Sciencemadness Discussion Board

Impure ZnCO3?

OctanitroC - 15-1-2013 at 14:48

Alright, so I made some Zinc Carbonate as follows:

I added impure zinc (melted pellets from the inside of pennies) to 40 mL water, added a small amount of CuCl2, brought to a boil (to speed up reaction on the bulk metal), and keep boiling, adding CuCl2 as necessary when the solution went clear. After adding a small amount, perhaps 3-5 g, of CuCl, which had precipitated enough "spongy" copper precipitate to just cover the bottom of a 50mL beaker, I hot filtered the crude ZnCl2 solution into an erlenmeyer and obtained a crystal clear solution. I then added concentrated Na2CO3 solution in quantity until no more gelatinous precipitate was formed, and filtered this precipitate out. I pressed out water from the precipitate and obtained a still very wet looking cake of ZnCO3. This was air dryed until slightly damp and then stirred with ethanol to dehydrate it, which I presumed would be easy to evaporate off after decanting.

This is where my problem comes in; after decanting almost all EtOH from the thick suspension of carbonate in the alcohol, I figured that burning off of the remaining solvent would be nondetrimental, as the decomposition point of ZnCO3 is 300 degrees C or thereabouts. However, as the carbonate started to appear dry, the leading edge of the burning ethanol changed the carbonate to a dark black color, which faded to lemon yellow and then light grey as it cooled. After the ethanol finished burning off, I broke up the clumps of dry carbonate and mixed the powder. This does still seem to be carbonate, however when mixed with vinegar, only the white powder fizzes and dissolves. the grey powder is left behind. So what is this other off-white substance? I don't think I decomposed the carbonate, it doesn't look at all like the snow-white color of ZnO. Maybe copper remains? Other metal contamination? Dunno, do you guys have any ideas?

99chemicals - 15-1-2013 at 16:18

It brings this to mind but yours is persistent. It could be a CuO impurity. I would repeat the experiment with using HCl instead of CuCl2

OctanitroC - 15-1-2013 at 16:44

That is the exact shade it turned when hot. I think it could also be carbonized impurities, the EtOH I used is only 95%, denatured. I'll heat it all strongly, it doesn't matter which salt I have. I just wanted a reasonably pure zinc reagent to explore its chemistry. I don't have any high-quality HCl, although I'm ordering some from elemental soon. Is there an acid that will preferentially dissolve ZnO over CuO?

12AX7 - 15-1-2013 at 19:37

Burning is a process which generally occurs well over 300C. The edges, light and powdery with their newfound absence of solvent, were dutifully heated by combustion. IIRC, the stuff turns yellow over 500C or thereabouts. Not sure about the gray -- maybe some impurities from the ethanol itself (incomplete combustion within the powder?), or Cu or other residues (Cu, Al, Mg, Fe and Sn are common impurities of Zn). Maybe even trace metallic Zn.

I see two issues with your process: you didn't mention a washing step, and the precipitation was done hastily, resulting in extremely small particles ("gelatinous"), which tends to adsorb a lot of whatever's in solution. Certainly, you'll have a huge sodium impurity -- when the ethanol burned off, was the flame especially yellow around the edges?

To address this, I would suggest using dilute, rather than intensely concentrated, solutions, and mixing them slowly. Remember, you're in no hurry to precipitate -- not only do you want it done slowly, but the solubility of ZnCO3 is low enough that you don't have to worry about leaving a lot behind.

In fact, you probably should leave some zinc in solution, which will help make sure that the last thing precipitated had the (in)solubility of ZnCO3, and not something otherwise more soluble which would stay in solution if you stopped early.

By extension, you should ideally pre-precipitate the fresh solution, just to expunge particularly insoluble materials, Fe(OH)3 for instance. In other words, after dissolving the substrate in acid, titrate to a murky neutral pH, then filter and discard the precipitate. Then proceed with the main precipitation; stop when the pH starts to rise again. Finally, you can cook the solution, to sort-of "finish" precipitation and get the particles to consolidate; the suspension should settle faster and filter easier.

Come to think of it, carbonate is a good anion to use: at higher pH, it forms a complex with copper, which will tend to keep it in solution rather than co-precipitating. When filtering, it might not hurt to do the second wash with very weak Na2CO3 solution. Then finish up with lots of distilled or RO water.

