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Silver compounds must be kept away from:

Acetylene or acetylides  (e.g., calcium carbide)

• Ammonia or ammonium salts (prolonged contact)

Oxalic acid or oxalates

CR Scientific

Silver and Silver Nitrate

WARNING!  This procedure involves nitric acid and hydrochloric acids, which are extremely corrosive to skin and other organic materials.  Silver nitrate itself is also caustic to skin and eyes.  The experiment also involves generation of toxic nitrogen oxides, which can easily kill you.  If you choose to attempt the following procedure, you do so entirely at your own risk.

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Over the course of studies in mineralogy and jewelry-making, it's easy to accumulate a pile of silver scraps.  While not nearly as expensive as gold, these bits of silver are still far too valuable to throw away.  At some point it makes sense to consolidate the scrap silver and purify it.  The silver metal can then be used in subsequent alloying, plating, or re-casting.
The purification of silver as detailed herein tends to consume large amounts of nitric and hydrochloric acids, depending on how much silver is involved;  the operation's economy, of course, depends also on the current market price of silver.  The following procedure is based largely on traditional qualitative analysis methods of differential solubility.
Silver compounds should not be left in contact with ammonia or ammonium compounds for prolonged periods.  Some of the more obscure recipes for silver-plating baths may call for ammonium salts;  these are alright for immediate use but should not be stored. 
Traditional silver plating formulations involve cyanide (a deadly poison).    However, the metals as plated out in our experiment are not intended to form durable, attractive coatings;  we wish only to recover them in a more pure state than they were.  Therefore we can forego the use of cyanides.  Plating for aesthetics and durability should be a separate operation anyway.


This article is not a recommendation that the reader attempt the experiment.  It must not be attempted by beginners or anyone else untrained in working with highly corrosive materials and toxic gases.
Wear safety goggles at all times.  Acids can make you blind almost  instantly.  Silver nitrate is equally dangerous in this respect, possibly more so.
Do not heat acids without adequate ventilation.  The vapors should be collected by a fume hood which carries them to the outdoors.  Test your fume cupboard with a stream of smoke to make sure vapors are being removed, not lingering in your laboratory.  You can also use the NH4Cl "smoke" that is produced from the vapors of hydrochloric acid as they combine with gaseous ammonia.  Watch where the stream goes.
Nitrogen dioxide is very poisonous and corrosive to the lungs.  Do NOT inhale it under any circumstances.  It can kill you quite easily, with death occurring sometimes a day or two after the inhalation.  Also avoid inhaling acid vapors.
The silver, copper, and other metals that may be present are recycled, so the generation of hazardous waste is minimal.  Do NOT pour silver, copper, or nickel solutions down the sink, out on the ground, etc.  Recycle and reclaim all compounds via crystallization or precipitation.
If lead is present in the scrap silver (not likely, but conceivable depending on where you got the scrap silver), there will of course be lead compounds present in the solutions.  The amounts should be no more than a few ppm, but they should nevertheless be treated as if Pb is present.   The aim in such circumstances is either to plate out the Pb2+ as lead metal or to precipitate it as lead sulfide (PbS).  PbS has the lowest aqueous solubility product constant (Ksp) of all common lead compounds;  the Ksp is on the order of 10-28 or 10-29, depending on what source is quoted.
Keep silver and its compounds away from oxalic acid, oxalates, acetylene,  calcium carbide, strong reducing agents, or prolonged exposure to ammonia. 
Silver nitrate is caustic.

Materials & Equipment: 

Safety goggles and face shield;
Neoprene gloves;
Heavy lab coat and rubberized apron;
Silver, scrap
Nitric Acid, concentrated
Hydrochloric Acid, concentrated
Sodium Thiosulfate
Distilled Water
Copper metal (wire or strip)
Graphite electrodes & Lab power supply

Methods & Observations:

A piece of scrap sterling silver was set aside for the experiment.  The piece was stamped ".925", indicating 92.5% purity in silver.  There had been small, faceted "marcasite" stones set in the piece.  Removal was facilitated by heating the bracelet in a torch flame and then quenching the bracelet in cold water.  Most of the stones fell out, but a few became stuck and melted into a hopeless mess of black residue that would not come loose from the silver.
The cooled scrap silver was placed in a porcelain crucible.  This was covered and heated with an acetylene torch until the crucible assumed a bright orange glow;  heat was applied for at least two to three minutes more. 

