Sciencemadness Discussion Board - Metathesis reaction to make copper (II) nitrate?

Metathesis reaction to make copper (II) nitrate?

Romain - 23-12-2013 at 07:20

Hi everybody,

I'm new to the forum and English isn't my mother tongue so if I make any mistakes, please correct me.

I would like to know if the following reaction is viable:
2NH₄NO₃ + CuCO₃ ⇋ (NH₄)₂CO₃ + Cu(NO₃)₂

You just need to boil off the resulting solution to decompose the ammonium carbonate and you are left with copper (II) nitrate (hopefully).

I believe boilling the solution also drives the equilibrium to the right.

I don't want to try and see by myself because as I live in switzerland, it's extremely complicated to obtain chemicals.

Just to give you an idea, I have access to 13% ammonium nitrate fertilizer which I need to recrystalize multiple times to get a reasonably pure product.
I don't have a dessicator so it takes several weeks total for 100 grams of ammonium nitrate!

The copper carbonate is difficult to produce too:
I start by electrolyzing a sodium bisulfate solution with copper electrodes to get copper a mixture of copper and sodium sulfate. Then I add some sodium carbonate to precipitate copper carbonate which I still need to dry...

And I won't talk about NaOH, HCL, H₂SO₄, ...

So please help me with that reaction,
thanks a lot,
R.

bismuthate - 23-12-2013 at 08:05

Nope it won't work. Although using chemicals you have you could:
Option 1 heat sodium bisulfate with NaCl and lead the HCl gas into water. add this to amonium nitrate and distill the HNO3 or put copper in it and lead the gas into water.

Option 2 mix NaHSO4 with NH4NO3 and distill the HNO3.

Option 1 variant B take the HCl and add H2O2 and copper to it. Then mix the CuCl2 with NH4NO3 anb decompose the NH4Cl.

Option 3 react CuSO4 with NH4NO3 and boil it down. Then separate the Cu(NO3)2 from the mix by dissolving it in ethanol.
I know that some reactions use HCl sorry.

[Edited on 23-12-2013 by bismuthate]

[Edited on 23-12-2013 by bismuthate]

Romain - 23-12-2013 at 08:50

Thanks for the fast reply!

Just wondering: why doesn't it work?

Also I unfortunately don't have a distillation apparatus

. And I can't even make the HCl because I don't have adequate glassware (yet). Your option 3 seems rather easy, so I will try it as soon as I can.
Though since I will need to make copper sulfate, it will take some time.

And I didn't mention H2O2: It costs 50$ a liter here. (for 3%)!

bismuthate - 23-12-2013 at 08:59

The reason is that ammonia is more reactive than copper so it will keep the strong nitrate ion and the copper will keep the less reactive carbonate ion.
http://en.wikipedia.org/wiki/Reactivity_series
Read this. (ammonia would be much higher up than copper) As you go up the series the metals (ammonia behaves like a metal) want stronger anions and the nitrate anion is much stronger than carbonate.

barley81 - 23-12-2013 at 09:04

Bismuthate, that's false. You cannot say that carbonate is "less reactive" in the context of metathesis reactions. Such reactions are driven by solubility and other equilibria, not by metals "wanting stronger anions." Ammonia does not behave much like a metal, either. The only setting where that is true is Isaac Asimov's short story The Magnificent Possession [which is pretty interesting]. In some situations, the ammonium ion does behave somewhat like alkali metal ions. Its size is close to that of a caesium ion. Wikipedia: "In terms of size, the ammonium cation (r<sub>ionic</sub> = 175 pm) resembles the caesium cation (r<sub>ionic</sub> = 183 pm)."

I am not sure if the reaction between ammonium nitrate and copper carbonate actually works. Try and see, though the ammonium cation may not be acidic enough to dissolve the CuCO3 at an acceptable rate. If it does work, it will probably involve boiling the mixture for a long time.

[Edited on 23-12-2013 by barley81]

Romain - 23-12-2013 at 09:15

Thank's for the reactivity series page. I'll read about it, it looks interesting.
And if what Bismuthate said is incorrect, is it still impossible to make copper nitrate via metathesis?

bismuthate - 23-12-2013 at 09:20

Ammonium behaves like a metal my appologies. Geeze I suck at explaining. Also I'm sorry about the mess up so everyone ignore my post as it is wrong. I'm going to leave the explaining to the people who are much more qualified.

By the way check out this article ammonia could behave like a metal

(although that wasn't what I meant)
http://adsabs.harvard.edu/full/1954MNRAS.114..172B

[Edited on 23-12-2013 by bismuthate]

blogfast25 - 23-12-2013 at 10:14

Quote: Originally posted by barley81

I am not sure if the reaction between ammonium nitrate and copper carbonate actually works. Try and see, though the ammonium cation may not be acidic enough to dissolve the CuCO3 at an acceptable rate. If it does work, it will probably involve boiling the mixture for a long time.

