Sciencemadness Discussion Board - Copper carbonate - controlling the colour

A year ago I prepared copper carbonate, one is blue, the other green.
The blue is 2Cu(CO₃)(OH)₂
The green is Cu(CO₃)(OH)₂
This mimics azurite and malachite respectively.

[snip]

Duh! of corse it did not oxidise, they both have the same amount of valence electrons, so this means that there are Two copper carbonates to one hydroxide in the blue, but how?

2Cu(CO₃)(OH)₂ doesn't make any sense, actually. I can see from the page code you're using the rectangular brackets, not the proper angular ones. Try again? I think it is 2CuCO3.Cu(OH)2.

Cu(CO3

(OH)2 does also not make sense. Must be Cu2(CO3

(OH)2, I think...

Correct about the oxidation though: no further oxidation possible, all credit cards maxed out!

HINT: the bluest one is the one with the highest OH to CO3 ratio and that reflects preparation conditions, I believe. Try to use a solution of NaOH + Na2CO3, in the ratio that OH/CO3 occurs in the mineral, as precipitating agent? Just a thought, you know?

[Edited on 26-11-2014 by blogfast25]

Texium - 26-11-2014 at 15:16

I've made a few different shades of CuCO3 too. I've found that using bicarbonate leads to a greener product than using carbonate does. I think using a more dilute solution would help produce a greener color too. It all seems to be dependent on the pH of the solution. At some point I think I'll run an exhaustive test tube scale experiment on the different shades of this compound, and maybe make a video about it to be posted on the Rador Labs channel.

CHRIS25 - 27-11-2014 at 05:31

zts16 - well I will be doing some tests and experiments so expect a post back here at he end of the week. With pictures hopefuly.

Blogfast26 - I see by the chemical formulas for each rock the following:

Malachite = Cu₂(CO₃)(OH)₂
Azurite = Cu₃(CO₃)₂(OH)₂
source: http://www.galleries.com/Azurite
So the azurite has the same amount of hydroxides, but one more copper and two more carbonates - I have to admit I am confused now. No problems about experimenting, I just don't understand the mechanism here.

[Edited on 27-11-2014 by CHRIS25]

blogfast25 - 27-11-2014 at 05:57

No problems about experimenting, I just don't understand the mechanism here.

[Edited on 27-11-2014 by CHRIS25]

When I think about it, like marble, these minerals are almost certainly metamorphic rocks. They may have started out from the same basic copper carbonate and then altered in geological processes into the two azurite/malachite variations. If this is so then simple lab synthesis may not be, erm... simple! During their metamorphosis different conditions of temperature and (high!) pressure would have caused the different chemical compositions to arise. That's not something that can be controlled by precipitation conditions alone.

http://en.wikipedia.org/wiki/Metamorphic_rock

Try writing the stoichiometrical equations for the formation of azurite and malachite from CuSO4 and mixtures of NaOH and Na2CO3 solutions.

Hint: Malachite can be re-written as CuCO3 + Cu(OH)2, Azurite as 2 CuCO3 + Cu(OH)2.

[Edited on 27-11-2014 by blogfast25]

blogfast25 - 27-11-2014 at 06:04

Quote: Originally posted by zts16

At some point I think I'll run an exhaustive test tube scale experiment on the different shades of this compound, and maybe make a video about it to be posted on the Rador Labs channel.

That would be an interesting thing to do. What matters of course is the dry colour of the stuff, not the wet precipitate.

Texium - 27-11-2014 at 07:12

Quote: Originally posted by zts16

At some point I think I'll run an exhaustive test tube scale experiment on the different shades of this compound, and maybe make a video about it to be posted on the Rador Labs channel.

That would be an interesting thing to do. What matters of course is the dry colour of the stuff, not the wet precipitate.

Yes, the samples will be thoroughly dried before making any conclusions. I think I'll also compare the samples to some copper hydroxide as a sort of control, if I can dry it without having it decompose.

Also, I have a natural copper ore chunk that I got at a rock shop which consists of a mixture of malachite and azurite (mostly azurite), so it seems like both are formed under the same conditions to a certain extent.

CHRIS25 - 27-11-2014 at 08:49

CuSO₄+2NaHCO₃ = Cu(HCO₃)₂ + Na₂SO₄

But the same equation above can also yield: CuCO₃ + Na₂SO₄ + H₂O + CO₂

YET CuSO₄ + equally same amount of NaHCO₃ = CuCO₃ + NaHSO₄ +H₂O + CO₂ (actually the bisulphate here would be nice to have as by-product).

As far as the 2 moles to one mole ratio is concerned looking at the first equation that yields the copper bicarbonate is not possible since I read that it can not exist? so this is essentially a copper hydroxide mix whereas the second equation yields......yep, too many questions. Anyway my interest deepens, better get up and get moving and get mixing.

Out of interest: I painted some green copper carbonate after grinding and onto watercolour paper, (using linseed oil), it worked beautifully, and after abrasive action and wetting the paper after 24 hours to see if it would come off - it does not budge.

[Edited on 27-11-2014 by CHRIS25]

subsecret - 27-11-2014 at 08:55

Out of interest: I painted some green copper carbonate after grinding and onto watercolour paper, (using linseed oil), it worked beautifully, and after abrasive action and wetting the paper after 24 hours to see if it would come off - it does not budge.

[Edited on 27-11-2014 by CHRIS25]

At least the carbonate only decomposes in the presence of heat or water. If you painted an ocean, you'd come back the next day to continue and you'd see the most morbid thing ever.

CHRIS25 - 27-11-2014 at 09:02

Quote: Originally posted by Awesomeness

At least the carbonate only decomposes in the presence of heat or water. If you painted an ocean, you'd come back the next day to continue and you'd see the most morbid thing ever.

No idea what you mean.

blogfast25 - 27-11-2014 at 09:11

Stoichiometries, for what it's worth:

Malachite:

Na2CO3 + 2 NaOH + 2 Cu2+ == > CuCO3.Cu(OH)2 + 4 Na+

Azurite:

4/3 Na2CO3 + 4/3 NaOH + 2 Cu2+ == > 2/3 (2CuCO3.Cu(OH)2) + 6 Na+

So for Malachite the ratio of carbonate to hydroxide would be 0.5, for Azurite it would be 1. Worth a try...

It would not be hugely difficult to determine approx. Cu content in these precipitates, as these compounds decompose to CuO on calcining quite easily.

CHRIS25 - 27-11-2014 at 09:28

((Try writing the stoichiometrical equations for the formation of azurite and malachite from CuSO4 and mixtures of NaOH and Na2CO3 solutions.

Hint: Malachite can be re-written as CuCO3 + Cu(OH)2, Azurite as 2 CuCO3 + Cu(OH)2.))

If that is the answer to what you asked me to do above then no way could I have worked that out.

Also ((So for Malachite the ratio of carbonate to hydroxide would be 0.5, for Azurite it would be 1.)) Don't understand, what is the ratio here? 1 carbonate to 0 hydroxide?

[Edited on 27-11-2014 by CHRIS25]

blogfast25 - 27-11-2014 at 10:20

Also ((So for Malachite the ratio of carbonate to hydroxide would be 0.5, for Azurite it would be 1.)) Don't understand, what is the ratio here? 1 carbonate to 0 hydroxide?

For malachite we found Na2CO3 + 2 NaOH, so that is a molar ratio of carbonate:hydroxide of 1:2 (' 1 divided by 2') or 0.5.

For Azurite we found 4/3 Na2CO3 + 4/3 NaOH , so that is a molar ratio of 4/3:4/3 of 1:1 or 1.

These would be interesting ratios to use as precipitant solutions: prepare a solution that contains 1 mol (or multiple of) Na2CO3 and 2 mol (or the same multiple of) NaOH and add it to a solution of CuSO4. Observe.

Now do the same with solution that contains 1 mol (or multiple of) Na2CO3 and 1 mol (or the same multiple of) NaOH and add it to a solution of CuSO4. Observe and compare.

[Edited on 27-11-2014 by blogfast25]

CHRIS25 - 27-11-2014 at 11:16

Ok, I understood my own but I was thrown by the way you wrote yours ie 4/3 (1.33:1.33 is how I saw that) and I was seeing 2:1 but seeing yours as 0.5. Anyway, phew, sometimes goes over my head, I actually am getting ready to do exactly what you had suggested. A 2:1 and a 1:1. Out of de-ionized water so will have to wait till tomorrow.

Is it me or did we just take the long windy country road to Dublin?

What is defined as strong glass in this situation, (since I have no idea how to relate to the concept of 5 atm's. 1 = where we live that's it.
"""A solution of copper nitrate is mixed with an excess of pieces
of chalk, and the mixture is placed in a large-diameter tube of
strong glass connected to a mercury manometer. The tube is then
sealed. The azurite forms at room temperature when the liberated
CO2 creates a pressure of 5-8 atm."""

[Edited on 27-11-2014 by CHRIS25]

blogfast25 - 27-11-2014 at 13:17

Is it me or did we just take the long windy country road to Dublin?

What is defined as strong glass in this situation, (since I have no idea how to relate to the concept of 5 atm's. 1 = where we live that's it.
"""A solution of copper nitrate is mixed with an excess of pieces
of chalk, and the mixture is placed in a large-diameter tube of
strong glass connected to a mercury manometer. The tube is then
sealed. The azurite forms at room temperature when the liberated
CO2 creates a pressure of 5-8 atm."""

ALL roads to Dublin are country roads. When we drove from Holyhead to Dublin a few years back we were promised a motorway. It didn't come, not until Dublin either. Or rather, what we had been on, a dual carriage way, was IT.

1 atm is the pressure you'd experience when you dive to 10.33 m, very tolerable. Most ordinary glass wouldn't flinch at that pressure. 5 - 8 atm: multiply accordingly. At 50 m you already need to decompress to avoid the 'bends'.

Well, actually at 10.33 m you experience 2 atm, one from the air above the water, one from the water itself.

Does that make it more 'real'?

[Edited on 27-11-2014 by blogfast25]

CHRIS25 - 27-11-2014 at 13:49

Holyhead? That must have been an amazing trip, a dual carriageway across the irish sea. They only just built a motorway from Cork to Dublin 5 years ago!

So, on the definition of strong glass in the above for 5 atm. My guess is that a sealed douwe egberts (you know what that is) jar would have to have a lot of pressure generated by carbon dioxide before it succumbed?

Preliminary Results of Tests

CHRIS25 - 28-11-2014 at 03:17

3 Jars received 150 mLs water each + 0.15mol of Na₂CO₃ + 0.15mol CuSO₄ each. Then each Jar received a 2, 1, and none NaOH respectively.

Images need no explanation, though the jar with 1mol NaOH began to release CO₂ a few minutes after the photos were taken, and the jar with no NaOH exploded immediately (not literally). The colours are zeer opmerkelijk (noteworthy), Blue to blue green to green.

Will filter and dry, and do some other tests but wanted to post these images straight away.