For further purification, I suggest dissolving in sulfuric acid and recrystallizing the ZnSO4.7H2O product. Keep iron (specifically Fe(II)) and magnesium away, these will form solid solutions in the crystal. IIRC, it has a crystal habit similar to Epsom Salts; as purity improves, you should be able to grow reasonably sized, colorless, columnar crystals.

Tim

violet sin - 15-1-2013 at 20:17

hey just a note. easier than destroying penny's ( or for me at least) you can find "1.5 ton crab pot zinc's" that are ~1# of fairly pure Zn (or so I'm told) for 4.65$ each. mine came in a large gumdrop shape, with 4 stainless whiskers sticking out the flat side. they are to tie to stainless crab pot netting so it isn't "hot". guess the dungeoness crab here on west coast(cali) don't like the pots if they are actively fighting the ocean, but have no prob if the sacrificial anode is easing the stress. my lill brother is out fishing now and that was his explanation. sooooo, long story short, 1# Zn for <5$ in shiny, easy to stock form. I plan on melting some to shot. oh ya, be sure to ask for "crab pot zinc's" b/c they had some for boat motors, props, prop shafts etc. that were easily 80$ for <1/2#

OctanitroC - 16-1-2013 at 03:40

Thanks everyone. I attempted redissolving the precipitate and large amounts of solid black material were left behind. I think I'm going to retry the dissolution with purer zinc and an acid instead of CuCl2, and make sure to wash the precipitate adequately.

12AX7: for slower addition of the carbonate, should I have to add dropwise, or will a dilute solution added in ~10 mL amounts be enough to increase precipitate particle size? I know ZnCO3 is very slightly soluble.

violet sin: Thanks! I actually live right near a boating supply store, I'll check there this afternoon. Have you tested the purity of the zinc any?

Edit:I don't know what I was thinking about ZnCO3 solubility, I'd hardly lose any at 0.21 g/L...

[Edited on 16-1-2013 by OctanitroC]

OctanitroC - 16-1-2013 at 07:26

I was browsing some other threads and saw it mentioned that Zn2+ precipitates carbonate with CO2. If the ZnCl2 solution was reasonably pure, could bubbling CO2 from bicarb+acid through the solution be a viable method of precipitation? It seems that would be better than Na2CO3 because it eliminates risk of contamination. However, I can't figure out the balance of that reaction...It seems that it would have to produce HCl, which would turn around and redissolve the ZnCO3. Would this work?

AJKOER - 16-1-2013 at 09:15

OK, perhaps this problem can benefit from a hydrometallurgical ore processing viewpoint. For example, per (Australian) patent 20100180728, to quote:

"A zinc ore containing primarily hemimorphite was found to contain goethite which coated the hemimorphite restricting leaching of zinc to 43% after 24 h. Curing the ore in oxalic acid solution for 24 h prior to leaching resulted in >90% dissolution during otherwise identical subsequent leaching. Without being bound by theory, the oxalate anions formed a soluble complex with the iron from the goethite thereby removing it from the surface of the zinc mineral which was subsequently accessible to the ammoniacal-ammonium carbonate leach solution."

Apparently, the leaching, curing and oxidizing agents vary with the target metal and metal mix (Zn in our case, or is it Zn with Fe and/or Cu from filing/melting of the penny?) with respect to various solutions' concentration. In addition to concentration for curing, the author states "the temperature at which the curing step occurs, the pH at which the curing step occurs and the time for which the ore is exposed to the curing agent may all be varied in response to the composition, mineralogy, texture and pore volume of the ore (with low pore volumes necessitating higher concentrations). "

Read more: http://www.faqs.org/patents/app/20100180728#ixzz2IBOJcGJq

Note this definitional paragraph:

"As would be understood by a person skilled in the art, the term curing is fundamentally distinct from leaching. Leaching describes a process by which a solution containing a leaching agent is contacted with an ore, the solution recovered and valuable metals extracted therefrom. The curing step of the present invention renders the ore to be leached more amenable to the leaching process, improving both the extent and rate of recovery of the target metal. Without wishing to be bound by theory, this may arise from one or more of the oxidation or reduction of the target metal or otherwise refractory ores containing the target metal, the complexation of metal (target or non-target) and the mobilisation of metal (target or non target)."

