Sciencemadness Discussion Board

copper silicate?

vampirexevipex - 30-3-2012 at 14:15

Hello today i been surfing in the internet and came across a particular compound, Copper Silicate. I searched for its preparation and found little to no information about it. So, is it possible to make copper silicate from copper and silicate? If it is, then how?

[Edited on 22-02-12 by vampirexevipex]

weiming1998 - 30-3-2012 at 16:52

Quote: Originally posted by vampirexevipex  
Hello today i been surfing in the internet and came across a particular compound, Copper Silicate. I searched for its preparation and found little to no information about it. So, is it possible to make copper silicate from copper and silicate? If it is, then how?

[Edited on 22-02-12 by vampirexevipex]


1, What is silicate? Do you mean silicon dioxide?

2, Copper silicate cannot be made from copper and silicon dioxide because copper cannot react with non-oxidizing acids

3, The way to make copper silicate is to probably precipitate it (if possible) with a soluble copper salt, like CuCl2 and sodium silicate. If that doesn't work, then the reaction of Cu(OH)2 with hydrated SiO2 probably will.

[Edited on 31-3-2012 by weiming1998]

AJKOER - 30-3-2012 at 17:56

I would try reacting Sodium Silicate, Na2SiO3, which is available per Wiki as an aqueous solution, with CuCl2:

Na2SiO3 (aq) + CuCl2 (s) --> 2 NaCl (aq) + CuSiO3 (s)

where I am assuming that CuSiO3 will have low solubility.

If you are able to work with molten salts, per Wiki:

"Sodium carbonate and silicon dioxide react when molten to form sodium silicate and carbon dioxide:[1]

Na2CO3 + SiO2 → Na2SiO3 + CO2"

weiming1998 - 30-3-2012 at 18:32

A non-molten salt way of making sodium silicate is to react silica gel (found in food packaging, etc) with a concentrated solution of NaOH.

vampirexevipex - 30-3-2012 at 18:53

Thanks in advance! Copper acetate will work aswell right?
Cu(OAc)2 + NaSiO3 -> CuSiO3 + Na(OAc)2

[Edited on 22-02-12 by vampirexevipex]

jamit - 30-3-2012 at 20:34

@vampirexevipex

Can I ask why you want to make copper silicate? There's not much info on this chemical but what's its use in a lab or commercially?

vampirexevipex - 31-3-2012 at 04:24

Quote: Originally posted by jamit  
@vampirexevipex

Can I ask why you want to make copper silicate? There's not much info on this chemical but what's its use in a lab or commercially?


That's for me to know, and you to find out... shhhhhhh...

barley81 - 31-3-2012 at 04:45

Mix solutions of sodium silicate and copper sulfate. Both are cheap and readily available. Use dilute solutions because the solutions don't mix very well if they're concentrated (a layer of copper silicate will form between the two solutions). The insolubility of copper silicate and other metal silicates allows the chemical garden experiment to work.

bbartlog - 31-3-2012 at 04:49

It sounds none too stable to me. Should fall apart into CuO and SiO2 at the drop of a hat, assuming it exists at room temperature at all.

blogfast25 - 31-3-2012 at 05:52

Whether or not a copper silicate will precipitate from a copper (II) solution and waterglass (Na2SiO3) will depend on the relative solubilities of Cu(OH)2 and 'CuSiO3' (assuming the copper silicate is a simple one) because a Na2SiO3 solution is strongly alkaline. It's worth a try.

Alternatively you might want to try adding waterglass to a solution of cuprate. Cu(OH)2 is slightly amphoteric and dissolves in strong alkali to form Cu(OH)<sub>4</sub><sup>2-</sup> (cuprate) anions (cobalt blue). If adding waterglass to a cuprate solution yields a precipitate then that is likely to be 'a' silicate (but it could be a basic silicate).

[Edited on 31-3-2012 by blogfast25]

Nicodem - 31-3-2012 at 06:08

Quote: Originally posted by AJKOER  
I would try reacting Sodium Silicate, Na2SiO3, which is available per Wiki as an aqueous solution, with CuCl2:

Na2SiO3 (aq) + CuCl2 (s) --> 2 NaCl (aq) + CuSiO3 (s)

where I am assuming that CuSiO3 will have low solubility.

And just what would prevent the expected product, Cu(OH)<sub>2</sub>(s), to precipitate instead? If anybody has references for anything like this then post them please. Otherwise stop spreading misinformation.
Copper silicate (which one anyway?) is most likely prepared the way such compounds are prepared, by the solid-solid reaction between CuO and SiO<sub>2</sub> (in the form of silicagel). It probably requires several millings and repeated calcinations before a reaction with an acceptable conversion is achieved. I will not even bother looking for references, as I see that nobody even cares for actual data.

blogfast25 - 31-3-2012 at 10:37

Nicodem:

A tad harsh, I feel.

I did look for some references and just Wiki mentions "CuSiO3·H2O (also reported as CuSiO2(OH)2)", strongly suggesting synthesis from watery solution is possible. Barley81's reference to the 'chemical garden' is also relevant, as the formation of these colourful 'stalagmites' is believed to be due to the formation of insoluble d-block metal silicates.

