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FrickinA
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[*] posted on 23-10-2010 at 23:07
KCl and HCl questions (New member here)


Hey everybody,
New amateur chemist here. I have always liked chemistry and I would like to start some synthesis and lab experiments of my own. I've got classroom lab experience and have taken 1.5 years of regular chem, so my experience is little but enough to get started. (I know basic lab safety and have worked with concentrated acids/bases before but nothing too serious).

I have 2 questions....
1) I'm planning on using KCl in various syntheses, and my main source is salt substitute (kinda pricy for the amount you get, anyone know a better source? KCl seems like a very important chemical to have for any synthesis of potassium salts). The stuff I have has several impurities, these are fumaric acid, tricalcium phosphate, and monocalcium phosphate. Based on some stoichiometric calculations using the nutrition info (1.2g total serving and contains 610mg K), this is 97% pure KCl, I assume this is pure enough for most lab use but just want to make sure with some experienced guys if these impurities need special considerations.

Second, I'm planning on making some hydrochloric acid using the reaction of sodium bisulfate and sodium chloride. I have an electric pH reader and I'm wondering if in the lab its okay to use the pH to determine molarity of the acid using the reverse log function, would this be a good way to avoid having to do a titration or is that a better method (I don't have any phenolphthalein on hand)?

I've got some good experience in "theoretical" chemistry, however these days classroom chemistry is rarely applied hands on in basic chem courses. Any help is appreciated.
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[*] posted on 24-10-2010 at 06:49


Quote:
anyone know a better source


Can't speak to the purity, but in the USA various big-box stores like Home Depot sell large bags as a water softener supply. Might well be better than the 97% in the salt substitute.

Quote:
I'm wondering if in the lab its okay to use the pH to determine molarity of the acid using the reverse log function


No. At the very least, it's not as simple as using the reverse log function, and rather than try to map out the exact relationship between pH and concentration for your solution you're better off just doing the titration. I expect it would be more precise anyway except maybe for very dilute strong acids and bases.
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[*] posted on 24-10-2010 at 07:57


Quote: Originally posted by FrickinA  

Second, I'm planning on making some hydrochloric acid using the reaction of sodium bisulfate and sodium chloride. I have an electric pH reader and I'm wondering if in the lab its okay to use the pH to determine molarity of the acid using the reverse log function, would this be a good way to avoid having to do a titration or is that a better method (I don't have any phenolphthalein on hand)?

I've got some good experience in "theoretical" chemistry, however these days classroom chemistry is rarely applied hands on in basic chem courses. Any help is appreciated.


The inverse pH method isn't perfect but more usable than bbart makes out, IMHO. In fact for quite dilute solutions, i.e. 0.1 M or less, it works quite well.

But I'm guessing that for your synthesis you'll need more concentrated solutions, at least 1 M (otherwise you'll be processing bucket loads of water to obtain the reaction product of your synthesis) and then simple pH measurements become unreliable because the relation between acid concentration and [H3O+] concentration no longer holds so theoretically.

Titration then becomes the real option. With a functioning pH meter you won't need indicators for end-point determination: just construct a 'volume added v. pH' titration curve, the end-point is unmistakable.

With the NaHSO4 - NaCl dry distillation method, getting really concentrated solutions of HCl is rather difficult. But hydrochloric acid is cheaply available as 'patio cleaner', 'toilet cleaner' and various products around the 20 w% mark from hardware store type outlets...


[Edited on 24-10-2010 by blogfast25]
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FrickinA
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[*] posted on 24-10-2010 at 08:38


Interesting, that's good to know about the pH meter. I've heard that the distillation process for HCl is complicated. Does anyone know how this would be done to up the concentration of an HCl solution? Im assuming boiling wouldn't work because that would drive HCl out of solution at the same time as boiling of the water?
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[*] posted on 24-10-2010 at 09:21


HCl-H2O form a constant boiling mixture - azeotrope - at standard pressure about 20,2% HCl or roughly 6 M. To get higher you need to mess with the pressure, which it's convenient and not all that effective, or soak up water in some fashion, or generate HCl gas and bubble that into water or the constant boiling acid. There are several threads in which making HCl gas were discussed, there's the standard NaCl+H2SO4, using CaCl2 to displace some of the HCl from hydrochloric acid (unfortunately this works best with more concentrated HCl), and others.

Note that constant boiling hydrochloric acid is sufficiently concentrated for many uses, and has the advantage that it doesn't "fume" with HCl gas sneaking out past the bottle cap and corroding all the nearby metal.
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[*] posted on 24-10-2010 at 09:22


Quote: Originally posted by FrickinA  
Interesting, that's good to know about the pH meter. I've heard that the distillation process for HCl is complicated. Does anyone know how this would be done to up the concentration of an HCl solution? Im assuming boiling wouldn't work because that would drive HCl out of solution at the same time as boiling of the water?
The distillation is "complicated" because HCl forms a constant boiling azeotrope with water. Solutions of HCl weaker than the azeotrope (about 20%) can be concentrated up to 20% by distillation because the water will distill off first. Then the vapor temperature rises and the azeotrope distills. But this is not really complicated, and can be quite useful. Google "constant boiling hydrochloric acid".
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[*] posted on 24-10-2010 at 09:30


HCl and water form what's called an azeotropic mixture, at about 20.2 w% HCl:

http://www.qvf.com/en/processsystems_3/Mineral%20Acids/Conce...

