Chemical Forums
Chemistry Forums for Students => Inorganic Chemistry Forum => Topic started by: Bushka on October 18, 2010, 02:27:31 AM
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Hi - great forum here! I have a couple of bottles of silver chloride that have been sitting in the basement for literally 30 years! I decided to try and extract the silver through electrolysis. I’m dissolving the silver chloride in ammonia and then electrolyzing it. I had a couple of questions if anybody would be good enough to comment.
Question 1: When I dissolve the silver chloride in ammonia I have silver and ammonium cations and chloride and hydroxide anions in solution. When I electrolyze this solution, I precipitate silver at the cathode and evolve chlorine gas at the anode. I assume that I will also evolve hydrogen at the cathode and oxygen at the anode as time passes and there is less and less silver chloride in solution. Assuming this is correct, what am I left with in solution when I’ve gotten most of the silver plated out? I’m thinking some of the ammonia would have evaporated as NH3, but that the remaining NH4 and OH balance would be maintained in solution. Given this, I should be able to reuse this “spent” solution to dissolve some more silver chloride. Does it make sense that I can reuse this solution, or am I depleting the strength of the ammonia somehow during the electrolysis? Are there other reactions of significance that I need to be aware of?
Question 2: The silver chloride salt has a very pale greenish tint. When I put an excess of it into an ammonia solvent and let it sit for awhile, the solution turns very blue. I’m thinking there may be some copper compounds in the salt. Would that explain the color? I don’t see anything that looks like copper during electrolysis, but it could be there in small amounts.
Thanks!
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If there was copper, it would form Tetraamminecopper(II) chloride
http://en.wikipedia.org/wiki/Tetramminecopper%28II%29_sulfate
which would have a deep blue copper.
For silver, it would form diamminesilver(I) chloride, which should be colorless.
http://en.wikipedia.org/wiki/Tollens%27_reagent
also, it mentions risk of formation of silver nitride (explosive) as it ages. WARNING WARNING WARNING.
Wouldn't it be better to heat the AgCl to some high temp, perhaps 1000 C and you get molten silver?
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What are the electrodes made of?
Evolving chlorine is so corrosive it will bring nearly any electrode metal into solution.
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Electrolysis would be a very slow way to recover the silver. You can reduce silver from the diammine complex with an aldehyde like formaldehyde and I suppose glucose would work too (though you might need some heat to help the reaction along, depending on pH and such). This is just the old Tollen's reagent reaction. There are definitely better ways to reduce the silver than electrolysis.
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Hi everybody. I guess about a year has passed and I haven't done anything with my silver chloride. Part of the reason is that I didn't want to blow myself up after reading comments here, and part is because I just haven't had time/priority to try and extract the silver from the silver chloride powder. I would like to take a shot at getting this done now, though.
To recap, I have dissolved what I believe to be AgCl with ammonia bought at the local supermarket. In the past I have heated this solution to improve the solubility of the AgCl, and yield of silver. I electrolyze the solution with a 2/10 amp battery charger and collect what looks like silver at the cathode.
I've included some photos of what I get:
1. dark blue solution (looks lighter in photo) after mixing AgCl with ammonia. This appears to be a copper compound based on replies I've received here.
2. I am using stainless steel for the cathode and anode, and while there is some corrosion, it seems minimal. The second photo is from the cathode. Black/grey film grows quickly and nearly fills the entire beaker if I let it. The film collected in this photo took about 15 seconds. I just wanted to show what the cathode deposit looked like.
3. When the precipitate is dry it forms a grey powder. If I take the side of a screwdriver and compress the powder as best I can, I do get a shiny silvery surface.
At this point I am assuming the blue color indicates that I am plating out some copper with the silver, but it may be in small enough quantities as not to see visually. Also, I'm not sure how great the risk of silver nitride is. I looks to me like I need to have silver nitrate and silver oxide involved in the reactions in order to get to the explosive silver nitride. I don't believe I have nitrides in any appreciable quantity. Also, I've electrolyzed enough in the past to have a small cup full of "silver" and nothing has blown up yet. While not 100% comforting, at least I'm probably not forming anything explosive.
Any comments are greatly appreciated. I'd like to avoid buying more and more reagents because I don't even know if the final product will be worth anything. I mainly wanted to see if I was on-track in getting a blue solution (which it seems I'm not), and to be sure I wasn't going to blow up by doing this.
In answer to one other comment, it isn't practical for me to melt the AgCl at 1000C. I don't even have enough heat to melt the silver powder to see if I really have silver pure enough for sale. I'll have to buy a small acetylene torch or figure out some cheap way to melt this stuff, but that's a subject for down the road.
I would appreciate any comments!!
Dan
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Hi there, I have not done anything with silver chloride or ammonia, but have made batches, as needed, of silver nitrate from sterling silver (92.5% Ag and 7.5% Cu), I use the fact that copper sits higher on the reactivity scale and displaces the silver immediately before it in order to eliminate the copper. However one tip about melting, since I also do that. Either, buy a cheap jewelers torch that heats beyond 1100c, available from rio grand in USA, or, buy a weed burner, they usually come around 1200c, the weed burner I have was only 20 Euros. Take a piece of wood, burn the surface thoroughly, until there is just a charcoaled area (keeps the heat very high as opposed to trying to melt silver on wood itself), place your clump of silver mud on the piece of wood and burn away - the wood will not catch fire if you aim at the charcoaled area, I have done this so many times. Hope something here is of use.
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Thanks Tittywahwah. I've been googling around and there seems to be a lot more stuff on the internet about silver than there was even a year ago. There are several videos (maybe yours) on charcoaling wood and using it for a mold. Anyway, I saw one video where they just put a copper pipe in silver nitrate solution due to the electronegativity difference. I tried that about an hour ago and I was surprised to see the silver precipitating out pretty nicely! I don't know how the yield would compare to electrolysis. My fear with electrolysis was that I might be precipitating out other things like copper in with the silver. It would appear that wouldn't happen as long as there was enough silver still in solution. I suppose if I electrolyzed beyond the point where silver was still in solution, I could start plating out copper on the cathode at that point. If the simple copper pipe trick gives a near complete yield, then copper contamination is not an issue.