Read more: http://www.faqs.org/patents/app/20100180728#ixzz2IA0OH63A


[Edited on 16-1-2013 by AJKOER]

12AX7 - 16-1-2013 at 18:14

Quote: Originally posted by OctanitroC  

12AX7: for slower addition of the carbonate, should I have to add dropwise, or will a dilute solution added in ~10 mL amounts be enough to increase precipitate particle size? I know ZnCO3 is very slightly soluble.

violet sin: Thanks! I actually live right near a boating supply store, I'll check there this afternoon. Have you tested the purity of the zinc any?

Edit:I don't know what I was thinking about ZnCO3 solubility, I'd hardly lose any at 0.21 g/L...

[Edited on 16-1-2013 by OctanitroC]


Ksp is something like 10^-11... shouldn't be a problem. :)

I'd figure doing the addition over 10 minutes or so, by hand, stirring enough to keep it moving. If you could put it on a drip for hours, over a stir bar, that'd be fine too, but you don't need to be babying it, that's a pain.

Tim

violet sin - 17-1-2013 at 01:58

I have not had a chance to check the Zn purity. the metal is stamped "Canada". I specifically asked the store owner if it was pure zinc, and he said ohhhh yeah (not completely convinced lol). funny he used the term "martyr" instead of sacrificial anode. and martyr is also a brand from a canadian supplier. but I will have to look into it.......... ..... .....
http://www.canadametal.com/zinc-anode-zinc-balls.php
based on the name stamped on them and a few min of searching.. I found the above pic from the canadian supplier that look EXACTLY like mine, except mine have the whiskers sunk in for fastening to the pots.
and this one is "martyr" brand alloy composition, but the form is not similar.
http://www.martyranodes.com/content/martyr-resources/alloy-s...
it's not the same company, and I can't seem to find an identical product online. the 4 s.s tie wires being the only missing aspect. though most of these manufactures seem to have nearly pure Zn. likely byproduct from Ni, Cr, Pb, etc. mining and processing. can't really think of any ways to check the purity of my sample off the top of my head either. suggestions? I did mix some fillings with sulfur and got some green flame bursts, but it was unreliable.

12AX7 - 17-1-2013 at 17:41

Well... if nothing else, they have no reason to put junk in them. Just commercial grade zinc, which is probably distilled, and separated from values (Ag, Cd, Se, Te, etc.), so a reasonable purity (>99.9% I think??). Pure zinc is stable enough as-is, and they don't need to alloy it for strength, moldability, corrosion resistance (lol!) or anything like that.

If you wanted to verify yourself, you'd be best served taking a field trip to a university with SEM or XRF equipment. If you can find a receptive professor (geology or metallurgy?), they might do it for free.

The chemistry department tends to do things one-at-a-time, so sure, you could check the concentration of Pb using AAS, but you have to go to the trouble of prepping a sample, in various concentrations, plus a reference, just to see if something isn't even there. Then repeat for 40 other impurities or until you get fed up. I'm not sure how popular ICP-MS and such systems are getting; you might get a reasonably complete one-shot analysis from this, if they have it.

Tim

watson.fawkes - 17-1-2013 at 21:12

Quote: Originally posted by violet sin  
though most of these manufactures seem to have nearly pure Zn. likely byproduct from Ni, Cr, Pb, etc. mining and processing.
An interesting question, so I looked it up in Kirk-Othmer. There's no need for high-grade zinc in a sacrificial anode, and there are two relevant, common commercial grades. One is 98% grade zinc, with Pb added: Pb 1.4%, Cd 0.20%, Fe 0.05% and Al 0.05%. If Pb is cheaper than Zn at the source, expect this alloy. The other is 99.9% grade, with Pb 0.03%, Fe 0.02%, Cd 0.02%. All impurity percentages are maximum level permitted to meet grade requirements. No guarantees on what grade is actually there (or even if it's one of these two), but that should give you an idea.

violet sin - 17-1-2013 at 21:47

that's exactly the type of labor I had imagined. not for me. though I may ask at the local university. only 20 min away and I have to do some errands there soon. CSUC. years back my physics major room mate and I hung out in the labs a bit. at least back then the science professors were quite accommodating to the curious, so couldn't hurt. I fully agree that leavings of more valuable commercial metals or alloying on purpose is unlikely.

the martyr brand I linked above had roughly the same contaminants + a 0.1 - 0.5% Al impurity. I looked for about 45 min and I had trouble finding anything less than 95% Zn anodes. I think they were the machined version meant to fit shafts and specific motors if memory serves. probably had to had to be less brittle in low temp rotating and vibrating precision applications. I don't have any references to that on hand, and I'm tired of reading about it for the time being.