This means that as you distill a mixture of HCl/water of less than 20.2 % HCl, the liquid in your boiler will increase in HCl concentration until 20.2 % is reached: then that mixture boils over. So, yes, you can concentrate weak solutions of HCl by distillation but only to 20.2 w%.

Distilling HCl solutions isn't a particularly difficult thing to do, especially with those low boling points but HCl gas is of course both quite toxic and corrosive, not to mention very pungently smelling. You need a glass distillation kit for that and a way to deal with HCl gas, something like a NaOH wash bottle to lead your exhaust fumes through to capture any HCl. Otherwise you'll be choking on the HCl.

But there is a 'simpler' way of getting to reasonable concentrations and that's by repeating several NaHSO4/NaCl dry distillations and leading the HCl into the same solution multiple times. In theory about 40 % concentration can thus be obtained.

Considering how time consuming all this is, getting 5 l of 'patio cleaner' from your garden centre or hardware store is really the preferred 'method'...

Looks like entropy51 beat me to it! AND not_important! Sigh!


[Edited on 24-10-2010 by blogfast25]
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[*] posted on 24-10-2010 at 09:51


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[*] posted on 24-10-2010 at 10:29


Thats interesting, I had not considered the azeotrope before but it makes sense (same thing happens with 94%ish ethanol right), I had erroneously assumed HCl behaved like other dissolved gases and would bubble out of solution on heating. Thanks for the info guys. I will probably end up buying HCl when i need it for more synthesis but I find synthesis more satisfying when I make all the reagents I can :P
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[*] posted on 24-10-2010 at 10:52


I used NoSalt for KCl. It is HORRIBLE. I had to filter out tons of sticky gritty slime they used for anti-caking agents. If you don't have a vacuum filter setup I wouldn't bother.



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[*] posted on 24-10-2010 at 11:24


Quote: Originally posted by FrickinA  
Thats interesting, I had not considered the azeotrope before but it makes sense (same thing happens with 94%ish ethanol right), I had erroneously assumed HCl behaved like other dissolved gases and would bubble out of solution on heating. Thanks for the info guys. I will probably end up buying HCl when i need it for more synthesis but I find synthesis more satisfying when I make all the reagents I can :P


Gasses that are dissolved in solvents behave very much like any other solution. As it so happens most gasses are only sparingly soluble in water, HCl being one of the exceptions because it dissociates strongly when it dissolves. As a result the overwhelming majority of HCl in a HCl solution is present there as solvated H3O+ ions and solvated Cl- ions. Only a tiny minority is present as the diatomic molecule H-Cl... That improves solubility in water no end. In solvents where no dissociation can take place (alkanes, for instance) HCl solubility tends also to be piss poor, like most other gasses in water...

[Edited on 24-10-2010 by blogfast25]
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[*] posted on 24-10-2010 at 12:03


Quote: Originally posted by FrickinA  

1) I'm planning on using KCl in various syntheses, and my main source is salt substitute (kinda pricy for the amount you get, anyone know a better source?


If you live in an area that gets snow and ice you may find it as an ice melter. The best deal I ever found on KCl was 40 lbs for about $18. This was "Vaporizer" brand. The front of the bucket says "100% Potassium Chloride Granules" but the fine print on the back says 98% Potassium Chloride. [I suspect the remaining 2% is an anti-caking agent.]

I was able to purify this with a simple crystallization and the KClO3 I made from it produced a clean violet flame with no trace of any sodium yellow.




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[*] posted on 24-10-2010 at 13:17


Quote: Originally posted by NeutralIon  
Quote: Originally posted by FrickinA  

1) I'm planning on using KCl in various syntheses, and my main source is salt substitute (kinda pricy for the amount you get, anyone know a better source?


If you live in an area that gets snow and ice you may find it as an ice melter. The best deal I ever found on KCl was 40 lbs for about $18. This was "Vaporizer" brand. The front of the bucket says "100% Potassium Chloride Granules" but the fine print on the back says 98% Potassium Chloride. [I suspect the remaining 2% is an anti-caking agent.]

I was able to purify this with a simple crystallization and the KClO3 I made from it produced a clean violet flame with no trace of any sodium yellow.


Had you looked at the flame even with a primitive R-DVD spectroscope you would almost certainly still have seen the Na D line(s)... It would get feebler and feebler with each crystallisation. That's not to say that your product wasn't quite pure, it probably was. KCl isn't that easy to purify that way because it has quite a flat solubility-temperature profile... KClO3 should be easier.

