# Series of reactions



## TomVader (May 7, 2013)

Hello everybody. I have a couple of questions. I had two gallons of silver nitrate solution, I dropped the silver with copper and recovered the silver. Now I'm dropping the copper with iron. I'm using old nails which are made of steel. What will happen to the carbon in the steel? Will it combine with oxygen and escape as a gas or will it remain as a powder with the copper? After that, I want to drop the iron with aluminum. Once I have the aluminum in solution, can I drop it with table salt? Will this result in sodium nitrate solution, aluminum powder and free chlorine gas, or sodium nitrate solution and aluminum chloride? if it results in aluminum chloride, can I simply heat it to drive off the chlorine? And the sodium nitrate solution, can I evaporate that down to a powder and re-use it to make nitric acid ? I don't have a practical reason for doing this other than the experience. Thank you in advance for your knowledge and patience. ( I guess I had more than a couple of questions)


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## butcher (May 7, 2013)

The steel nails will form an iron nitrate solution,and replace copper from solution, the steel would not have much carbon, formula for one type of carbon steel, 0.026% S, 0.06% P, 0.4% Mn, 0.1% C, balance Iron, I doubt the nitrate would oxidize carbon to CO2 gas, so my guess is it would be in the copper powders.

I do not see any benefit in dropping Iron with aluminum, although it will.

No you would not drop out aluminum with salt, you would just have sodium, and chlorides mixed in with your aluminum nitrate solution.

if you want to make sodium nitrate here is a suggestion:
Use your copper nitrate solution, slowly and carefully add sodium hydroxide (with stirring, reaction is exothermic so do not let it heat up too much), watch for copper hydroxide precipitant it will be gelitinous and solution should go clear, you do not want to add too much hydroxide so take your time, check for color of solution (clear, no blue in solution) after giving time for copper hydroxide to settle, the pH should be slightly basic when you test pH.

Cu(NO3)2 (aq) + 2NaOH (s)--> Cu(OH)2 (s) + 2NaNO3 (aq)

After the copper hydroxide settles you can decant the clear sodium nitrate solution.

You can then add a little water to the copper hydroxide and heat it gently with stirring until the color changes to form a copper oxide, let this settle and dry it can be used to make other copper compounds later like copper sulfate, copper nitrate, copper chloride...

Cu(OH)2 (s) --heat--> CuO (s) + H2O (l)

or you could make it back into copper metal, using several different processes if you wished.

CuO (s) + H2SO4 (aq) --> CuSO4 (aq) + H2O (l)

CuSO4 (aq) + Zn (s) --> Cu (s) + ZnSO4 (aq)


There are other ways you could reuse your copper nitrate, it is more useful than iron or aluminum nitrate solutions in my opinion.


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## TomVader (May 7, 2013)

I would like to be able to re-use these solutions when possible. Going up the reactivity chart is just for fun, but a practical application is better. OK, NaOH to drop copper as CuOH; NaNO3 in solution remains. I assume that NaOH could drop FE or Al as hydroxides and still leave NaNO3... Thank you once again for your knowledge and effort


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## solar_plasma (May 7, 2013)

> NaOH could drop FE or Al as hydroxides



http://de.wikipedia.org/wiki/Hydroxide:

_ 
AgOH 8,3–11,3

Ca(OH)2 12,4–13,9
Mg(OH)2 9,6–11,1
Fe(OH)2 8,3–9,8
Ni(OH)2 8,1–9,6
Cd(OH)2 8,1–9,6
Mn(OH)2 7,9–9,4
Pb(OH)2 7,2–8,7
Co(OH)2 7,2–8,7
Zn(OH)2 6,6–8,1
Be(OH)2 5,7–7,2
Cu(OH)2 5,1–6,6
Sn(OH)2 2,4–3,9

Cr(OH)3 4,6–5,6
Al(OH)3 3,8–4,8
Fe(OH)3 2,2–3,2
Sb(OH)3 0,9–1,9_

...depends on the pH


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## butcher (May 7, 2013)

Laser Steve posted some good ways to reuse your copper nitrate solution, using graphite anodes to deplete copper from solution, where he used the solution to recover silver from sterling.

note: you will not get back as much nitric, or nitrates from your solution as you put in as acid to dissolve the original metal, because a portion of the acid goes up in gases of NO or NO2 in the reaction, but we can recover a portion of our nitric acid.


