# Looking for a chloride salt solubility chart



## ericrm (Apr 1, 2013)

Im looking for a chloride salt solubility chart that include température below freezing point of water. If it exist...


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## Palladium (Apr 1, 2013)

I seen one somewhere that showed the solubility of silver chloride in hcl and can't find it to save my life. Been looking for over a month. It seems that the numbers surprised me when i looked at them though. I to would like to know to.


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## kkmonte (Apr 1, 2013)

I found this one but it doesn't have AgCl listed..


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## butcher (Apr 1, 2013)

http://www.saltlakemetals.com/Solubility_Of_Silver_Compounds.htm
http://www.saltlakemetals.com/Solubility_Of_Silver_Chloride.htm


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## solar_plasma (Apr 2, 2013)

Such tables and graphs I am looking for,too. Does any of the chemists on the forum know a book that contains soluability data for all usual ion compounds? I know the soluablity product, but that is always given for a standard temperature. I know that the soluability is related to the enthalpy change of solution, since the soluability at higher temp. will raise, if the dissolving reaction is endotherm and will fall, if it is exotherm. Is it possible to *calculate *the soluability at any temperature by any function or can those data only be found empirically?

If you can give hints or keywords, I will try to find out by myself.


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## butcher (Apr 2, 2013)

I have two books I use:

Handbook of chemistry and physics (thirty first edition), editor in chief Charles Hodgman M.S.
Published by Chemical Rubber Publishing Co. 1948

Handbook of chemistry (Lange), compiled and edited by Norbert Adolf Lange, PH.D. 1946

GSP posted a link to these and some other books at a fair price some time ago, I purchased them then, they are filled with data of all types, some day I would like to go to a big book store, in the Big City and spend a day or two searching for more books like these.

GSP has made some very good references to books, and showed the forum some of his collection of books he had in his library, (try a search --Books and library--using GSP as author of the topic).


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## shaftsinkerawc (Apr 2, 2013)

Is the Kl in the upper left hand corner of the chart a misprint or what is it?


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## kkmonte (Apr 2, 2013)

shaftsinkerawc said:


> Is the Kl in the upper left hand corner of the chart a misprint or what is it?



Think its a KI (i) for potasium iodide.


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## ericrm (Apr 2, 2013)

What im mostly hopping is for below freezing point. Wikipédia is a very good start but they are all from 0deg celcius an over...


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## butcher (Apr 2, 2013)

Try looking through these:

http://books.google.com/books?id=H0IrAAAAYAAJ&pg=PA272&lpg=PA272&dq=solubility+of+silver+chloride+below+freezing+point&source=bl&ots=lrU1uNdsg0&sig=u64P7zdIbyTBNg4gjiYbz20-nbI&hl=en&sa=X&ei=2zJbUbqvIMLhiAKTy4H4Cg&sqi=2&ved=0CC0Q6AEwAA#v=onepage&q=solubility%20of%20silver%20chloride%20below%20freezing%20point&f=false

http://www.google.com/#hl=en&gs_rn=8&gs_ri=psy-ab&gs_mss=International%20Critical%20Tables%20of%20Numerical%20Data%2C%20Physics%20...%2C%20Volume&pq=solubility%20of%20silver%20chloride%20below%20freezing%20point&cp=70&gs_id=iu&xhr=t&q=International%20Critical%20Tables%20of%20Numerical%20Data%2C%20Physics%20...%2C%20Volume%206&es_nrs=true&pf=p&sclient=psy-ab&oq=International+Critical+Tables+of+Numerical+Data,+Physics+...,+Volume+6&gs_l=&pbx=1&bav=on.2,or.r_qf.&fp=4e273446afd35bd0&biw=939&bih=545


http://www.google.com/#hl=en&gs_rn=8&gs_ri=psy-ab&gs_mss=solubility%20of%20silver%20chloride%20be&pq=polyethylene%20chemical%20resistance&cp=50&gs_id=zo&xhr=t&q=solubility+of+silver+chloride+below+freezing+point&es_nrs=true&pf=p&sclient=psy-ab&oq=solubility+of+silver+chloride+below+freezing+point&gs_l=&pbx=1&bav=on.2,or.r_qf.&fp=4e273446afd35bd0&biw=939&bih=545

http://chla.library.cornell.edu/c/chla/browse/title/2944761.html


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## goldsilverpro (Apr 2, 2013)

I find these 2 handy. You can download them by clicking on the arrow next to the gear in the upper right corner.