It must be said that it appears as if not all the copper sulphate has dissolved in the jar with 2mol NaOH. Adding more water made no difference it appears that precipitating copper carbonate here is suppressed. It definitely seems that the jar with 2mol NaOH Prevents copper carbonate from precipitating and instead we have a blue-green Copper Hydroxide layer sitting on top of undissolved copper sulphate. Anyway I can filter this and allow the copper hydroxide to react with the air CO₂ to form copper carbonate. See what colour we get.

[Edited on 28-11-2014 by CHRIS25]

[Edited on 28-11-2014 by CHRIS25]

[Edited on 28-11-2014 by CHRIS25]

blogfast25 - 28-11-2014 at 05:28

That journey did indeed involve a Seacat ferry, traveling on water, not an A-type road. My bad.

I'm a bit at a loss to explain the CO2 evolutions in 2 and 3. It appears that in those cases a hydroxycarbonate was formed but why does some of the carbonate convert to CO2? It's not clear to me right now.

I can only come up with:

Cu2+(aq) + CO32-(aq) + H2O(l) ===> Cu(OH)2(s) + CO2(g)

That must be it. But why no bubbles in case 1?

Oh, and I wouldn't recommend coffee jars or jam jars at 5 atm! You'd need at least thick walled borosilicate glass to be on the safe side. Do you have a link for that Azurite preparation you referenced? (Edit: found it. From Brauer's 'Inorganic Preparative Chemistry'. A Very Authorative source, indeed)

http://books.google.co.uk/books?id=Pef47TK5NfkC&pg=PA102...

[Edited on 28-11-2014 by blogfast25]

[Edited on 28-11-2014 by blogfast25]

CHRIS25 - 28-11-2014 at 06:16

I suppose by looking at the reaction equations the 2mol NaOH solution just makes the copper hydroxide and the sodium and carbonate ions do not participate at all?

When NaOH is less, or absent the it seems that the reaction proceeds. It was noticeable that the 1mol NaOH solution took about 2 minutes before there was any liberation of carbon dioxide.

Magpie - 28-11-2014 at 06:49

FYI, here's an old thread that relates some experience on this subject:

http://www.sciencemadness.org/talk/viewthread.php?tid=5529&a...

CHRIS25 - 28-11-2014 at 08:53

Quote: Originally posted by Magpie

FYI, here's an old thread that relates some experience on this subject:

http://www.sciencemadness.org/talk/viewthread.php?tid=5529&a...

This thread where you said: "...the more carbonate the more green.." is what I have seen. The solution without the NaOH gave more green to the filtrate. The solution with more NaOH gives more blue. Also the solution without the NaOH yielded plenty of CO2 but with the NaOH no CO2 was visible, on the contrary I have copper hydroxide layer on top of very deep blue, I presume copper sulphate.

The formula for malachite has less carbonate than the blue azurite. So an apparent contradiction appears which I am certain is no contradiction - I am unable to understand any of this due to the usual lack of knowledge.

[Edited on 28-11-2014 by CHRIS25]

blogfast25 - 28-11-2014 at 09:20

Chris:

That Azurite preparation method by Brauer could be carried out successfully in a PET C*ke bottle.

This here fellow pumps up a PET drinks bottle to 120 psi with CO2 without blinking: that’s about 8 bar. Other references confirm these values.

https://www.youtube.com/watch?v=X6o561T7N6I

And this one:

https://www.youtube.com/watch?v=ZBJ-Wd5RVGI

180 psi (12.5 bar) and beyond:

https://www.youtube.com/watch?v=D78K3sW-1fA

The reaction equation is:

3 Cu(NO3)2 + 3 CaCO3 + H2O === > 2 CuCO3.Cu(OH)2 + 3 Ca(NO3)2 + CO2

(reaction equation edited: H2O added)

So with each mol of Azurite is produced also 1 mol of CO2. As in the Brauer method we can use that to build up pressure to the needed level.

A 2 L bottle (at STP) contains about 2/24 = 0.08333… mol air. By adding some moles of CO2 we can increase pressure. According the Ideal Gas Law p V = n R T (with p pressure, V volume, n number of moles, R the Ideal Gas Constant and T the temperature (in Kelvin)) and since as we add the CO2 without changing neither the volume nor the temperature of the bottle we can say:

p1 / p2 = n1 / n2, with p1 = 1 bar and n1 = 0.0833.

If we set the final pressure p2 at a safe 5 bar, we can calculate n2 = 5 x 0.0833 = 0.417. Subtract from that n1 and the number of moles of CO2 to be added to the bottle to obtain 5 bar pressure is 0.417 – 0.0833 = 0.333 mol CO2.

That also corresponds to 0.333 mol Azurite or about 0.333 mol x 344.6 g/mol = 115 g of Azurite. The reaction would require 3 x 0.333 = 1 mol Cu(NO3)2 and the excess of limestone.

Quite neat!

The solution without the NaOH gave more green to the filtrate.

You meant the filter cake, not the filtrate, right?

[Edited on 28-11-2014 by blogfast25]

CHRIS25 - 28-11-2014 at 11:23

That's a whole lot of unfamiliar maths, but the idea is something I would like to try. Just need to find fizz giz tops and a CO2 extinguisher for dry ice. Thankyou for the links Gert, something new.

DraconicAcid - 28-11-2014 at 11:35

Chris:

That Azurite preparation method by Brauer could be carried out successfully in a PET C*ke bottle.

This here fellow pumps up a PET drinks bottle to 120 psi with CO2 without blinking: that’s about 8 bar. Other references confirm these values.

I have had such bottles pop from wine fermentation, believe it or not, so while they will probably take the pressure, you may want to put some duct tape around the bottle just in case.

I wonder if a champagne bottle would take more pressure than that (of course, if it shatters, you get high velocity broken glass, so wrapping it with tape is a must, and even with tape, terrifying).

blogfast25 - 28-11-2014 at 13:02

That's a whole lot of unfamiliar maths, but the idea is something I would like to try. Just need to find fizz giz tops and a CO2 extinguisher for dry ice. Thankyou for the links Gert, something new.

What would you want the dry ice for? Here the CO2 and pressure are generated in situ, by the reagents. No extra CO2 needed.

Quote: Originally posted by DraconicAcid

Where these genuine OEM PET bottles? I had some delamination problems with bottles from a ‘starter’ beer making kit but these clearly weren’t OEM (nor did they 'pop'). And was that CO2 from primary or secondary fermentation?

Glass, when well utilised can take much more pressure than flimsy PET carbonated drinks bottles but as you point out, exploding glassware is no joke.

CHRIS25 - 28-11-2014 at 13:07

That's a whole lot of unfamiliar maths, but the idea is something I would like to try. Just need to find fizz giz tops and a CO2 extinguisher for dry ice. Thankyou for the links Gert, something new.

What would you want the dry ice for? Here the CO2 and pressure are generated in situ, by the reagents. No extra CO2 needed.

My thinking was to test any bottle to see how much pressure it could take first.

DraconicAcid - 28-11-2014 at 13:30

The one incident that I remember was a batch of plum wine that my mother was making- the batch was slightly larger than the carboy, so she put the remainder into a pop bottle, intending to let off the pressure frequently (which I do all the time). She forgot and went on vacation for a couple of days- when she came back, the bottle had ruptured, knocking a great chunk out of the carboy next to it. What a horrible mess, and the loss of gallons of wine....

blogfast25 - 29-11-2014 at 06:36

Quote: Originally posted by DraconicAcid

For 1 L of 13 ABV wine, about 125 g of CO2 are produced. If you tried to contain that you'd exceed the recommended carbonation levels of 8 g CO2/L enormously and things would definitely go pop. I had a water lock blown off a glass demijohn in my debutante brewing days, presumably due to partial blockage of the lock.

But anything in the range of 8 to 12 g/L CO2 must be safe, taking into account bottle manufacturers will build a safety factor into their products.

[Edited on 29-11-2014 by blogfast25]

Images from tests

CHRIS25 - 29-11-2014 at 06:46

Maybe this thread has run dry but I thought I would at least post some images, comments would be helpful if anyone has anything further to add:

The Paint trial: Left hand watercolour paper painted with the blue copper carbonate and the right hand painted with the green. (Powdered down and mixed with linseed oil). As opposed to the right side which is stable, fixed and does not wash or rub away, the left side blue turned to powder and came off (after 15 hours) leaving a slightly darker green underneath?

The Blue and yellow containers (copper sulphate and sodium carbonate): Blue container contains the 'cake' from the "No NaOH added" (bad photo but is greener) solution and the yellow container contains the "Added 1 mol NaOH" (clearly blue) solution

The Sodium BIcarbonate and copper sulphate:
No explanation, colour obvious.

blogfast25 - 29-11-2014 at 06:58

Chris:

Since as you are looking for greenish copper based pigments, presumably for chalk paints, there's another one that's easy to prepare: dicopper chloride trihydroxyde, or Cu2Cl(OH)3:

http://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide#...

I prepared this from CuSO4, NaCl and NaOH:

2 CuSO4(aq) + NaCl(aq) + 3 NaOH(aq) === > Cu2Cl(OH)3(s) + 2 Na2SO4(aq)

Add the NaOH solution slowly to the 2 CuSO4 + NaCl solution, with constant stirring. Add only a stoichiometric amount of NaOH: overshooting converts the greenish blue dicopper chloride trihydroxyde precipitate to blue Cu(OH)2.

I didn't actually characterise the obtained precipitate but I did test the carefully washed product for bound chloride and there was plenty of that.

[Edited on 29-11-2014 by blogfast25]

blogfast25 - 29-11-2014 at 07:06

The Paint trial: Left hand watercolour paper painted with the blue copper carbonate and the right hand painted with the green. (Powdered down and mixed with linseed oil). As opposed to the right side which is stable, fixed and does not wash or rub away, the left side blue turned to powder and came off (after 15 hours) leaving a slightly darker green underneath?

Again, I suspect inadequate drying may cause the poor adhesion on the left hand side. Better drying of ANY of these products should also reinforce the green hues, not the blue tones.

I suggest after washing with water and careful dripping 'dry', to wash the filter cake a few times with small amounts of acetone or clear methylated spirits (or even clear methanol). Then dry at RT (or below 50 C in any case) to drive of the solvent, then grind to required fineness and give it a final dry for a few days in a CaCl2 desiccator.

'Small amounts of moisture + linseed oil = problems', I think, because oil and water don't mix.

[Edited on 29-11-2014 by blogfast25]

CHRIS25 - 29-11-2014 at 08:40

Actually I totally agree, plus I predicted that a copper carbonate with water molecules would not work. But amazingly, the green carbonate from a year ago was quite moist and the blue carbonate from a year ago was totally dry due to it being easily powdered and well, (non-scientifically) certainly more dry than the green which was clearly "wet". Anyway the green was what turned out to be quite good. So I am very thankful for the above links, and will get onto that right away. (Actually I am experimenting with all this just for the paint pigments this time, not for the chalk paint which was the calcium carbonate from last week). I appreciate the links and information. (At least the yellow pigment works), though could be better with a buchner flask

blogfast25 - 7-12-2014 at 10:18

I’ve prepared some of what I believe to be Cu2Cl(OH)3, from a mixture of CuSO4, NaCl and NaOH.

