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K12Chemistry
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Making Hydrochloric Acid
Hi,
I want to make hydrochloric acid by mixing sulfuric acid and sodium chloride but I only have 64% sulfuric acid(titrated by me). Do I HAVE to
concentrate the acid or can I use the dilute version as long as I get the right stoichometry.
Can someone tell me the stoichometry because I have no Idea how to calculate it. I have 250ml of 64% sulfuric acid. How much salt?
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elementcollector1
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You don't have to concentrate the acid, but yes, stoichiometry is heavily recommended.
http://theodoregray.com/PeriodicTable/MSP/BalanceReactions
Enter the equation, with HCl and K2SO4 as the products.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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Diablo
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I don't know about the concentration but I can help with the stoichiometry.
First 0.64 * 250=160 ml of acid The reaction NaCl + H2SO4-> HCl + NaHSO4 Now we multiply the volume of acid by its density of 1.84 to get
160*1.84=294.4 grams acid. Molar mass of sulfuric acid is 98 grams per mole 294.4/98= about 3 moles so 3 moles of nacl=58.44 grams per mole 58.44 * 3
=175.32 grams sodium chloride.
Edit: what he said about the concentration.
[Edited on 2-15-2013 by Diablo]
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K12Chemistry
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thank you diablo and elementcollector,
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blogfast25
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Actually, it really works well only with at least 95 % H2SO4.
If you can't get that (Draino!), try anhydrous sodium bisulphate (NaHSO4, aka sodium hydrogen sulphate): fusing NaCl with NaHSO4 gives you dry HCl:
NaCl + NaHSO4 === > Na2SO4 + HCl
But some fumes of SO3 are likely to come over too, if you overdo it on the temperature...
[Edited on 15-2-2013 by blogfast25]
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Adas
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Quote: Originally posted by blogfast25  | Actually, it really works well only with at least 95 % H2SO4.
If you can't get that (Draino!), try anhydrous sodium bisulphate (NaHSO4, aka sodium hydrogen sulphate): fusing NaCl with NaHSO4 gives you dry HCl:
NaCl + NaHSO4 === > Na2SO4 + HCl
But some fumes of SO3 are likely to come over too, if you overdo it on the temperature...
[Edited on 15-2-2013 by blogfast25] |
I tried this with normal table salt, and the salt did carbonize, lol. It must have been contaminated with organics and the SO3 dehydrated them. BIG
FAIL  but at the same time a good experience. Now I know that table salt is not
just NaCl...
Rest In Pieces!
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DraconicAcid
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| Quote: Originally posted by Adas | | Quote: Originally posted by blogfast25 | Actually, it really works well only with at least 95 % H2SO4.
If you can't get that (Draino!), try anhydrous sodium bisulphate (NaHSO4, aka sodium hydrogen sulphate): fusing NaCl with NaHSO4 gives you dry HCl:
NaCl + NaHSO4 === > Na2SO4 + HCl
But some fumes of SO3 are likely to come over too, if you overdo it on the temperature...
[Edited on 15-2-2013 by blogfast25] |
I tried this with normal table salt, and the salt did carbonize, lol. It must have been contaminated with organics and the SO3 dehydrated them. BIG
FAIL but at the same time a good experience. Now I know that table salt is not
just NaCl... |
I'm surprised that table salt would contain enough other stuff to carbonize with sulphuric acid. You might get iodine from the iodized salt...
Next time, recrystallize the table salt first. That should get rid of any organics (along with the iodides), and give purer HCl.
You didn't take it right out of the salt shaker, did you? Sometimes people add rice to keep it from clumping. Or maybe you accidentally grabbed the
sugar container?
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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blogfast25
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Table salt is often adulterated with anti-caking agents, driers etc. Use lab grade or with kitchen grade, dissolve, filter and recrystallise as pure
NaCl...
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Endimion17
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holy shit
http://www.sciencemadness.org/talk/search.php?token=&srchtxt=hydrochloric+acid&srchfield=subject&srchuname=&f[]=all&srchfrom=0&
;filter_distinct=yes&searchsubmit=Search
[Edited on 15-2-2013 by Endimion17]
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K12Chemistry
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I am going to try it with the dilute sulfuric acid anyway. Why wouldn't it work? Also how much water should I dissolve the HCl gas in to get
concentrated hydrochloric acid.
Is it worth boiling the sulfuric acid to concentrate and then making HCl or should I just use dilute?
thanks
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Vargouille
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If you try it with dilute sulfuric acid, you might get the 20.2% HCl/H2O azeotrope. Higher concentrations of sulfuric acid will result in higher
concentrations of HCl coming over. If you calculate it just right, you can get 37% coming over at somewhere between 61 and 48C. You can check the
physical properties section of the Wiki page for hydrochloric acid for more information.
