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turd
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Quote: Originally posted by blogfast25  | | So your precipitate does have the same molar ratio CuO:SiO2 as ‘CuSiO3’ but it isn’t actually CuSiO3 because it contains no actual
SiO<sub>3</sub><sup>2-</sup> anions. |
SiO<sub>3</sub><sup>2-</sup>? Your ideas are intriguing to me, and I wish to subscribe to your newsletter. 
But seriously, you certainly meant (SiO<sub>3</sub><sub><i>n</i></sub><sup>2<i>n</i>-</sup> with <i>n</i>>=3.
PS: Dioptase (anhydrous and monohydrate) which seems to be the topic of this thread (is it?) is actually a rather simple hexacyclo-silicate known
since the late 18th century. 
PPS: I doubt hydrothermal treatment and calcination of the precipitate would lead to the same results. Also 700°C in a test tube? This is a bit
dubious.
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blogfast25
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@Unionised:
You think the structure you linked to can arise in these conditions described by PH11?
I think this merits a test of my own, just to see this precipitate and fire it.
Conditions in which colloidal silica gel arises are quite particular: get it wrong and you end up with precipitated silica, also a commercial product.
If it is a co-precipitate of copper hydroxide and hydrated silica then that might well end up grey after drying. It didn't retain blue as expected
either...
@turd:
No one is claiming that 'hydrothermal treatment and calcination of the precipitate would lead to the same results. Also 700°C in a test
tube?'.
[Edited on 14-9-2013 by blogfast25]
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unionised
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I still think "It's probably a complex copper silicate something like this"
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Nicodem
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| Quote: Originally posted by unionised | This observation
"Upon heating at approximately 700oC, the blue solid expands and turns white/gray. water condenses on rim of test tube." proves that it's not copper
hydroxide which would dehydrate to the oxide which is black.
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This only proves that it if there is copper hydroxide, then it is not only copper hydroxide. Yet, this is no evidence that there is
no copper hydroxide there. However, the evolution of water upon heating and the visual transformation of the solid is an excellent indication that the
product is not any of the copper silicates (or at least not just a copper silicate).
| Quote: | | It's probably a complex copper silicate something like thishttp://www.ncbi.nlm.nih.gov/pmc/articles/PMC2960335/and the assertion that it's a mixture
of silica gel and copper hydroxide is baseless. |
Interestingly, the experimental of this article demonstrates that the precipitate which forms upon mixing waterglass with aq. copper sulfate, not only
is not a copper silicate, but is not even a compound at all. Just like the article up-thread that analyses the "chemical garden" precipitate from
copper nitrate and waterglass (and find it to be "copper hydroxide nitrate, together with amorphous silica"), this one also indicates the material is
a mixture of silicagel with copper and sodium products (they only did the elemental analysis of the precipitate - they did not characterize the
components of the mixture). The article also describes the product of the calcination of this precipitate which is again relevant to the recent posts.
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blogfast25
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The way in which the material studied in the article was obtained was also very specific. I quote:
"Experimental
Chemicals were purchased from commercial sources and used without further purification. An alkaline solution was prepared by mixing 13.86 g of a
sodium silicate solution (Na2O 8 wt%, SiO2 27 wt%), 16.13 g H2O and 4.11 g NaOH, and a second solution was prepared by mixing 17.87 g H2O with 7.60 g
of Cu(SO4).15H2O. These two solutions were combined, stirred thoroughly during 2 h and the resulting gel, with a molar composition of CuO: 3.1SiO2:
1.4Na2O: 94.5H2O, was autoclaved for 10 days at 503 K. A crystalline material was obtained [Na2(Cu2Si4O11).2H2O], filtered and treated thermally at
573 K for six hours leads to the removal of the crystallization water molecules."
Not exactly just mixing a bit of CuSO4 solution with some waterglass and keeping fingers crossed...
(I also presume "Cu(SO4).15H2O" is a simple typo)
[Edited on 14-9-2013 by blogfast25]
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unionised
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| Quote: Originally posted by Nicodem |
Interestingly, the experimental of this article demonstrates that the precipitate which forms upon mixing waterglass with aq. copper sulfate, not only
is not a copper silicate, but is not even a compound at all.