[Edited on 23-12-2013 by barley81]

If it had anything to do with acidity of the ammonium ion then ammonium carbonate could never exist, instead it can exist both as a solid and in solution.

The reason why this can't work is that ammonium carbonate is fairly soluble, much more than copper carbonate. In short, nothing will happen.

Romain - 23-12-2013 at 11:04

Ok thanks, I'll try the copper sulfate + ammonium nitrate method as soon as I can.

So in fact a metathesis reaction works if products are less soluble than reactants?

TheChemiKid - 23-12-2013 at 11:10

How do you plan to separate the copper nitrate and ammonium sulfate?

bismuthate - 23-12-2013 at 11:12

Copper nitrate is soluble in ethanol and ammonium sulphate is not.

TheChemiKid - 23-12-2013 at 11:40

Ok, that is a good method. I was just wondering if he knew how to separate them.

WGTR - 23-12-2013 at 12:56

Quote: Originally posted by bismuthate

Copper nitrate is soluble in ethanol and ammonium sulphate is not.

Copper sulfate is not particularly soluble in ethanol either, but I couldn't tell you whether CuSO4 or (NH4)2SO4 is more or less
soluble than the other.

You could try boiling ammonium nitrate solution with quicklime to form calcium nitrate. The ammonia would boil away (watch
the fumes). Filter off the leftover slaked lime (it's caustic, but not very soluble).

2NH4NO3 + Ca(OH)2 --> Ca(NO3)2 + 2NH3 + 2H2O

Then mix the Ca(NO3)2 and CuSO4 together in water, and filter off the insoluble CaSO4:

Ca(NO3)2 + CuSO4 --> CaSO4 + Cu(NO3)2

Romain - 23-12-2013 at 13:24

The quicklime idea looks pretty good too, thanks! The only problem is I don't have quicklime, I don't know where to find some and if it is regulated somehow.

But so far it's the best method as it also produces calcium sulfate, a useful dessicant.

WGTR - 23-12-2013 at 13:57

I suppose you could buy it from a garden store. I couldn't imagine it being restricted. But why buy it when you could make it?

Plante1999 published a short writeup on how to do this at this link.

I showed briefly how to make CaO in this thread where I built a small homemade electric kiln.

I've even done it in a small charcoal furnace on the front porch. It's best to get a pure CaCO3, but you can get it from seashells,
limestone rocks, some forms of chalk, etc. Of course, you can make CaCO3 from CaCl2 (desiccant) and Na2CO3 (washing
soda), both of which are available here in every grocery store.

blogfast25 - 24-12-2013 at 05:11

Quote: Originally posted by Romain

So in fact a metathesis reaction works if products are less soluble than reactants?

In essence, yes, one of them:

AX(aq) + BY(s) → AY(aq) + BX(s)

Quicklime is easy to get and not really worth making yourself, IMHO. Get it from eBay. Slaked lime (Ca(OH)2, even easier to get) would work here too.

Bear in mind that CaSO4 is poorly soluble but not completely insoluble. Your Cu(NO3)2 will be slightly contaminated. Better would be to use Ba(NO3)2 but that's much harder to get and much more expensive.

Cu(NO3)2 is also hard to crystallise because it so damn soluble (83.5 g Cu(NO3)2 per 100 g of water @ 0 C, Wiki). The solid is deliquiescent, at least in my experience.

[Edited on 24-12-2013 by blogfast25]

Romain - 24-12-2013 at 05:37

I may order Ca(OH)2 on ebay but I'll first try to find some at my garden store.
Here in switerland, a lot of chemical are regulated or unavailable so I don't think it's the solution though.

I may also try to make some Ca(OH)2 but I don't want to get too fancy and build a frunace.

And for the solubility porblem of Cu(NO3)2, if I order some Ca(OH)2 on ebay, I will also order some silica gel to dessicate the copper nitrate.

Or maybe a few hours in my oven at 200°C is enough to dry it?

Anyway, thanks a lot to everybody for your answers, I didn't expect to get so many!

blogfast25 - 24-12-2013 at 05:52

Quote: Originally posted by Romain

And for the solubility porblem of Cu(NO3)2, if I order some Ca(OH)2 on ebay, I will also order some silica gel to dessicate the copper nitrate.

Or maybe a few hours in my oven at 200°C is enough to dry it?

Neither of these things will work. Silica gel isn't a very strong desiccant (more for keeping things dry than for drying things, there's a difference).