What would prevent the formation of cupric hydroxide instead of cupric silicate? If the solubility product of the silicate is much lower than that of the hydroxide, the silicate will preferentially precipitate. It's a plausible hypothesis looking for evidence.

barley81 - 31-3-2012 at 11:14

A long time ago, I mixed waterglass and copper sulfate solutions. A hard, crunchy precipitate formed. The texture of it was vastly different from copper basic sulfate (made by mixing copper sulfate and sodium hydroxide). This is evidence that a different compound was formed, and likely contained silicate ions.

Nicodem - 31-3-2012 at 13:18

Quote: Originally posted by blogfast25  
Nicodem:

A tad harsh, I feel.

The harshness is on place. I keep seeing more and more threads like this, starting from zero effort and being immediately derailed into idle speculation based on zero references and zero thought. Always by the same usual suspects. I hate it when this occurs to a potentially interesting topic.

Anyway, to keep on the topic... What barley81 describes above is probably the only thing one can expect from such a reaction. If you throw a crystal of CuSO4.5H2O into a diluted solution of waterglass it starts growing blue gelatinous worm-like structures as the diffusion of ions progresses trough them. I did this as a kid and it was quite entertaining. Other coloured similarly acidic salts do the same. I would not call that insoluble stuff copper silicate though. It would be expected to be silicagel entraining copper(II) hydroxide if considering the reaction equation. Doubtlessly some Cu-O-Si bonds might be there at the silicagel surface, but that is far from a compound. Just consider the isoelectronic carbonates. Na2CO3 in water is less basic than Na2SiO3 yet as far as I know nobody ever succeeded to prepare CuCO3 by its reaction with CuSO4(aq). All you get is Cu(OH)2×CuCO3. I don't even know how CuCO3 is prepared, if it is a known compound at all, but I'm quite sure it is not made by aqueous precipitation. Same goes for other salts of acidic metal ions with anions of poor acids. Acids generally do not like to combine with acids to give salts in aqueous media (unless coordination prevails). They require different, more forcing conditions.

Eddygp - 31-3-2012 at 13:32

Although this looks more like a strange hypothetical reaction to form a possible compound, it is possible to look at if from another point of view.

It of course depends on which silicate you want to do ( "copper(I) silicate", "copper (II) silicate" ). As this is not the case, I believe, it would be useful to look at the formation of other metal silicates, to have ideas.

However, I think the best way would be to put copper(II) nitrate or copper(II) chloride in silicic acid.

EDIT: Wouldn't it be copper silicide? It is far more documented.

[Edited on 31-2012-3 by Eddygp]

Nicodem - 3-4-2012 at 08:20

Barley81 brought to our attention the reaction of acidic colored metal salts with waterglass. I was unaware that it is actually a common demonstration and has a name: chemical gardens reactions. Those who ever tried it, know that it is a nice visual demonstration with "colored things growing", especially suitable for kids (it was impressive to me when I was a kid).

I checked a little bit of the literature on it. The products are not metal silicates, so this is not really directly related to this thread, but nobody seems interested in copper silicates anyway, so I'll stay off topic. There is an educational and review article on the topic that is worth reading (J. Colloid Interface Sci. 2002, 256, 351–359; freely available here). The mechanism of growth of these structures is quite interesting, relying on diffusion through semi-permeable membranes that the products form (this strongly influences their complex microscale morphology). The composition of the precipitates depend on the reactants: "The compositions of two silicate-garden precipitates have recently been studied in detail: those formed from aluminium nitrate [15,16], and from copper nitrate [17]. The former precipitate is a material with a hierarchical structure on the nanoscale, consisting of silica nanotubes clustered together over several orders, surrounded by aluminosilicate and aluminium hydroxide. The latter precipitate is more crystalline, being formed in part of crystalline copper hydroxide nitrate, together with amorphous silica."

Read also:

http://www.rsc.org/Publishing/ChemScience/Volume/2007/01/che...

For pictures see http://chcscienceandmath.blogspot.com/2009/08/crystal-garden...

... a weird one from DOI:10.1209/0295-5075/89/44004

Zephyr - 11-9-2013 at 20:06

I've found a reference to copper silicate in The Golden Book of Chemistry Experiments
"Making Silicates:
1. Dilute 5ml of water glass (Na2SO3) with 5ml of water.
2. dissolve small crystal of copper sulfate in water.
3. Add a few drops of the water glass to get a blue precipitate of copper silicate."

It's strange that there isn't any other references to copper silicate, but after reproducing this experiment I got a light blue precipitate a promised :)
the golden book of chemistry
http://openmaterials.org/cache/The%20Golden%20Book%20of%20Ch...
(pg 55, lower left)

bbartlog - 11-9-2013 at 21:00

Although it is not copper silicate but rather calcium copper silicate, the ancient pigment known as Egyptian Blue seems relevant to the original question: http://en.wikipedia.org/wiki/Egyptian_blue

blogfast25 - 12-9-2013 at 04:46

Quote: Originally posted by Pinkhippo11  
3. Add a few drops of the water glass to get a blue precipitate of copper silicate."