Also, since I don't really know what I'm starting with as far as purity (silver chloride), I could redissolve any silver I get with nitric acid and use one of the methods to improve purity. However, do you have any idea of the yield in this kind of thing? For example, if I dissolve 1 oz of pure silver in nitric acid, what percentage of that silver could I expect to have back after precipitating and melting it?
Thanks a lot for your post!
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<<<<<<<<<<If the simple copper pipe trick gives a near complete yield, then copper contamination is not an issue. Also, since I don't really know what I'm starting with as far as purity (silver chloride), I could redissolve any silver I get with nitric acid and use one of the methods to improve purity. However, do you have any idea of the yield in this kind of thing? For example, if I dissolve 1 oz of pure silver in nitric acid, what percentage of that silver could I expect to have back after precipitating and melting it?>>>>>>>>>>
Firstly dissolving silver in nitric acid will yield silver nitrate. You need to heat it. Once dissolved allow to cool, the silver nitrate crystals will appear within seconds, but you will need to allow to evaporate for some days depending upon how much solution you made. I advise you to follow the strict stoichemetry here. If you want to turn the Ag(NO3) crystals into silver just heat this to a very high temperature to get the elemental silver. As for the copper issue, the copper will displace every bit of silver out of solution. Once you filter the solution and you have silver mud in the filter - take a few drops of the filtrate into a test tube and add a single drop of HCL; if it turns cloudy then there is still silver in your copper/silver nitrate solution, if it remains clear then all the silver is out. There is no contamination of copper mixed with silver, as far as my limited chemical knowledge is concerned this is impossible with a displacement reaction, unless it happens on a microscale which is not important for me since I never need super high purity reagents or products. As far as loss is concerned I would expect always some loss, my loss seems always to be just a few grams each time, but this could be also my maths, or even the amount of copper alloy in the original scrap silver that I use has been miscalculated. I can not help you with electrolysis here I am sorry to say. Hope something here helps you.
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Bushka, actually forgot something obvious, too focused on one thing. Since you have AgCl, you might try this: Add NaCl and take into natural light, or UV source if you have it. Watch the silver being separated from the Cl and the solution turn black slowly. This black is actually silver's natural colour when it is so small, particle size. But this black is actually silver metal. Just giving you more info, sorry if you already know this.
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Hi TWW: Can you take it back a step when you mention to add NaCl to AgCl? What solution are you referring to? I'm not sure how I'm precipitating silver by adding NaCl to a solution of AgCl.
Also, I'm finding that the solubility of AgCl isn't great in ammonia, even when boiled for awhile. I'm going to need gallons of ammonia to dissolve everything, and how much Ag I will lose in solution is unknown to me. Of course, if the copper pipe really will convert 100% of the Ag in solution back to solid silver, then having lots of ammonia isn't such a problem I guess.
I thought I read someone who said you can just melt AgCl crystals at high temp, and the Cl will burn off, leaving pure silver. Do you know anything about this? I assume I will need something like a MAPP torch from Home Depot. Any other options (ie, the cheapest way) to melt silver, or melt AgCl into silver?
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Ok, firstly I am not a student of chemistry, all my knowledge is self taught. Here is the best I can offer, I have never done this experiment, but parts only, as for ammonia i have never tried that either. I simply use silver nitrate for different things and make it as needed from scrap silver alloy. So here is my best: Adding NaCl to water with AgCl mixed in increases the Cl- ion in solution. NaCl is highly soluble as you know but AgCl is not, however the Cl- ions will dissociate in the water so that you have Ag+, Na+ and Cl- ions all having a party and not dancing together. The solubility of the AgCl will naturally go right down because of all this, add the UV light to the equation and you will see the silver particles turn dark brown to black, these will be your precipitates, not the Cl ions and not the Na ions, these will remain dissolved. Now all this is theoretical and not based on any observation I have made. I do know that when you add silver nitrate to tap water it turns milky white. This is because there are chloride ions in the tap water and the silver nitrate is losing its silver ions and becoming silver chloride precipitate, which as you know is insoluble. All that I have said is my best knowledge on this, it would need a student chemist or teacher to take it any further. But Light will force the silver to separate from it's chloride bond in solution - though on this I have never experimented, I have seen a couple photographs of this some time ago.
HAve you seen: http://www.youtube.com/watch?v=zJd0EnLwt44
http://en.wikipedia.org/wiki/Silver_chloride
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TWW: Thanks for the reply. I'm kind of the opposite of you. Although I've been out of it for about 15 years, my background is in chemical engineering. So I understand some/many of the interactions if I brush off my brain, but I have no practical experience in refining silver. I don't want to reinvent the wheel, nor create something that can be dangerous.
I understand how adding Cl- ions can shift the equilibrium of AgCl solution back to solid Ag. However, the problem is that AgCl has low solubility in water, and from what I can tell, limited solubility in ammonia. The other issue I'm running into is that everybody says something different and contradictory to each other. I've seen the links you provided (thanks). I've seen other links that say not to heat AgCl, not to put copper pipe into your silver solution, and so on. I decided to take a step back and see how much silver I'm likely to recover, and if any of this is worth the effort. (Part of my effort is for the fun of it, part is the financial gain). So I weighed all the silver and silver chloride I have with an electronic scale and I have around 770 grams. AgCl is 75% Ag by weight, so 770 works out to 578 grams. I can't expect 100% yield, but some of that 770 grams is already pure and/or dirty silver. So if I go with the 578 gram yield, this equates to about 20 ounces of silver. Not bad and more than I thought. The spot price for silver is about $34/oz.. How close to spot can I expect to get for homemade silver ingots?
So with the above in mind, I'm thinking which way to go. I have to spend $50 for a MAPP torch from Home Depot (you know any cheaper way to get hot enough?), need borax and then possibly some chemicals other than ammonia if melting the AgCl is found to be the best bet. If I can get anywhere near the spot price for silver, it makes sense to spend a some money on getting the right chemicals to get the silver. The AgCl I have appears to have some copper contamination because it turns light blue when I add ammonia. I'm thinking that even if I just melted the AgCl with a torch, I would need to dissolve away the copper first. So maybe I should first wash my dirty AgCl with ammonia to dissolve the copper, and flush it until there is no more blue color (I suppose I'm not really dissolving the Cu, but more dissolving some of the AgCl, which then sets up the solid Cu to be oxidized by the silver in solution). Then I can recover the silver from the spent solvent by dropping in a copper pipe and precipitating the silver. That would give me silver from the dirty copper solution, plus clean AgCl precipitate that maybe could be melted directly.