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AJKOER
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Preparation of HCl
In place of H2SO4, consider using Oxalic acid, H2C2O4. React with any convenient chloride, as for the most part, there are all not very soluble. Also,
can use solid chlorides (like NaCl per Watt's) to produce HCl gas.
---------------------------------------
Don't have pure H2SO4 or H2C2O4, generate Chlorine gas (at least one old thread on paths). To a vessel containing the Cl2, add a little cold distilled
H2O and let seat in direct sunlight. Reactions:
Cl2 (g) + H2O <--> HOCl + HCl
2 HOCl --uv--> 2 HCl + O2 (g)
Note, in a closed vessel, the Oxygen formation will increase pressure. An idea I had, insert an open beaker in the vessel containing the chlorine
water, with say aqueous FeCl2 (conveniently but slowly prepared by the action of H2CO3 on NaClO to which Fe is added forming a pale green solution).
The O2 will increase pressure moving the 1st reaction to the right. Slowly over days, however, the O2 will react with the aqueous FeCl2 (changing
color of the solution) removing O2 from the system. The reaction (see "Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K.
Mukherjee at
http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fiWLs05y8f&am
p;sig=mi-pV94woVj7JABKBBzLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=onepage&q=FeCl2%20%2B%20O2%20%3
D%20FeO(OH)%20%2B%20FeCl3&f=false ):
6 FeCl2 + 3/2 O2 + H2O --> 2 FeO(OH) + 4 FeCl3
and, upon raising the pH of the FeCl3 (aq), say upon dilution, the same source (equations 87 and 88) gives:
4 FeCl3 + 8 H2O --> 4 FeO(OH) + 12 HCl
FeCl3 + 3 H2O --> Fe(OH)3 + 3 HCl
so dilute HCl can be recovered/prepared. Note, a potential complication of the FeCl2 synthesis occurs if the Chlorine Bleach product contains Na2CO3
in excess of the NaOCl.
[Edited on 16-2-2013 by AJKOER]
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mr.crow
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Can't you get HCl from the hardware store?
The azeotrope of HCl and water is 20.2% at 110 deg C. Water and weak HCl will distill over until this temperature is reached. Your weak acid should be
fine if used in excess.
You need a proper distillation aparatus. HCl is very unpleasant for something so common. I haven't done it myself so I can only give general advice.
If you have high school chemistry education you already know enough about stoichiometry calculations. PM me if you need any more help 
Double, double toil and trouble; Fire burn, and caldron bubble
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K12Chemistry
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If I do get weak HCl then can I concentrate it?
I also tried to boil down my sulfuric acid but when I added it to sugar nothing happened. Do I need to boil more?
In NurdRage's method he uses 30ml of sulfuric acid. So even though my sulfuric acid will be about 98% at 100ml, I am going to boil to 30ml to
GUARANTEE that I will get conc. sulfuric acid and I will have enough to make HCl
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AJKOER
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Actually and surprisingly, you can make dilute HCl act like concentrated hydrochloric acid using a chloride system (and no, I am not on drugs). See
Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K. Mukherjee, page 15 at
http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fiWLs05y8f&am
p;sig=mi-pV94woVj7JABKBB
zLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=snippet&q=Magnesium%20chloride%20MgCl2&f=false .
The trick is used by hydrometallurgists as dissolving metal ore with conc acid is expensive. The author says data confirms that a 2M HCl in 3M CaCl2
or MgCl2 (or FeCl3) behaves like 7M HCl.
So you may wish to try this in place of concentrating depending on the application. Note, there is an apparent more limited improvement in the
presence of a monovalent chloride (like NaCl).
--------------------------------
The apparent increase in ionic strength of HCl in the presence of divalent chloride may be an explanation of why Cl2 is observed to be produced by the
action of CO2 on moist Bleaching powder (Ca(OCl)2 and CaCl2). The action of CO2 in a small amount of water on Ca(OCl)2 could form unstable HOCl (in
the conc range of 30% to 50% or more). Some of Hypochlorous could decompose/disproportionate, especially in light, to HCl, O2 and perhaps, a small
amount of HClO3, as well. The HCl is assumed to be, in the presence of the divalent chloride, active enough to drive the reaction HCl + HOCl to Cl2
(g) and H2O, which in the presence of CO2...and the reaction chain continues, well at least in my speculation. Note, this Chlorine production (in
accordance with the mechanics of the chloride system) is not similarly observed with NaClO or KClO (where more conc HCl is required to liberate
Chlorine).
[Edited on 17-2-2013 by AJKOER]
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blogfast25
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| Quote: Originally posted by AJKOER |
So you may wish to try this in place of concentrating depending on the application. Note, there is an apparent more limited improvement in the
presence of a monovalent chloride (like NaCl).