It's strange that there isn't any other references to copper silicate, but after reproducing this experiment I got a light blue precipitate a promised :)


What this proves is that you get a blue precipitate in the conditions described. But copper hydroxide is blue as well. At a very minimum you'd have to prove there is silica in your precipitate, as evidence the blue precipitate is a silicate of sorts.

Filter off the precipitate and wash it carefully. The following tests would be useful:

1. Heat it: if it goes brown/black it's likely to be Cu(OH)2, which dehydrates quickly by heat to CuO (black). If it goes blue it's likely to be a silicate of copper

2. Carefully add weak acid to it: it should split into a blue solution (the copper salt of the acid used) and a gelatinous whitish hydrate of silica, if it is a copper silicate of sorts. If it dissolves completely to a blue solution, it's likely to have been Cu(OH)2

The Golden Book of Chemistry is interesting but not really an authorative source on chemistry...

[Edited on 12-9-2013 by blogfast25]

Zephyr - 12-9-2013 at 21:09

Quote: Originally posted by bbartlog  
Although it is not copper silicate but rather calcium copper silicate, the ancient pigment known as Egyptian Blue seems relevant to the original question: http://en.wikipedia.org/wiki/Egyptian_blue

bbartlog, in the wiki article, the procedure for making Egyptian Blue(calcium copper silicate) calls for silica sand,copper carbonate,and calcium carbonate and temperatures of 800-1000oC!
Cu2CO3(OH)2 + 8 SiO2 + 2 CaCO3 → 2 CaCuSi4O10 + 3 CO2 + H2O

When making copper silicate you simply mix sodium silicate and water glass
CuSO4+Na2O3Si -> CuO3Si+Na2SO4
this doesn't involve calcium and does not require heat initiation like when making the said "Egyptian Blue".
Quote: Originally posted by blogfast25  
What this proves is that you get a blue precipitate in the conditions described. But copper hydroxide is blue as well. At a very minimum you'd have to prove there is silica in your precipitate, as evidence the blue precipitate is a silicate of sorts
in answer to your question blogfast25, the blue solid I described didn't decompose upon heating so it isn't copper carbonate or copper hydroxide.
Also in the "conditions described" what other blue precipitate could form?
Almost all copper compounds are blue/green so the color isn't a reliable point on which to base an argument.

Nicodem - 13-9-2013 at 08:31

Quote: Originally posted by Pinkhippo11  
When making copper silicate you simply mix sodium silicate and water glass
CuSO4+Na2O3Si -> CuO3Si+Na2SO4
this doesn't involve calcium and does not require heat initiation like when making the said "Egyptian Blue".

That equation makes no sense chemically. You can't just pretend that reactions in aqueous solutions don't involve water. You can't ignore pKa values either. I suggest you to read the whole thread and the cited references. All this was already discussed up-thread. There are good reasons for such products being synthesized by calcination based solid-solid reactions.
Quote:
Also in the "conditions described" what other blue precipitate could form?

Firstly, copper silicate is the last thing one would expect. Secondly, plenty of other blue precipitates could form, like the materials described in the articles cited above. You would need a powder XRD analysis to prove the presence of copper silicate in that precipitate. Just wishing or expecting it to be what you want is not enough.

blogfast25 - 13-9-2013 at 11:56

PH11:

Whether or not a copper silicate can exist is a different question from to whether such a product can be made from mixing a (watery) solution of sodium silicate and a (watery) solution of a soluble copper salt.

Solutions of waterglass are quite strongly alkaline, so they contain a lot of hydroxide (OH<sup>-</sup>;) ions. With Cu<sup>2+</sup> ions these form a blue, insoluble, flocculant precipitate:

Cu<sup>2+</sup>(aq) + 2 OH<sup>-</sup> === > Cu(OH)<Sub>2</sub>(s) (very simply put)

The hydroxide ions come from reaction of the silicate with water. Since as there's plenty silicate there's in principle also enough hydroxide to cause Cu(OH)<Sub>2</sub> to precipitate out.

You say you heated up your precipitate – what were the heating conditions? Did you at a very minimum drive off most of the water?

By all means also add a weak solution of acid to your precipitate and see what happens. In acid conditions a freshly prepared silicate should yield hydrated silica.

You're quite a long way from proving what you obtained is a silicate of sorts (and which one).



[Edited on 13-9-2013 by blogfast25]

Zephyr - 13-9-2013 at 21:35

i apologize for not explaining my procedure more thoroughly.
here is my procedure:
1. 0.5 g of copper sulfate if dissolved in 5 ml of water.
2. 1 ml of sodium silicate is added to the copper sulfate solution, a thick blue precipitate is formed.
3. The precipitate is filtered and left to evaporate.

Upon heating at approximately 700oC, the blue solid expands and turns white/gray. water condenses on rim of test tube.

When 1.5 ml HCl is added, the blue solid becomes a bright green liquid. copper chloride maybe?

again sorry for the obscurity the first time, any help is appreciated!

blogfast25 - 14-9-2013 at 05:34

Quote: Originally posted by Pinkhippo11  

Upon heating at approximately 700oC, the blue solid expands and turns white/gray. water condenses on rim of test tube.