Does anybody know of any solvent that will dissolve AgCl better than ammonia does?
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Have you looked into reducing AgCl to Ag using NaOH and dextrose? This is one common way to get the reduction done and it is fast and exothermic. I've done it and the yield was good despite the clumpy AgCl I started with. I think the yield would be improved the more fine AgCl particles are.
After you have that silver 'cement' it's just a matter of melting it in a crucible. Some sodium carbonate with borax/boric acid flux will help to reduce any left over silver chloride in the cement during the melt.
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Have you looked into reducing AgCl to Ag using NaOH and dextrose? This is one common way to get the reduction done and it is fast and exothermic. I've done it and the yield was good despite the clumpy AgCl I started with. I think the yield would be improved the more fine AgCl particles are.
After you have that silver 'cement' it's just a matter of melting it in a crucible. Some sodium carbonate with borax/boric acid flux will help to reduce any left over silver chloride in the cement during the melt.
Hi. Thanks for the reply. If you check TWW's last post, he provides a link to this method. It looks like the way to go if I can clean up the copper first. The problem is that instructions are pretty vague. How much NaOH to use - stoichiometric or an excess? how much dextrose? sodium carbonate, borax and so on. If you or somebody could help me with those questions then I think I'd be all set!
Regards.
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Hi Bushka:
Ammonia is not then the good way to go at all....Found this
Dissolve the silver chloride in a solution of ammonia. This is a complexation reaction.
AgCl(s) + 2NH3(aq) --> [Ag(NH3)2]+ + Cl-
Add copper metal or almost any metal above silver in the activity series. This is a redox reaction. The silver is reduced and the copper is oxidized.
2[Ag(NH3)2]+ + Cu(s) --> 2Ag(s) + [Cu(NH3)4]+
============= Follow up ==============
Placing copper metal into solid silver chloride will not result in any appreciable reaction. In order for the replacement reaction to occur, silver ions must come into contact with copper metal, but the silver ions are tied up in an insoluble compound.
The above was something that I found on the net. HOWEVER, very important, copper works very very well at displacing silver from nitrate solutions where you have just dissolved silver alloy - I do it all the time. Now then, have you read this by any chance?
http://www.chemicalforums.com/index.php?topic=32471.0
also this: http://www.finishing.com/195/29.shtml
Sorry if I can not answer all of your questions, I wish I could.
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Hi. Thanks for the reply. If you check TWW's last post, he provides a link to this method. It looks like the way to go if I can clean up the copper first. The problem is that instructions are pretty vague. How much NaOH to use - stoichiometric or an excess? how much dextrose? sodium carbonate, borax and so on. If you or somebody could help me with those questions then I think I'd be all set!
Regards.
Sorry, I didn't see that. As to quantities to use for reduction, I used the info from this patent, end of page 3, beginning of page 4. Basically, you use an excess of both NaOH and dextrose.
http://www.google.com/patents/US3658510?printsec=abstract#v=onepage&q&f=false
As to the flux quantities, I don't think they are too critical. You want to coat your crucible with the borax/boric acid (1:1) mix and also perhaps cover a bit of the surface of the melt. The soda ash is what reduces left over AgCl though and you would need a 1:1 molar ratio there, if you knew how much left over AgCl you had, which you won't. So, you'd guess. It releases CO2 when it reacts with AgCl so if you see fizzing on adding annhydrous NaCO3, there's still AgCl left. I think. A patent regarding that reaction:
http://www.google.com/patents/US4306902?printsec=abstract#v=onepage&q&f=false
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Sorry, I didn't see that. As to quantities to use for reduction, I used the info from this patent, end of page 3, beginning of page 4. Basically, you use an excess of both NaOH and dextrose.
http://www.google.com/patents/US3658510?printsec=abstract#v=onepage&q&f=false
As to the flux quantities, I don't think they are too critical. You want to coat your crucible with the borax/boric acid (1:1) mix and also perhaps cover a bit of the surface of the melt. The soda ash is what reduces left over AgCl though and you would need a 1:1 molar ratio there, if you knew how much left over AgCl you had, which you won't. So, you'd guess. It releases CO2 when it reacts with AgCl so if you see fizzing on adding annhydrous NaCO3, there's still AgCl left. I think. A patent regarding that reaction:
http://www.google.com/patents/US4306902?printsec=abstract#v=onepage&q&f=false
That's interesting. I think I get it now.
Other questions:
* Everybody mentions Borax, but googling says borax isn't the same as soda ash. I assume real soda ash is what I want.
* Maybe you covered this before but what does the boric acid do?
* What's the cheapest way to get a hot flame that will work well. Home Depot has a MAPP torch but it's $50. Is that the best way to go?
* Do you know if the silver reduction can be done in a stainless steel container? I'm a little nervous about doing it in a plastic bucket.
* I have a crucible to melt the silver, but nothing to pour it into. TWW uses a block of wood carbonized. Do you also do that?
Thanks for all the help, and it would be great for anybody else to chime in. (Then I'll shut up with all the questions). :)
Dan
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Borax is Sodium borate. I use it as a flux and as an excellent shield against oxidation of silver and copper, it is also used to keep melting silver in a single globular shape and to stop it from splitting as you melt it, also acts to prevent silver oxide from forming during the heating process and does the same for copper.
Soda ash is Sodium Carbonate, a bleaching agent, a very strong alkali.
I use the carbonized wood, pouring it is not necessary unless you are actually wanting to pour it into a mold. Melt and allow to cool on the wood.
Hot flame? For small amounts of Silver under about 10 grams a normal butane pen flame from a jewellers supplier will do since the carbonized wood underneath holds and reflects back up at the bottom of the silver the heat necessary to melt the silver as you heat it from above. Otherwise as I said before a weed burner is much hotter, I use this for larger amounts. Maybe I should add about melting silver that it is not really fluid like say when you melt tin or lead. You can pour the silver onto concrete really, as long as it does not run off.
Silver reduction? I assume you mean the reaction itself? Neither stainless steal or plastic - Never. Glass for absolutely everything. I hope that I have not misunderstood what you meant by 'Silver Reduction'.