Heating a hydrate like copper nitrate (there are acc. Wiki three hydrates: tri, hexa and hemipenta) will cause it to melt and dissolve in its own crystal water. Trust me, I've been there...

Your best bet is probably to start from a quite concentrated solution of Cu(NO3)2 and gently boil this in until there is almost no liquid left. The solid crystals are probably Cu(NO3)2.3H2O (hard to be sure without a two phase temperature diagram). Remove what liquid is left while still hot. After cooling try storing in a CaCl2 desiccator.

I am of course assuming it is the hydrate you want, not the anhydrous salt.

[Edited on 24-12-2013 by blogfast25]

Romain - 24-12-2013 at 07:01

The goal of all this is to make copper nitrate flash powder so as to get a green flame and also to have a sample of copper nitrate to add to my chemicals collection.

At first I thought of making some copper carbonate and add it to existing sodium nitrate flash powder but copper carbonate would slow down the reaction.

I believe hydrated copper nitrate works for flash powders so I won't need to dehydrate it completly, but it's just a guess.
I also heard that anhydrous copper nitrate is way too reactive to be used in flash powders.

blogfast25 - 24-12-2013 at 10:08

Depends what you call a flash powder. Best results would be obtained with an anhydrous nitrate. Cu(NO3)2 dehydrates well. Just keep heating until steam evolution stops.

Re. reactivity, it depends largely on composition. For example, if you use Mg powder the amount (among other things) will determine how violent it is.

Cu(NO3)2 + 2 Mg === > CuO + N + 2 MgO

But with extra Mg, also:

CuO + Mg === > Cu + MgO, a lot more violent if both happen.

[Edited on 24-12-2013 by blogfast25]

Marvin - 24-12-2013 at 10:35

Quote: Originally posted by blogfast25

Best results would be obtained with an anhydrous nitrate. Cu(NO3)2 dehydrates well. Just keep heating until steam evolution stops.

My understanding is that the anhydrous salt cannot be made by drying the hydrate, it decomposes.

There are reports of solutions containing ammonium nitrate and copper salts spontaneously detonating during crystallisation. (Mellor). Might be best to make it a different way.

Romain - 24-12-2013 at 14:24

When I say it's not suitable for flash powder, I mean it would make a very sensitive mixture that may deflagrate (or even detonate) upon impact or friction.

I sure do not want that, I want a stable mixture that won't ignite spontaneously. I think I read somewhere that it's a very potent oxidized and I thought it may make very reactive mixtures.

I just read on this thread that dehydrating copper nitrate is extremely difficult.$

And about solutions of ammonium nitrate and copper salts, perhaps it was copper chlorate in the solution? I know ammonium chlorate is extremely unstable and explosive even in solution according to Wiki..

And now that I think of it, making copper chlorate could be another way of making green flash powder though if copper nitrate is reactive, copper chlorate may be even more. What do you think?

blogfast25 - 25-12-2013 at 05:38

Acc. to my CRC (just got it, best Crimbo present ever!) copper chlorate is highly soluble in water and forms a hexahydrate. Possibly even harder to dehydrate than the nitrate...

Have you considered anhydrous CuSO4 as oxidiser?

Calcium sulphate (anh.) works very well for oxidising Al powder for instance: CaSO4 + 8/3 Al === > CaS + 4/3 Al2O3. Mixtures like that burn fiercely hot when lit.

[Edited on 25-12-2013 by blogfast25]

PHILOU Zrealone - 25-12-2013 at 09:01

1°) The reaction you thought about of mixing CuCO3 and NH4NO3 should work but not the way you all think about... because:
a) NH4(+) is mildly acidic and so H2CO3 (H2O and CO2(g)) will be produced in some extend and thanks to heating, the gaseous CO2, will be expelled out of the system.
b) NH3 is leading to tetraamino complexes with Cu(2+) what is also a driving force for the reaction

So
2CuCO3 + 4NH4NO3 --> Cu(NO3)2 + (NH3)4Cu(NO3)2 + 2CO2(g) + 2H2O

The same holds true for Cu(OH)2 and CuO...
2Cu(OH)2 + 4NH4NO3 --> Cu(NO3)2 + (NH3)4Cu(NO3)2 +4H2O
2CuO + 4NH4NO3 --> Cu(NO3)2 + (NH3)4Cu(NO3)2 +2H2O

The reactivity of the oxydation layer of Cu (CuCO3, Cu(OH)2 and CuO) towards NH4NO3 is the main reason for the uncompatibility of Cu and NH4NO3. Tetraamino copper nitrate being a sensitizer of decomposition of NH4NO3.

2°) A flash powder with NaNO3 would be extremely detrimental to the integrity of the blue color from copper because Na produces a strongly remanent yellow colour that will hide any blue or green shade. Better use KNO3!