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Where?
As far as I can see, they just cooked the stuff- they didn't check for Si-O-Cu bonds.
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turd
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| Quote: Originally posted by Nicodem | The article also describes the product of the calcination of this precipitate which is again relevant to the recent posts.
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The Acta Crystallogr. article describes the product of a hydrothermal reaction followed by drying. 230°C is typically used in Teflon lined
autoclaves. Or are you talking about the Chem. Comm. one?
This is a classic hydrothermal silicate synthesis under basic conditions from silica gel. Also note that this is only an addendum to the actual
article, which is this one: https://www.ncbi.nlm.nih.gov/pubmed/15724175 (hopefully they have a more thorough characterization like phase purity in there). This is just a
case of "Hey, we got the same crystal structure as before, just without H2O. *collective yawn* Let's spend a lazy afternoon and spice up our
publication list with a new entry."
Since I had Na2SiO3 at hand, a quick experiment: 5 g Na2SiO3 dissolved in 30 ml H2O, filtered. 10 g CuSO4.5H2O dissolved in 40 ml H2O. Add 10 ml of
the "water glass". Stir 5 min. Filter off bright blue precipitate. Heat in test tube over Bunsen burner. Precipitate turn first green, then brown,
then black. Quelle surprise!
Maybe if I had let it incubate for a longer time to get a gel... But why bother? This will certainly give an undefined gel with undefined stuff
entrapped in it. That's how sol/gel "chemistry" works. 
Edit: CuSO4, not CaSO4 :p
[Edited on 15-9-2013 by turd]
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Nicodem
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Blogfast, the NaOH is added to make a "Na2SiO3" from the waterglass they used. They used the elemental analysis of their
waterglass and calculated how much NaOH needs to be added to obtain the correct composition. Waterglass has a variable composition where
Na2O vs. SiO2 is usually not 1 : 1 unless you compensate it somehow. Don't forget that waterglass is not a solution of a
compound, but the formal solution of silica in aq. sodium hydroxide.
They did enough. They did the elemental analysis of the precipitate (gel) which demonstrates this is not a single compound as the elemental ratio does
not fit the known valences. They also mention the precipitate is a gel, hence amorphous (copper silicates alone would have to be crystalline). This
indeed does not demonstrate there is no component with Cu-O-Si bonds, but if you claim there is, you encounter a pretty big problem: the reaction
pathway that would lead to such a product. If you think that copper sulfate, for some unknown reason behaves so dramatically different than copper
nitrate in this precipitation reaction and for some reason forms some compound having Cu-O-Si bonds, then please explain how this could occur. I'm not
aware of any reaction mechanism that would give compound with a Cu-O-Si bond by a precipitation reaction from an aqueous solution. I believe
this would be a well known reaction, if it gave such an unexpected result. If it would be so easy then what would be the point of using calcination
reactions or hydrothermal treatments?
| Quote: Originally posted by turd | | The Acta Crystallogr. article describes the product of a hydrothermal reaction followed by drying. 230°C is typically used in Teflon lined
autoclaves. Or are you talking about the Chem. Comm. one? |
You are quite right, the Acta article is just a rehash of the (attached) Chem. Comm. which is barely more detailed in regard to the
experimental:
| Quote: | | Synthesis: AV-23 was synthesised in Teflon-lined autoclaves under static hydrothermal conditions. Chemicals were purchased from commercial sources and
used without further purification. Typically, an alkaline solution was prepared by mixing 13.86 g of a sodium silicate solution (Na2O 8
wt%, SiO2 27 wt%, Merck), 16.13 g H2O and 4.11 g NaOH (Merck). A second solution was prepared by mixing 17.87 g H2O with 7.60 g
of Cu(SO4).15H2O (Merck). These two solutions were combined and stirred thoroughly. The resulting gel, with a molar composition
of CuO : 3.1SiO2 : 1.4Na2O : 94.5H2O, was autoclaved for 10 days at 230 °C. |
Funnily enough, they even copy-pasted the "Cu(SO4).15H2O" typo into the Acta article.