When 1.5 ml HCl is added, the blue solid becomes a bright green liquid. copper chloride maybe?



The second results strongly points to the precipitate being Cu(OH)2:

Cu(OH)2 + 2 HCl === > CuCl2 + 2 H2O. Solutions of CuCl2 are indeed green.

With a copper silicate, say hypothetically 'CuSiO3', the very weak silicic acid would become displaced by the strong HCl, very simply put:

CuSiO3 + 2 HCl === > CuCl2 + H2SiO3

The silicic acid should precipitate as a voluminous gel of SiO2.nH2O. You didn't observe that.

The second test result is more ambiguous. You claim to have heated at 700 C. But did you also heat to 700 C? In other words, what temperature did the precipitate reach, approximately? Cu(OH)2 should definitely yield black CuO on dehydration. But it's possible that your precipitate contained 'trapped' excess Na2SiO3. Then, on heating you might end up with something grayish.

I would suggest to repeat the test using equivalent amounts of CuSO4 and Na2SiO3, not "0.5 g and 1.5 ml" (apples and oranges!) So calculate how many moles is 0.5 g of CuSO4 (presumably you use the pentahydrate? CuSO4.5H2O) and then work out how much Na2SiO3 that is equivalent to. Make solutions of these and carefully mix them together.

Edit:

I’m probably barking up the wrong tree myself here. If, as I strongly suspect, mixing solutions of a copper salt with waterglass results in Cu(OH)2 to precipitate then silica MUST also precipitate at the same time, according to the following stoichiometry:

Cu<sup>2+</sup>(aq) + SiO<sub>3</sub><sup>2-</sup> (aq) + 2 H<sub>2</sub>O(l) === > Cu(OH)<sub>2</sub>(s) + H2SiO3, the latter manifests itself as hydrated silica (SiO2.nH2O)

This is due silicic acid being an extremely weak acid (pKa1 = 9.84, pKa2 = 13.2, Wiki), even weaker than carbonic acid (*). Hydrolysis of the SiO<sub>3</sub><sup>2-</sup> anion to HSiO<sub>3</sub><sup>-</sup> and subsequently to H<sub>2</sub>SiO<sub>3</sub> with release of OH<sup>-</sup> ions causes both Cu(OH)2 and hydrated silica to precipitate out in a 1:1 molar ratio.

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. It’s possible that when carefully washed (sulphate!), dried and subjected to prolonged heat and mechanical action such a mixture would fuse to something akin to ‘CuSiO3’ but only advanced measuring techniques like X-ray diffraction crystallography could shed light on that structure. Silicates can be excruciatingly complicated in structure, as the myriad of different silicate based minerals shows.

(*) in your 'Golden Book of Chemistry' is described an interesting little experiment that shows how silicic acid can be displaced from waterglass by carbonic acid, proving in a simple way that the latter is a stronger acid than the former.

[Edited on 14-9-2013 by blogfast25]

unionised - 14-9-2013 at 07:28

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.

It's probably a complex copper silicate something like this
http://www.ncbi.nlm.nih.gov/pmc/articles/PMC2960335/
and the assertion that it's a mixture of silica gel and copper hydroxide is baseless.

A colloidal "solution" of silica gel is a commercial product used in wine making, so there's no problem with a lack of a silica precipitate.

turd - 14-9-2013 at 08:54

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. :o

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.

blogfast25 - 14-9-2013 at 09:12

@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]

unionised - 14-9-2013 at 10:16

I still think "It's probably a complex copper silicate something like this"

Nicodem - 14-9-2013 at 12:27

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.

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.

blogfast25 - 14-9-2013 at 13:22

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]

unionised - 15-9-2013 at 03:51

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.


Where?
As far as I can see, they just cooked the stuff- they didn't check for Si-O-Cu bonds.

turd - 15-9-2013 at 06:21

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.

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. :P

Edit: CuSO4, not CaSO4 :p

[Edited on 15-9-2013 by turd]

Nicodem - 15-9-2013 at 07:21

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.
Quote: Originally posted by unionised  
Where?
As far as I can see, they just cooked the stuff- they didn't check for Si-O-Cu bonds.

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)
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unionised - 15-9-2013 at 08:30

"! 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.

blogfast25 - 15-9-2013 at 08:37

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]

unionised - 15-9-2013 at 12:02

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?

blogfast25 - 15-9-2013 at 12:31

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?


I couldn't find any reference to elemental analysis, I just kind of superficially assumed that they'd done it anyway. :(

turd - 15-9-2013 at 14:09

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.

AJKOER - 15-9-2013 at 15:18

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."

unionised - 16-9-2013 at 10:40

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?

unionised - 16-9-2013 at 11:02

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?

blogfast25 - 16-9-2013 at 11:25

One thing I think we can ALL agree on: PH11's precipitate is almost certainly not a defined silicate.

unionised - 16-9-2013 at 11:59

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.

turd - 16-9-2013 at 12:59

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.

AJKOER - 16-9-2013 at 19:00

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]

turd - 17-9-2013 at 03:52

Quote: Originally posted by AJKOER  
OK, more physical chemistry

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:
Quote: Originally posted by Pinkhippo11  
Upon heating at approximately 700oC, the blue solid expands and turns white/gray.