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OK, for anybody still paying attention. I think the lye/corn syrup method is easiest for me to get silver cement. However, I have some confusing issues. Any comments on my saga below are appreciated:
I’ve had these bottles of silver chloride in my basement for about 30 years, and finally got around to doing something with them. A year or so ago, I dissolved some in ammonia and electrolyzed it. I didn’t get very far, though. After revisiting the subject recently on the internet, I have found out much more information than seemed available even a year or two ago. So I’ve been playing around with techniques such as copper replacement.
So now I am serious about converting this stuff to silver ingots and being done with the issue. The problem is that I’m having a little trouble finding a good way to take what I have now, and turn it into pure, clean silver. Here’s what I mean: I have about 750 grams total of dry material. Some is pretty clean AgCl, some is dirty AgCl (meaning unknown amounts of copper contamination, and/or other metals, as well as some small bits of foreign debris caught up in the AgCl). I have another jar of lumpy AgCl that seems to have partially converted itself over to silver over the years. There are also some chunks of green solid in the bottle. In addition, I have several jars of ammonia with copper in solution, with some amounts of silver and likely silver chloride in solution as well as precipitated. I have a couple of other flasks with a small amount of clear ammonia covering a mix of AgCl and silver precipitate. Basically, I’ve got all these different “cats and dogs” in these bottles and I want to get all the silver out of them.
My goal was to get everything into clean AgCl, and then use the lye/corn syrup method to get silver cement. One concern is how do I get everything (silver, AgCl, copper compounds?) into solution so that I can filter out any bits of debris? I guess any further questions would depend on that answer.
I washed one of the bottles of AgCl with some ammonia to remove any copper, and then I found that if I left copper pipe in the solution for a few days, I get a surprising amount of silver precipitate. Is this precipitation free of copper after washing. Should I use just distilled water to wash?
The other concern is that I was originally electrolyzing the AgCl ammonia solution in the presence of copper. I’m thinking that if I left the electrolysis going too long, I could start plating out copper and contaminate the silver. If so, then would I have to dissolve the silver in nitric acid, and then add HCl (or salt I suppose) to get clean AgCl and copper in solution?
That’s mostly it. It seems a bit scatter brained but I hope my general question is understood.
Thanks.
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AgCl dissolves in ammonia.
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AgCl dissolves in ammonia.
Have you tried to dissolve AgCl in household ammonia? It takes a ton of ammonia. I read that you can theoretically dissolve 1 kg of AgCl in 1 liter of concentrated (something like 25%) ammonia. I don't know the concentration of household ammonia, but it would take gallons to dissolve what I have.
How about dissolving the silver in nitric acid, pouring it off, and then dissolving the AgCl in concentrated lye?
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I think from what you describe your best bet is going to be to sell the crude, contaminated, clumpy AgCl to someone who knows how to do the extraction. Or pay them a fee to do the extraction and send you the purified silver. As you are finding out, it is not easy. IF you had pure AgCl that was fresh (finely divided) it would be quick and easy to do the lye/glucose reduction and then cast pure ingots.
But with what you have, there will be large amounts that don't get reacted. Clumpy AgCl doesn't react or dissolve easily at all and is a b$*%( to grind. It will reduce with sodium carbonate in a melted state but that isn't an easy process either since you need a furnace and proper crucible and sodium carbonate tends to dissolve crucibles. Once you've reduced it you will have some copper contamination. Unfortunately, copper I oxide is soluble in a silver melt.
If you want pure silver you would need a two step process, reduce the crude metal, cast into an ingot and then electrolyze in silver nitrate bath to purify and melt and cast again pure. Unless you are looking for the fun of doing the above, just sell it to a small refiner.
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eprigge: Selling off doesn't sound like much fun! Honestly, though, I'm half in it for the money, and half for the fun of it. I may have overstated the "mess" that the AgCl is in. I'll have to take a pic of what I have at this stage. Does the ingot have to be .9999 pure silver to have value? Coins are 90% silver and they have good value. My understanding was that the ingot buyer tests the silver and pays you based on that test. Not having ever sold an ingot, I could be way off on that.
For now I'm going to get the rest of the AgCl out of the second bottle (prolly have to bust the glass to get it out), and then see exactly what I have to work with.
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eprigge: Selling off doesn't sound like much fun! Honestly, though, I'm half in it for the money, and half for the fun of it.
That problem afflicts us both. If you want to sell a 'fine' silver ingot it better be .999+. I can pretty much guarantee you won't get 3 9s from from crude AgCl reduction, maybe 2 9s. In some cases impurities flux out and float off on top of a melt but copper oxide will dissolve in and contaminate a silver melt. So it has to be electrolytically removed. If you sell an ingot where you can't characterize the purity, you can expect a lot less money. You can always have an a$$ay done I guess.
You could electrolyze out of an AgCl bath directly (I would use thiosulfate instead of ammonia to dissolve it) but it is *really* hard (slow) to get old, clumpy AgCl to dissolve in anything AND you need an inert anode (platinized or perhaps carbon). And, as you pointed out, you need to know when your bath is depleted or else the copper/impurities plate out too. AND if you don't get it all out of the bath you are wasting it...
http://www.saltlakemetals.com/Solubility_Of_Silver_Chloride.htm
If you want fun, you got it. :-)
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I was trying to avoid Salt Lake Minerals because their procedures for getting pure silver were depressingly complicated and expensive if you don't already have a real lab. ::) But thanks for the link to solubilities.
There's something I don't understand, though. You talk about having to electrolyze silver and guess at endpoints etc etc to get anything really pure and worth selling, but what about all these guys online plating out silver from copper wire. (as in http://www.youtube.com/watch?v=znByBb6wf7g ). In that video, he says he plated out most of the silver from AgNO3 with copper wire, and made some AgCl in the video from what was left over. If people are starting with sterling silver (loaded with copper) and then adding copper wire, then what are they doing with the silver they cement or precipitate out? If they wash it and then melt it into ingots, are you saying it won't be 999 purity, or something that will bring top dollar?
Also, how do you prove you have achieved 999? I assumed the buyer would confirm this with some kind of simple jewelers test.
Regards.