Attachment: b410731d.pdf (561kB)
This file has been downloaded 857 times
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unionised
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"! If you think that copper sulfate, for some unknown reason behaves so dramatically different than copper nitrate in this precipitation reaction and
for some reason forms some compound having Cu-O-Si bonds, then please explain how this could occur."
Do you mind if I do that the other way round?
Rather than the difference between the sulphates and nitrates of copper, can I ask you to explain why magnesium sulphate is well known to be
precipitated from solutions of sodium silicate and magnesium silicate but yet, you say, copper won't do the same thing?
Now, I appreciate that there will be differences in detail, but why would they not form similar products?
Also they say
" An alkaline solution was prepared by mixing 13.86 g of a sodium silicate solution (Na2O 8 wt%, SiO2 27 wt%), 16.13 g H2O and 4.11 g NaOH, and a
second solution was prepared by mixing 17.87 g H2O with 7.60 g of Cu(SO4).15H2O. These two solutions were combined, stirred thoroughly during 2 h and
the resulting gel, with a molar composition of CuO: 3.1SiO2: 1.4Na2O: 94.5H2O, was autoclaved "
Can someone check my maths on this bit?
7.6 grams of copper sulphate ( with 5 H2O as I think we all agree) at 249.7 g/mol is 0.0304 moles of copper.
16.13 grams of water in the first solution, plus the stuff in the silicate solution to begin with: 8+27=35 % of that is sodium silicate so the other
65% is water. That's another 9.009g of water
The first solution has 9.009+16.13=25.14 grams of water
Then there's 17.87 grams in the second and 2.74g from the copper sulphate (assuming the pentahydrate)
So the mixture has 45.75 grams of water in it. That's 2.54 moles
Then there's the water produced by the reaction of sodium hydroxide with polysilicates.
4.11g is 0.103 moles which gives 0.0513 moles of water . That's another 0.925 grams.
So, the grand total is 46.67 grams of water or 2.59 moles
And it also has 0.0304 moles of copper
That's 85 moles of water for each mole of copper
But their "analysis" indicates 94.5 moles of copper for each mole of copper.
Where did it come from?
OK, you can't "magic" water into the mixture.
What if they really believed the copper sulphate hydrate had three times as much water as it really did?
That gives another 5.48 grams of water so the grand total would be 52.15 grams in total 2.897 or moles
and that gives us 94.3 moles of water for each mole of copper.
On that basis I strongly suspect that the "analysis" of the gel is actually just a calculation (with the wrong hydration for copper sulphate).
Anyone care to check the other elements (Na, Si)?
Incidentally, since they haven't included the sulphate ions in that mixture it isn't surprising that it doesn't work chemically.
Now, if I'm right, they didn't analyse the product at all.
If they didn't analyse it the composition of the mixture can't be used as a basis for saying what the product was.
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blogfast25
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| Quote: Originally posted by Nicodem | Blogfast, the NaOH is added to make a "Na2SiO3" from the waterglass they used. They used the elemental analysis of their
waterglass and calculated how much NaOH needs to be added to obtain the correct composition. Waterglass has a variable composition where
Na2O vs. SiO2 is usually not 1 : 1 unless you compensate it somehow. Don't forget that waterglass is not a solution of a
compound, but the formal solution of silica in aq. sodium hydroxide.
|
The sodium metasilicate I sell is advertised to me as 'Na2SiO3.5H2O', so it would have that Na2O:SiO2 molar ratio of 1:1.
Turd and Nicodem: so this paper is essentially 'journal filler'?
| Quote: Originally posted by unionised | Rather than the difference between the sulphates and nitrates of copper, can I ask you to explain why magnesium sulphate is well known to be
precipitated from solutions of sodium silicate and magnesium silicate but yet, you say, copper won't do the same thing?
Now, I appreciate that there will be differences in detail, but why would they not form similar products?
|
That doesn't make any sense. Did you mean ‘why magnesium silicate is well known to be precipitated from solutions of sodium silicate and
magnesium sulphate but yet, you say, copper won't do the same thing?’