Heating the crucible Molten silver

On attempting to pour the silver into a miniature ingot, the molten metal rapidly solidified as it hit the cooler, upper edge of the crucible (side note: time to build a crucible furnace).  At one point the molten silver was accidentally spilled on the benchtop, leaving hot beads of freshly-solidified silver.  These were allowed to cool and were then collected.
Casting attempts were thus temporarily abandoned. 
Most of the cooled silver was placed in a clean 50 mL beaker labeled Solution A ;  this was then filled about 3/4 full with concentrated nitric acid.  This was covered with 1000-mL beaker and heated on a sand bath.  Nitrogen oxides were released as reddish-brown fumes;  these remained inside the large beaker until it was lifted, at which point the ventilation system carried the fumes to the outdoors (NO2 is highly toxic!).  The solution became yellow and then green upon standing. 
It took two treatments in this manner to dissolve all the silver;  the second treatment produced a solution with not quite as much green coloration.  A larger initial volume of nitric acid in a larger beaker would have eliminated the need for two treatments;  the theoretical amount could easily have been calculated based on the estimated number of moles silver versus the number of moles HNO3 required to dissolve it..

The silver dissolves, releasing NO2
Nitric acid is indispensable in the chemistry of minerals and metals.  However, extreme caution must be exercised when dissolving certain materials;  they cause the acid to release nitrogen oxides.
Nitrogen dioxide (shown at left) is an insidious and deadly poison.  Do not breathe it.  Like nitric acid, it can cause delayed pulmonary edema and death after a period when the victim seems to be alright. 
Lesser doses can cause severe lung irritation, asthma attacks, and other symptoms.

The remainder of the silver was placed in another beaker and covered with concentrated HNO3.  To this beaker was added an equal volume of concentrated HCl, forming a mixture that lacked the proportions of true aqua regia (3 HCl : 1 HNO3).  As the acid attacked the silver, AgCl precipitated out as white clumps.  The amount of precipitate steadily increased, peaking at about four hours.  It redissolved after prolonged standing, however.  The final color was green.  This solution was labeled Solution B.

In both the solutions A and B, the color was deep green.  It was not certain at this point whether the green tint was due to Ni or to Cu compounds.  Although aqueous Cu++ is usually blue, it can form green complexes.

Solution A was diluted to about three times its volume with distilled water.  Reddish-brown NO2 was released during dilution, but not as heavily as it had been when the metal was dissolved in the acid.  The liquid's green color promptly turned to a pale blue.  The solution was filtered using a Buchner funnel and vacuum filtration flask.  The filtrate was saved. 
The dilution was necessary for two reasons:  1.) to ensure that the nitric acid wasn't strong enough to attack the filter paper, and 2.) to make sure any solid silver nitrate was fully re-dissolved.
The filtrate was evaporated slowly on a hot plate, the process being done with ventilation to carry off the nitric acid vapors.  Heating was discontinued when the first silver nitrate crystals began to form.  The AgNO3 crystals grew larger upon cooling.  The blue solution, which appeared to retain most of the copper nitrate and presumably any other salts, was decanted.  The process was repeated to get progressively purer AgNO3.

Solution B was diluted to about three times its volume with distilled water as well, causing the silver chloride to re-precipitate.  The total amount of AgCl was estimated at less than 0.5 gram.  This was filtered on a Buchner funnel and the precipitate washed and saved.  It was placed in a clean beaker and covered with 25 mL of distilled water.  The silver chloride was then solubilized by adding three or four crystals of sodium thiosulfate to the water and stirring with a glass rod.  A preliminary attempt at plating out the silver was made using a cleaned copper wire.  Upon standing for about 10 minutes in the solution, the wire began to assume a very fine, silvery coating.  The wire was left in the solution for about five hours;  it developed a black, scaly coating of silver.  The solution, however, remained clear.
The silver metal was scraped off with wooden tongue depressor.  This metal, mostly black and crumbly, was saved for making AgNO3 solutions of higher purity.

Crystals of silver depositing on a copper wire
Silver crystals such as the ones shown at left can be grown from concentrated solutions. 
As aqueous Ag+ becomes more dilute, the deposited silver tends to be black rather than its normal silvery-gray. 
Copper metal goes into solution as cupric ion, while the silver ion comes out of solution as silver metal.  Copper is oxidized and silver ion is reduced.  This particular reaction will occur spontaneously-- that is, without applied voltage.

Conclusions & Discussion:

The two solutions, A and B, represent two different routes to silver recovery.