0.1 mol of CuSO4.5H2O and 0.2 mol of NaCl were dissolved in 150 ml water. While stirring magnetically 0.1 M NaOH was slowly added. A light green precipitate formed but also a few specs of dark blue presumed to be Cu(OH)2. On further stirring these disappeared completely. I didn’t add the full stoichiometric amount of NaOH to avoid further Cu(OH)2 formation (previous experience showed that the precipitate converts to Cu(OH)2 with excess alkali). Here it is after settling:

The remainder of Cu2+ can be seen in the supernatant.

The precipitate filtered and washed with difficulty, slow even with a Buchner. A final rinse with a bit of acetone removed most of the wash water.

The acetone was evaporated on the hot plate, low setting. Here’s the product:

It’s already much greener than the original precipitate.

I’m now drying it in an oven at a modest 90 C, hopefully to constant weight. A preliminary test showed it’s much more resistant to heat than Cu(OH)2 but it’s still possible I end up with a black product of mainly CuO.

If it holds up this green colour I’ll probably determine Cu content to compare it to the theoretical value of Cu2Cl(OH)3.

aga - 7-12-2014 at 13:43

I sense a range of colours looming - a Chemical Palette ...

CHRIS25 - 8-12-2014 at 02:34

So far I have prepared Cu carbonate in 4 different ways, and four different shades ranging from complete blue, to blue-green turquoise and dark green. The good news is that when the dried compound is mixed with linseed oil it always turns to the same colour green and paints on water colour paper very well indeed. I wish I could grind it a bit more but seems impossible without a mechanical device. I will be doing Blogfasts's recipe today and then I want to be able compare the copper and chloride amounts with the copper oxychloride I bought from a farm supply.

So using Blogfast's mole ratios: I had quite a fair amount of blue globulars spinning around with the magnetic stirrer and although the precipitate was turquoise when I had finished adding the NaOH, I decided to leave it for a few hours and now it is quite clearly green, I will be interested to see whether or not there is any copper hydroxide precipitate mixed in with this.

(NOTE: THE 'SUPER' TEXTS BUTTON ARE NOT WORKING? NO IDEA WHY TRIED 4 TIMES)

So: Cu⁺²SO^-2 + 2Na⁺Cl^- + Na⁺OH^-

can provide us with either: Cu(OH)₂ or Cu₂(OH)₃Cl all dependent upon how fast one adds the NaOH, assuming that there are enough free Cl^- in solution?

I see that Na⁺ replaces Cu²⁺ in solution to give sodium sulphate, (reactivity series), but I still want to understand why the Copper and the chloride and the hydroxide come together as they do, and what prevents CuCl₂ forming instead? - stupid not thinking, because one needs H ions
an acidic environment for this.

[Edited on 8-12-2014 by CHRIS25]

[Edited on 8-12-2014 by CHRIS25]

[Edited on 8-12-2014 by CHRIS25]

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blogfast25 - 8-12-2014 at 13:25

My product reached constant weight after about 1 h @ 90 C. I kept drying for another 1/2 h and another 1/2 h, getting identical weight readings. The colour is now very much like my last photo: green, not green/blue or blue/green. It's waiting in a CaCl2 desiccator for the Cu determination.

I also tested a small piece at about 200 C and it fell apart to CuO + HCl + H2O and turned black. This is in agreement with Wiki's entry (even though it wrongfully mentions an MP of 250 C!).

(NOTE: THE 'SUPER' TEXTS BUTTON ARE NOT WORKING? NO IDEA WHY TRIED 4 TIMES)

So: Cu⁺²SO^-2 + 2Na⁺Cl^- + Na⁺OH^-

can provide us with either: Cu(OH)₂ or Cu₂(OH)₃Cl all dependent upon how fast one adds the NaOH, assuming that there are enough free Cl^- in solution?

Chris, you need to use the angular brackets ( < and >

, not the rectangular ones ( [ and ] ) to embed the sub and sup tags.

A mixture of CuSO4 and NaCl in solution is an 'ion soup' in water. Besides water molecules, slightly simply put it contains Cu2+, Na+, SO42- and Cl- ions. The reactivity series doesn't enter into it because dissolving the salts causes no reduction or oxidation anywhere.

The Cu(OH)2 seems to form briefly when locally high OH- concentration is reached during the NaOH solution addition.

But as long as the supernatant solution is neutralish AND still contains Cl- ions, the Cu(OH)2 is converted to Cu2Cl(OH)3:

2 Cu(OH)2(s) + Cl-(aq) ===> Cu2Cl(OH)3(s) + OH-(aq)

This makes me believe the chlorohydroxide is even more insoluble than the straight hydroxide.

CuCl2 cannot form because its (quite high) solubility limit has not been reached.

[Edited on 8-12-2014 by blogfast25]

CHRIS25 - 9-12-2014 at 08:46

Another case of Aluminium sulphate - darn tough to drive off the water. Anyway since each Cl ion requires 2 Cu ions to make Cu₂(OH)₃Cl; the 1 mol Copper sulpate supplied is the limiting reagent in terms of Cu ions and therefore can only yield 0.5 mol of the Cu₂(OH)₃Cl. Expected yield was therefore 10.68g. Actual yield is 5.96g, but am allowing for 1g loss on the filter paper (very sticky and tough to scrape off when so wet) plus maybe a 2g loss through incomplete reaction since there was some copper hydroxide mixed in with the precipitate. So a 3 gram loss is quite feasible here, close then to the 10gram expected yield.

You did not mention your yield Gert?

Re the HTML text, I always use the buttons and never use the actual HTML code when inserting the sub and super scripts.

So now to compare this with copper oxychloride which both Wiki and ChemSpider say are exactly the same.

blogfast25 - 9-12-2014 at 10:50

Another case of Aluminium sulphate - darn tough to drive off the water. You did not mention your yield Gert?

[snip]

Re the HTML text, I always use the buttons and never use the actual HTML code when inserting the sub and super scripts.

Hmm. The difficulties of getting rid of water in aluminium sulphate hydrate and wet Cu2Cl(OH)3 are very different in nature. Cu2Cl(OH)3 fresh precipitate e.g. isn't a hydrate, it's simply a wet crystalline substance that's insoluble in water.

I didn't bother with yield because of too much product loss during filtration difficulties. Next time I'll use mainly decantation to wash the product. What did you do?

Interesting observation: those bits that had gone black at approx. 200 C have regained their original colour, pointing to thermochromism rather than decomposition at 200 C. That males the chloridehydroxide much more resistant to dehydration than Cu(OH)2.

I didn't even know there were such buttons, haha!

[Edited on 9-12-2014 by blogfast25]

CHRIS25 - 9-12-2014 at 11:10

I allowed it to filter overnight - 12 hours. Then scraped the single solid wet mass into a dish, put it into my homemade teapot oven, kept the temperature between 55 to 90c, (can't control the temp very well in a teapot! After regular checking and weighing, a total of 8.52g of water was driven away after 2 hours 45 mins interspersed with occasional breaking up of the mass.
....yes those buttons save time

blogfast25 - 9-12-2014 at 12:59

12 hours? Ouch. There has to be a way to improve the filtrability of that stuff...

Azurite in a PET bottle?

blogfast25 - 13-12-2014 at 09:05

Following Chris' link to a Brauer referenced method of preparing Azurite, here:

http://books.google.co.uk/books?id=Pef47TK5NfkC&pg=PA102...

... I've started a small experiment to that effect.

In a former carbonated drinks PET bottle of 600 ml volume, 0.3 mol of Cu(NO3)2 dissolved in about 100 ml of water and 90 g of small (0.5 - 1 cm) chunks of clean limestone (about 3 times the stoichiometric requirement) were loaded:

Fizzing started immediately but I believe this is mainly due to small amounts of residual nitric acid (I had to prepare the Cu(NO3)2 somewhat ad hoc). So for the next couple of hours I'll release pressure every now and again.

Then I will allow reaction to take place for about a week (it's really cold here and the bottle has to stay outside for obvious reasons).

Calculations show that on stoichiometric conversion pressure in the bottle should be about 5 - 8 bar.

[Edited on 13-12-2014 by blogfast25]

CHRIS25 - 13-12-2014 at 10:44

On that link just scroll down a couple of times.

@Gert: Calculations show that on stoichiometric conversion pressure in the bottle should be about 5 - 8 bar. Something I need to understand but am too scared to ask.......

[Edited on 13-12-2014 by CHRIS25]

blogfast25 - 14-12-2014 at 06:45

On that link just scroll down a couple of times.

@Gert: Calculations show that on stoichiometric conversion pressure in the bottle should be about 5 - 8 bar. Something I need to understand but am too scared to ask.......

[Edited on 13-12-2014 by CHRIS25]

I'm no sure I follow? Scroll down to what?

Re. pressure. When you're topping up a sagging tyre you're adding gas without substantially changing the volume (or the temperature) of the tyre. The pressure increase that results is governed by the Ideal Gas Law: http://en.wikipedia.org/wiki/Ideal_gas_law

In the case of our reaction we are adding gas to the bottle because the reaction generates CO₂, namely 1 mol per every 3 mol of Cu(NO₃)₂, at least stoichiometrically speaking.

Since as I know how much Cu(NO₃)₂ I used (0.3 mol) and what the volume of the bottle is (600 ml, but strictly speaking I need to account for the volume taken up by the reagents: water, Cu(NO₃)₂ and limestone) I can use the Ideal Gas Law to predict the end pressure in the bottle. I designed the experiment to achieve 5 - 8 bar, as prescribed in the method you linked to.

I'm increasingly convinced that Azurite and Malachite are part of one equilibrium:

2 Cu3(CO3)2(0H)2(s) [Azurite, blue] + H2O(l) <===> 3 Cu2CO3(OH)2 [Malachite, green](s) + CO2(g) .... Eq.1

... so that one can be obtained from the other, depending on actual conditions.

The bottle now feels 'hard' due to pressure and there is unmistakably some blue precipitate already formed.

I believe the formation of Azurite may proceed as follows:

1st step:

Formation of Malachite according to:

6 Cu(NO3)2 + 6 CaCO3 + 3 H2O === > 3 Cu2CO3(OH)2 + 6 Ca(NO3)2 + 3 CO2 .... Eq.2

2nd step:

Conversion of the Malachite to Azurite, basically as per Eq.1 (but in reverse):

3 Cu2CO3(OH)2 + CO2 === > 2 Cu3(CO3)2(OH)2 + H2O

Add this to Eq.2 for the overall reaction equation:

6 Cu(NO3)2 + 6 CaCO3 + 2 H2O === > 2 Cu3(CO3)2(OH)2 + 6 Ca(NO3)2 + 2 CO2

and divide by 2:

3 Cu(NO3)2 + 3 CaCO3 + H2O === > Cu3(CO3)2(OH)2 + 3 Ca(NO3)2 + CO2

So the conversion of the first Malachite to Azurite indeed uses the CO2 generated in step 1 because we don't allow it to escape.

Quite neat...