White/gray? Sounds more like the dehydration of CuSO4, than anything.

Also it seems somewhat dubious that the silicate dissolves so easily in HCl.

AJKOER - 17-9-2013 at 04:42

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]

blogfast25 - 17-9-2013 at 09:31

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]

AJKOER - 17-9-2013 at 10:42

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]

unionised - 17-9-2013 at 11:12

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]

blogfast25 - 17-9-2013 at 12:30

AJ:

The notation ‘Na2O.Al2O3’ is just that: an antiquated notation. Look at older texts: the formulas of many double salts that are known now to be coordination complexes were written in that format because at least it reflects molar ratio. But it says nothing about structure or bonding.

Mixing a solution of sodium aluminate and ammonium chloride solution precipitates flocculant Al(OH)3.nH2O. The aluminate ion in solution is likely more accurately described as:

[Al(OH)<sub>4</sub>(H<sub>2</sub>O)<sub>2</sub>]<sup>-</sup>, which may lose that water quickly to become the tetrahedral [Al(OH)4]- anion.

It’s likely to be surrounded by more loosely bound water (solvation), causing a highly hydrated gel to precipitate when a proton is absorbed.

Anhydrous sodium aluminate would react with such an ammonium chloride solution much in the same way, in the sense that it would have to re-hydrate and solubilise at the surface first. The results would be gel-like Al(OH)3 too. Nowhere does Al2O3 come into it.

‘Concise Enyclopedias’ have their uses but are generally not so great in dealing with ‘the devil that’s in the detail’.



[Edited on 17-9-2013 by blogfast25]

blogfast25 - 17-9-2013 at 12:38

Unionised:

That's very interesting. I assume the molar ratio CuSO4.5H2O:Na2SiO3.5H2O is 1:1 (don't have time to check right now)?

Those solutions are fairly strong (0.4 M), did the whole mixture 'solidify'?

It's quite intriguing that a bit of the gel in water doesn't change appearance on boiling.

Interesting result with the ammonia too!

I want to replicate your experiment, perhaps with some photos too...

[Edited on 17-9-2013 by blogfast25]

unionised - 17-9-2013 at 13:07

I was aiming for 1:1 though the silicate isn't exactly "analytical reagent".
http://www.amazon.co.uk/CHILTERN-CONNECTIONS-DEVELOPER-UNIVE...
The mixture was viscous, but not solid.

turd - 17-9-2013 at 13:15

Quote: Originally posted by unionised  
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.

Did you miss the part where I had my precipitate analyzed and it was brochantite? I don't know if this forms ammonia complexes, but it is stable up to 250°C.

Why all the mumbo-jumbo about the precipitate not being Cu(OH)2, when it's not expected to be Cu(OH)2?

AJKOER - 17-9-2013 at 16:19

OK, I was hoping that the chemistry surrounding silicates wasn't that complex or extensive.

However, after reviewing the material at Atomistry,com (link: http://silicon.atomistry.com/silicates.html ), I am clearly way off the mark.

ElectroWin - 17-9-2013 at 18:19

YES! should we be surprised, though? (since Si is in the same group as C)

blogfast25 - 18-9-2013 at 05:07
Quote: Originally posted by turd  
Why all the mumbo-jumbo about the precipitate not being Cu(OH)2, when it's not expected to be Cu(OH)2?


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.


Hmmm...


turd - 18-9-2013 at 07:17

Hey now, that was posted before I got the analysis. Expectation can change...

But your criticism is right - I was an idiot for suggesting copper hydroxide, since Nicodem's paper said the precipitate is a basic nitrate, so one would likewise expect a basic sulfate. I sugarcoat it by pretending that I meant a copper hydroxide, not copper hydroxide. :P

Anyway I made a further experiment which convinced me even more that PinkHippo didn't actually obtain a silicate. The first time I tried to simulate PinkHippo's experiment by pouring silicate in an excess copper sulfate. This time I wanted to do it unionised-style: copper sulfate in an excess/stoichiometric amount of silicate.

The silicate solution was made by dissolving 3 g NaSiO3 in 30 ml warm H2O. This was stirred for a few h at ~50°C for complete hydrolysis. Then 3 g CuSO4.5H2O in 30 ml H2O was added (still at ~50°C in the hope that this would speed up gelling). Immediate precipitation of a nasty colloid (note: PH11 got a filterable precipitate). The blue "gel" was suction filtered (took quite long), washed with H2O and dried at 60°C. Filtrate clear, colorless.

A part of the precipitate was heated to red heat over a Bunsen burner. It turned first green, then black (as I had expected for a copper silicate, but see below). Contrast PH11: white/gray.

Digestion with 25% NH3 gave the characteristic copper complex (in contrast to unionised)?

Analysis of the precipitate: completely amorphous.

Analysis of the heated precipitate: not well defined, but definitely contains CuO. Also very likely tridymite/SiO2 and at least one other unidentified phase, presumably a copper silicate.

Conclusion 1: just because your gel is completely amorphous doesn't mean that it's a copper silicate.