A:  crystallizing silver nitrate to separate it from copper nitrate, etc.;  the AgNO3 is either saved as a reagent or redissolved and used in plating.

B:  precipitating silver chloride to separate it from copper chloride, etc.;  the AgCl is redissolved by thiosulfate, and then the silver is deposited out from the resulting solution by either spontaneous or electrolytic deposition.  Eventually the concentration of aqueous silver ions (or silver-thiosulfate complex ions) drops to the point that no more silver will plate out spontaneously.  This "spent" solution yields mostly copper sulfide on evaporation.  This material, however, should be presumed to contain at least trace amounts of silver;  it may be saved for future recovery operations (e.g., ion exchange).

Method A is obviously much better when the starting silver and other reagents are already 99.9+% pure.  This eliminates lossy purification steps.
In Method B, note that the HCl could be added to the solution after the silver is dissolved completely;  the method of using mixed HCl-HNO3 was done in this experiment to illustrate a curious situation in which Ag seems to go into solution via the nitric acid and comes right back out again as AgCl because of the hydrochloric.  If we could make a diagram to show what we think is happening, it would show newly-mobilized Ag+ ions coming off the metal surface, only to form insoluble AgCl as soon as they meet up with nearby Cl- ions.
It is well-known that silver chloride is highly insoluble, even in nitric acid;  however, beginning chemistry texts often neglect to mention that silver chloride will in fact form a soluble complex ion (AgCl2-) in the presence of excess chloride. AgCl is therefore soluble in excess NaCl or concentrated HCl, especially when some AgNO3 is present  (Merck Index, 1983).  
In this experiment, much of the silver came right out of solution as AgCl.  However, some remained in solution as AgCl2-.  Dilution of this liquid caused more AgCl to precipitate.  This tends to confirm, by the way, that [Cl-]aq must be high in order to allow significant [AgCl2-]aq.  Dilution favors solid AgCl over dissolved AgCl2- ions.
We know that AgCl will not dissolve in HNO3 alone.  On the other hand, what would happen if Ag metal were treated with HCl only?  Would the metal  dissolve and give a precipitate of AgCl, or would it give a solution of AgCl2-?  Could one deposit silver metal from a solution of AgCl2- without applying voltage?  These questions might prompt another, simple experiment.
Method B (as done here) uses spontaneous deposition of silver metal from solution onto a piece of copper, but a better scheme is probably to use graphite electrodes at some carefully-regulated voltage;  the silver would come out of solution on the graphite cathode.  A graphite anode would prevent unwanted metal ions from entering the solution;  a graphite cathode would present no acid-soluble contamination when the Ag was redissolved.  The carbon particles from the graphite would stay behind as the silver dissolved in HNO3.

Update:  If the AgCl precipitate is "aged" for too long, it does not dissolve very well in the thiosulfate.  Upon prolonged standing in thiosulfate solution, the aged silver chloride is slowly replaced with what appears to be silver sulfide.  It leaves a brown, slightly iridescent coating on the glass. The solid Ag2S will slowly settle to the bottom where we can collect it, filter it, and dissolve it in nitric acid;  just keep in mind the significant hazard of H2S that may be released.  Silver sulfide will also dissolve in cyanide solutions.
It is therefore best to use the AgCl precipitate shortly after it forms, minimizing exposure to light the entire time.  

Each route has its relative merits depending on the intended use;  Method A is a more direct route to silver nitrate, a very useful laboratory reagent which happens to be quite expensive when purchased directly from a supply house. 
There is nothing to prevent portions of each method from being used in a combined scheme.  For example, the silver plated out from method "B" could be washed with distilled water and redissolved in HNO3 to give even purer silver nitrate;  the AgNO3 crystals obtained from "A" could be redissolved and treated with HCl or NaCl to precipitate purer AgCl;  and so forth.  For very pure AgCl, Parder et al. (1979) suggest using calcium chloride in dimethyl sulfoxide (DMSO) to solubilize the silver chloride, then precipitating the AgCl again by adding water to the resulting extract.  They suggest roasting the precipitate at 1100 °C with excess sodium carbonate to give pure silver metal (1979). 