http://books.google.com/books?id=ZnsMAQAAIAAJ&printsec=frontcover&dq=intitle:solubilities&hl=en&sa=X&ei=zDVbUdntFoOo2gWzuYC4CQ&ved=0CEgQ6AEwBA#v=onepage&q&f=false

http://books.google.com/books?id=uTmc-BeVbZoC&printsec=frontcover&dq=intitle:solubilities&hl=en&sa=X&ei=zDVbUdntFoOo2gWzuYC4CQ&ved=0CF8Q6AEwCA#v=onepage&q&f=false


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## FrugalRefiner (Apr 2, 2013)

I learn something new here every day. I never used the gear icon to initiate a download. I always mouse over the EBOOK-FREE button to the left and it drops down a menu, then mouse down to Download PDF and click. Now I know another way. :lol: 

Dave


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## butcher (Apr 2, 2013)

ericrm,
I have looked for the data for below freezing point but have not found it yet, I am beginning to wonder if it would even exist as I think water from solution would just ice up... 
I am unsure why you are looking for this.

AgCl silver chloride is very insoluble already, unless it is silver chloride of a higher oxidation state like AgCl2 which is slightly soluble in high acid high chloride mix (like concentrated aqua regia), but when diluted converts back to insoluble AgCl when diluted 3 to 4 parts water.

There is also what they call a common ion effect, basically if an fairly insoluble ionic salt like silver chloride has another ionic salt similar added to it like NaCl sodium chloride (common salt), the silver chloride is much more insoluble in the solution, noteur aqua regia is loaded with NaCl and other chloride salts which help with this common ion effect.

The freezing points change for a solution when ionic salts are added to a solution, changing the freezing point of the solution, more salts more change,(not sure how to say this) water freezes easily, salt water is almost impossible to freeze, I think the silver chloride in solution would be very hard to freeze, I do not know but I do not think you could freeze it easily, but you may be able to freeze out some of the water involved, pushing silver chloride to the side in less water to saturate (thus pushing some more silver chloride out of remaining solution, and if you got it that cold, you would only freeze the water involved, leaving less solution for the silver chloride to saturate, so in essence it would be similar to evaporating the solution to drive off water, so that it would hold less salt.

solubility silver chloride in 100g water,
@10 deg C = 8.9 x 10 (-10)
@20 deg C = 1.5x 10 (-4)
@50 deg C = 0.0005
@100 deg C = 0.002


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## ericrm (Apr 2, 2013)

Butcher and Gsp, thanks for the links. It is not much about silver chloride as i dilute and cool to drop the agcl, what i look for is to know if it is possible to precipitate out of hcl with cooling (lowering saturation point) enuf of salt(mixed chloride), to make it dirty but reusable again . Dealing that way would allow me to reduce my liquid waste exit.


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## butcher (Apr 3, 2013)

Basically depending on what process you use you may not be left with HCl that can be reused, or recovered.

Lets look what happens when we dissolve a metal in HCl;
First lets state what happens:
Metal + acid = salt of the metal

Example 
If the metals are higher than hydrogen in the reactivity series, Zinc for example, then hydrogen is gassed off from the reaction making a salt of the metal, zinc chloride ZnCl2

Zn + 2HCl --> ZnCl2 + H2 (gas)

We see here there is no HCl left, unless we added more acid than we had metal to react with then all we would do would make the solution more acidic, we could evaporate this salt to powders, if we heated the dry ZnCl2 powders strongly we could drive off chlorine gas making zinc oxide powders.

When we dissolve metal below hydrogen they are more noble, and need an oxidizer to dissolve in HCl acid, but the reaction is similar but hydrogen gas does not leave solution, here the hydrogen and oxygen normally forms water, as the metal becomes a chloride salt.
I believe the reaction would look similar to this: 
Cu + 2HCl + H2O2 --> CuO + 2HCl + H2O + O 
Then:
CuO + 2HCl + H2O + O --> CuCl2 +2H2O

So here we have copper chloride salt dissolved in water, note no acid left (unless we acidified solution with excess acid).
Again we could evaporate water making a copper chloride salt, if we kept and stored, if kept acidic the crystals could saved as CuCl2 crystals, if not enough acidity they would make CuCl losing a chlorine in the process, the CuCl2 could be stored slightly wet in a few drops of acid, and reused if mixed these with acid to dissolve more copper at a later date.