[Edited on 16-2-2013 by AJKOER] |
And contaminating everything you use it for with Ca, Mg or Fe(III)? No but thanks!
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AJKOER
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But Blogfast, if you have to choose between using Muriatic acid and a dilute, but otherwise pure, HCl with an introduced Calcium impurity (from CaCl2)
not expected to be overly problematic for the synthesis in question, which would you select?
Also, if acid strength has to be increased to generate a gas (Cl2, ClO2, SO2, H2S...) does a Ca impurity in the HCl matter?
What if the reaction mixture already has a metal X salt, and there exist a not monovalent soluble chloride of metal X, which you have or can easily
prepare, would you consider a chloride system to increase acid strength?
What if you have to prepare a certain quantity quickly, but do not have sufficient HCl, but could employ a chloride system, what would you do?
Perhaps yes to some of the questions. But, after all, it is just an option, not a mandate.
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Poppy
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| Quote: Originally posted by blogfast25 | Actually, it really works well only with at least 95 % H2SO4.
If you can't get that (Draino!), try anhydrous sodium bisulphate (NaHSO4, aka sodium hydrogen sulphate): fusing NaCl with NaHSO4 gives you dry HCl:
NaCl + NaHSO4 === > Na2SO4 + HCl
But some fumes of SO3 are likely to come over too, if you overdo it on the temperature...
[Edited on 15-2-2013 by blogfast25] |
Blogfast25, as for that retort which HCl come by, is it stainless steel pan enabled?
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zed
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Use rock salt.
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repo1030
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Just use kosher or popcorn salt. This should be pure NaCl.
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blogfast25
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Huh? Please rephrase, I don't know what you mean.
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Poppy
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Errh, sorry Blogfast.
Is this set of NaHSO4 and NaCl you proposed enabled to be produced with stainless steel apparatus? Or would it corrode the pan so bad? As far as it
look the pH would be buffed! 
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elementcollector1
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Glass works fine.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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| Quote: Originally posted by Poppy | Errh, sorry Blogfast.
Is this set of NaHSO4 and NaCl you proposed enabled to be produced with stainless steel apparatus? Or would it corrode the pan so bad? As far as it
look the pH would be buffed! |
As EC1 said: glass is fine. SS will be attacked and some FeCl3 may come wafting over with the HCl. Having said that, H2SO4 + NaCl + heat in cast iron
kettles is how they used to produce HCl industrially. And it was a bit yellow because of the FeCl3 in it...
[Edited on 19-2-2013 by blogfast25]
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AJKOER
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| Quote: Originally posted by AJKOER |
Actually and surprisingly, you can make dilute HCl act like concentrated hydrochloric acid using a chloride system (and no, I am not on drugs). See
Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K. Mukherjee, page 15 at
http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fiWLs05y8f&am
p;sig=mi-pV94woVj7JABKBB
zLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=snippet&q=Magnesium%20chloride%20MgCl2&f=false .
The trick is used by hydrometallurgists as dissolving metal ore with conc acid is expensive. The author says data confirms that a 2M HCl in 3M CaCl2
or MgCl2 (or FeCl3) behaves like 7M HCl.
So you may wish to try this in place of concentrating depending on the application. Note, there is an apparent more limited improvement in the
presence of a monovalent chloride (like NaCl).
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However, if you still wish to concentrate HCl via distillation, apparently even a very dilute HCl solutions will work if you know how. Basically add
MgCl2 and if very dilute, a little FeCl3. For more details see http://www.neoferric.ca/documents/Harris%20et%20al%20Iron%20... . To quote:
"Figure 4 shows the OLI simulation of the pure HCl-H2O system (it should be noted
that virtually identical values are obtained from the Chemical Engineers Handbook [35] and
the commercially available modelling software AspenPlus® [36]). The data predict that it is
not possible to distill HCl from dilute solutions (i.e. solutions below the concentration of the
azeotrope, 20% HCl). On the other hand, when magnesium chloride is added, the situation
changes quite dramatically as shown in Figure 5. In this case, it is clear that it should be
possible to distill HCl at any liquid phase concentration from magnesium chloride solutions
of concentration greater than about 2.5 m. When ferric chloride is also present, as shown in
Figure 6, then it is apparent that at high concentrations of magnesium chloride, relatively
concentrated HCl should be recovered even from brine solutions containing very little free
HCl."
Also:
"The results (Figure 7) confirmed the OLI simulation, showing
that in strong magnesium chloride solution, the free HCl was flashed off very quickly at
close to the boiling point. Interestingly, the initial concentration in the distillate was close to
36% HCl, and overall, approximately 80% of the free HCl in the solution was recovered at a
cumulative concentration in excess of the azeotrope."
[Edited on 22-2-2013 by AJKOER]
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