It seems quite a waste to put KCl on ice when CaCl2 does such a good job...
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[*] posted on 24-10-2010 at 16:03


Quote: Originally posted by blogfast25  



It seems quite a waste to put KCl on ice when CaCl2 does such a good job...


KCl is not even a great ice melter, works only down to +20°F. But it is sold as safer for plants -- about the same time I found this as ice melter Agway was selling a small 5 lb box of KCl as a fertilizer for about $8 -- much more expensive and had so much filler that it was only about 70% KCl




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[*] posted on 24-10-2010 at 21:14



Quote:

I used NoSalt for KCl. It is HORRIBLE. I had to filter out tons of sticky gritty slime they used for anti-caking agents. If you don't have a vacuum filter setup I wouldn't bother.


The stuff I have is Morton Salt Substitute. It was about 3$ for 88.6g which seems a little pricey considering the apperant availability of the compound. It does contain tricalcium phosphate which is an anticaking agent and I would assume might be what gunked it up. I dont live in a snowy area but i might be able to find it as ice melter anyways.


Quote:

Gasses that are dissolved in solvents behave very much like any other solution. As it so happens most gasses are only sparingly soluble in water, HCl being one of the exceptions because it dissociates strongly when it dissolves. As a result the overwhelming majority of HCl in a HCl solution is present there as solvated H3O+ ions and solvated Cl- ions. Only a tiny minority is present as the diatomic molecule H-Cl... That improves solubility in water no end. In solvents where no dissociation can take place (alkanes, for instance) HCl solubility tends also to be piss poor, like most other gasses in water

Alright, that makes sense, didn't consider the complete ionization of the strong acid. Thanks for the info. So in theory I could make an HCl solution by NaHSO4 and NaCl reaction and bubble it through water then boil that until for awhile until it stops "boiling off" and then I would have about 20% HCl?
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[*] posted on 24-10-2010 at 22:13


I also purchased a can of "nosalt" with the idea that I'd burn off the tartaric, fumaric and adipic acids with manganese dioxide- but that made a bigger mess. I'll be looking for the ice melter now.

As for a K source for making salts, K2CO3 is available at the pottery supply and maybe the farm shop too. Metal carbonates are a good parent stock for making other salts- the C02 will evolve and leave the conveniently oxidized metal for the acid.

I also like to make some of my own reagents, but I'll let the economy of purchase weigh in on that decision- eg. make NH4VO3 but buy Ca(NO3)2.
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[*] posted on 26-10-2010 at 20:37


I just set up my whole lab in the garage and performed the HCl generation with NaHSO4 and NaCl. The process seemed to work okay (except im not sure my hoses are okay to use..., they seem to be off color now). I got acid with a pH of 2.6 and boiled it down to .9 then tried boiling it down but ended up with 1 drop. Except the last drop the strength was too weak for visible neutralization with sodium bicarbonate. I think i can improve my process but it seems like you guys said... this is a fairly inefficient way of making HCl and im better off buying it.
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[*] posted on 27-10-2010 at 04:53


Quote: Originally posted by FrickinA  
I just set up my whole lab in the garage and performed the HCl generation with NaHSO4 and NaCl. The process seemed to work okay (except im not sure my hoses are okay to use..., they seem to be off color now). I got acid with a pH of 2.6 and boiled it down to .9 then tried boiling it down but ended up with 1 drop. Except the last drop the strength was too weak for visible neutralization with sodium bicarbonate. I think i can improve my process but it seems like you guys said... this is a fairly inefficient way of making HCl and im better off buying it.


If you used silicone tubing, it goes from opaque to white very easily: you're turning the silicone into silica through hydrolysis. pPVC should be OK. Neoprene tooish. Been there, done that...
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[*] posted on 27-10-2010 at 06:29


Quote: Originally posted by FrickinA  
I got acid with a pH of 2.6 and boiled it down to .9 then tried boiling it down but ended up with 1 drop. Except the last drop the strength was too weak for visible neutralization with sodium bicarbonate.
That's because once the concentration is raised to 20% by boiling, the 20% HCl boils off.

You should distill it so that you know when the vapor temperature rises and the 20% concentration has been reached.

Did you by any chance calculate the expected concentration based on the amount of NaCl used and the volume of water in which you dissolved the HCl(g) ? If you dissolve sufficient HCl(g) in the water to reach 20%, you can distill good 20% HCl right out.

Admittedly it's easier to buy muriatic acid, but this is a nice experiment in terms of learning chemistry as well as technique. This is especially true if you titrate the acid and follow the HCl concentration as the process proceeds.
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[*] posted on 27-10-2010 at 14:03


The tubing was vinyl. If i distilled it how would i know when to put the HCl mixture through the condenser, IE how do i ensure I am not distilling excess water at any given point to ensure the max concentration of HCl
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