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I was trying to avoid Salt Lake Minerals because their procedures for getting pure silver were depressingly complicated and expensive if you don't already have a real lab. ::) But thanks for the link to solubilities.
There's something I don't understand, though. You talk about having to electrolyze silver and guess at endpoints etc etc to get anything really pure and worth selling, but what about all these guys online plating out silver from copper wire. (as in http://www.youtube.com/watch?v=znByBb6wf7g ). In that video, he says he plated out most of the silver from AgNO3 with copper wire, and made some AgCl in the video from what was left over. If people are starting with sterling silver (loaded with copper) and then adding copper wire, then what are they doing with the silver they cement or precipitate out? If they wash it and then melt it into ingots, are you saying it won't be 999 purity, or something that will bring top dollar?
Also, how do you prove you have achieved 999? I assumed the buyer would confirm this with some kind of simple jewelers test.
Regards.
I always begin with sterling silver because I do not have any pure silver. After the displacement reaction with copper, I then filter and wash with a few rinses while still in the filter bag, I use distilled water after one treatment with 0.1 mole Nitric acid to help flush out any copper ions that might be left. After allowing to dry it's just a simple case of heating on carbonized wood. The silver is, I believe, at the minimum 98% no less. More steps can be achieved I believe to purity by re-dissolving in Nitric acid, getting the silver nitrate crystals then heating these back into silver. (This seems to be the best method for obtaining max purity in a home situation/backyard chemistry so to speak)
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I found this in one of my articles that I have collected: Now It may be of no use to you but I thought that I would post it anyway - I have also given a link to the original site where I found it.
""A third method of reduction of the chloride, is one which is very convenient for those who do not possess a furnace, or have the convenience of fusing ores or residues. Moisten the chloride with dilute hydrochloric acid, and immerse a plate of zinc in the moistened mass for several hours. Decomposition will gradually take place, the silver being deposited, whilst the soluble chloride of zinc is formed. After the chloride has been thus completely decomposed, the remaining zinc is withdrawn, and the precipitate is washed with dilute hydrochloric acid, until there is no longer any precipitate formed in the decanted fluid by means either of ammonia or of sulphide of ammonium. The precipitate is next well washed with warm water. It is now in a condition for being dissolved in nitric acid.""
http://albumen.conservation-us.org/library/monographs/sunbeam/chap13.html
Link:
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Thanks for the comments TWW. Can I inquire as to what you do with 98+% silver ingots?
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Thanks for the comments TWW. Can I inquire as to what you do with 98+% silver ingots?
I use it to make silver nitrate anhydrous for use in old photographic processes, I mean 1840's style, not black and white film.
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Thanks for the comments TWW. Can I inquire as to what you do with 98+% silver ingots?
I use it to make silver nitrate anhydrous for use in old photographic processes, I mean 1840's style, not black and white film.
That's very interesting. I guess according to eprigge I'm going to need to get 999 purity in order for it to be sale-able.
The AgCl I'm dealing with was made by high school students a long time ago. There are some green flakes and specs in the powder, which I attribute to some kind of copper salts. I've washed all the silver/AgCl powder in dist. water to start, and next I thought I'd soak everything in an excess of muriatic acid to dissolve the copper salts. I realize it won't dissolve copper, but I don't have any reason to think there is pure copper in the powder. I have to pick up some muriatic, but I tried dissolving a little in some of that CLR I had sitting around (HCl plus sulfamic acid). It seemed to dissolve the copper salts pretty well.
From there we'll see... but at least I can take a stab at cleaning up, drying and pulverizing the AgCl.
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""muriatic acid to dissolve the copper salts. I realize it won't dissolve copper,""
Just as a side note, Copper will dissolve in HCL. One of two things you need, (if you want to do this ever), add H2O2. or a cheaper way, insert a tube from an aquarium pump. I could give you all the copper ion interchange details on why if you want. But suffice it to say that all you need is a source of oxygen in the HCL to dissolve copper. I often etched copper with copper chloride solution. I just used an aquarium pump to keep the oxygen level at a certain level and regularly titrated to make sure my copper chloride was always at an appropriate level, ionically speaking. Anyway that was a side note.
As for selling, I have no idea, never done that, pity you are in USA, I would purchase some silver from you - sure we could have come to an amicable arrangement. As it is I am running out of sterling silver bits, so may well have to purchase a silver coin soon. Pity, I enjoyed complicating my life with the impure stuff.
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""muriatic acid to dissolve the copper salts. I realize it won't dissolve copper,""
Just as a side note, Copper will dissolve in HCL. One of two things you need, (if you want to do this ever), add H2O2. or a cheaper way, insert a tube from an aquarium pump. I could give you all the copper ion interchange details on why if you want. But suffice it to say that all you need is a source of oxygen in the HCL to dissolve copper.
Interesting that you've done that. I was reading some stuff on needing an "oxidizing" acid, and that peroxide would help. I didn't realize an air pump would do it, too. I think the problem with an oxidizing acid like nitric, is that it will dissolve the silver, too. I'm assuming the HCl/peroxide or air method would tend to dissolve silver as well? Maybe the silver would dissolve and immediately precip out as AgCl? I think I read that somewhere.
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Yes, 2Ag+2HCl=2AgCl+H2
Notice that you need to do this stoichemetry, accurately. Although I have to add that if you don't then Ag+HCl will not react. You need two moles of silver and two moles of HCl for a reaction. However I have never done this so I actually can not tell you why you need two moles of each in order to start a reaction - as I said I am no chemist and this frustrates me sometimes - logic dictates that if Two moles react then why not one? I confess a lack of insight here. But if anyone is reading this please tell me why.
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Yes, 2Ag+2HCl=2AgCl+H2
Notice that you need to do this stoichemetry, accurately. Although I have to add that if you don't then Ag+HCl will not react. You need two moles of silver and two moles of HCl for a reaction. However I have never done this so I actually can not tell you why you need two moles of each in order to start a reaction - as I said I am no chemist and this frustrates me sometimes - logic dictates that if Two moles react then why not one? I confess a lack of insight here. But if anyone is reading this please tell me why.