[Edited on 14-12-2014 by blogfast25]

CHRIS25 - 15-12-2014 at 10:33

Thought maybe you might find this interesting: What I find interesting are the 2 different oxidation states of copper and the way they are bonded. But I am probably stupid for asking (who cares) would this be the same difference in the blue and the green carbonate made synthetically. (don't mean the turquoise), I mean literally the green I made and the totally blue that I made?

Azurite Copper carbonate hydroxide

Screen shot 2014-12-15 at 18.30.29.png - 5kB

Screen shot 2014-12-15 at 18.30.29.png - 5kB

Malachite
Carbonic Acid copper hydrate

Screen shot 2014-12-15 at 17.44.12.png - 6kB

[Edited on 15-12-2014 by CHRIS25]

aga - 15-12-2014 at 12:24

So the conversion of the first Malachite to Azurite indeed uses the CO2 generated in step 1 because we don't allow it to escape.

Does that mean that the CO₂ mostly escapes in the Malachite reaction due to sheer Kinetics, or does the increase in pressure also affect the thermodynamics enabling the Azurite reaction to happen ?

Presumably Malachite would become Azurite due to atmospheric CO₂ over time otherwise.

[Edited on 15-12-2014 by aga]

CHRIS25 - 15-12-2014 at 12:42

Aga, from what I deduce so far (and am always wrong in my deductions...) the azurite blue is the Pure copper carbonate while the green is impure, so to speak. Also the green does not turn blue and the blue does not turn green, otherwise all my three samples from a year ago would have changed colour by now. All I can say is that whenever I mix the blue carbonate with Linseed oil, wet brushes on paper, it immediately turns green, even when the blue is applied to the paper it does not take long for it to change colour to green. These are only my observations though. I still need to know whether the molecular structures above that I found apply to the rocks only or also the synthetically made product at home.

Should also have added that when I used Sodium carbonate with copper sulphate AND NaOH I got a blue copper carbonate.

When the same mix was done without the NaOH I got a green carbonate. What is interesting is the fact that I noticed Half as much Carbon dioxide being driven off with the NaOH add into the equation. This then at first glance looks like with less carbon dioxide being driven off, (Since I used exactly the same quantities in all the experiments, half of it must have been used and remained in the reaction, I don't know enough chemistry at the moment to understand all this.

[Edited on 15-12-2014 by CHRIS25]

aga - 15-12-2014 at 12:45

Is the Paper treated with anything like a Sizing agent ?

blogfast25 - 15-12-2014 at 13:33

OK. Lots of interesting questions here, so let me try.

Thought maybe you might find this interesting: What I find interesting are the 2 different oxidation states of copper and the way they are bonded. But I am probably stupid for asking (who cares) would this be the same difference in the blue and the green carbonate made synthetically. (don't mean the turquoise), I mean literally the green I made and the totally blue that I made?

The oxidation state of Cu in BOTH Azurite and Malachite is +2, so that cannot explain the difference in colour. Cu(I) is only really quite stable in Cu2O, copper(I) oxide, which is kind of Bordeaux.

BTW, that second structure labelled 'carbonic acid copper hydrate' is simply baloney. Plain wrong. The real structure involves Cu2+ ions, CO32- ions and OH- ions and no water molecules. Going by that structure the copper wouldn't even carry any charge at all!!!

So the conversion of the first Malachite to Azurite indeed uses the CO2 generated in step 1 because we don't allow it to escape.

The formation of Malachite (from copper nitrate and calcium carbonate) generates lots of CO2, as per the equation. If we allowed that CO2 to escape (open container) then nothing further would happen but if we ‘keep’ the CO2 then:

3 Cu2CO3(OH)2 + CO2 === > 2 Cu3(CO3)2(OH)2 + H2O

… occurs. I can kind of see the green Malachite being converted to the blue Azurite, just looking at the bottle.

Azurite at atmospheric conditions is probably so-called meta-stable: thermodynamically it SHOULD release that CO2 again (at least in the presence of water) but it may only do this very slowly, so it APPEARS to us to be stable.

But atmospheric CO2 is of far too low pressure to cause that conversion (Malachite to Azurite) and indeed all copper hydroxycarbonates we see on statues and other copper structures are always green, never blue (even though it’s subjective, colours always are).

Aga, from what I deduce so far (and am always wrong in my deductions...) the azurite blue is the Pure copper carbonate while the green is impure, so to speak. Also the green does not turn blue and the blue does not turn green, otherwise all my three samples from a year ago would have changed colour by now. All I can say is that whenever I mix the blue carbonate with Linseed oil, wet brushes on paper, it immediately turns green, even when the blue is applied to the paper it does not take long for it to change colour to green. These are only my observations though. I still need to know whether the molecular structures above that I found apply to the rocks only or also the synthetically made product at home.

It’s interesting that your blue turns green so quickly on wetting. I’ll be checking that with the Azurite, once it's come off the high pressure CO2 diet.

[Edited on 15-12-2014 by blogfast25]

aga - 15-12-2014 at 13:45

My hypothesis is that the Thermochemistry is OK, in that Malachite => Azurite in a CO2 environment reaction Will happen at STP.

Your experiment is using augmented pressure as well as a saturated CO2 environment, implying that the Azurite reaction might only take place if the Kinetics are agreeable at 5+ atm.

If Malachite + CO2 => Azurite at 1 atm, then case proven.

So a simple experiment would be to make some Malachite, then take half of the sample and subject it to a stream/atmosphere of CO2 for a while.

OK. I think i can do that (famous last words).

blogfast25 - 15-12-2014 at 14:05

Aga:

As indicated in my earlier post I think Malachite and Azurite are in equilibrium according to:

3 Cu2CO3(OH)2(s) + CO2(g) < === > 2 Cu3(CO3)2(OH)2(s) + H2O(l)

According to the Le Chatelier Principle, if you increase the pressure in that system it will respond by trying to undo that change. Here it does that by absorbing some of the CO2. The overall pressure decreases because the right hand side of the equation is much less voluminous than the left hand side.

What I observe is in agreement with Le Chatelier and that means that applying CO2 in atmospheric conditions will not convert Malachite to Azurite because there is no driving force.

But by all means prove me wrong, empirically.

[Edited on 15-12-2014 by blogfast25]

aga - 15-12-2014 at 14:48

I'll give it a go and see what happens.

Whether one is right, wrong or simply misled shall be known by said go-ing-happeneing.

[Edited on 15-12-2014 by aga]

I very much like CHRIS25's posts.
So much to actually DO !

[Edited on 15-12-2014 by aga]

CHRIS25 - 16-12-2014 at 01:26

@Gert

Quote: ""BTW, that second structure labelled 'carbonic acid copper hydrate' is simply baloney. Plain wrong. The real structure involves Cu2+ ions, CO32- ions and OH- ions and no water molecules. Going by that structure the copper wouldn't even carry any charge at all!!!"" Unquote

Then I am truly puzzled, would you mind having a look at these, because I do not understand this at all then: Plus the above molecular structure is listed clearly with the same formula on chemspider, the one you that you use for malachite copper carbonate. Although I do admit they are MISSING a diagrammatic Carbon atom somewhere.
http://pubchem.ncbi.nlm.nih.gov/compound/3081961?from=summar...
http://www.chemicalregister.com/MALACHITE_CU2_CO3_OH_2/Suppl...

blogfast25 - 16-12-2014 at 05:56

Chris:

BOTH references are WRONG: they start off from the wrong molecular formula of: CH6Cu2O5

But it actually is Cu2(CO3)(OH)2!

CH6Cu2O5 has 4 H too many! That explains why in their faulty structure model the Cu would be chargeless, which is clearly IMPOSSIBLE.

It's good practice to always include a link to any information you're unearthed, so people can check context.

Don't be too surprised by this: science isn't perfect and scientific literature contains errors, often of the 'cut 'n paste' type.

Later today: the suspected role of water in the Malachite - Azurite transformation.

[Edited on 16-12-2014 by blogfast25]

CHRIS25 - 16-12-2014 at 07:20

Very surprised by this. But thanks Gert. I have used chemspider quite a lot and thought it a completely trustworthy source for learning.

Ok, I tested both by copper carbonates for copper content. One is thoroughly Blue, the other thoroughly Green. These are the results and the Blue results really surprised me and have not had time to work out why yet. I tested the blue carbonate a second time believing that I made a mistake somewhere, no I can not see a mistake.

Everything in Grams, saves an extra key punch!

GREEN:
Empty Beaker and test tube = 110.89
Carbonate weight = 1.02
Beaker/Test tube filled with conc sulphuric plus Carbonate = 126.71
Weight after reaction is completed = 110.98
Copper content = 0.09

BLUE: Test A

Empty Beaker and test tube = 110.85
Carbonate weight = 1.03
Beaker/Test tube filled with conc sulphuric plus Carbonate = 125.44
Weight after reaction is completed = 109.64
Copper content? Something has been added? = 1.21 or not all the acid was consumed?

Blue: Test B

Empty Beaker and test tube = 111.35
Carbonate weight = 1.16
Beaker/Test tube filled with conc sulphuric plus Carbonate = 125.53
Weight after reaction is completed = 109.24
Copper content? Something has been added? = 2.11 or not all the acid was consumed?

The most noticeable thing about these reactions with sulphuric acid was that the Blue carbonate liberated a huge amount of Carbon dioxide, whereas the Green liberated just a gentle fizz. Compare a firework banger to a firework rocket

Also the Green upon completion, stayed green. the Blue however turned a milky white upon completion, whereupon addition of water turned the white into a blue again.

[Edited on 16-12-2014 by CHRIS25]

[Edited on 16-12-2014 by CHRIS25]

aga - 16-12-2014 at 08:52

My hypothesis is that the Thermochemistry is OK, in that Malachite => Azurite in a CO2 environment reaction Will happen at STP.

No. It doesn't.

I made some copper carbonate today and bubbled CO₂ through it for a couple of hours.

Nothing happened.

blogfast25 - 16-12-2014 at 09:44

Chris:

Interesting what you're trying to do there and decent reasoning but with flaws. I really don't see how the copper content can be deduced from it, though.

I'm still at work so just wanted to post my water spiel quickly. Will look at your numbers tonight.

The role of water in:

3 Cu2CO3(OH)2(s) + CO2(g) < === > 2 Cu3(CO3)2(OH)2(s) + H2O(l)

Left we have 6 OH ions, 3 CO3 ions and one CO2, right we have 4 OH ions, 4 CO3 ions and one H2O. So it looks like we’ve combined 2 OH ions with a CO2, to form another CO3 ion and 1 water.

Imagining how 2 OH- + CO2 = H2O + CO32- is really not easy.

Enter… the water.

As we know CO2 is much more soluble in water when under pressure. CO2 in water forms carbonic acid acc. the equilibrium:

H2O(l) + CO2(aq) < === > H2CO3(aq)… (I)

Carbonic acid is a weak acid and its two deprotonations can be summarised as:

H2CO3(aq + 2 H2O (l) < === > 2 H3O+(aq) + CO32-(aq)… (II)

And now we have a way to eliminate 2 OH ions:

2 OH- + 2 H3O+(aq) === > 4 H2O(l)… (III)

Add up (I), (II) and (III) and eliminate redundancies on both sides:

CO2(aq) + 2 OH- === > CO22-(aq) + H2O(l)… (IV)

The water used in (I) and (II) is fully returned in (III), so it act as a catalyst, making (IV) possible.