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## TomVader (May 8, 2013)

Yeah, that's a shame. How cool would it be to re-use everything and set up a cycle where all you add is energy (heat or electricity), and recover and refine ad infinitum?


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## mitchd (May 8, 2013)

butcher said:


> No you would not drop out aluminum with salt, you would just have sodium, and chlorides mixed in with your aluminum nitrate solution.
> 
> if you want to make sodium nitrate here is a suggestion:
> Use your copper nitrate solution, slowly and carefully add sodium hydroxide (with stirring, reaction is exothermic so do not let it heat up too much), watch for copper hydroxide precipitant it will be gelitinous and solution should go clear, you do not want to add too much hydroxide so take your time, check for color of solution (clear, no blue in solution) after giving time for copper hydroxide to settle, the pH should be slightly basic when you test pH.
> ...




Will this still work if NaCl is in the mix with the Cu(NO3)2? I add salt to recover the silver then add scrap iron the recover the copper but would like to recover some of the NaNO3 for reuse but with the salt mixed in, though it might make poor mans AR.

Mitch.


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## solar_plasma (May 8, 2013)

There should not be much NaCl in it then, but NaNO3....since you have precipitated the Cl-

In fact are no salts at all, since they form ions in water, so there are Na+ ions and NO3- ions.


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## butcher (May 8, 2013)

As solar Plasma stated, you will have no reaction as nothing will precipitate, you just have all of the ions floating around in solution.

Copper nitrate with hydrochloric acid:
Cu(NO3)2 + 2HCl --> CuCl2 + HNO3 (no reaction nothing precipitates) (here copper would just be dissolved in a form of aqua regia).

Copper nitrate with sodium chloride salt:
Cu(NO3)2 + 2 NaCl --> CuCl2 + 2NaNO3 (here again no reaction nothing precipitates) so we have copper and sodium cations and nitrate and chloride anions all floating around in solution.

If you have silver nitrate and you add salt
AgNO3 + NaCl --> AgCl (s) + NaNO3 
(here we have sodium nitrate) (and have a reaction because silver precipitates).

If we used HCl to precipitate silver from silver nitrate solution to form silver chloride powders,we make nitric acid as a byproduct.
AgNO3 + HCl --> AgCl (s) + HNO3
(here if we had some copper from previously dissolving sterling silver, we would be left with some copper in this solution but we could still use this nitrate to dissolve base metals or even another batch of sterling, if we did not use excess HCl we would not be left with chlorides in solution (which could make a form of aqua regia), one thing we could do is leave a little silver in solution or remove chlorides from our remaining nitric by using a solution of silver nitrate solution (the silver grabs the chlorides from solution and precipitates the chlorides out of solution as silver chloride.


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## solar_plasma (May 9, 2013)

For completeness (correct me, if I'm wrong): Some ions, though none of those above, will interact with water molelyles in an equilibrium reaction. Therefore some salts react acidic or alkalic in water solution.


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## danogarvin123 (May 9, 2013)

Gentlemen.......may I commend you on an awesome thread....much learned. May I say it connected quite a few dots for me..or should I say "ions". (warped humor)


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## TomVader (May 10, 2013)

I like the idea of dropping copper with sodium hydroxide. You remove the copper hydroxide, and have sodium nitrate to reuse later in your next batch of nitric. Is lye pure enough to use or should I look for sodium hydroxide from a chem supply co.? Thanks. ( I hope someday to answer questions instead of always asking them!)