Aqua regia not only uses up the acids making gases, and dissolving metals to make salts of those metals (gold chloride AuCl3), after the process we not only have very little free HCl, but we actually add a precipitant that can will change the remaining solution to a different acid, or salt of metal, or mixture of these as we precipitate the gold:
Ferrous sulfate (copperas):
2AuCl4 + 6FeSO4 --> 2Au + 2Fe(SO4)3 + FeCl4
Here we are left with salts of iron.

Sodium metabisulfite:
SMB in acidic solution to produce sulfur dioxide gas:
3Na2S2O5 + 6HCl --> 6NaCl +3H2O + 6SO2
Then SO2 gas in the gold chloride solution
2AuCl3 + 3SO2 + 6H2O --> 2Au + 3H2SO4 + 6NaCl 
Here NaCl + H2SO4 could also react:
6NaCl + 3H2SO4 -->3 Na2SO4 + 6HCl


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## solar_plasma (Apr 3, 2013)

> Sodium metabisulfite:
> SMB in acidic solution to produce sulfur dioxide gas:
> 3Na2S2O5 + 6HCl --> 6NaCl +3H2O + 6SO2
> Then SO2 gas in the gold chloride solution
> ...



If you only would use SO2-gas for reduction, you would still have your HCl contaminated with H2SO4. The only easy way, I can see, to get rid of that, was resin. Here for example they have a resin, that is designed for taking NO3-ions an give Cl- AND as a "disadvantage" it also takes SO4--  
http://www.jat-service.de/Ionenaustauscher/entnitratisierung.html
But that is made for water, not for diluted acids.

But I am not so sure, that SO2 only generates SO4-- in that solution, sulfur is as tricky as tin is.


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## ericrm (Apr 3, 2013)

I usualy and volontary use too much acid and water because it remove uncertainty ( why did it stop reacting, i know it is not because i miss acid...) youve made à good point , if i canot reuse the acid i could simply évaporate it to salt, it is not à perfect solution but at least it will reduce the size of the waste.


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## Geraldo (Apr 3, 2013)

You most certainly can purify a solution by freezing. The first thing to solidify is water. This concentrates the salt solution, further depressing the freezing point, so more and more pure water gradually freezes out and the salt solution becomes more and more concentrated. 

Conceptually, it is an ancient technology - molten metal can be purified the same way (you can look at "zone refining").

In aqueous systems, there is a limit of course - how cold you can make your solution. A normal freezer is only a few degrees below 0 C.

A difficulty - you generally have to freeze it fairly slowly, so that salt solution doesn't become trapped in the ice mass. Also, the outside of the ice mass is contaminated with salt solution, and has to be rinsed off.

There are industrial plants using freeze purification for aqueous solutions (in hydrometallurgy). Some years ago there was a journal dedicated to the topic but I cannot find it right now.

It really comes down to how much solution you have, the local cost of energy and your willingness to design a proper system. Evaporation is often easier to design and can sometimes be more economical (it is very energy-inefficient to cool stuff off below freezing, compared to heating, even though the enthalpy of boiling is much higher than the enthalpy of freezing). If you have access to waste heat, or cobble together a solar system, you have essentially free evaporation potential, but freezing requires a lot more equipment and solar freezers (yes, they are a thing) are tricky and expensive to produce.

Just as an aside, I have concentrated acetic acid (vinegar) using the simple freeze purification approach - put it in a plastic container, stick it in the freezer, wait a few hours, skim off the ice on top, the liquid is concentrated acetic acid...eventually (might have to do it a few times depending on how concentrated you want the acid).

For general salt solution concentration, it is really easy to make an evaporator.

Just my thoughts.

Best Regards, Geraldo


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## butcher (Apr 4, 2013)

One way I evaporate the salt solution, which has already been treated for waste, is to put it in my ceramic casserole dish,
Now I am using the same heat to do another job also.

In this same casserole dish I have a canning jar sitting in the salt solution with values inside the jar, the values in solution are heated, to dissolve base metals or getting water rinses...

The solution to be evaporated is used for a water bath.
The salt water concentrates, once it concentrates down to a thicker solution, I replace it with fresh solution to be evaporated down, at the same time I am heating my solution in the jar I am working on, when I do this, I do not get the solution too hot, but enough to evaporate the salt water slowly.

Now this is the salt water solution that has already been treated for waste, almost all metals have been removed and it is mostly a sodium chloride solution, which can easily be evaporated down further by itself.