Conclusion 2: PH's experiment resembled more closely my first run where I got well defined basic copper sulfate. (If it isn't completely fake.)

Outlook: I'm tempering the precipitate for 18 h at 900°C - let's see if the finely dispersed SiO2 and the CuO react to a silicate and whether the other phases are better defined to allow for a definite assignment. Unfortunately I cannot do this much longer, since I don't have infinite access to the analytics.

blogfast25 - 18-9-2013 at 09:15

Turd:

I still think it could be worth looking at the filtrate too. Does it contain the sodium sulphate one would expect? Any 'free ' silica', perhaps?

On 'digestion with 25 % NH3', what precisely happened?

I'm preparing a 'Na2SiO3.5H2O' 0.4 M solution tonight for an experiment tomorrow.

Lucky you to have at least finite access to the analytics! It'd be great if you could spend a few words on the type of analysis that were performed.

[Edited on 18-9-2013 by blogfast25]

unionised - 18-9-2013 at 10:40

Quote: Originally posted by turd  
Quote: Originally posted by unionised  
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.

Did you miss the part where I had my precipitate analyzed and it was brochantite? I don't know if this forms ammonia complexes, but it is stable up to 250°C.

Why all the mumbo-jumbo about the precipitate not being Cu(OH)2, when it's not expected to be Cu(OH)2?


"I don't know if this forms ammonia complexes"
Well, you should.
Practically every kid who has done high school chemistry has taken a solution of copper sulphate and added ammonia solution gradually. The initial reaction produces a ghastly mess of "basic copper sulphate" and / or copper hydroxide.
That material then dissolves when an excess of ammonia is added.
If brochantite is formed by the reaction of copper sulphate and an alkali then it's definitely soluble in ammonia solution. If it isn't produced then nobody cares.

Incidentally, re "Analysis of the precipitate: completely amorphous." What sort of analysis?

[Edited on 18-9-13 by unionised]

turd - 18-9-2013 at 11:08

I should know a bazillion things that I have forgotten. Fortunately I don't have to - I just ask smartasses on the internet.
Quote: Originally posted by unionised  
Incidentally, re "Analysis of the precipitate: completely amorphous." What sort of analysis?

XRD, I thought that was completely obvious.

So, are you still sure that PH11 got a "copper silicate" after I am able to crash out crystalline Cu salts and grow amorphous gels which seem not to be copper silicates?

blogfast: On digestion with NH3, the precipitate became blue and slowly migrated into the liquid phase.

[Edited on 18-9-2013 by turd]

blogfast25 - 18-9-2013 at 12:27

I meant elemental analysis.

I've never been convinced that PH11's was anything like a copper silicate and expressed much scepticism about that. Now I'm not sure about anything with regards to copper silicates anymore...

unionised - 18-9-2013 at 13:09

Well, I still have a test tube with a blue precipitate at the bottom of a nearly colourless solution that smells strongly of ammonia.
It's clearly not copper hydroxide and it's not basic copper sulphate.
It was prepared by reaction of a solution of a silicate with a solution of copper sulphate.
What do you think it is?

turd - 18-9-2013 at 13:14

Hm? I think we're talking at cross-purposes...

I asked unionised who seemed to be dead set that PH11 got a silicate and yet couldn't give any convincing argument. To be honest I had from the beginning the suspicion that PH11 was fake - why dig up such an old thread? And the turning white on heating thing is dubious. OTOH I don't want to insinuate anything.

You are right that I should have had a closer look at the filtrate, but my time is very limited at the moment. And I can't do elemental analysis, at least not with an acceptable time effort.
Quote:
Now I'm not sure about anything with regards to copper silicates anymore...

Then simply don't call these undefined and uncharacterized gels copper silicates. ;) It's certainly not what the original poster had in mind. And I object to the comparison with glasses - both are amorphous but glasses are single-phase (in principle) and can be perfectly defined (think silica or some perfectly pure organics that form glasses).

turd - 18-9-2013 at 13:21

Quote: Originally posted by unionised  
Well, I still have a test tube with a blue precipitate at the bottom of a nearly colourless solution that smells strongly of ammonia.
It's clearly not copper hydroxide and it's not basic copper sulphate.
It was prepared by reaction of a solution of a silicate with a solution of copper sulphate.
What do you think it is?

I have no idea what this. But why do you insist that PH11 got exactly what you got and not what I got in two different experiments? What happens when you heat it to ~700°C? For all I know, a copper silicate should be dark green/black.

What do you think was the perfectly amorphous gel, that did dissolve in aq. NH3 and calcinated into SiO2 and CuO (+ other phases)?

[Edited on 18-9-2013 by turd]

turd - 19-9-2013 at 01:23

Quote: Originally posted by unionised  
What do you think it is?

Now that I managed to get the same thing I'm sure that this is a gel containing CuO and SiO2. It is certainly not anything like the Na/Cu silicate of the Acta Crystallogr. paper. This would not form such a nice tetramine complex. Relating the two would be like relating diamond and graphite because they're both composed of carbon.

[Edited on 19-9-2013 by turd]

blogfast25 - 19-9-2013 at 04:52

Quote: Originally posted by turd  
Quote:
Now I'm not sure about anything with regards to copper silicates anymore...

