RaD
Harmless
Posts: 4
Registered: 24-4-2015
Member Is Offline
Mood: No Mood
|
|
Copper Carbonate Synthesis
If the formula is: 2CuSO<sub>4</sub> + 2Na<sub>2</sub>CO<sub>3</sub> + H<sub>2</sub>O -->
Cu<sub>2</sub>(OH)<sub>2</sub>CO<sub>3</sub> + 2Na<sub>2</sub>SO<sub>4</sub> +
CO<sub>2</sub>. Does that mean that I would need to use 499.37g of CuSO<sub>4</sub>, 211.98g of
2Na<sub>2</sub>CO<sub>3</sub>, and 18.015g of water to get a yield of ~221.116g of
Cu<sub>2</sub>(OH)<sub>2</sub>CO<sub>3</sub> + 284.08g of Na<sub>2</sub>SO<sub>4</sub> +
CO<sub>2</sub>?
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
Yes and no. Those are the stoichiometric amounts, but you are assuming that you have copper sulfate pentahydrate and anhydrous sodium carbonate, are
you sure you do? In addition, the reaction needs a solvent(water is the best choice) and both reactants are preferably dissolved prior to mixing
slowly in a large container to prevent overflow from the CO2 foam. Mixing the two solid chemicals and 18ml of water is not going to get
you your product. In addition, be aware that while yields will be good, they will not be quantitative, likely if you are inexperienced they will be
around 90% due to mechanical losses/purity issues.
|
|
|
RaD
Harmless
Posts: 4
Registered: 24-4-2015
Member Is Offline
Mood: No Mood
|
|
Thanks! Would using more water affect the amount of reactants needed for the equation to remain true?
|
|
|
Amos
International Hazard
Posts: 1406
Registered: 25-3-2014
Location: Yes
Member Is Offline
Mood: No
|
|
Nope, an excess of water is just fine. Both sodium carbonate and sodium bicarbonate can be used as the base. I recommend that you either use sodium
bicarbonate, or use cold solutions when using sodium carbonate. Because copper(II) carbonate is amphoteric, it can react with hot sodium solutions to
partially re-dissolve as cuprate, which will yield a blue solution that cannot be further precipitated using more carbonate. And since the dissolution
of sodium carbonate in water is exothermic, it will get pretty warm.
Last thing, be careful of fizzing; the bicarbonate reaction gives off a lot of carbon dioxide.
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
Nope, water in this case is used in large excess as a solvent. This is common in many reactions which require a solvent, if one of the reactants is a
relatively inexpensive liquid, it is simply used in a large excess to dissolve everything. Technically, you could use 2.5L of water here as that is
the minimum amount required to dissolve everything, but using a larger amount such as 3L will lower the number of copper carbonate particles suspended
in solution and make it easier to filter. When making magnesium hydroxide, if I use concentrated solutions of both magnesium sulfate and sodium
hydroxide I get an unfilterable gel. Copper sulfate takes a long time to dissolve in water, warming the solution will help to speed up the process.
To get an idea of what the process looks like, there is a video here : https://www.youtube.com/watch?v=Uorlv05KLX4 . This reaction may be an exception to the case I mentioned above, you should be able to use 2L of
hot water to dissolve the copper sulfate, then use a minimum of hot water to dissolve the sodium carbonate. The video doesn't do this, but this can
result in sodium contamination in the final product.
EDIT: I didn't think of that Amos, have you had any personal experience with redissolution to cuprate? I've never had that issue.
[Edited on 4-25-2015 by gdflp]
|
|
|
Amos
International Hazard
Posts: 1406
Registered: 25-3-2014
Location: Yes
Member Is Offline
Mood: No
|
|
You might want to check this, and the rest of the thread as well: http://www.sciencemadness.org/talk/viewthread.php?tid=55588#...
Yes, cuprate does form if you use excess sodium carbonate, I've had it happen a couple of times(once on purpose, so I could bottle some solution; it's
really rather pretty); It's not necessarily a bad thing if all you desire from the reaction is basic copper carbonate, especially given that the
reaction with sodium carbonate is twice as efficient per mole compared to sodium bicarbonate. But if you want a true double replacement reaction and
to collect sodium sulfate as well, you might not want cuprate hanging around; I believe it would convert back to copper sulfate if you attempted to
acidify the leftover solution as a means of collecting sodim sulfate.
[Edited on 4-25-2015 by Amos]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by gdflp  |
EDIT: I didn't think of that Amos, have you had any personal experience with redissolution to cuprate? I've never had that issue.
|
A little perspective here. In the thread that Amos linked to the poster obtained a slightly alkaline, (quote) "weakly blue" supernatant, after the
precipitation. The copper(II) hexaaqua ion cannot exist in alkaline conditions, the only soluble copper form then is cuprate, which is a deep, cobalt
blue. The poster obtained a light blue solution, indicating that cuprate was at very low concentration.
Get the stoichiometry more or less right and you're unlikely to get any significant quantities of cuprate when preparing Malachite.
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
As mentioned above, a common mistake is forgetting to account for crystal water in your calculations. If your copper sulfate is blue, it is actually
the pentahydrate: CuSO<sub>4</sub>*5H<sub>2</sub>O. This has a molecular weight of about 249.5 g/mol, as opposed to 159.5
g/mol for the white anhydrous variety. A pretty large difference that will throw off your stoichiometry.
Making copper compounds is a lot of fun, and in my experience colors can be quite different from experiment to experiment. I've made copper carbonate
that appeared as a range of colors from baby blue to light green. I think it has something to do with chloride impurities (if you start with
CuCl<sub>2</sub> or the variable nature of what we call "copper
carbonate".
I've also encountered the weakly blue solution after precipitation. Stubbornly adding more carbonate does not solve the problem  If this is indeed from cuprate, the very light blue color says that it's a very
dilute solution.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
MrHS:
The perceived ‘variability’ is in my opinion due to various and often conflated subjectivities. Moisture content, granulometry and
lighting all contribute to small variations in perceived colour which tend to level out if the product has been prepared properly and washed, dried
and ground in comparable conditions. The ‘true’ colour is also hard to capture photographically, contributing to the general confusion. Copper
basic carbonate has a very well defined chemical composition, despite its ‘reputation’ for variability, which is further reinforced by erroneous
but persistent claims about 'CuCO<sub>3</sub>'.
The influence of chloride could be worth investigating because chloride contamination is so easy to determine. But personally I doubt that it
contributes much.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