Blue color can be obtained only at low temperature and flash powder would produce a very hot and bright flame inducing a shift to the white side of the emission spectra (silvery light) this will hide the blue expression. Better find a colder flame!

Usually the blue color of copper is enhanced by the presence of chloride in the powder mix (PVC, CuCl2, Cu(OH)Cl, Cu(ClO4)2, K chlorate - perchlorate or chloride, ...). Use sodium free ingredients in your powder (suggar and natural charcoal contains a lot of Na...)!

[Edited on 25-12-2013 by PHILOU Zrealone]

blogfast25 - 25-12-2013 at 09:27

PH Z:

I doubts if the acidity of ammonium is really enough. As stated above, it that was the case, (NH4)2CO3 solutions would 'self destruct'. Yet only on heating do they lose CO2.

[Edited on 25-12-2013 by blogfast25]

PHILOU Zrealone - 25-12-2013 at 10:11

Never underestimate the slow processes or low concentration processes...

-Acidity of NH4NO3 solution is enough in solution to chew Al foil in a few days while Al stands HNO3.
-HNO3 69% can effectively nitrate toluen in monthes at 20°C ... you get nitrobenzoic acids mixed with o-/p- nitrotoluen.

Sometimes the only way is to try and find out!

AJKOER - 25-12-2013 at 12:36

OK, I recall the report of an experiment where Al foil was used to cover a flask containing NO2. Crystals of Aluminum nitrate were formed. This is not an isolated instance as per Wikipedia (http://en.wikipedia.org/wiki/Nitrogen_dioxide ) to quote:

"NO2 is used to generate anhydrous metal nitrates from the oxides:[6]

MO + 3 NO2 → 2 M(NO3)2 + NO"

So, first generate NO2 and react with CuO (or Cu, see below) to create anhydrous Copper nitrate.

Now, as to why an aqueous path is not wise, per Wikipedia on Copper Nitrate to quote (see http://en.wikipedia.org/wiki/Copper_nitrate ):

"Anhydrous Cu(NO3)2 forms when copper metal is treated with N2O4:

Cu + 2 N2O4 → Cu(NO3)2 + 2 NO

Attempted dehydration of any of the hydrated copper(II) nitrates by heating instead affords the oxides, not Cu(NO3)2. At 80 °C, the hydrates convert to "basic copper nitrate" (Cu2(NO3)(OH)3), which converts to CuO at 180 °C.[2]"

So, one could also use either Cu or CuO to prepare anhydrous Cu(NO3)2 with N2O4.

I doubt if I one could call the Copper metal based preparation, however, a metathesis reaction. The reason being that such a reaction is technically define as occurring when cations and anions exchange partners (see discussion at http://chemistry.osu.edu/~woodward/ch121/ch4_metathesis.htm ).

[Edited on 25-12-2013 by AJKOER]

blogfast25 - 25-12-2013 at 13:48

Quote: Originally posted by PHILOU Zrealone

Sometimes the only way is to try and find out!

Amen to that.

TheChemiKid - 26-12-2013 at 03:53

The metal oxide method works, I just tried it yesterday.

blogfast25 - 26-12-2013 at 06:00

Quote: Originally posted by TheChemiKid

The metal oxide method works, I just tried it yesterday.

Aren't you going to share what you did with us? From 'it works' very little can be learned.

Romain - 30-12-2013 at 14:31

@PHILOU Zrealone for you hints about flash powders, I'll try to avoid sugar for color tests and use Al or Mg instead!

I'll also try to make some pure KNO3 (without Na) but that will be costly since KCl is 5CHF (~6$) for 100g.. I may order some chemicals on ebay sometime. I'll try PVC too as it's readily available.

@blogfast25 I saw a video on youtube of some CuSO4/Mg flash powder. Impressive!
I'll porbably try that soon, I have some CuSO4 crystallizing at the moment, it should be ready in a few weeks.

The NO2 route to make nitrate salts is interesting though you need some way of generating NO2.
Perhaps via the Birkeland-Eyde process?
You could put some metal oxide at the bottom of a 3 neck round bottom flask with a stir bar. 2 of the necks would be used for the high-voltages electrodes and the third would be stoppered for quick access and to add some oxygen and nitrogen when needed. The stir-bar would be used to allow the gas to come in contact with the finely powdered oxide.

If that works that may be useful to make exotic anhydrous nitrates

blogfast25 - 31-12-2013 at 06:02

Quote: Originally posted by Romain

@blogfast25 I saw a video on youtube of some CuSO4/Mg flash powder. Impressive!

Wow. Big bangs. But no significant flame colouration though... Adding some anhydrous CuCl2 (easy to prepare) could increase colouration.

[Edited on 31-12-2013 by blogfast25]