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## Don in Mindanao (May 10, 2013)

when i was in high school chemistry class we did an experiment to remove silver metal from silver nitrate solution we just put a coil of clean copper wire in the silver nitrate solution over nite and the next day silver crystals were all over the copper wire. then we wiped the crystals of metallic silver off of the wire and melted the silver and had small silver buttons . and later when i went to work for a photo lab we extracted silver from the fixer using zinc oxide powder and then strained and dried the grey precipitate and melted it into silver buttons about an ounce each. and i still use the zinc oxide to precipitate silver out of my aqua regia. i think i use it for every thing I just use different strength mixtures for different metals. Ag Au Pt


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## butcher (May 10, 2013)

Don in Mindanao, 
we use these methods also for recovery of metals, we recover silver from silver nitrate with copper, (the silver will then need to be refined, or melted for a less pure silver, we can also recover PGM from solution with zinc, (we will normally use the metal as it is plenty reactive without being an oxide of the zinc).

I would not use zinc for recovery of gold from aqua regia, for recovery of gold from a dirty aqua regia I would prefer using copper, this way my gold would have less base metals to deal with, if the gold in the aqua regia is fairly pure I would precipitate with other reducing agents such as copperas, sodium metabisulfite or other chemical reducing agents, this way I can convert the gold to metal with out precipitating other metals with my gold.

study the forum you will learn much better methods than you are using now, also to get a better understanding of all of the other metals you are also precipitating with your gold when you use zinc oxide, Google and study the reactivity series of metals, you are also reducing every metal below zinc in the reactivity series, with your gold, so if you also dissolved these other metals besides pure gold, then you just about precipitated back out almost all of the metals with your gold, which really in my opinion was just a waste of time and acids, this could include these other metals like, chromium, iron, cadmium, cobalt, nickel, tin, lead, antimony, arsenic, bismuth, copper, mercury, palladium...
so basically your are not really precipitating gold from aqua regia using zinc oxide, you are reducing many metals from solution with your gold, this really makes no common sense, because the primary reason we use aqua regia is to refine gold, and to dissolve almost pure gold and leave the other small amount of metals in solution when we precipitate our gold, very seldom will aqua regia be used as a recovery method, and in the very rare instance it is, then copper metal would be the best choice to recover your gold, by a replacement reaction, we will normally do this from a dirty solution or a mess, copper will not precipitate all of the other metals listed above, copper would only reduce the more valuable metals from solution.

Hope this helps.




TomVader,
lye is just another name for caustic soda, sodium or potassium hydroxide, it should work.
http://en.wikipedia.org/wiki/Lye


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## TomVader (May 20, 2013)

Tried what butcher suggested; added sodium hydroxide to copper nitrate, copper hydroxide precipitated out as a light blue gel, then with time, settled into a more dense powdery layer. This happened with each incremental addition of sodium hydroxide, until the solution turned clear. The first time I added the sodium hydroxide as a solid, very slow reaction. Further additions were sodium hydroxide dissolved in water, reaction was instantaneous.
My second container of copper nitrate I added nails, solution turned green over a few days. I decanted the liquid today and found the nails were mostly intact, with a small amount of copper powder in the bottom of the container. I believe some of the iron went into solution and some of the copper dropped out, but the reaction is not complete. Either the acid is too weak to dissolve any more iron, or the nails have been passivated. Tomorrow I'll add more clean iron and see if the reaction continues.


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## butcher (May 20, 2013)

The copper hydroxide will be fluffy and most hydroxides are voluminous and do not settle well as powders, this precipitate could look like there is more in solution than if it was a powder or metal.

The elemental copper metal powders produced fro the reaction with iron will be more dense as metal, and may not look like the same amount of copper as the reaction above.

This might work to see if you can test for copper in solution try few drops of suspect solution in a spot plate, or white plastic spoon, with a few drops of ammonia seeing if you get the deep blue from the reaction with copper amines.

Iron can passivate in strong nitric, but can dissolve easier in dilute nitric acid, the iron may be getting an oxidized layer even though this is a replacement reaction, you might try it more dilute.