You have some reference for this? If correct (i.o.w. an Mg silicate of sorts can be simply precipitated) then maybe that could be due to higher
solubility of Mg(OH)2 with respect to Cu(OH)2? In a simpler situation (no polysilicates), if ‘MgSiO3’ (of sorts) was much more insoluble than
Mg(OH)2 the silicate (or a basic silicate) would precipitate rather than the hydroxide.
[Edited on 15-9-2013 by blogfast25]
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unionised
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Oops, typo.
I mean that this stuff
http://en.wikipedia.org/wiki/Magnesium_trisilicate
is made from the sulphate by precipitation.
It must be less soluble than the Mg(OH)2
Similarly, since the complex silicate in that paper is made by precipitation, it must be less soluble than copper hydroxide.
Also, do you agree with my assessment that they didn't actually analyse the gel and that you can't use that composition as evidence of anything?
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blogfast25
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| Quote: Originally posted by unionised | Also, do you agree with my assessment that they didn't actually analyse the gel and that you can't use that composition as evidence of anything?
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I couldn't find any reference to elemental analysis, I just kind of superficially assumed that they'd done it anyway. 
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turd
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| Quote: Originally posted by unionised |
Similarly, since the complex silicate in that paper is made by precipitation, it must be less soluble than copper hydroxide. |
In which paper? Not in the one you posted. There it is made by a classical hydrothermal reaction (10 <i>days</i> at 230° under autogenous
pressure - this is not a precipitation reaction). Hydrothermal in alkaline medium is the method to grow these huge 30 cm quartz crystals. With
hydrothermal you can't simply argue via solubility. It's a crazy world of its own.
God knows what the gel looks like - it certainly is not a defined compound and highly depends on the gel growth conditions, as usual with sol/gel
chemistry. Not that it matters once you autoclave it in alkaline medium.
| Quote: | | Also, do you agree with my assessment that they didn't actually analyse the gel and that you can't use that composition as evidence of anything?
|
Yes indeed, they put the <i>whole</i> thing in an autoclave and give the molar ratios for reproducibility. Again: this is NOT a
precipitation reaction. I'm pretty sure that if you introduce solid CuSO4, NaOH, silica gel and some water in an autoclave you would likewise grow
these crystals (or other Cu/Na-silicates).
Also what does "complex" silicate mean? At least structurally this thing looks very simple.
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AJKOER
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On sources relating directly to the formation of copper silicate, I did find a limited mention of it on the atomistry.com website (link: http://nickel.atomistry.com/nickel_ore_smelting.html ). To quote:
"The ferric oxide present is reduced in the furnace by the sulphur of the pyrites to form ferrous oxide, which, in the presence of silica, forms a
slag: FeS + 3Fe2O3 + nSiO2 = SO2 + 7FeO.nSiO2. Any nickel monoxide which may be present reacts with an equivalent amount of ferrous sulphide to form
nickel sulphide and ferrous oxide, which in turn passes into slag.
The copper and nickel in the slag range up to about 0.4 per cent. The slags may be rejected, or part may be used again in similar smeltings, or in
later stages of the concentration process. The matte must contain sufficient iron to prevent nickel passing into the slag.