[Edited on 18-1-2013 by violet sin]

OctanitroC - 18-1-2013 at 05:12

So once I can get HCl and H2SO4, I'll retry this with the sacrificial zinc. Depending on the grade, redissolving the carbonate in sulfuric and crystallizing most, but not all, of the zinc out should yield reasonable purity ZnSO4, as PbSO4 is nearly insoluble. Would any other metal try to cocrystallize (ie Cd, Fe, etc) during recrystallization, or do they have sufficiently different crystal habits?

12AX7 - 18-1-2013 at 06:48

Not sure about Cd, but Fe(II), Mg, Mn and others will form a solid solution -- recrystallization won't help separate these contaminants. (Fe ans Mn can be oxidized with H2O2, then precipitated by raising the pH until zinc just starts to come out of solution, leaving Fe(OH)3 and Mn2O3 or MnO2. This is probably not necessary as iron contamination will be apparent.)

Tim

annaandherdad - 20-5-2013 at 12:24

Hi, OctanitroC, we've been doing something similar lately with American pennies, post 1982. Our object was to make some zinc sulfate. The unknown factor is the purity of the zinc in the pennies; of course there is the copper cladding, but I don't know about the rest.

The basic strategy was to dissolve the pennies in sulfuric acid and evaporate to crystallize the zinc sulfate.

In more detail, we first melted the pennies in about 50 gram batches, which is about 20 pennies. This wasn't necessary but was interesting, because the copper dissolves in the melted zinc, so when you pour the melted metal out and let it cool you get nice, silvery slugs, with no evidence of copper. Then we dissolved these in sulfuric acid. I'm fortunate to have access to lab grade H2SO4, so I don't have to worry about impurities in the acid. Gizmodium has a video on youtube where he used hardware grade acid and does a nice job of removing impurities.

As the acid eats away at the zinc, it makes dark spongy masses of copper. Of course these do not dissolve in the acid, and if there are any copper ions in the solution (eg from copper oxides) they are nicely removed by the zinc metal itself, as long as it's in excess. I didn't test the final solution for copper, but there was no color visible in it. The copper masses can be removed by filtering.

When washed and dried, the spongy masses of copper convert to copper oxide in the air. I assume this is because they have so much surface area and the material is wet. When the masses are dried, you can touch them with your finger and they just crumble. In this way we got some nice samples of copper oxide, and a nice separation of the copper from the zinc in the pennies. If we had just dumped the pennies into the acid (scratching or sanding to give the acid access to the zinc) would have gotten empty copper shells, which wouldn't have been as interesting as the copper oxide we actually did get. This was one reason for melting the pennies first.

The zinc sulfate has some excess acid (there's always some amount left), and it won't evaporate. So to get rid of it, we drained off 10% or so of the clear solution and precipitated all the zinc as carbonate, using Na2CO3. A little excess carbonate neutralizes the excess H2SO4 in the sample drained off. The ZnCO3 is very fluffy and holds a lot of water, as you noted. This makes it hard to wash. But if you dry it out (without bothering to wash it), then you can add it back to water and it readily settles to the bottom of the beaker. Somehow its propensity for sucking up water has been destroyed by the first drying. You really can't wash it before doing this first drying, in fact if you leave it over night it takes on the consistency of oatmeal. I'm not sure what is going on there but I assume it's a well known phenomenon. Anyway, the zinc carbonate is now easy to wash by decanting, after which a filter and drying session gives a nice sample of clean zinc carbonate as very fine, white powder.

Now adding some zinc carbonate back to the ZnSO4 + residual H2SO4, you can neutralize the excess H2SO4 (add ZnCO3 until no more dissolves or fizzes). Then filter off the excess ZnCO3 and evaporate the result to get ZnSO4. We easily made 300 gr of ZnSO4.H20 this way. The impurities must be mostly from the impurities in the original Zn, and, as I said, I don't know what these are.

We got into this because I had some ZnSO4 I bought off ebay that was bad. It had brown, insoluble impurities in it, don't know what they were, but I threw the whole lot away. Then it seemed like it would be more fun to make our own ZnSO4 instead of just ordering some.

