Basic Copper Carbonate - a Copper Series

Hi all,

I've been fairly interested in doing a series on copper compounds and complexes for a while now(my first attempt at precipitating copper metal failed initially). So I thought I'd revisit this and try to post my experiments with proper documentation, and for the first lab I decided on basic copper carbonate.

There are a wide variety of copper compounds that can be made through a number of different ways. Many of them can be synthesized from extremely easy to obtain chemicals, and basic copper carbonate is one of the easiest to make from off the shelf products. This makes it a good starting point.

Making Basic Copper Carbonate

Step 1: The Ingredients

We really only need 3 chemicals to make this compound. Water (also known as H2O), Copper (II) Sulfate (also known as root killer), and Sodium Bicarbonate (also known as that thing my mom puts in the fridge a lot for some reason).



All of these were purchased at local stores(the root killer was bought at my hardware store, and the water and bicarb at my supermarket).

Step 2: The Stoichiometry

There are two very simple routes we can take to synthesizing our compound. We can either react an aqueous solution of Copper (II) Sulfate with Sodium Carbonate, or Sodium Bicarbonate. In my case I chose to go the bicarbonate route since it's just not necessary adding an extra step for this reaction in making Sodium Carbonate. The balanced reaction is:

2 CuSO4 + 4 NaHCO3 → Cu2(OH)2CO3 + 2 Na2SO4 + 3 CO2 + H2O

And if you feel like doing your experiments in 10-gallon buckets because you've got a giant lab, you absolutely can do that. For me however, I chose a much smaller .1 scale for my tiny glassware, which corresponds to the following:

49.9 g CuSO4(aq) + 33.56 g NaHCO3 → 12.35 g Cu2(OH)2CO3 + 13.2 g CO2 + 1.8 g H2O

Step 3: The Procedure

First we measure out about 50 grams of our Copper (II) Sulfate, and get it in about 200mL of water. In my case, my copper sulfate crystals are so gigantic because I did a recrystallization from my root killer. Doing so is absolutely not necessary for this lab, and you can simply use the ground up pellets from the root killer directly.



It'll take a while for everything to dissolve into solution. Simply keep it on the hot plate and stirring, eventually it will become clear. Note: it's not necessary to use a hot plate, you can also simply just microwave the solution for a few seconds at a time to get the water near boiling and the copper sulfate will dissolve. It's just a little bit stubborn. Make sure to turn off the heat afterwards.



Next, measure out 33.56 grams of Sodium Bicarbonate, and EXTREMELY SLOWLY and with fast stirring, and I cannot emphasize this enough EXTREMELY SLOWLY NO REALLY YOU'RE GOING TOO FAST, begin adding the bicarbonate to the sulfate solution. Use no more than a teaspoon per addition until the bubbling dies down.

We are producing CO2 gas very quickly, and it is so easy to become overconfident and add too much of the bicarbonate too fast. This will cause it to bubble out of the container. In my case, all I had to work with was a small 250mL flask. For a .1 scale of this reaction, I would recommend a 1-liter flask or beaker just to make it easier to add the bicarbonate. It took me about 30 minutes to add 33.56 grams slowly.

In the event of an overflow and spill, it is important not to panic. Simply wipe up the liquid with some paper towels, wipe off the exterior of the reaction vessel, and put the paper towels in a waste container for disposal.



The most important step in any lab procedure is safety. Although the copper compounds we are working with are not very harmful, other chemicals and copper compounds can be very toxic and must be handled with care. Again, slow addition of the bicarbonate and patience would have prevented this.

Once the addition is finished, let the mixture continue stirring and reacting for a good half hour or so, and the CO2 bubble production should wind down near the end. You should wind up with a blue-ish green precipitate that looks something like this:



Let the precipitate settle to the bottom of the flask, then wait until the bubbling dies down completely. The liquid at the top layer of the flask should be very nearly clear and colorless once it's completely reacted.

Finally, decant off the top layer of liquid, and filter out the blue-green precipitate. Use a small amount of water to wash out the flask and filter everything through. You should wind up with a greenish-blue wet paste. This can be left to dry in open air for a day or two until it becomes brittle enough to powderize (I was not able to take a picture of these final steps since my phone decided then would be a good time to form a labor union and go on strike against me ;_; ).

Step 4: So What Happened?

We are essentially performing an acid-base reaction, where the Copper (II) Sulfate solution (which is acidic) reacts with the Sodium Bicarbonate, which is a mild base. In common acid-base reactions, typical byproducts will be CO2 gas and water. If you have these byproducts in a reaction, that can be a fairly good indicator of an acid-base reaction. By adding the base we lower the acidity of the solution, precipitating our desired product.

Instead of creating a neutral compound however, our precipitate is actually a base, characterized by the addition of an OH- (hydroxide) molecule attached to the compound, as well as the CO3- (carbonate) molecule. In this form, our final copper salt can be reacted fairly easily with a wide variety of different chemicals to produce many different compounds in different oxidation states.

Note: There are different types and definitions of the terms 'acid' and 'base'. In the case of the classic acid-base reactions where CO2 and water are produced, it's usually assumed that the acid in question is an H+ ion donator, and the base is an OH- recipient for the H+ ion. These are known as Arrhenius acids and bases.

Step 5: In Conclusion

This is a fun little experiment which teaches some basic chemistry concepts and can be done with kids under adult supervision and with proper safety gear. Clean up is fairly easy and all solid waste is tossed, all liquid waste can generally be safely poured down the drain with copious amounts of water depending on your local sewage waste disposal laws and guidelines.