This is another reason why I don’t believe converting Malachite to Azurite in a dry stream of CO2 will work: presence of water is essential.

blogfast25 - 16-12-2014 at 10:37

GREEN:
Empty Beaker and test tube = 110.89
Carbonate weight = 1.02
Beaker/Test tube filled with conc sulphuric plus Carbonate = 126.71
Weight after reaction is completed = 110.98
Copper content = 0.09

BLUE: Test A

Empty Beaker and test tube = 110.85
Carbonate weight = 1.03
Beaker/Test tube filled with conc sulphuric plus Carbonate = 125.44
Weight after reaction is completed = 109.64
Copper content? Something has been added? = 1.21 or not all the acid was consumed?

Blue: Test B

Empty Beaker and test tube = 111.35
Carbonate weight = 1.16
Beaker/Test tube filled with conc sulphuric plus Carbonate = 125.53
Weight after reaction is completed = 109.24
Copper content? Something has been added? = 2.11 or not all the acid was consumed?

Chris, you're going to have to describe in more detail what exactly you've done, to try and make sense of these numbers.

The way I see it, you are actually trying to determine the CO2 content of these samples by driving off the CO2 with conc. H2SO4 and weighing before and after. But without further explanation the numbers don't add up.

For Malachite the reaction would be:

Cu2CO3(OH)2 + 2 H2SO4 === > 2 CuSO4 + CO2 + 2 H2O

For Azurite:

Cu3(CO3)2(OH)2(s) + 3 H2SO4 === > 3 CuSO4 + 2 CO2 + 4 H2O

[Edited on 16-12-2014 by blogfast25]

CHRIS25 - 16-12-2014 at 11:02

Ah sorry. Still learning to be more meticulous and in presenting data properly. (untrained in that but training myself as I go along); Ok I weighed an empty beaker with an empty test tube in it. I then weighed out the copper carbonate. I then poured in to the test tube 11 mL 98% sulphuric acid and weighed the beaker again that now contains the test tube with acid and the copper carbonate. Poured the acid directly into the carbonate in the beaker, waited for completion and then weighed the beaker with the completed reaction inside.

The following: CuCO₃ + H₂SO₄ = Cu + H₂O + CO₂ +SO₄

Theoretically since all the dioxide has been driven away we are left with a solution of sulphate ions, water and copper 2⁺ ions.
Yes this does give us the amount of Copper, and here is how:
We know that in GREEN Test that 126.71 - 110.89 = 14.80 is the weight of the acid I put in. We also know that 126.71 - 110.98 = 15.73 is the weight AFTER the reaction is completed; logically then 15.73 (After completion) - 14.80 (weight of acid) = 0.93 can only be the Copper content. And I did this all by myself - pride comes before a knock on the head with a stoichiometric sledgehammer?

PS, I like what you have deduced and demonstrated above, it makes sense to me (that is a miracle you know).

[Edited on 16-12-2014 by CHRIS25]

blogfast25 - 16-12-2014 at 13:28

Unfortunately the above is wrong in many ways. I'll try and keep it short.

1. You start from a false premise: your carbonate is a hydroxycarbonate, NOT CuCO3. It makes a WHOLE lot of difference.

2. "CuCO3 + H2SO4 = Cu + H2O + CO2 +SO4" is NOT a reaction equation at all. The correct ones are in my last post (bottom).

3. "We know that in GREEN Test that 126.71 - 110.89 = 14.80 is the weight of the acid I put in"..

Nope. Above that you wrote: "Beaker/Test tube filled with conc sulphuric plus Carbonate = 126.71"

So that difference (126.71 - 110.89) is acid PLUS carbonate, not just acid.

You are, slightly clumsily, trying to determine the CO2 content, by driving it off and doing a weight difference (before - after). That's fine. But do not believe that from that alone you can determine copper content.

I suggest you run it again, as follows.

1. weigh beaker empty (W1)

2. add about 1 g carbonate and weigh again (W2)

3. Add empty test tube to beaker containing the carbonate and weigh again (W3)

4. Add 11 ml conc H2SO4 to test tube and weigh again with the beaker and the carbonate in it (W4)

5. Empty test tube in beaker, set empty test tube aside (don't clean or wash it!), allow to react. Then add empty test tube back to the beaker with the reacted mixture and weigh the ensemble (W5).

From these weights can be calculated how much CO2 the carbonate sample has lost, and thus the CO2 content.

DraconicAcid - 16-12-2014 at 13:38

Following that, if you allow the solution to evaporate to give copper(II) sulphate pentahydrate, you can weight that and estimate the copper content of the original sample.

aga - 16-12-2014 at 14:12

Currently i have a high pressure experiment running to see if Azurite will be made.
(essentially a copy of blogfast25's, but in copper piping.)

As a test i tried heating the part of the apparatus containing the carbonate using a blowtorch.

Pretty quickly the result was copper/copper oxide as a brown/black residue.

It would appear that Natural Azurite is not formed under great temperatures, so chemical mix and pressure are what remain as the variables.

Tests Re-done

CHRIS25 - 17-12-2014 at 03:50

So, the sledgehammer was not that bad considering my infantile reasoning. Anyway, here are the new tests, with the same Pattern of results and same sort of consistency. The same beaker and test tube was used, washed only after Completion of First test and not in between. The same 10.5 mL 98% acid used in both tests. Each stage was weighed minimum of 3 times by removing off the scales and placing back on to allow for minor discrepencies and to double check absolute accuracy. Plus both reactions were allowed to stand for exactly 10 minutes each before final weight was taken.

----------------------------------------------------------------BLUE-----------GREEN
Weight Beaker empty---------------------------------------95.4-------------95.4
Weight Beaker + CO₃---------------------------------------96.9-------------96.6
Weight Bker + empty TestTube + CO₃----------------112.4-----------112.2
Weight Bker + TestTube + Acid + CO₃----------------131.3-----------131.0

Weight Bker + empty TestTube After reaction------131.3-----------115.1

OBSERVATIONS:

The Blue Carbonate produced a plume of smoke in an instant, after completion of reaction a milky white solution which precipitated a white fine powder, this I deduce to be Anhydrous copper sulphate (which happens to be white and by the fact that sulphuric acid is very hygroscopic).

The Green carbonate produced mild fizzing and clearly this was carbon dioxide. The solution stayed green and it also precipitated a green fine powdery residue.

Don't ask me to interpret the results - bit fed up with being the forum idiot! But I find the above both interesting and unexpected result from the Blue. Interestingly I made the above blue with NaHCO₃ and the green was made with Na₂CO₃ but without NaOH. I also made another blue with just Na₂CO₃ + NaOH

I made a Blue-green Carbonate by doubling the amount of Na₂CO₃ but no NaOH added.
Whether any of the above is useful to know I don't know to be honest.

blogfast25 - 17-12-2014 at 03:54

Chris, thanks very much. I will interpret these data tonight. And no, you're not the 'forum idiot', just still a beginner. Rome wasn't built in a day either.

Quote: Originally posted by DraconicAcid

Following that, if you allow the solution to evaporate to give copper(II) sulphate pentahydrate, you can weight that and estimate the copper content of the original sample.

Copper sulphate pentahydrate isn't great for gravimetry and he's also got a lot of unreacted conc. H2SO4 in there which would be difficult to eliminate. It's a recipe for huge errors.

Currently i have a high pressure experiment running to see if Azurite will be made.
(essentially a copy of blogfast25's, but in copper piping.)

As a test i tried heating the part of the apparatus containing the carbonate using a blowtorch.

Pretty quickly the result was copper/copper oxide as a brown/black residue.

It would appear that Natural Azurite is not formed under great temperatures, so chemical mix and pressure are what remain as the variables.

All these copper hydoxycarbonates and hydroxychlorides appear thermo-labile from about 200 - 250 C onwards. But they're more stable than Cu(OH)2 which starts dehydrating below 100 C, even when still wet.

---------------------

I've found a way to determine the formula of a well-defined copper hydroxycarbonate, say xCuCO3.yCu(OH)2 with x and y integers, from an accurate determination of CO2 content. I will give that a whirl tonight or tomorrow.

[Edited on 17-12-2014 by blogfast25]

blogfast25 - 17-12-2014 at 09:43

Chris, once again I can’t make head or tails of your numbers: it seems in the first case NO CO2 was lost and in the second case far, far too much. Something is wrong and I don’t know what it is. You’re right about the white product being anh. CuSO4. I'll be using 50 % H2SO4.

I conducted my own experiment, using this basic copper carbonate.

A 100 ml empty volumetric flask plus funnel was accurately weighed (W1 = 71.00 g). About 4 g of alleged Cu2CO3(OH)2 was added, this gave W2 = 75.04 g. About 7 ml of water was added to create a slurry and the total weight recorded as W3 = 82.58 g.

Separately 10 ml of 50 v% H2SO4 had been prepared in a measuring cylinder, total weight W4 = 23.65 g. The acid was then bit by bit added to the flask plus funnel and its content and allowed to react. A blue solution resulted after the fizz and some CuSO4.5H2O crystallised out. The final weight was W5 = 95.76 g. The weight of the empty measuring cylinder was W6 = 9.66 g.

Photo post-experiment:

Calculations:

From W2 – W1 = 4.04 g we know the weight of copper basic carbonate.

From W6 – W4 = 13.99 g we know the weight of acid added.

From W3 + (W6 – W4) – W5 = 0.81 g we know the amount of CO2 that was released.

The weight percent CO2 in the product was thus 0.81/4.04 x 100 % = 20.0 w%. The theoretical value for Cu2CO3(OH)2 is 19.9 w%.

Now for the determination of x and y.

0.81 g CO2 is 0.81 / 44 = 0.0184 mol CO2, that corresponds to 0.0184 x 123.54 = 2.27 g CuCO3.

The remainder of the 4.04 g of product, i.e. 4.04 – 2.27 = 1.77 g must be Cu(OH)2, or 1.77 / 95.54 = 0.0185 mol.

Gotcha! Because 0.0185 mol Cu(OH)2 / 0.0184 mol CuCO3 = 1 and thus x = y.

So that product does indeed correspond to the formula CuCO3.Cu(OH)2 or Cu2CO3(OH)2

If that Azurite ever finishes (it’s still bubbling away!) I will use this method to estimate composition.

[Edited on 17-12-2014 by blogfast25]

CHRIS25 - 17-12-2014 at 10:31

Chris, once again I can’t make head or tails of your numbers: it seems in the first case NO CO2 was lost and in the second case far, far too much. Something is wrong and I don’t know what it is. You’re right about the white product being anh. CuSO4. I'll be using 50 % H2SO4.

