 with CuO) and pictures of
minerals I have seen recently are green.
There are many articles about copper basic sulfates, for example
Thermal analysis, X-ray diffraction and infrared spectroscopic study of synthetic brochantite (DOI: 10.1007/BF01913606) and
it is written there: Sample SO is pale-green in colour and changes to brownish-black when heated beyond 300 C
(sample dried in air at 30 C. It is designated "SO" in this paper).
Results/products from sodium silicate are likely more complex than it seems. |
On prolonged boiling it would almost certainly also turn green: the dried (as in 'sucked dry') precipitate starts turning green on drying on a hot
plate quite quickly.
A fully fledged elemental analysis of the product would be in order to obtain certainty beyond reasonable doubt with regards to its primary structure.
[Edited on 22-9-2013 by blogfast25]
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turd
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As I posted upthread: I had the blue product analyzed. It was undoubtedly nicely crystalline brochantite. Colors are treacherous!
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miken277
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Not to resurrect a dead thread on the subject of (possible) copper silicate, I recently got a very interesting blue product from a reaction between
basic green copper carbonate and sodium silicate solution, when the two were heated together at very moderate temperatures for hours (a rice warmer).
Upon filtering off the unreacted copper carbonate, two crystalline products resulted from evaporation of the bright blue liquid: a clear crystalline
solid (actually very slightly blue from inclusion of small amounts of the blue compound) that upon drying deliquesced to a white powder (presumably
unreacted hydrated sodium silicate), and a crunchy sand like blue precipitate that upon examination was entirely clear and had a reasonably high index
of refraction. Upon washing this product in warm and cold water multiple times, one noted the lack of easy solution, and upon later examination, the
transformation to an opaque turquoise solid instead of a transparent light blue glassy solid. They were later compared under light magnification and
found to be in different crystal states.
The degraded blue solid was subsequently washed with acetone and allowed to dry, whereupon it was found that the deliquescence could be partially
rubbed off, resulting in a minority of the crystals regaining some transparency. This solid was decomposed in a pyrex dish with the aid of focused
sunlight. The result was the expected black copper oxide plus a fusible white solid. Further tests need to be performed to confirm the presence of
silica in this remaining sample. I will do these at some point and post pictures if people are interested.
--Mike
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blogfast25
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Hello Mike,
The only reaction I can see possible between copper basic carbonate and sodium silicate is the formation of small amounts of sodium cuprate,
Na<sub>2</sub>Cu(OH)<sub>4</sub>, due to the alkalinity of sodium silicate solutions. Even dilute solutions of cuprate are
significantly blue.
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Nicodem
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A Chemical Reviews article on chemical gardens was recently published. It has an open access policy and contains beautiful examples of this
demonstration experiments as well as interesting explanations of the phenomenon:
From Chemical Gardens to Chemobrionics (open access)
Laura M. Barge et al.
Chem. Rev., 2015, 115(16), 8652–8703.
DOI: 10.1021/acs.chemrev.5b00014
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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