Then simply don't call these undefined and uncharacterized gels copper silicates.


Nowhere have I done that. I was very sceptical about PH11's results (I think he did actually do an experiment) and thought the precipitate to be an undefined mixture of Cu(OH)2 and hydrated silica gel.

It's unionised's result that had me wondering, when on heating his precipitate didn't turn black so easily. You claim it's a basic copper sulphate but other than your assertion we have precious little evidence for that either. Also, why would a basic copper sulphate form in these conditions? Right now we don't even know where the sulphate anions have gone because neither of you checked the filtrate for anything.

I'm very short of time too but hope to run a replica of unionised's run this week end, looking also at the filtrate if I can get some.

turd - 19-9-2013 at 05:08

When quoting you removed the smiley. I was kidding.

I admit the last posts where a bit chaotic because I'm posting besides real work. :o

To summarize:

- I tried to simulate PH11 conditions - I got basic copper sulfate. 100% certain
- unionised got a mixed CuO/SiO2 gel. He heated to <100°C, where basic copper sulfate does not decompose. But the ammonia test proves that there is none.
- I tried to reproduce unionised gel and got a 100% amorphous precipitate that *did* dissolve to good parts in NH3. I think this is remarkable. Google gives a few hits on amorphous Cu(OH)2, but so far I haven't had time to check them out.
- In another try I did get the CuO/SiO2 gel. But not as cleanly. I still get a distinctly blue solution on treatment with 25% NH3 but also a very nice amine-complex precipitate (or maybe blue solution in the gel?). To check I will have to wash multiple times with aq. NH3.

unionised - 19-9-2013 at 11:43

Puts on suit of armour.
Quote: Originally posted by unionised  
Well, I still have a test tube with a blue precipitate at the bottom of a nearly colourless solution that smells strongly of ammonia.
It's clearly not copper hydroxide and it's not basic copper sulphate.
It was prepared by reaction of a solution of a silicate with a solution of copper sulphate.
What do you think it is?


Well, it's possible that the answer is copper phosphate.
:mad:
I got suspicious when the results I got weren't repeatable. I'd not expect a perfect match because silicates are often a bit (well, a lot) variable.
But I thought I'd better check the sodium "silicate" I was using.
It doesn't give a precipitate with acid so, whatever it is, it's not a silicate.
It's a strong base, it was fairly cheap and it's not hygroscopic so it's not NaOH. It doesn't fizz with acid so it's not carbonate.
My guess is phosphate.
I have been ripped off.

Now, just as soon as someone answers my point about hydroxide being a strong enough base to react with silica ( to produce silicate ions) and the fact that copper hydroxide, while not very soluble, is a source of hydroxide ions to the extent that it dissolves and, given that there is at least one copper silicate species that has a very low solubility- lower than that of the hydroxide since it's formed in solution where hydroxide would be an alternative, then they will have ruled out the idea that you get a copper silicate by reacting copper sulphate solution with sodium silicate solution.

The amorphous product proves that it's not the crystalline one, but it doesn't rule out a silicate.

Raman spectrum anyone?
That should show up any Cu-O-Si groups


[Edited on 19-9-13 by unionised]

turd - 19-9-2013 at 18:58

Maybe PH11 got the same Na2SiO3 batch? That would explain the (green)grey color on heating?
Quote: Originally posted by unionised  
Now, just as soon as someone answers my point about hydroxide being a strong enough base to react with silica ( to produce silicate ions) and the fact that copper hydroxide, while not very soluble, is a source of hydroxide ions to the extent that it dissolves and, given that there is at least one copper silicate species that has a very low solubility- lower than that of the hydroxide since it's formed in solution where hydroxide would be an alternative, then they will have ruled out the idea that you get a copper silicate by reacting copper sulphate solution with sodium silicate solution.

The amorphous product proves that it's not the crystalline one, but it doesn't rule out a silicate.

You will find I never ruled out such a thing. On the contrary, I'm convinced such gels exist and if you don't get them by precipitation/ageing of water glass, then you will probably get them by the classical alkoxide route.

I object to your logic that because a highly crystalline silicate is obtained under hydrothermal condition you will get a homogenous SiO2/CuO gel (As opposed to a SiO2 gel with a few Si-O-Cu bonds at the surface) by precipitation. The hydrothermal experiment is no proof since autoclaving solid Na2SiO3, Cu(OH)2 and water will provide the same/similar silicates.

And I object to the equation of SiO2/CuO gels with the Na/Cu-silicate in the Acta paper. It's not about amorphous/crystalline, but about degree of condensation.

PS: What's up with the double posting?

[Edited on 20-9-2013 by turd]

[Edited on 20-9-2013 by turd]

unionised - 20-9-2013 at 09:39

He might, but I doubt it, for a start he said he was using water glass. The stuff I was using was a solid.

Since I have always accepted that the gel is amorphous it's hard to see why you think that I believe"you will get a homogenous SiO2/CuO gel (As opposed to a SiO2 gel with a few Si-O-Cu bonds at the surface) by precipitation. "

"The hydrothermal experiment is no proof since autoclaving solid Na2SiO3, Cu(OH)2 and water will provide the same/similar silicates. "
All it needs to prove is that there is a silicate with a low solubility.