[Edited on 17-12-2014 by blogfast25]

I can't either ! But this I do know, I did not make any errors in the weights, and I followed a very slow procedure double checking as I went along. So I have no idea, except that I must have made an obviously very unique copper carbonate! The difference in the reaction dynamics upon adding the acid give a clue? As I said earlier, the blue one gave an immediate puff of smoke, though probably steam and I will add it gave quite a lot for such a small amount, whereas the green fizzed away quietly. This is all I can say.

blogfast25 - 17-12-2014 at 10:45

One BIG problem is for the green one: you lose 15 g, yet used only about 1 g of carbonate!

One word of advice: use 50 % H2SO4, not conc. H2SO4.

[Edited on 17-12-2014 by blogfast25]

CHRIS25 - 18-12-2014 at 03:11

The pressure increase that results is governed by the Ideal Gas Law: http://en.wikipedia.org/wiki/Ideal_gas_law

In the case of our reaction we are adding gas to the bottle because the reaction generates CO₂, namely 1 mol per every 3 mol of Cu(NO₃)₂, at least stoichiometrically speaking.

Since as I know how much Cu(NO₃)₂ I used (0.3 mol) and what the volume of the bottle is (600 ml, but strictly speaking I need to account for the volume taken up by the reagents: water, Cu(NO₃)₂ and limestone) I can use the Ideal Gas Law to predict the end pressure in the bottle. I designed the experiment to achieve 5 - 8 bar, as prescribed in the method you linked to.

[Edited on 14-12-2014 by blogfast25]

I've tried finding out the pressure exerted inside a 2 L carbonated water bottle, since I just bought one that actually says 'PET' on the bottom. I have also tried to understand the ideal gas law and all that maths so that I can work out, as you did, a pressure of 8 bar and how much reactant to use. Well, needless to say, I can't find answers or work things out here. I simply can not understand how to do those sort of calculations. So for the first time ever I am asking for the answer: How much copper nitrate could I use in a 2 L bottle without waking up the neighbours in the middle of the night with a loud bang. How much water should I use and lastly what kind of pressure do you think a 2 L bottle could take anyway?

Also I will be using Limestone flour, this removes any dubious casualties from the equation just in case my own calcium carbonate has other things in it.

Will re-do that copper carbonate as you suggested in the coming days, this is a week where nothing seems to be going right.

Third Round Tests

CHRIS25 - 18-12-2014 at 05:43

Same Procedure used as previous except Sulphuric acid diluted to 50%, using approx 17 mL (8.5 : 8.5) in each test tube.

Results:

----------------------------------------------------------------BLUE-----------GREEN
Weight Beaker empty-------------------------------------95.8-------------95.8
Weight Beaker + CO₃-------------------------------------96.8-------------96.8
Weight Bker + empty TestTube + CO₃--------------112.3-----------111.97
Weight Bker + TestTube + Acid + CO₃--------------134.65-----------135.22

Weight Bker + empty TestTube After reaction-----134.54-----------135.21

------------------------------------------------CO₂ loss -------0.11----------0.01

Ok my measurements and procedure is spot on. No Mistakes not this time. All I can say is that either I should patent my copper carbonate or my scales are moody. However I must add that I double and re-double the weighing to see if there is any difference. I do notice that just 2 drops of acid added to the beaker with everything still in there, (I did this at the end to test it), 2 drops change the scale reading from a 0.35 to a 0.67, very sensitive.

Observations:

The Blue carbonate stayed blue and the precipitate that slowly dissolved stayed blue, the green however surprised me, with dilute acid the green very slowly (after 5 mins) changed to blue, and the green precipitate also changed from green to blue as it dissolved. The blue did not yield a plume of steam this time with the addition of acid, both green and blue yielded their dioxide at the same rate.

The whole point of all this is that I am trying to account for the different behaviours between the blue and green carbonates. Both with concentrated acid and dilute acid, and of course the results. Even if the Green figures are slightly off, the fact is that the blue yielded a lot more carbon dioxide than the green. And if one looks at the formulas for both blue and green one clearly sees that the Blue has twice as much carbon dioxide to generate. But why, with concentrated sulphuric acid did the blue not do this? I will do another test using about 5 g of carbonate and compare concentrated with dilute again just to see whether there is any consistency with those skew-whiff results.

[Edited on 18-12-2014 by CHRIS25]

blogfast25 - 18-12-2014 at 06:17

Chris, you're definitely getting closer but I think your main problem is one of scale (too small) and possibly some losses you're overlooking. Try just the blue pigment at a 5 g scale. Make sure none of the fizz (a mist of solution/slurry droplets with CO2) escapes, as that would lead to grave errors: that's why I used the semi-closed shape of a volumetric flask. You've already experienced the sensitivity of this type of quantitative experiment, so learn from it.

Can you remind me of the main differences in preparation of the blue and green carbonates? Long thread...

As regards quantities for a 2 L Azurite experiment:

In a former carbonated drinks PET bottle of 600 ml volume, 0.3 mol of Cu(NO3)2 dissolved in about 100 ml of water and 90 g of small (0.5 - 1 cm) chunks of clean limestone (about 3 times the stoichiometric requirement) were loaded:

... simply multiply the quantities I used by 2000/600 = 3.3333. That will develop the same pressure as in my 600 mL bottle experiment (assuming the method works and the reaction runs more or less to completion). Note that the limestone really does have to be as small chunks, not powder. Using powder any excess limestone would end up mixed in with the produced Azurite, inseparable from it. By contrast excess limestone in chunks can simply be sieved out of the reaction products, without contaminating the Azurite precipitate. That's if my reaction ever finishes because it's still bubbling...

The bottle can withstand approx. twice the pressure the experiment is designed to achieve.

May I ask what Cu(NO3)2 you will use? The Cu(NO3)2 solution should be reasonably neutral (pH no lower than 4 - 5).

[Edited on 18-12-2014 by blogfast25]

CHRIS25 - 18-12-2014 at 07:19

Ok Gert thankyou for all that above. Here are my notes:

Na₂CO₃ + 2NaOH + CuSO₄ = Nothing, did not work at all.
Na₂CO₃ + NaOH + CuSO₄ = Blue
Na₂CO₃ + CuSO₄ = Green
CuSO₄ + 2NaHCO₃ = Blue-Green
CuSO₄ + NaHCO₃ = Blue (I think, not the colour I am doubting more the method, notes are becoming more organized as time goes by) And these red lines under my words are annoying, Engalnd infented the spellling and USA should not re-infent itss owne.

Ah copper nitrate, I made it last year, I was even more ignorant then than I am now, so maybe I should re-do another batch?

blogfast25 - 18-12-2014 at 08:13

Ah copper nitrate, I made it last year, I was even more ignorant then than I am now, so maybe I should re-do another batch?

What physical form is it?

What's important is that you can control the amount of actual Cu(NO3)2 added to the bottle fairly well. Purity is not important. Also, if in doubt, rather use a little less than too much: it's the amount of Cu(NO3)2 that determines how much CO2 will be developed and thus how much pressure. Too much could lead to a bursting bottle, not enough and no Azurite may form.

I'm convinced after all this reading and experimenting that all copper hydroxycarbonates formed at atmospheric pressure are of the Malachite type, no matter how you produce precipitate them. Only CO2 pressure (and water) can lead to the more carbonaceous Azurite.

Another thing that makes me believe this, is in your Brauer (google) reference:

http://books.google.co.uk/books?id=Pef47TK5NfkC&pg=PA102...

… scroll down to where it says:

”Alternate methods [Azurite] a) From a precipitated green basic carbonate under a CO2 pressure of 4 atm.”

In other words, as noted above repeatedly, a slurry of Cu2CO3(OH)2 can be converted with CO2 at 4 atm to Cu3(CO3)2(OH)2 (Azurite), at least acc. this reference. This weekend I will initiate a test using that method too.

Your CO2 measurements could prove me wrong on my views here (if you get the measurements right, of course!) and I'm almost hoping so.

I think the colour differences you are observing may be due to granulometry differences and other effects, not to fundamental differences in composition.

For your CO2 measurement woes, using a small chunk of shop bought Malachite, crushed and ground to a powder, could serve as a Standard.

[Edited on 18-12-2014 by blogfast25]

aga - 18-12-2014 at 09:51

Seeing as large quantities are formed naturally, would it not make sense to look up the geological locations in which Azurite forms and deduce the parameters from there ?

"Found largely in the oxidized portions of copper deposits, it is a secondary mineral formed by the action of carbonated water acting on copper-containing minerals, or from Cu-containing solutions, such as CuSO4 or CuCl2 reacting with limestones."

Source : http://www.mindat.org/min-447.html

Judging by that data, blogfast25's experiment is bang on the money.

blogfast25 - 18-12-2014 at 10:05

Judging by that data, blogfast25's experiment is bang on the money.

Yes, I know about the natural formation process. The experiment may be on the money but it's very slow. Each day I have a look at the bottle and I can see the supernatant liquid is still quite blue due to unreacted Cu(NO3)2. Bubbles continue to form slowly. I daren't stop the experiment because I believe as long as there is unreacted Cu(NO3)2 there will also still be unconverted Malachite. Patience is a virtue...

Your reference throws up two more interesting species:

Georgeite: Cu2CO3(OH)2.6H2O, a hexahydrate of Malachite

Chalconatronite: Na2Cu(CO3)2.3H2O, a double salt of CuCO3 and Na2CO3 trihydrate, it seems... This composition is also mentioned in the Google reference. It would have a significantly higher CO2 content: 31.1 w% (Malachite 19.9 w%).

[Edited on 18-12-2014 by blogfast25]

New Measurements

CHRIS25 - 18-12-2014 at 11:29

Firstly I have to apologize. The blue that I kept measuring was the Blue-Green; good news is that the green is indeed green, and the Real Blue I have now added to the tests. So now some comprehensive data.

All weights of carbonate aimed at 6 g
Sulphuric 8 mL + 8 mL water (this is 6.5 times more than the stoichiometric requirements)

Photos demonstrate how I did the reactions and the second photo you can see the completed reaction After all measurements taken - naturally.

----------------------------------------------------------------BLUE-------------Blue-Green---------GREEN
Weight Beaker empty----------------------------------138.61-----------------133.81-----------138.9
Weight Beaker + CO3----------------------------------144.61-----------------139.87-----------144.98
Weight Bker + empty TestTube + CO3--------------159.63---------------155.03-----------160.06
Weight Bker + TestTube + Acid + CO3--------------186.7-----------------182.61-----------188.05

Weight Bker + empty TestTube After reaction-----183.6-------------------181.75----------187.68

------------------------------------------------CO2 loss -------3.1----------------------0.86------------0.37

OBSERVATIONS:
The blue produced a huge amount of fizzing so much so that I could not add the whole test tube in all at once like the other two. Also the blue turned into this gel-like precipitate.

Adrian kindly sent me a PDF: A Renaissance of Color: Particle Separation and Preparation of Azurite for Use in Oil Painting by
Michael Price
"......Azurite is sky
blue below 30 μm, becoming paler as
the particle size decreases. It is dark blue
with a slightly greenish hue from 40 to
60 μm, similar to the hues of modern
cobalt blue....."