ElectroWin - 21-5-2013 at 11:41

i've seen lots of zinc-aluminum alloys. any chance there was some aluminum in there?

annaandherdad - 21-5-2013 at 19:39

I don't know. Do you mean, aluminum could explain the gelatinous consistency of the carbonate precipitate? Pennies are usually described as 97.5% Zn, 2.5% Cu (post 1982 American pennies, I mean), but I don't know what else is in there or in what percentage.

blogfast25 - 22-5-2013 at 05:07

Quote: Originally posted by annaandherdad  
I don't know. Do you mean, aluminum could explain the gelatinous consistency of the carbonate precipitate? Pennies are usually described as 97.5% Zn, 2.5% Cu (post 1982 American pennies, I mean), but I don't know what else is in there or in what percentage.


Al doesn't from a carbonate and would precipitate as gelatinous Al(OH)3. Carbonate solutions aren't alkaline enough to hold Al in solution as aluminate. So that would fit.

From what I've read, pure ZnCO3 isn't easy to prepare and several commercial products appear to be either basic zinc carbonate (zinc hydroxy carbonate) or mixtures of ZnCO3 and basic zinc carbonate(s ?). Not sure whether that could explain the gelatinous nature.

annaandherdad - 22-5-2013 at 08:19

I did a little googling on the zinc in American pennies, including some web sites from the US Mint. They use an alloy of zinc called A190, which is manufactured by Jarden Zinc Products. The JZP web site says the A190 alloy is 0.7 to 0.9% copper, with the balance being 99.9995% pure zinc. So I guess there's no aluminum.

References:
http://jardenzinc.com/techdata/Tech_Brief_Zinc_Alloys.pdf
http://www.usmint.gov/about_the_mint/PDFs/United_States_Mint...

ElectroWin - 17-6-2013 at 10:37

hmm... maybe you're getting gelatinous zinc hydroxide?

"Zinc(II) ion reacts with aqueous ammonia to precipitate white gelatinous Zn(OH)2:"
-- http://www.public.asu.edu/~jpbirk/qual/qualanal/zinc.html

blogfast25 - 17-6-2013 at 12:31

Quote: Originally posted by ElectroWin  
hmm... maybe you're getting gelatinous zinc hydroxide?

"Zinc(II) ion reacts with aqueous ammonia to precipitate white gelatinous Zn(OH)2:"
-- http://www.public.asu.edu/~jpbirk/qual/qualanal/zinc.html



Annaandherdad are using Na2CO3 to precipitate the Zn2+, they should at least get something that contains CO32- ions because ZnCO3 really does exist.

But the Ksp for ZnCO3 is 1.46×10-10 and the Ksp for Zn(OH)2 is 3×10-17. That would indicate Zn(OH)2 to precipitate preferentially over ZnCO3, at least in strongly alkaline conditions. NaHCO3 may be a more appropriate base here.

To complicate matters further, Zn(OH)2 is amphoteric, forming zincate anions [Zn(OH)42-] when there's plenty OH- present in solution. Then it would become a 'struggle' between the Kf of Zn(OH)42- and the Ksp of ZnCO3...



[Edited on 17-6-2013 by blogfast25]

AJKOER - 5-2-2014 at 19:55

Quote: Originally posted by OctanitroC  
Alright, so I made some Zinc Carbonate as follows:
...................
I then added concentrated Na2CO3 solution in quantity until no more gelatinous precipitate was formed, and filtered this precipitate out.......



You cannot use Na2CO3 to form the neutral ZnCO3 as it forms basic zinc carbonates of varying compositions. Source:"Elementary Chemistry: With Special Reference to the Chemistry of ...", Volume 1, by Harry Mann Gordin, page 322. Link: http://books.google.com/books?pg=PA322&lpg=PA322&dq=...

To quote:

"Zinc Carbonates. Neutral zinc carbonate ZnCO3 is found in nature (calamine), and can be made by adding a bicarbonate to a solution of a zinc salt ZnSO4 + 2NaHCO3 = ZnCO3+ Na2SO4 + CO2 + H2O. If instead of a bicarbonate a neutral alkali carbonate is used, a basic zinc carbonate results (cf. magnesium carbonate). The composition of the basic carbonate varies with the conditions of the reaction."

The good news is that to form the neutral Zinc carbonate just use Baking Soda (NaHCO3).

Interestingly, the action of a bicarbonate on a Magnesium salt is not precisely the same, as an unstable Mg(HCO3)2 is created. It, however, decomposes to the insoluble basic magnesium carbonate on boiling or on standing to MgCO3 (reference, for example, same source as above, page 317).

[Edited on 6-2-2014 by AJKOER]