If you feel like I missed something important about this procedure or got something wrong, please let me know and I'll make an update. Thanks for reading

Have a look at the SM Copper Carnival

Quote: Originally posted by Sulaiman

http://www.sciencemadness.org/talk/viewthread.php?tid=62608

After about a day of open-air drying this is what it'll look like (the shade of blue/green will be a little different since my lighting and camera are awful):



You might find white specs and discoloration on the powder. This is likely either hydroxide decomposition or left over bicarbonate. There are some ways to avoid this, but in general it should be fine for most copper complex demonstrations.

Well done, Sciencebum! BigD is correct; the pigment ultramarine blue was renowned by Renaissance artists for its beauty, though many such works have turned a rather dull black copper (II) oxide. I produced some basic copper carbonate not long ago, rather a sticky product, much like copper hydroxide! Included are photos azurite and of the beautiful verdigris-colored pigment, I dried the material in a 'double boiler' essentially a steam bath produced by a pyrex bowl over a vessel of boiling water. Then I sifted the product to a medium fine consistency.

Quote: Originally posted by Syn the Sizer

Quote: Originally posted by Sulaiman   http://www.sciencemadness.org/talk/viewthread.php?tid=62608 Nice, that would have been a fun contest, I wish I was on the forum then.

Quote: Originally posted by Big D

Do you or any other chemists know how to make the azurite basic copper carbonate

Quote: Originally posted by Tsjerk

1. Azurite can be made by slowly adding a copper sulfate solution to a solution of carbonate. The order of addition seems to be key here, my guess is that azurite needs basic conditions to form. 2. Or you could do it via electrolysis of a saturated bicarbonate solution with a copper anode. Electrolyse until black copper oxide starts forming on the anode and filter the solution. Leave the solution until crystals start to form. [Edited on 5-5-2020 by Tsjerk]

Azurite, as a pigment, is also known as Blue Verditer.

1. Yes, Azurite needs alkaline conditions but whether the order of addition under high agitation really matters I don't know.

I mixed some 6 g Na₂CO₃ (anh.) in 100 g DIW, then with high stirring I added drop by drop sat. CuSO₄, until no more precipitate formed and the supernatant was light blue:



This looks suitably blue for Azurite.

I filtered off the precipitate and washed it:




Finally washed with MeOH and dried, I weighed 0.33 g of product, which is very low compared to the stoichio amount of 9.75 g.

I checked the supernatant and it was strongly alkaline and bubbled strongly on addition of HCl.

I'll determine the CO₃^2- of the product tomorrow.

2. You're referring to the method described here:

https://youtu.be/UsaANlJo_OU

I've tried that 3 times no, w/o success, despite following 'AScannerClearly' guidelines.



Someone in the comment section of that UToob suggested the product might be Chalconatronite (Na₂Cu(CO₃)₂.3H₂O). Chalconatronite is slightly soluble, so formation of crystals from a mother liquor is quite probable.

Another method using Cu(NO₃)₂ and CaCO₃ is described here:

https://youtu.be/Y9uZpZBy76w



[Edited on 4-2-2023 by blogfast25]

Here's my ad hoc 'gazometer', 25 mL capacity:



Samples were weighed to 1 mg precision (typically 0.1 g). Gas readings were adjusted for hydrostatic pressure and temperature.

Reaction is banal:

Carbonate(s) + HCl(aq) ---> CO₂(g) + Chloride(aq)

Using two samples of basic copper carbonate described above I obtained values of CO₂ w% of 14.2 and 14.7.

This doesn't really match the CO₂ contents of either Azurite or Malachite, of resp. 25 w% and 19.9 w%.

But during development of the gazometer I repeatedly measured the values for NaHCO₃ and Na₂CO₃ and found systematic under estimation of 4 - 5 w%.

Adding that error to the values for the basic copper carbonate would then put its value in the 20 w% range, consistent with Malachite, not Azurite.

And after drying the basic carbonate looked more like the blue green (green blue?) of Malachite than the azure blue of Azurite.

&&&&&&&&&&&&&&&&&&&&&

Next attempt will be to add either CO₂ or CO₃^2- to a solution of cuprate (Cu(OH)₄^2-.

[Edited on 10-2-2023 by blogfast25]

Quote: Originally posted by maldi-tof

My experience says that what you need to control in order to get azurite or malachite is the temperature of the reaction. What I do not recall is what needs to be at higher temperature and which one at lower temperature.

Quote: Originally posted by blogfast25

My latest attempt with CuCl2 at about 18 C failed. I got Malachite (again!) [Edited on 11-2-2023 by blogfast25]

Quote: Originally posted by Amos

You might want to start from a different water-soluble salt. [Edited on 2-11-2023 by Amos]

So I tried 'TheAlchemicalArts' method as shown in his long winded video (above), minus some of the obvious BS designed to make TAA look smart and... erm... alchemical!

I used 6 g of copper nitrate trihydrate (which arrived a few days aqo) and 2.5 g of calcium carbonate, making a molar ratio of 1:1, in 80 ml of DIW.

I kept the temperature strenuously below 12 C, stirring every hour or so, for about 2 days.

All I got was more Malachite, while the supernatant was colourless.

The Azurite Gods do not favour me...



[Edited on 19-2-2023 by blogfast25]