Just a tiny quote that confirms what you said earlier about particle size. He goes into detail describing how he controls the colour of the pigment by controlling the partical size of this copper carbonate and even has electron microscope images detailing everything.

RE the copper nitrate: Does this answer your question? It is quite crystalline but looks slushy here, I forgot to sharpen the image in order to replicate the actual nitrate.

aga - 18-12-2014 at 12:13

Erm, is it just me, or do all the three samples look equal, and Turquoise ?

The nitrate looks Blue.

[Edited on 18-12-2014 by aga]

CHRIS25 - 18-12-2014 at 12:29

Yep and no it's not you, Green turns to blue with sulphuric and blue turns to green in linseed, (also what that article said). And here is another fascinating piece of data from http://www.galleries.com/malachite
".....Malachite is an impostor of its own. It frequently pseudomorphs the closely associated mineral azurite. A pseudomorph is a mineral specimen where the original mineral has been chemically replaced by another mineral, but the outward appearance is still retained. The transformation is fascinating and sometimes leaves a nearly perfect azurite crystal shape that is actually malachite. Often the transformation is incomplete and leaves a blue/green mineral specimen unlike any other".

aga - 18-12-2014 at 12:36

So the Carrier oil is the Key then ?

If i dunk my batch of Malachite (which is slowly going green) into Linseed oil, it should be a Green colour ?

Does that take Time or is it pretty quickly obvious ?

blogfast25 - 18-12-2014 at 14:07

Chris:

First off. Like aga, I have difficulty distinguishing the colours of these products. Very similar to me. I expected larger differences.

Second. Your Cu(NO3)2 looks perfect. No problems at all. Very nice, indeedy.

Third, I’m convinced the good people of galleries.com have got the wrong end of the stick: it’s Azurite that can morph into Malachite, not the other way around! For Malachite to ‘pseudomorph’ into Azurite 3 mol of Malachite would have to absorb at least 1 mol of CO2 (at STP, that’s 24 L!) at ATMOSPHERIC pressure. The CO2 pressure in the air however is still much lower than 1 atm because air only contains 0.039 % of CO2! With all due respect but these people may be excellent mineralogists but not chemists.

The weight percentages now come out respectively as 51.7; 14.2 and 6.08 w% CO2. The first is way too high, the two others really a bit too low. Let me think about this some more before I try and make sense of it. You’ve put too much work into this to shoot it down without due consideration.

So the Carrier oil is the Key then ?

If i dunk my batch of Malachite (which is slowly going green) into Linseed oil, it should be a Green colour ?

Does that take Time or is it pretty quickly obvious ?

Yes, the carrier makes a lot of difference and that should become apparent immediately after mixing. Sunflower oil can be used for a test, if you haven't any linseed oil.

And how come this *.pdf wasn't also sent to moi?

[Edited on 18-12-2014 by blogfast25]

CHRIS25 - 19-12-2014 at 00:53

Ok Ok Guys, there seems to be some communication mis-wired, or you have colour management issues on your monitors, but I doubt the latter.

The solution in the jars is DURING the reaction and had changed colours to the point that they were the same colour except for the Gel like reaction which was clearer and had lost its colour. I only showed these because dear old Gert advised me that I should not allow any spray to escape and use a volumetric, and after removing the lids I was surprised to see how much spray would have been lost, especially from the far left.

The second photo shows the reactions AFTER completion, and the two on the right are ALMOST/VERY CLOSE to the same colour, the far right having changed from GREEN to BLUE, but with a tiny tiny hint of green hue visible

These new photos show the three different carbonates that I used for the tests: THEY ARE IN EXACTLY the same order as all the above images. IE: Green far right, green-blue middle, blue far left.

[Edited on 19-12-2014 by CHRIS25]

blogfast25 - 19-12-2014 at 10:00

That's much better, Chris, thank you.

Tomorrow I'm wrapping up my Azurite experiment although there is still some blue copper nitrate left. It's gone on for long enough now.

Rummaging through the recesses of my dark mind yesterday, I had an idea.

The overall reaction equation for the Azurite experiment was:

3 Cu(NO3)2 + 3 CaCO3 + H2O === > Cu3(CO3)2(OH)2 + 3 Ca(NO3)2 + CO2

What if we were to replace all nitrate by sulphate and all calcium by sodium?

The reaction equation would then be:

3 CuSO4 + 3 Na2CO3 + H2O === > Cu3(CO3)2(OH)2 + 3 Na2SO4 + CO2

The advantages of using these reagents are clear:

1. CuSO4 is cheaper than Cu(NO3)2. With calcium, Cu sulphate could not be used because CaSO4 is insoluble.

2. The source of carbonate ions/CO2 is now water soluble, instead of CaCO3, which is not.

I think the insolubility of limestone is what makes this reaction very slow (at least at these low temperatures).

I expect the (modified) reaction to proceed similarly, with Malachite being formed first, then the Malachite converts to Azurite because of the high CO2 pressure, but faster.

Replacing the ingredients mol for mol should lead to the same pressure build up.

So tomorrow I will load solid CuSO4.5H2O and a saturated solution of Na2CO3 in the correct ratio into the PET bottle and hopefully Azurite will also form but faster this time.

[Edited on 19-12-2014 by blogfast25]

CHRIS25 - 19-12-2014 at 10:45

Reference the brain-wave above - Did you not advise against using a powder?

Not that beforehand I know what difference this makes, but re-reading Bauer I was struck by the fact that he mentions Chalk. Now I did some searching and discovered that Chalk is in fact a product of biological secretion type calcium carbonate, whereas Limstone (generally speaking there are exceptions in Fossilecerous Carbonate), anyway limestone from a quarry is normally known as Chemical calcium carbonate, this is over-simplified, but on further reading I got the impression that the two are more different than I first thought. As I said probably nothing, but at least this is new to me.

As for your suggestion above, since I have just made a huge amount of Copper nitrate, I will be getting a second and a third bottle now, one for Seashells, one for limestone, and one to follow what you are recommending above.

The PDF list on the rock compositions gave the following:
Azurite: 25% Carbonate, 69% copper oxide and Water 5%
Malachite: 20% carbon dioxide, 72% Copper oxide and water 8 %

Why would they list the carbonate composition for azurite yet replace that with the carbon dioxide composition for malachite? According to the same source Azurite metamorphs to malachite at the loss of Hydroxyl groups, I don't see how this is so when one looks at the formulas for both, especially since it appears that it is the composition of carbon dioxide that has the strongest influence on its character?

[Edited on 19-12-2014 by CHRIS25]

[Edited on 19-12-2014 by CHRIS25]

blogfast25 - 19-12-2014 at 14:06

Chris:

I advised against powder in the case of limestone (CaCO3) because I was using an excess (with respect to stoichiometry) of CaCO3. Had I used powdered CaCO3, the unreacted CaCO3 would have ended up with the Azurite, impossible to separate the two.

But Na2CO3 is highly water soluble, so any excess will be washed out of the Azurite when I clean it up.

Using CuSO4 and Na2CO3 I expect the Malachite to form very quickly, then it will convert to Azurite at unknown rate when CO2 pressure is achieved. Wait and see, I guess…

Re. limestone/seashells etc, what is important in the preparation re. Brauer is that one uses CaCO3, not where that CaCO3 comes from. Ultrapure synthetic CaCO3 would work the same as more natural forms like the ones you listed.

The theoretical amount of CO2 in pure, anhydrous Azurite, Cu3(CO3)2(OH)2, is as follows: MM of Azurite is 3 x 63.5 + 2 x 12 + 8 x 16 + 2 = 344.5 g/mol. CO2 content of 1 mol Azurite is 2 x 44 g/mol = 88 g/mol. W% CO2 = 88 / 344.5 x 100 = 25.5 w%.

Not sure why they revert to ‘carbonate’ either.

The pseudomorphing reaction of Azurite to Malachite is the following reaction in reverse (left is Malachite, right is Azurite)

3 Cu2CO3(OH)2(s) + CO2(g) < === > 2 Cu3(CO3)2(OH)2(s) + H2O(l)

After that absorption of H2O/loss of CO2 the 'new' molecule then rearranges itself to a new structure (that of Malachite, that is where these stories of Azurite mysteriously turning into Malachite come from).

I can’t put it more objectively than that...

[Edited on 19-12-2014 by blogfast25]

blogfast25 - 20-12-2014 at 08:19

Well, the Azurite experiment was an epic failure.

I opened the bottle and it fizzed like mad, including the slurry: lots of CO2, so far, so good and as expected.

I sieved off the chunks of limestone, washed and dried and weighed them but there was hardly any weight loss at all: not good.

Then the precipitate was Buchnered off which appeared blue but lost all colour on washing: the blue had simply been caused by unreacted copper nitrate in the supernatant. I washed this filter cake with two portions of acetone and dried it at 100 C for 1 h. I got about 8 g of this white powder.

I performed the CO2 determination as above and… found no fizz whatsoever, not even a bubble!

My working hypothesis is that it is calcium sulphate, due to residual sulphate in my copper nitrate. I may test this hypothesis. With so much CO2 around, calcium would otherwise stay in solution as Ca(HCO3)2.

So that was a damp squib, sadly.

In the next experiment, 75 g of CuSO4.5H2O powder (0.3 mol) and 32 g of Na2CO3 powder (anh.) (0.3 mol) were loaded into a 600 ml colourless and clean PET bottle and dry-mixed by shaking. 100 ml of water was added and the bottle quickly capped and then shaken. Green Malachite and CO2 (the bottle stiffened up considerably) were formed immediately:

2 CuSO4 + 2 Na2CO3 + H2O === > Cu2CO3(OH)2 + 2 Na2SO4 + CO2

I now live in hope that the captured CO2 and resulting pressure will convert the Malachite to Azurite, over time. At least Malachite did certainly form this time...

[Edited on 20-12-2014 by blogfast25]

CHRIS25 - 20-12-2014 at 09:09

Well, the Azurite experiment was an epic failure.
My working hypothesis is that it is calcium sulphate, due to residual sulphate in my copper nitrate. I may test this hypothesis. With so much CO2 around, calcium would otherwise stay in solution as Ca(HCO3)2.

[Edited on 20-12-2014 by blogfast25]

How did you make your copper nitrate then? you must have used copper sulphate to make the nitrate? I am just boiling down some copper nitrate at this moment, used copper and nitric acid. So I will still go ahead and repeat your attempt but will think about what other parameters might be changed, if any. Not that I can improve on this, except that I won't get any sulphate.

[Edited on 20-12-2014 by CHRIS25]

blogfast25 - 20-12-2014 at 11:09

Yes, I started from CuSO4, precipitated it as Cu(OH)2, filtered and washed (but obviously not enough!), then dissolved it in nitric acid and boiled down to a known final volume.

The identity of the mystery non-fizzing white precipitate as CaSO4 was just confirmed as follows. The material was treated with an excess of saturated Na2CO3 solution at about 70 C for an hour with magnetic stirring. This would cause the CaSO4 to convert to the much less soluble CaCO3:

CaSO4(s) + Na2CO3(aq) === > CaCO3(s) + Na2SO4(aq)

The slurry was then filtered and the filter cake washed with copious amounts of hot DIW, then a smidge of acetone. It was dried on filer at 100 C for an hour.