I can tell you. You don't really need 2 moles of each, but to balance the equilibrium equation you provided above, you need 2 moles each of Ag and HCl if you want to represent 1 mole of H2. Notice it takes 2 moles of HCl to liberate one mole of H2. You could also write:
Ag + HCl <=> AgCl(s) + 1/2H2
The (s) indicates a solid or precipitate and the 1/2H2 shows a half mole of H2 liberated. It's simply a matter of how you want to show the equilibrium. Of course in this example you said you are also providing oxygen through either the bubbler or peroxide, so that should really be in the equation as there will be no reaction between Ag and HCl without the oxygen.
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Funny, I was just searching for some insight on this and found this article after a reference stated that the above reaction is impossible:
http://www.sciencedirect.com/science/article/pii/S0021979706006710
I will digest your answers though. Yes because in any reaction with HCl hydrogen HAS to be a product, is that why?
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Funny, I was just searching for some insight on this and found this article after a reference stated that the above reaction is impossible:
http://www.sciencedirect.com/science/article/pii/S0021979706006710
I will digest your answers though. Yes because in any reaction with HCl hydrogen HAS to be a product, is that why?
Not sure exactly what you are asking/saying. However, and I'm going back 30 years, electronegativity is what determines whether a metal will dissolve in a strong acid like HCl (I'm sure there are other factors, and someone will correct me). In order to dissolve the metal, the H+ ion from the HCl solution has to be able to strip away an electron from the metal, thereby forming H2 and an ionized atom of metal. For metals like gold and platinum, their electronegativies are higher than hydrogen, so they are able to hold onto their electrons. If you look up an electronegativity table like here: http://tinyurl.com/alojvc8 you'll see that silver actually has a slightly lower number then hydrogen. So, it should dissolve, slowly, as I understand it. However I'm reading that since AgCl is so highly insoluble, it forms a protective shell over the silver, and prevents any further attack by the acid. Aluminum apparently forms an impenetrable shell of aluminum oxide under certain acidic conditions, and cannot be dissolved further. On the other hand, the H+ from nitric acid IS able to dissolve silver because silver nitrate is soluble and cannot form a protective shell, unlike AgCl.
Not sure if that's what you were getting at...
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electronegativity is what determines whether a metal will dissolve in a strong acid
You are confusing electronegativity with reactivity series.
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<<Not sure exactly what you are asking/saying.>> Sorry was getting slightly confused in my writing. What I meant was that this reaction: 2Ag+2HCl=2AgCl+H2 is often stated as being impossible; but that the article I read (link given above) proves otherwise when it comes to silver nano particles. A bit off tangent regards conversation, I was just investigating it myself since I have never used silver chloride or had a use for it yet and therefore needed to do some speed reading. Trouble is so many people say that silver and HCl do not react, as they often do with copper and HCl, well I get a little bemused at these people because clearly these chemicals do react with each other under certain conditions, so one should qualify these statements by adding this. I believe.
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electronegativity is what determines whether a metal will dissolve in a strong acid
You are confusing electronegativity with reactivity series.
I'm sure you're right. I did a bit of googling of the terms, but it seems to me that electronegativity and reactivity are almost the same thing. Intuitively aren't we really talking about how "badly" the metal wants to hang on to its electrons, vs hydrogen's ability to grab it? Electronegativity seems a more straightforward way to understand it. On the other hand I don't want to be spreading bad info, so correct me where I'm wrong.
Thanks!
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Trouble is so many people say that silver and HCl do not react, as they often do with copper and HCl, well I get a little bemused at these people because clearly these chemicals do react with each other under certain conditions, so one should qualify these statements by adding this. I believe.
Yes, you're right. Maybe that's why there is so much conflicting (to me) information on how to refine silver or silver chloride. Somebody says, "just dissolve it in ammonia." OK, but if you look at the solubilities I'll end up needing 10 gallons of ammonia.
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I did a bit of googling of the terms, but it seems to me that electronegativity and reactivity are almost the same thing.
No, they are not. They are not completely unrelated, but if you you will take a closer look at the table you have posted, electronegativities of gold, sulfur and iodine are pretty close to each other, yet sulfur and iodine are highly reactive, while gold is very inert.
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Just as a side note, Copper will dissolve in HCL. One of two things you need, (if you want to do this ever), add H2O2. or a cheaper way, insert a tube from an aquarium pump.
Then you are dissolving copper not in the HCl, but in the mixture of HCl and additional oxidizing agent. That's not the same.
I suppose you would not say calcium carbonate dissolves in water - but by your logic it does, you just need to add some HCl.
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Just as a side note, Copper will dissolve in HCL. One of two things you need, (if you want to do this ever), add H2O2. or a cheaper way, insert a tube from an aquarium pump.
Then you are dissolving copper not in the HCl, but in the mixture of HCl and additional oxidizing agent. That's not the same.
I suppose you would not say calcium carbonate dissolves in water - but by your logic it does, you just need to add some HCl.
Well, I agree with your point to an extent. No I would never say that CaCO3 dissolves in water - it does not. However my logic with regard to copper and silver is that what you are adding is already there in solution, just not enough that's all - Oxygen. Not being a chemist I see that Oxygen - yes can be an oxidizing agent - but on the other hand is not really an ingredient so to speak (obviously I have set myself up for a downfall here). But let's take gold, I would agree that gold will not react with any of the acids. The fact that you need to mix HCl and Nitric together to dissolve gold is not the same as saying that gold therefore does react with acids. My only argument is weak I know, but concerns the fact that extra oxygen is all you need to promote the reactions and therefore copper and silver will react with these acids.
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I did a bit of googling of the terms, but it seems to me that electronegativity and reactivity are almost the same thing.
No, they are not. They are not completely unrelated, but if you you will take a closer look at the table you have posted, electronegativities of gold, sulfur and iodine are pretty close to each other, yet sulfur and iodine are highly reactive, while gold is very inert.
I've been reading up some on reactivity (activity) series. All the sites seem to say pretty much the same thing, but don't explain WHY one element is more reactive than another. It seems to me that electronegativity is a physical explanation as to why. I can't speak to electronegativities between gold and iodine, but if we are looking only at transition metals, do elect. values not explain why silver reacts as it does relative to hydrogen, and so on with gold? In other words, why is it incorrect to say that hydrogen has a greater affinity for electrons than silver, and gold has a greater affinity for them than hydrogen, and that this explains their behavior?