The precious metals in the ore accumulate in the matte, and in the latter case, there were present 1.90 ozs. of silver, 0.35 oz. of platinum, and 0.35
oz. of palladium per ton. According to G. P. Schweder, the sulphur in the matte is present as mono-sulphides of silver, copper (ous), nickel, and
iron; and if insufficient sulphur is present to form NiS, and FeS, the excess of metal dissolves in the molten sulphide. Any nickel silicate which may
be formed is decomposed by the iron sulphide to form iron silicate and nickel silicate, and so long as enough iron sulphide is present, only a very
small proportion of nickel can pass into the slag - prills of matte may be imprisoned in the slag if its viscosity in the furnace is too great. If the
ores have been over-roasted nickel will appear in the slag, and in that case some unroasted ore is mixed with the furnace charge. Copper silicate
behaves like nickel silicate, but the cobalt silicate does not react so easily with the iron sulphide, and when cobalt silicate is produced, it will
pass into the slag. "
Per this statement, Copper silicate may actually form (but still not definite in my opinion) via the direct synthesis of the metal oxide and SiO2 in a
furnace as follows:
CuO + nSiO2 ---> CuO.nSiO2
This source ("Compositional analysis of copper-silica precipitation tube", link: http://www.ncbi.nlm.nih.gov/pubmed/17164892 ) does not cite its creation in Silica gardens just amorphous silica and copper(ii) hydroxide. To
quote:
"Silica gardens consist of hollow tubular structures that form from salt crystals seeded into silicate solution. We investigate the structure and
elemental composition of these tubes in the context of a recently developed experimental model that allows quantitative analyses based on
predetermined reactant concentrations and flow rates. In these experiments, cupric sulfate solution is injected into large volumes of waterglass. The
walls of the resulting tubular structures have a typical width of 10 microm and are gradient materials. Micro-Raman spectroscopy along with energy
dispersive X-ray fluorescence data identify amorphous silica and copper(ii) hydroxide as the main compounds within the inner and outer tube surfaces,
respectively. Upon heating the blueish precipitates to approximately 150 degrees C, the material turns black as copper(ii) hydroxide decomposes to
copper(ii) oxide. Moreover, we present high resolution transmission electron micrographs that reveal polycrystalline morphologies."
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unionised
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| Quote: Originally posted by turd | | Quote: Originally posted by unionised |
Similarly, since the complex silicate in that paper is made by precipitation, it must be less soluble than copper hydroxide. |
In which paper? Not in the one you posted. There it is made by a classical hydrothermal reaction (10 <i>days</i> at 230° under autogenous
pressure - this is not a precipitation reaction). Hydrothermal in alkaline medium is the method to grow these huge 30 cm quartz crystals. With
hydrothermal you can't simply argue via solubility. It's a crazy world of its own.
God knows what the gel looks like - it certainly is not a defined compound and highly depends on the gel growth conditions, as usual with sol/gel
chemistry. Not that it matters once you autoclave it in alkaline medium.
| Quote: | | Also, do you agree with my assessment that they didn't actually analyse the gel and that you can't use that composition as evidence of anything?
|
Yes indeed, they put the <i>whole</i> thing in an autoclave and give the molar ratios for reproducibility. Again: this is NOT a
precipitation reaction. I'm pretty sure that if you introduce solid CuSO4, NaOH, silica gel and some water in an autoclave you would likewise grow
these crystals (or other Cu/Na-silicates).
Also what does "complex" silicate mean? At least structurally this thing looks very simple. |
May I invite you to consider what would happen to the silicate if it was soluble and someone autoclaved it in water for days at high temperature?
The components of the product got there somehow. Are you asserting that they were not in solution while in transit?
Didn't they precipitate?
Re " With hydrothermal you can't simply argue via solubility. "
I think you will find that I can.
If it wasn't of lower solubility then it wouldn't form.
In this context I'm using the word "complex" to refer to the fact that the lattice contains sodium ions as well as copper ions. It's not a simple
copper silicate.
I agree with you about this bit
"God knows what the gel looks like - it certainly is not a defined compound"
That's also true of glass.
Are you saying glass isn't a silicate?
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unionised
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| Quote: Originally posted by turd | | Quote: Originally posted by unionised |
Similarly, since the complex silicate in that paper is made by precipitation, it must be less soluble than copper hydroxide. |
In which paper? Not in the one you posted. There it is made by a classical hydrothermal reaction (10 <i>days</i> at 230° under autogenous
pressure - this is not a precipitation reaction). Hydrothermal in alkaline medium is the method to grow these huge 30 cm quartz crystals. With
hydrothermal you can't simply argue via solubility. It's a crazy world of its own.