Treated with 50 v% H2SO4 this dry filter cake fizzed like mad, with white CaSO4 precipitating:

CaCO3(s) + H2SO4(aq) === > CaSO4(s) + CO2(g) + H2O(l)

So that proves it's mainly CaSO4.

I'm not convinced the sulphate in the copper nitrate was the cause of the failure of this first Azurite experiment, but can't rule it out either. I have much more faith in the newest experiment.

[Edited on 20-12-2014 by blogfast25]

CHRIS25 - 20-12-2014 at 11:47

Just found this: http://www.sciencemadness.org/talk/viewthread.php?tid=3623#p...

[Edited on 20-12-2014 by CHRIS25]

blogfast25 - 20-12-2014 at 13:03

Just found this: http://www.sciencemadness.org/talk/viewthread.php?tid=3623#p...

[Edited on 20-12-2014 by CHRIS25]

Interesting that others have been working on Azurite here on SM.

But there's little there we don't already know.

Interesting tidbit about the copper phosphate, presumably Cu3(PO4)2 or possibly a basic phosphate.

The thread confirms again that there is considerable confusion with regards as to what converts to what but it is now firm and clear in my mind that Malachite is the stable version and that by water and high pressure CO2 it can be converted into metastable Azurite. The latter then slowly or more rapidly converts back to Malachite in the presence of water and other conditions being 'right'.

[Edited on 21-12-2014 by blogfast25]

CHRIS25 - 20-12-2014 at 16:10

Yes I agree, and if these people can make it, http://www.kremer-pigmente.com/en/search.html?keywords=azuri... we can. I also homed in on the copper phosphate.

But I have this as well: Add copper sulphate to a Concentrated solution of sodium carbonate at 45 - 64c for an azurite, and copper sulphate to the bicarbonate at the same temp for a green. I will try this tomorrow or the next and post results, however the former will turn to malachite within 2 hours apparently. We'll see. In fact have just gathered a number of recipes, it's going to be a busy week I think. will post images and results as they are completed.

blogfast25 - 21-12-2014 at 05:52

Blimey! At those prices we should get cracking, presto!

It's unlikely that Kremer actually produce anything themselves though, I'm fairly sure they are traders only, not producers. Going by their website they may do some pigment mixing but I think that's about it.

I'm running a test with an excess saturated Na2CO3 solution, about 50 C, CuSO4 solution stirred in with high speed mixing. See what gives later on...

Update:

So 106 g anh. Na2CO3 (1 mol) was mixed with 265 g of water and magstirred and heated to about 50 C until all carbonate had dissolved. Similarly 25 g of CuSO4.5H2O (0.1 mol) was dissolved in about 200 g of water, to about 30 C.

The sodium carbonate solution stirring speed was then increased to maximum and the copper sulphate solution added to it slowly. The colour changed to a deep blue, so for about 5 ms I thought ‘Bingo!’ But it wasn’t to be: on filtering the filtrate ran a deep blue and it’s caused by cuprate anions: Cu(OH)4-(aq). I’ve seen this anion form many times and it’s caused by the strong alkalinity of the saturated sodium carbonate solution. The deeply intense blue of cuprate makes estimating its concentration difficult.

But there was much precipitate too, which Buchnered reasonably well, was washed a few times with DIW and two minimal rinses of acetone. This product looks far more green than blue to me and is now drying at 100 C for 1 h. I will determine CO2 content later on.

Update 2:

Ok, this is another failure: on drying this product went to a dark green, then developed black patches. These are due to Cu(OH)2 dehydrating to black CuO. Here’s a photo comparing it to a standard light blue circle printed with my inkjet:

This is crappy Malachite, in short. Not worth determining CO2 content on.

And it further confirms my opinion that Azurite cannot be prepared by atmospheric pressure precipitation from a cupric solution and a soluble carbonate/bicarbonate.

Here's also 'the real deal' (from Russia):

http://www.alibaba.com/product-detail/Azurite-pigment_106375...

[Edited on 22-12-2014 by blogfast25]

aga - 22-12-2014 at 14:28

The stuff i made is closer to the Russian True Blue, yet not exactly the same colour.

Copper Sulphate + Sodium Sulphate.

I've been waiting over a week for it to go Green, and it hasn't.

blogfast25 - 23-12-2014 at 05:02

Copper Sulphate + Sodium Sulphate.

Do use emoticons to underline your jokes!

Seriously though, determine CO2 content as I showed above, to establish identity of the material.

[Edited on 23-12-2014 by blogfast25]

blogfast25 - 23-12-2014 at 12:36

Chris:

What you should do is run one of your tests in such a way that you can measure what I wrote above:

Quote-

Malachite:

2 Cu2+(aq) + 2 CO32-(aq) + H2O(l) === > Cu2CO3(OH)2(s) + CO2(g)

So for every 2 mol of carbonate 1 mol is shed.

Azurite:

3 Cu2+(aq) + 3 CO32-(aq) + 2 H2O(l) === > Cu3(CO3

2(OH)2(s) + CO2(g)

So for every 3 mol of carbonate 1 mol of CO2 is shed. (Of course Azurite cannot be prepared that way).

Unquote-

By weighing everything carefully and measuring the weight loss after reaction, the molar ratio of carbonate used/CO2 lost can be established.

Mass balances don't lie: if the molar ratio was 2 (or close), what you precipitated was Malachite, it simply cannot be otherwise stoichiometrically speaking.

In any case, if that paper on M/A stability shows us one thing it's that in ordinary precipitation conditions pH will ALWAYS be (way) too high for Azurite to form.

Reading that paper reinforces my belief that Debray had the right idea and that it's me who is doing it wrong.

[Edited on 24-12-2014 by blogfast25]

DJF90 - 24-12-2014 at 00:51

So 106 g anh. Na2CO3 (1 mol) was mixed with 265 g of water and magstirred and heated to about 50 C until all carbonate had dissolved. Similarly 25 g of CuSO4.5H2O (0.1 mol) was dissolved in about 200 g of water, to about 30 C.

The sodium carbonate solution stirring speed was then increased to maximum and the copper sulphate solution added to it slowly. The colour changed to a deep blue, so for about 5 ms I thought ‘Bingo!’ But it wasn’t to be: on filtering the filtrate ran a deep blue and it’s caused by cuprate anions: Cu(OH)4-(aq). I’ve seen this anion form many times and it’s caused by the strong alkalinity of the saturated sodium carbonate solution. The deeply intense blue of cuprate makes estimating its concentration difficult.

Change the order of addition...

CHRIS25 - 24-12-2014 at 03:25

It would appear from further reading that Pressure does not play such an important role in the formation of Azurite - reading about the Delphi bronze statue confirms this (contrary to one of Brauer's reported methods, and we do not know what other variables were present in that situation). But also I found the following PDF on the stability of malachite and azurite, it seems also that precipitating azurite from solution requires a slightly acidic environment PH 6.5 yet the malachite can also precipitate at this PH but will definitely precipitate above this level. Furthermore the PDF demonstrates that it is carbonate/bicarbonate levels that are the key factor in bringing around azurite and not the CO2, CO2 certainly plays a role indeed, but it is not the main role as is demonstrated by the Delphi bronze statue where it was not until neighbouring stonemasons work with limestone began that the statue was covered in a thin film of calcite and a blue on the statue appeared, in the presence of CO2 from the grounds beneath. The CO2 alone did not cause the blue otherwise it would have remained that colour down through the years. Rather the calcite dust increased the pressence of the CO2 and influenced PH favourably.

These are just some very sketchy conclusions from a brief read of the two sources.

I will get round to further experiments after I have digested the information in more detail.

[Edited on 24-12-2014 by CHRIS25]

Attachment: Azurite Malachite CO2 effects.pdf (366kB)
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Attachment: Azurite Analysis Delphi.pdf (677kB)
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[Edited on 24-12-2014 by CHRIS25]

blogfast25 - 24-12-2014 at 07:21

Quote: Originally posted by DJF90

Change the order of addition...

Nonsense.

It would appear from further reading that Pressure does not play such an important role in the formation of Azurite - reading about the Delphi bronze statue confirms this (contrary to one of Brauer's reported methods, and we do not know what other variables were present in that situation). But also I found the following PDF on the stability of malachite and azurite, it seems also that precipitating azurite from solution requires a slightly acidic environment PH 6.5 yet the malachite can also precipitate at this PH but will definitely precipitate above this level. Furthermore the PDF demonstrates that it is carbonate/bicarbonate levels that are the key factor in bringing around azurite and not the CO2, CO2 certainly plays a role indeed, but it is not the main role as is demonstrated by the Delphi bronze statue where it was not until neighbouring stonemasons work with limestone began that the statue was covered in a thin film of calcite and a blue on the statue appeared, in the presence of CO2 from the grounds beneath. The CO2 alone did not cause the blue otherwise it would have remained that colour down through the years. Rather the calcite dust increased the pressence of the CO2 and influenced PH favourably.

These are just some very sketchy conclusions from a brief read of the two sources.

I will get round to further experiments after I have digested the information in more detail.

I'm studying the B.W. Vink one, great find BTW!!! In that pH < 7 lies the rub of course, as it is basically impossible to create carbonate rich conditions at that pH because of:

CO₃^2- + 2 H⁺ === > CO₂ + H₂O

Hence Debray's reasoning to increase carbonate concentration by high pressure CO2. Debray experiment is certainly on the correct side of pH = 7, i.e. slightly lower.

Chris, while it may well be true that Azurite may be formed at lower pressure, Le Chatelier's Principle still holds for:

3 Cu2CO3(OH)2 + CO2 < === > 2 Cu3(CO3)2(OH)2 + H2O

... because the 'consumption' of CO2 reduces the overall pressure of the system and that's how equilibria operate.

Re. Delphi: does the Delphi paper PROVE Azurite was formed? I've yet to read it.

Edit: Plutarch, wow, that's 1 century AD! It's like a testimony on the colour of Jesus' sandals!!! In short: unverifiable...

[Edited on 24-12-2014 by blogfast25]

CHRIS25 - 24-12-2014 at 07:39

I read "quickly' and have not yet gone through this slowly and digested, but I certainly get the sense that a big part of the statue colour turned to azurite, and that a blue patina may have been applied then lost again. Plutarch writing around 100 AD roughly comments on the statue's beautiful blue, and this was 400 years After it was made. The various analysis of this statue conclude (roughly, briefly) that Plutarch's observations were a clear sign that an azurite patina had formed as a direct result of building and restoration work during his time close by the statue, the position of the statue also giving it excellent opportunity to turn blue, something to do with CO2 vents from underneath in the limestone bedrock and the calcium carbonate dust that would have settled on the statue during the nearby works. But this is just my first impression of a quick read Gert.

What does B.W. Vink one mean?

Quote - In short: unverifiable...? The paper verifies his observations quite ingeniously.

[Edited on 24-12-2014 by CHRIS25]

blogfast25 - 24-12-2014 at 07:43