See what happens when you concentrate that clear solution from the sodium nitrate liquid, adding some sulfuric acid and test your nitric on copper again.


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## TomVader (May 22, 2013)

The copper powder in the copper nitrate/iron nitrate solution is VERY fine. A lot of it went right through the filter. The act of pouring it stirred it up and I think it will be days before it settles to the bottom. As I can't see into the container, I can't tell if the clean nail I added is reacting or not.
I decanted the liquid (sodium nitrate/sodium hydroxide in water) from my copper hydroxide gel. A little of the pale blue gel got stirred up and came with the solution, I think a little more sodium hydroxide should get it to gel up a little firmer. I don't have ammonia to test for copper in solution, but will get some. I plan on heating the copper hydroxide and I expect to see the hydrogen driven off as the hydroxide ion breaks up and then I expect the copper oxide to be a much more dense black powder. I think I can then decant the remaining liquid and rinse and recover the copper oxide. So that's what I THINK will happen. I'll post again and let you know what ACTUALLY happened. (Thanks Butcher)


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## solar_plasma (May 22, 2013)

> I plan on heating the copper hydroxide and I expect to see the hydrogen driven off as the hydroxide ion breaks up and then I expect the copper oxide to be a much more dense black powder.



Do you mean, you will heat the copper hydroxide while it still is in water? That will not work. It will first decompose at several hundred degrees celsius.


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## TomVader (May 23, 2013)

Wiki says copper hydroxide reduces to copper oxide at 80C. I haven't done it yet, but will post results when I do.


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## TomVader (May 25, 2013)

I heated my copper hydroxide solution on medium heat on a hot plate. The light blue gel changed into a fine black powder, but I didn't see any bubbles. Where did the hydrogen go? Could it have escaped as a gas without visible bubbles? Next I'll add water, wait for it to settle, decant and repeat, then dry. I should have relatively pure copper oxide. Thanks for listening.


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## TomVader (May 30, 2013)

Hand hits forehead; I just realized where the hydrogen went. Two hydroxide ions broke up, making one water molecule and leaving one oxygen atom to bond to the copper. No bubbles.


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## solar_plasma (May 30, 2013)

> I heated my copper hydroxide solution on medium heat on a hot plate. The light blue gel changed into a fine black powder,



That is very interesting. I didn't know that. Thanks for sharing!


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## TomVader (May 30, 2013)

The process was spelled out by butcher earlier in this thread, but I missed it. I'm still learning about chemical notation. (l) liquid. (aq) aqueous. etc.


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## butcher (May 30, 2013)

As you learn more about these chemical equations, and how the chemistry works and its laws, this opens up many possibility's of reactions you can create, or when reactions are created by the reactions of recovery and refining, it gives you Ideas of what you can use byproducts for, the understanding of the chemical equations and how they work can even help you in determining what is in solution, or ways to change one metal salt into a different metal salt, like for example using heat to drive off one acid from a metal salt dissolved in solution to form a different metals salt with an acid of different volatility or boiling point, or cementing a metal out of solution with another metal higher in the reactivity series, converting a metal salt to elemental metal with an acid and a metal higher in the reactivity series (like the chemistry we use in converting silver chloride to metal powders using iron and dilute sulfuric acid), you do not really need to understand the chemistry to be able to recover and refine metals, but with learning more about the chemistry side of this work it opens a whole new world of uses, understanding and possibility's, helping us to better understanding how things work, and how we can or cannot make something work and even in troubleshooting our problems.


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## TomVader (May 31, 2013)

Slowly but always forward. Reading time is limited for me, (work, kids, parents, maintaining the house, etc.) The chemistry is fascinating by itself, the PMs are an added bonus. This is the most interesting and challenging hobby I've ever tried. I started down this path with the idea of making money, that's still a possibility, but is now secondary. With the help of this forum, wikipedia and C.M. Hoke, I look forward to many years of learning this art/science.


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