Another way I may reuse a solution and evaporate it down.
Many times I will re-use the waste acidic solutions to dissolve base metals before it is treated for waste, usually using it with heat, it will dissolve base metals fairly well usually becoming pretty much an Iron chloride, or iron sulfate, by the time I am done with it. 

Old copper chloride solution which I have reused so much that has pretty well become a ferric chloride solution / copper II chloride solution, with heat I can dissolve a lot of copper with it, heating it fairly strongly it dissolves the copper like mad, as it concentrates to almost a syrup, I will put it into a cooling and settling jar, where copper I chloride precipitates out as a white powder as the syrup cools, the cooled liquid left from this is added back to the heating pot to dissolve more copper, and more solution added, (waste old copper II chloride / Ferric chloride solution).
more of the original waste solution which I am trying to lower the volume of and reuse to dissolve more copper at the same time, is added to replace liquid level in the heated vessel, eventually I have reduced a couple of 5 gallon bucket of waste down to a little more than a gallon, which is then treated for waste, I will also be left with a jar full of impure copper I chloride salt, if gold foils are involved, the copper I chloride powders are save to re-acidify back into copper II chloride later date,after re-acidifying the copper I Chloride powers back into solution, it is allowed to settle, and copper II chloride solution is filtered, the copper II chloride solution can then be used, and gold foils are refined.


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## solar_plasma (Apr 4, 2013)

Nice!

I wonder, if it would be possible to get the basemetals back by electrolysis on graphite electrodes? The gases could maybe be caught with Ca(OH)2 in a partly sealed system...


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## butcher (Apr 4, 2013)

Electrolysis with inert anode is a good option for recovery of metals and possibly a way to get a useable solutions, Laser Steve had a useful method for removing copper from his nitrate electrolyte.
I have used something similar but mainly just to push the metals out of solution, but I have been thinking of playing with it more.

I think with divided cells and salt bridges (or ion membranes) and possible multiple chambers or even with something like two power supply's and two sets of electrode a cell can be set up to make acids from some of our solutions, and possibly do some work on our metals at the same time, from time to time I try to sit and think about this, but I am usually so busy with other things I have not had enough time to even think.

Also before attempting we need to understand the dangers of the gases that can many times evolve from electrolysis reactions,and this is also an option for making acids or solution is capturing these gases from electrolysis.

If the solutions are known what they contain, distillation is very handy way to capture the gases and remake a dilute acid, which is concentrated to azeotropic acid, at the same time distilling off the solution will concentrate the metal in solution.


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## solar_plasma (Apr 4, 2013)

yes....but what happens in an electrolysis of blended metal chlorides, with lead, nickel, copper,iron or worst case beryllium, barium, cadmium (though it may be the best choice to leave the last three in solution...)? And how is the reaction related to the voltage? Is it possible to use the voltage to fractionate the metaltype output? I dont want to waste peoples time by questions, I could work out and read about by myself, but I'd like to know, if this would be a way to go, before I read some books, which will blow my head for some weeks and cost some sanity points (cthulhu RPG insider) :lol:


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## butcher (Apr 4, 2013)

As I see it if you have multiple metals in solution you will plate out or push out (reduce) multiple metals at the cathode.

If only a couple of metals there can be some what of separation of metal from each other, if the metals are far enough apart from each other in the electromotive series of metals, voltage is controlled and one metal was much more predominant in solution, or better yet if only one metal was compatible with the electrolyte chosen, best only one metal involved in the solution with minor trace of impurity.


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## solar_plasma (Apr 4, 2013)

So, you can never be sure that the metal output can be melted without generation of highly toxic fumes....tricky


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## ericrm (Apr 4, 2013)

it is not exactly in the main subject but does anyone have tryed to bubble h2 in the waste to see if it could precipitate copper and lower? i have tryed a sample with no good result but i had so many variable that i cant realy be sure that it didnt work...


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## butcher (Apr 4, 2013)

I believe hydrogen is in the reactivity series of metals as a reference point (when dealing with acids) and in the electrochemical series to determine voltage, I do not see how it would reduce metals, but then again I learn something new everyday.

Basically I do not know but I do not think it would.


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## solar_plasma (Apr 5, 2013)

Only if you heat the metal compound to red glowing under a hydrogene atmosphere, it will be reduced. I've done this with CuO.

I don't know what will happen, if you are bubbling the H2 over a platinum electrode, since only then you can measure its 0V relative electronegativity.


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