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Well, I agree with your point to an extent. No I would never say that CaCO3 dissolves in water - it does not. However my logic with regard to copper and silver is that what you are adding is already there in solution, just not enough that's all - Oxygen. Not being a chemist I see that Oxygen - yes can be an oxidizing agent - but on the other hand is not really an ingredient so to speak (obviously I have set myself up for a downfall here). But let's take gold, I would agree that gold will not react with any of the acids. The fact that you need to mix HCl and Nitric together to dissolve gold is not the same as saying that gold therefore does react with acids. My only argument is weak I know, but concerns the fact that extra oxygen is all you need to promote the reactions and therefore copper and silver will react with these acids.
I'm reading your discussion with Mr. Borek, and I think I'm seeing the difference between theory and real world application. I think it is wise to listen to Mr. Borek because he appears to be a chemist. In online forums people can say whatever they want, and others might come away with the wrong impression. So in your world, where you are interested in dissolving silver for commercial reasons or whatever, HCl WILL dissolve silver (or, let's say you can use HCl to dissolve silver) - just make sure you have enough oxygen. In Mr. Borek's world, there is no reaction between HCl and Ag. I think it's good and important on forums to be precise.
I think it comes down to semantics. It's like, "Yeah, you can get HCl to dissolve silver but it's not really just the HCl. You need enough oxygen." I wouldn't over argue the point, personally.
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I've been reading up some on reactivity (activity) series. All the sites seem to say pretty much the same thing, but don't explain WHY one element is more reactive than another. It seems to me that electronegativity is a physical explanation as to why.
It is not. It happens that both electronegativity and reactivity to some extent depend of the ionization energies, so in some cases there is an obvious correlation between one and the other. But it works only for some cases, so it has no predictive power which makes it useless in this context.
I can't speak to electronegativities between gold and iodine, but if we are looking only at transition metals, do elect. values not explain why silver reacts as it does relative to hydrogen, and so on with gold? In other words, why is it incorrect to say that hydrogen has a greater affinity for electrons than silver, and gold has a greater affinity for them than hydrogen, and that this explains their behavior?
Technically electron affinity is defined as amount of energy released when an electron is added to a neutral atom, so it is of no use here, as H+ is not neutral. I guess what you really mean is a "reversed ionization" - that is, amount of energy released when electron is added to a cation, and it is not difficult to address this case. Let's see the numbers:
Hydrogen - 13.6 eV
Silver - 7.58 eV
Gold - 9.22 eV
(see http://www.lenntech.com/periodic-chart-elements/ionization-energy.htm)
So it is incorrect to say that
hydrogen has a greater affinity for electrons than silver, and gold has a greater affinity for them than hydrogen
because it is not true. Hydrogen cation has the highest affinity for electrons of the three.
Note that part of the problem here is that you have not defined what you mean by the electron affinity - you just throw the term and add some hand waving hoping it will yield a result ;)
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I've been reading up some on reactivity (activity) series. All the sites seem to say pretty much the same thing, but don't explain WHY one element is more reactive than another. It seems to me that electronegativity is a physical explanation as to why.
It is not. It happens that both electronegativity and reactivity to some extent depend of the ionization energies, so in some cases there is an obvious correlation between one and the other. But it works only for some cases, so it has no predictive power which makes it useless in this context.
I can't speak to electronegativities between gold and iodine, but if we are looking only at transition metals, do elect. values not explain why silver reacts as it does relative to hydrogen, and so on with gold? In other words, why is it incorrect to say that hydrogen has a greater affinity for electrons than silver, and gold has a greater affinity for them than hydrogen, and that this explains their behavior?
Technically electron affinity is defined as amount of energy released when an electron is added to a neutral atom, so it is of no use here, as H+ is not neutral. I guess what you really mean is a "reversed ionization" - that is, amount of energy released when electron is added to a cation, and it is not difficult to address this case. Let's see the numbers:
Hydrogen - 13.6 eV
Silver - 7.58 eV
Gold - 9.22 eV
(see http://www.lenntech.com/periodic-chart-elements/ionization-energy.htm)
So it is incorrect to say that
hydrogen has a greater affinity for electrons than silver, and gold has a greater affinity for them than hydrogen
because it is not true. Hydrogen cation has the highest affinity for electrons of the three.
Note that part of the problem here is that you have not defined what you mean by the electron affinity - you just throw the term and add some hand waving hoping it will yield a result ;)
Borek - maybe it is a language thing (English your first language?), but you are taking me too literally. From the context you should realize that I am not talking about "affinity" as in eV's. I have not defined electron affinity because I wasn't talking about it. You assumed I was. I was actually talking about love. Hydrogen loves it's electrons more than silver, but not as much as gold does. :-* OK I think I'm safe because unless they've come up with some new chemical terms in the last few years, I don't think "love" is in the CRC handbook.
Let me get back to the original question, because I would like to understand better. Rather than you simply telling me what I'm saying wrong, help me understand the right answer. I originally said that electronegativity explains why sliver and gold react the way they do with hydrogen ions. You said I was confusing electro. with reactivity. Are you saying that reactivity is the correct explanation? If so, what causes one material to be more reactive than another (admittedly I think this can be a complicated answer).
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TWW if you're still reading: How much concentrated nitric acid do you need to dissolve silver? I read somewhere that it takes about 150 ml of acid to dissolve 1 ounce of silver. Does that sound about right, and what concentration do you use when you aren't "making your own"?
Also, I've taken all of my "dirty" silver chloride and soaked it in muriatic (HCl) acid. The solvent turned green/olive, which I presume might be copper sulfate. Interestingly, if I add ammonia to the acidic green solvent, and raise the pH above 7, the green turns bright blue. I assume it has something to do with copper valence. (?) Anyway, originally there were dark green flakes and crystals in the AgCl, and these are now all dissolved. I'm at the point where flushing the AgCl with additional acid does not yield a visibly green/yellowish liquid.
So now I have AgCl mixed in with some bits of silver. Some of it is a bit lumpy, so here's my plan: Dry out the AgCl and crush everything as finely as possible in a mortar and pestle. Resoak in muriatic acid to flush out any copper salts. Rinse and then add ammonia and soak. The idea is to dissolve anything else in the AgCl that might be soluble in ammonia. I would then reacidify the ammonia to drop out any AgCl that dissolved in the ammonia. I would discard the solvent, thinking that it might contain some contaminants that went into solution with the ammonia, and then did not precip. out like the AgCl does. Seems like this ammonia step wouldn't hurt anything.