God knows what the gel looks like - it certainly is not a defined compound and highly depends on the gel growth conditions, as usual with sol/gel
chemistry. Not that it matters once you autoclave it in alkaline medium.
| Quote: | | Also, do you agree with my assessment that they didn't actually analyse the gel and that you can't use that composition as evidence of anything?
|
Yes indeed, they put the <i>whole</i> thing in an autoclave and give the molar ratios for reproducibility. Again: this is NOT a
precipitation reaction. I'm pretty sure that if you introduce solid CuSO4, NaOH, silica gel and some water in an autoclave you would likewise grow
these crystals (or other Cu/Na-silicates).
Also what does "complex" silicate mean? At least structurally this thing looks very simple. |
May I invite you to consider what would happen to the silicate if it was soluble and someone autoclaved it in water for days at high temperature?
The components of the product got there somehow. Are you asserting that they were not in solution while in transit?
Didn't they precipitate?
Re " With hydrothermal you can't simply argue via solubility. "
I think you will find that I can.
If it wasn't of lower solubility then it wouldn't form.
In this context I'm using the word "complex" to refer to the fact that the lattice contains sodium ions as well as copper ions. It's not a simple
copper silicate.
I agree with you about this bit
"God knows what the gel looks like - it certainly is not a defined compound"
That's also true of glass.
Are you saying glass isn't a silicate?
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blogfast25
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One thing I think we can ALL agree on: PH11's precipitate is almost certainly not a defined silicate.
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unionised
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I think we can also agree that, while the solubility of copper hydroxide is small, the solution will contain hydroxide ions.
And I think we can also agree that hydroxide ions are a strong enough base to deprotonate silica.
So I think we can agree that at least some silicate ions will be present in the product.
Also, at least in principle, given time, if there's a copper silicate with a low enough solubility, it will be produced by the reaction, even if the
initial products are copper hydroxide and silica gel.
That reaction will happen more quickly if you autoclave the mixture.
I think we can also all agree that not all silicates are definite ones.
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turd
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| Quote: Originally posted by unionised | May I invite you to consider what would happen to the silicate if it was soluble and someone autoclaved it in water for days at high temperature?
The components of the product got there somehow. Are you asserting that they were not in solution while in transit?
Didn't they precipitate? |
Are you trying to play semantic games and tell me that hydrothermal reactions are precipitations because some ions traveled through water? Not only
would you be the only person to use it that way, you would also be wrong. If nothing changed in the last few years hydrothermal is not yet well
understood. But it is more akin to an aging than a precipitation.
| Quote: | Re " With hydrothermal you can't simply argue via solubility. "
I think you will find that I can.
If it wasn't of lower solubility then it wouldn't form. |
By hydrothermal treatment you get phases that you will never observe by simple precipitation, because in the latter some components crash out
and refuse to react (at least at appreciable rates during our lifetime). That's the whole point of this method.
| Quote: | | Are you saying glass isn't a silicate? |
I'm saying that your reasoning is not sound at all. I took some of the precipitate from my experiment above and calcinated it at 800°C over night.
Guess what? A fine black powder. This seems to support the more intuitive scenario: copper hydroxide is precipitated and then due to lowered pH, there
is slow gelling. Due to quick filtration I seem to have gotten not much silica. If you insist I can ask for an analysis of the precipitate as
well as the calcinated powder. Whether a putative gel includes copper in the framework or only as inclusion - who cares. It is irrelevant for a
subsequent hydrothermal treatment / calcination.
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AJKOER
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OK, more physical chemistry may help here. Per my research above, what we have with respect to copper silicate is a so called mixed oxide, CuO.nSiO2.
The problem here, as I see it, is the naming convention as historically mixed oxides are not expressly identified as such. As a common example, to
quote from Wikipedia (http://en.wikipedia.org/wiki/Aluminate ):
"for example the formula of anhydrous sodium aluminate NaAlO2 would be shown as Na2O.Al2O3."
Also, to quote another example:
"the mineral spinel itself, MgAl2O4 are mixed oxides with cubic close packed O atoms and aluminium Al3+ in octahedral positions.[7]"
which I would express as MgO.Al2O3. Cement is yet another example.