As far as I can tell, the double posting is a glitch with the editing.

Do you agree that silica gel and Cu(OH)2 would form a silicate on autoclaving in water?

Anyway, as I said, Raman would be really helpful here.
Any takers?

kmno4 - 20-9-2013 at 10:03

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

blogfast25 - 20-9-2013 at 12:28

Unionised:

Are you sure your sodium metasilicate is not a silicate? You see, mine (advertised as ‘Na2SiO3.5H2O’) also didn’t give a precipitate with acid. But I’m trying that again tomorrow because I’m sure I’m doing something wrong.

In the mean time I replicated unionised’s experiment, with some add-ons. 0.01 mole of sodium metasilicate was dissolved in 25 ml and stood overnight. 0.01 mole of CuSO4.5H2O dissolved in 25 ml of water was then added to it. A blue, gelatinous precipitate formed.

Boiling this in a test tube caused no change in appearance whatsoever. By contrast some freshly prepared Cu(OH)2 quickly turned black on boiling. Precipitate after boiling test:



Adding 33 % NH3 to the gel caused it to dissolve and the copper diammine complex blue colour to appear.

The precipitate was then filtered off on Buchner and washed once with 25 ml of deionised water (DIW). The filtrate ran clear and colourless, at a paper pH of about 5. To the 50 ml of filtrate was added 0.01 mole of Ba(NO3)2 dissolved in 75 ml of DIW. A lot of white precipitate formed, this is it after about 30 min of standing:



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.

A small teaspoon of the Buchnered and washed precipitate was then loaded into a nickel crucible and heated on a 10 inch hot plate on maximum setting, for about 1 h. The filter cake and crucible:




Some whitish material seemed to seep out and the blue substance gradually turned green. After a bit of drying, first formation of green material:



After about 1 h of drying:



The crucible was then further heated on a high, blue Bunsen flame for about 20 minutes with lid on. After cooling it was clear that the green material had turned black.



[Edited on 20-9-2013 by blogfast25]

turd - 20-9-2013 at 13:13

blogfast25: Your experiment is very consistent with my second run. And probably also my third run. There I probably just had a more condensed gel making leaching with ammonia harder.
Quote: Originally posted by blogfast25  
Are you sure your sodium metasilicate is not a silicate? You see, mine (advertised as ‘Na2SiO3.5H2O’) also didn’t give a precipitate with acid. But I’m trying that again tomorrow because I’m sure I’m doing something wrong.

Drop the silicate solution in HCl. That's apparently how those drying beads are made.
Quote:
Adding 33 % NH3 to the gel caused it to dissolve and the copper diammine complex blue colour to appear.

Exactly what I got: amorphous copper salt dispersed in amorphous silica gel. :(
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? Respect.
Quote:
A small teaspoon of the Buchnered and washed precipitate was then loaded into a nickel crucible and heated on a 10 inch hot plate on maximum setting, for about 1 h. [...]Some whitish material seemed to seep out and the blue substance gradually turned green. After a bit of drying, first formation of green material:
[...]
The crucible was then further heated on a high Bunsen flame for about 20 minutes with lid on. After cooling it was clear that the green material had turned black.

Exactly the sequence that I have seen, just with a Bunsen burner it took around a minute. The black residue, as I stated above, was CuO, SiO2 and an unidentified phase. Presumably a copper silicate, as one would expect from heating finely dispersed silica and copper to red heat.

Meanwhile I calcinated the residue over night. Got a hard shiny graphite-y residue. So far I couldn't be assed to stem it out of the crucible.

Seriously: I think it would be much more fruitful to do some literature work. The last two decades had a huge sol-gel fad (can you call it a fad if it lasts so long?). If there are mixed CuO/SiO2/H2O gels, there must be some literature - this is just too obvious a target. So far I only found things like Cu or CuO nano-particles in silica gel, Cu immobilized with organics grafted on silica gel, Cu immobilized with amines grafted on silica gel, etc.

unionised - 21-9-2013 at 00:47

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


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.

unionised - 21-9-2013 at 01:37

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


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.

blogfast25 - 21-9-2013 at 04:16

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]

Eddygp - 21-9-2013 at 08:59

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.

blogfast25 - 22-9-2013 at 08:20

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]

turd - 22-9-2013 at 11:00

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.

kmno4 - 22-9-2013 at 13:11

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.

blogfast25 - 22-9-2013 at 13:20

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]

blogfast25 - 22-9-2013 at 13:28

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]

turd - 22-9-2013 at 13:48

Quote: Originally posted by kmno4  
Everything is nice but......
This "copper basic sulfate" looks a little strange to me.

As I posted upthread: I had the blue product analyzed. It was undoubtedly nicely crystalline brochantite. Colors are treacherous!

miken277 - 13-6-2015 at 15:45

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

blogfast25 - 13-6-2015 at 16:28

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.

Nicodem - 26-8-2015 at 09:53

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

cr-2015-00014b_0023.gif - 117kB