Now the interesting part. I would like to add nitric acid to the pulverized AgCl powder in order to dissolve any free silver. If I poured off the solvent, I would have silver nitrate, which could then be converted to silver via the acidify/syrup method. Dissolving the silver with nitric acid in the first place would attempt to remove impurities from the silver. The next question is dealing with the remaining AgCl. I would like to add lye directly to the AgCl to see if it can be converted to silver, again, via the syrup method. The question is whether the AgCl is too "hard" to respond to the lye. If I can at least get all my silver chloride/silver powder turned into an ingot, then maybe I can analyze purity and go from there.
Any comments?
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Borek - maybe it is a language thing (English your first language?), but you are taking me too literally. From the context you should realize that I am not talking about "affinity" as in eV's. I have not defined electron affinity because I wasn't talking about it. You assumed I was. I was actually talking about love. Hydrogen loves it's electrons more than silver, but not as much as gold does. :-* OK I think I'm safe because unless they've come up with some new chemical terms in the last few years, I don't think "love" is in the CRC handbook.
There is a reason why "love" is not in the CRC handbook - it is not precisely defined, it can't be measured, it can't be used to predict properties of elements, so in this context it is useless.
It is similar problem like with dissolving silver in acids - you need to be precise with what you really mean, otherwise things you say don't mean anything.
English is my second language, but the way I see it is not the problem here. Problem is you are trying to use hand waving for the explanation. Sorry, it won't work. You need to be precise with the meanings, otherwise sooner or later we will run into communication problems - as we already did.
Let me get back to the original question, because I would like to understand better. Rather than you simply telling me what I'm saying wrong, help me understand the right answer. I originally said that electronegativity explains why sliver and gold react the way they do with hydrogen ions. You said I was confusing electro. with reactivity. Are you saying that reactivity is the correct explanation? If so, what causes one material to be more reactive than another (admittedly I think this can be a complicated answer).
Reactivity is a property that can be checked - basically it says which metal replaces which metal in the solution. As these reactions are in a way similar to the reaction between metals and H+ cations, H+ is also on the list. This is checked experimentally and the series can be used to predict what will happen when you add piece of a metal to the solution containing H+ or other metal cations. Does it explain "why"? Depends on how you look at it, I would say it doesn't.
But there is really no simple explanation. Basically reactivity of the metal in water solutions depends on its ionization energy and solvation effects. The lower the ionization energy, the more reactive the metal, the higher the solvation enthalpy, the more reactive the metal, but in some cases both can counteract each other, plus often we are talking about several ionization steps, each with its own ionization energy and each cation with its own solvation enthalpy. This can become quite complicated. Reactivity series ignores all these considerations, just gives a final (and because of the fine details not always precise) answer.
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Reactivity series ignores all these considerations, just gives a final (and because of the fine details not always precise) answer.
I understand better. However, it seems your original answer isn't correct, either. I'm not trying to nitpick, just trying to get it right in my mind. I said electronegativity explains, or let's say predicts, the behavior of gold and silver in HCl. The electronegativity of gold is higher than hydrogen, and it does not dissolve. The elect. of silver is lower than that of hydrogen, and silver DOES dissolve in HCl. Apparently it is the hard shell of AgCl that prevents any appreciable dissolution of Ag. On the other hand, the reactivity, to which you referred me, does not explain this same behavior. Since both gold and silver are shown lower than hydrogen on the activity series, this would predict that silver does NOT dissolve in HCl.
So maybe electronegativity just happens to predict what happens by chance, but by the same token reactivity does not predict what happens in this case, either. It seems that silver is the culprit, or "odd-man out." Do I have this right? According to activity, silver does not dissolve, but in reality it does (until it is blocked by AgCl). There must be a physical explanation for that, no?
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Sigh. Silver doesn't dissolve in HCl. It dissolves in HCl after adding an oxidizing agent. When we speak about dissolving in acid we speak about oxidizing a metal with H+, not with anything that you happen to add to the solution.
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Sigh. Silver doesn't dissolve in HCl. It dissolves in HCl after adding an oxidizing agent. When we speak about dissolving in acid we speak about oxidizing a metal with H+, not with anything that you happen to add to the solution.
No need to sigh. I understand the difference you talk about, and if you read my previous posts I defended you on that. However, I may be guilty of believing things I read on the internet. ::)
From this source: http://www.ehow.com/how_8254040_dissolve-silver.html I got the following:
Acids react with and dissolve most metals, but to achieve full dissolution, the resulting compounds must also exhibit solubility in water. Silver, for example, will dissolve in hydrochloric acid, or HCl, to form silver chloride, or AgCl. Silver chloride, however, is insoluble in water, which means a white solid of AgCl crystals will form in the resulting solution. The full dissolution of silver requires nitric acid, or HNO3, which reacts with silver to form silver nitrate, a water-soluble compound
This guy wrote the article so you'd think he'd have some knowledge. It wasn't just an anonymous person responding to a post. I could swear I found a post by a chemist saying about the same thing, but I can't seem to find it.
In googling more and more sources, it seems the vast majority say that silver does not dissolve in dilute HCl. Given that, I no longer see any discrepancy in the activity series for silver vs hydrogen.
I think this case is closed, amigo. Thanks for the refresher course.
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Can anyone help me explain what is going on here? I'm mostly curious about the color change. I washed some AgCl that had some green flakes in it in water, and a little ammonia. The solution turned dark blue. I assume this is due to copper sulfate or other copper salts. Can silver turn a solution blue? See photo 1.
Then I added HCl and the solution turned milky white/bluish. The milkiness I'm sure is caused by AgCl precipitating out. When the precip. settles , the solvent is clear. What causes the blue color to change to clear?
Thanks!
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Copper (Cu2+) is light blue in the solution (unless really concentrated). Plus ammonia it created a complex - dark blue tetraamminecopper. When you added strong acid ammonia got protonated, which is equivalent of removing the ammonia from the solution, and complexation reaction equilibrium shifted far to the left (reactants side) - so you got the Cu2+ back.
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OK, thanks!