Now, in the instance of the Silica garden, we have, I suspect, Cu(OH)2.SiO2. One might view this as an example of a mixed oxide and hydroxide (or, in
the case of some metals like Iron in place of Copper where Fe(OH)3, for example, is better express as Fe2O3.xH2O, reference: "Concise Encyclopedia
Chemistry" by DeGruter, as a mixed oxide hydrate).
In my opinion, the term "silicate" can be misleading name as to quote Wikipedia (http://en.wikipedia.org/wiki/Silicate ) on Silicates:
"A silicate is a compound containing an anionic silicon compound. The great majority of silicates are oxides"
Also, "Silicates are well characterized as solids, but are less commonly observed in solution. The anion SiO4 4- is the conjugate base of silicic
acid, Si(OH)4, and both are elusive as are all of the intermediate species. Instead, solutions of silicates usually observed as mixtures of condensed
and partially protonated silicate clusters"
So, in my opinion, what we are more likely addressing with respect to solids are mixed oxides (to which the term silicate has been applied in the case
of Copper silicate) or mixed oxide/hydroxides for compounds formed at more ambient temperatures (for which the term 'silicate' is apparently not
employed).
[Edited on 17-9-2013 by AJKOER]
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turd
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Your post has nothing to do with physical chemistry. WTF?
| Quote: | | Per my research above, what we have with respect to copper silicate is a so called mixed oxide, CuO.nSiO2. |
Everybody with the most superficial education in inorganic chemistry knows that.
| Quote: | | Now, in the instance of the Silica garden, we have, I suspect, Cu(OH)2.SiO2. One might view this as an example of a mixed oxide
hydrate. |
Wrong. Read the article above - what we have is something like Cu(OH)2 + SiO2.nH2O.
Anyway... I got my precipitate and the calcinated product analyzed: The former is practically single phase
Cu4(SO4)(OH)6 a.k.a brochantite and some amorphous unidentified product (presumably silica gel). The latter is CuO
with minor amounts of unidentified crystalline material. Not really surprising.
Now the stubborn will say: but Pinkhippo11 used completely different concentrations! Still there is nothing whatsoever suggesting that he got a
silicate. Rather the opposite:
Na2Cu2Si4O11: Hydrated green, anhydrous black (references above).
Dioptase: Natural green, synthetic blue, anhydrous black (Z. Kristallogr. 187 (1989), 15-23)
Shattuckite: blue / green (wikipedia)
Now contrast this to:
White/gray? Sounds more like the dehydration of CuSO4, than anything.
Also it seems somewhat dubious that the silicate dissolves so easily in HCl.
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AJKOER
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To quote a reference (http://en.wikipedia.org/wiki/Physical_chemistry) "Predicting the properties of chemical compounds from a description of atoms and how they bond is
one of the major goals of physical chemistry" and, in the current context, I personally am better able to predict the reaction of say an anhydrous
sodium aluminate with a compound (like cold aqueous NH4Cl) when it is represented as Na2O.Al2O3, rather than as NaAlO2.
And, what is the difference between "Cu(OH)2 + SiO2.nH2O" and "Cu(OH)2.SiO2.nH2O" ?
Also, I suspect Na2Cu2Si4O11 is better represented as Na2O.2CuO.4SiO2, a triple mixed oxide! Further, if you accept that a solid 'silicate' here is
more likely a mixed oxide, then by your own reference, one could call this compound a sodium copper 'silicate', or a mixed salt of sodium silicate and
copper silicate (although I still prefer calling it a mixed triple oxide of Sodium, Copper and Silicon).
In other words, the aqueous silicate anion is unlikely and on heating most likely just mixed oxides.
[Edited on 17-9-2013 by AJKOER]
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blogfast25
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| Quote: Originally posted by turd |
This seems to support the more intuitive scenario: copper hydroxide is precipitated and then due to lowered pH, there is slow gelling. Due
to quick filtration I seem to have gotten not much silica. If you insist I can ask for an analysis of the precipitate as well as the calcinated
powder. Whether a putative gel includes copper in the framework or only as inclusion - who cares. It is irrelevant for a subsequent hydrothermal
treatment / calcination. |
It could have been useful to check the filtrate for silica.
| Quote: Originally posted by AJKOER | […], I personally am better able to predict the reaction of say an anhydrous sodium aluminate with a compound (like cold aqueous NH4Cl) when it is
represented as Na2O.Al2O3, rather than as NaAlO2
[Edited on 17-9-2013 by AJKOER] |
How? Neither notations (Na2O.Al2O3 or NaAlO2) shed much light on actual structure or chemical bonds.