In our experiment it was at first troubling that some insoluble, white precipitate was present in the ordinary HNO3 solution made from the silver bracelet.  Had the reagent-grade nitric acid become contaminated with chloride ions?  It was doubtful;  a sample of this nitric acid gave no precipitate with silver ions, even after dilution and long standing.  One source of chloride ions was probably the general-purpose flux used on the silver bracelet during an unsuccessful attempt at silver-soldering it.  The bracelet had never been cleaned prior to melting in the crucible, and it's doubtful the molten silver had been hot enough to volatilize all the flux.  Furthermore, any piece of jewelry worn against the skin will have NaCl residue on it.  This was not washed off prior to melting and dissolving the bracelet.  The greatest source of Cl-, however, was probably the wooden benchtop;  when the molten beads of silver had contacted the wood during the accidental spillage of the silver, they caused charring of the wood and undoubtedly picked up some salts from it.

There is a big difference between laboratory electroplating and industrial electroplating.  The goal in the laboratory is typically to recover the metal from solution for weighing, analysis, or further experiments.  If the metal deposits spontaneously (i.e., with no applied current) and is easy to scrape off, so much the better.  Industrial platers, on the other hand, do a great deal of preparation to ensure the plated metal finish will be durable and attractive;  few industrial processes rely on spontaneous deposition ("dipping"), because such finishes are usually grainy, crumbly, etc.  Industrial plating is a mixture of art and science all its own.
Industrial silver plating usually uses cyanide, which school demonstrations must NOT use (with the possible exception of university labs)-- a mistake with cyanide can kill you almost instantly.  Silver-cyanide plating baths are unsurpassed for durable, bright finishes, but using and properly disposing of cyanides is too advanced for the chemistry novice. 
It is worth additional investigation, however, to see whether a passable finish could be obtained from silver-thiosulfate or other non-cyanide baths.  There are at least a couple of patents involving non-cyanide methods of silver plating.  For example, US Patent #4003806 (1975) claims an improved iodide-based plating bath that uses between 0.12 and 0.20 M Ca++ to help make the silver plate out smoothly.  Silver iodide is normally quite insoluble, but aqueous KI will dissolve it appreciably.

As a final note, one must never forget the danger of silver nitrate solutions;  they are highly destructive to living tissue.  Strong AgNO3 solutions can cause severe burns.  Superficial silver stains on the fingers and hands will eventually grow off as the dead skin is shed;  however, you must NEVER get silver nitrate solutions in or near the eyes.  They can cause permanent blindness.

Challenge Questions:

1.  Once in a while you may spot a person at a flea market testing a piece of silver with nitric acid to see whether it's "real or plate".  First they put a drop of nitric acid on the piece in an inconspicuous spot, then they treat it with ammonium hydroxide to look for a blue coloration (tetraamminecopper (II) or cuprammine ion).
Why, theoretically speaking, might this traditional test for silver vs. silver plate be flawed?  The test relies on the presence of copper.  It is assumed that the copper could only arise from the base metal layer beneath the silver plate.
Antique dealers and metalsmiths have used this test for years.  Decide whether the test works, and why or why not.

2.  We know that silver salts are incompatible with ammonia.  Why, then, do textbooks often recommend dissolving silver chloride precipitates in excess ammonia solution?
This method is part of the qualitative analysis routine taught in most chemistry labs, and it has been for probably a century or more.  Is the method unsafe?  Why or why not?

3.  Write out the equilibria that are happening in a system containing silver ions, chloride ions, and silver chloride.  What equilibria are present when excess chloride is added to dissolve the AgCl?  How does this alter the common-ion effect and the solubility product constant for silver chloride? 
Could you calculate the concentration of each ionic species in a system that contains, for example, 0.1 mole of silver chloride treated with 100 mL of 2 M aqueous sodium chloride?  What constants would you need to know?  What assumptions would you have to make about the ionic species?  Are there experiments you could devise to help you?

4.  How much co-precipitation or inclusion of unwanted ions, if any, might you expect when precipitating silver chloride from a solution containing copper and nickel ions?

5.  Propose chemical equations for the roasting of silver chloride with sodium carbonate.  Is silver metal really one of the products?  What else forms?  Are they reversible reactions?

6.  When a copper wire is immersed in aqueous AgNO3, why does the silver come out of solution bright silvery-gray at first but then black later on?  Propose a reason or search the literature to find out why this might be.


Works cited:

Merck Index, 10th edition.  Rahway, New Jersey:  Merck & Co., Inc.., 1983.

Parder, A. J., B. W. Clare, and R. P. Smith.  "A Novel Process for the Recovery of Pure Silver from Silver Chloride".  Hydrometallurgy 4(3): 233-245 (1979).

United States Patent 4003806.  "Silver Plating Bath".  Invented by R.W. Etter, assigned to RCA Corporation.  Filed May 30, 1975.

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