‘Cold aqueous NH4Cl’ isn’t a compound: it is two compounds that form a system with its own equilibria (deprotonation of water, dissociation of
ammonium chloride, deprotonation of ammonium ions).
Anhydrous aluminate, when mixed with this system would quickly revert to its hydrated from: Na<sup>+</sup>(aq) +
Al(OH)<sub>4</sub><sup>-</sup>(aq). The latter would in turn neutralise the weakly acidic ammonium ions:
Al(OH)<sub>4</sub><sup>-</sup>(aq) + NH<sub>4</sub><sup>+</sup>(aq) === >
Al(OH)<sub>3</sub>(s) + NH<sub>3</sub>(g, aq) + H<sub>2</sub>O(l).
How does representing sodium aluminate as NaAlO2 make you “better able to predict the reaction of say an anhydrous sodium aluminate with a compound
(like cold aqueous NH4Cl) when it is represented as Na2O.Al2O3, rather than as NaAlO2”??
The notation (e.g.) Na2O.Al2O3 serves stoichiometry more than anything else.
[Edited on 17-9-2013 by blogfast25]
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AJKOER
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Blogfast:
Agree to some extent with your comment, but with anhydrous sodium aluminate prepared by heating NaAl(OH)4, I could expect something different keeping
Na2O.Al2O3 in mind. The aluminum oxide may have become more resistance, with direct heating, to dissolving in a weak base (see 'Concise Encyclopedia
Chemistry' by deGrupter on Al(OH)3 and Al2O3).
Now, your expectations are to quote:
"The latter would in turn neutralise the weakly acidic ammonium ions: Al(OH)4-(aq) + NH4+(aq) === > Al(OH)3(s) + NH3(g, aq) + H2O(l)"
while I would not be surprised if I immediately saw a precipitate of Al2O3, which in time may only slowly dissolve into the clear jelly like Al(OH)3
in a solution of ammonia and NaCl.
[Edited on 17-9-2013 by AJKOER]
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unionised
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The discussion had become longer than the experiment, so I checked.
2.12 grams of nominally Na2SiO3.5H2O were dissolved in 25 ml of deionised water and the solution left to stand overnight in order for any
polysilicates to hydrolyse.
2.49 grams of crystalline copper sulphate were dissolved in 25 ml of deionised water.
The two solutions were mixed and formed a blue gel.
Aproximately 0.5 ml of that gel were mixed with 5 ml of water in a test tube and the liquid heated to boiling.
No visible change took place.
By way of comparison a similar precipitate of copper hydroxide was prepared, diluted, and heated.
As expected it turned black on heating (before the solution reached boiling point) due to the production of copper oxide.
Copper hydroxide turns black on boiling.
The blue ppt from Cu++ and "SiO3--" ions doesn't
That ppt isn't Copper hydroxide and , if it contains copper hydroxide, the quantity is too small to affect the colour on boiling.
So, for example, the reply to Ajoker's question "what is the difference between "Cu(OH)2 + SiO2.nH2O" and "Cu(OH)2.SiO2.nH2O" ?" might be that only
one of them is still blue after you boil it.
Furthermore, a second portion of the diluted blue gel was treated with aqueous ammonia.
The material became more blue.
This was left to settle.
The supernatant was almost colourless, but the precipitate was deep blue.
If copper hydroxide had been present it would have dissolved in the ammonia solution and the solution would have been blue, but the precipitate would
have been pale or white.
God knows what the structure of the blue gel is, but it's not copper hydroxide and I think it has as good a claim to be called a silicate as my
windows have.
[Edited on 17-9-13 by unionised]
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