Sciencemadness Discussion Board - Concentrating hydrochloric acid

Concentrating hydrochloric acid

xfusion44 - 13-10-2014 at 17:42

Hi,

I would like to know, what happens if you boil 20% HCl? Does concentration drop, or does it first drop and then rise? As far as I know, you can't concentrate HCl with just boiling water off, but 20% HCl has a boiling point of 108°C, so what happens then? Does it start to loose water or does it start loosing HCl gas, or both? I'm confused. Also 10% HCl has 103°C boiling point, but 30+% has lower boiling point than water, isn't that weird? So, can I concentrate it by boiling or not? If not, does that also hold true, if I use 31% HCl instead of 20%?

Thanks for your help

[Edited on 14-10-2014 by xfusion44]

elementcollector1 - 13-10-2014 at 18:10

If you boil 20% HCl, it boils off with the water as an azeotrope. Past that 20% threshold, heating the acid will release HCl gas.

subsecret - 13-10-2014 at 20:40

You could concentrate it by heating it and leading the gas into chilled water. If you can get your hands on HCl that is stronger than 20%, you could use this method to produce concentrated HCl, and you'd be left with the 20% stuff, which is still very useful.

xfusion44 - 14-10-2014 at 05:26

Thanks!

So, I could use 31% HCl, boil it and lead gas into water, to produce 38% HCl? But How do I know, when to stop? By measuring temperature? And how much chilled water should I use, to get 38% concentration with certain amount of 31% HCl?

Thx

PS: How about adding some concentrated H2SO4 to it and distilling it?

[Edited on 14-10-2014 by xfusion44]

subsecret - 14-10-2014 at 18:24

You wouldn't want to boil the acid too strongly, as this would put considerable water into the clean material.

If you're working outside or in a fume hood, you could use a graduated cylinder (extra height for efficiency) on a scale, and the mass increase would tell you how much HCl gas has been absorbed.

If you used a closed absorber, you could put a bubbler on the output end, and wait for the bubbling here to speed up, which would indicate saturation. The accuracy of this is entirely dependent on the efficiency of the absorber.

The NaCl / H₂SO₄ method would allow you to produce as much HCl as you wanted... If you live in the US (I believe this because you referred to 31% HCl), hopefully you can find large amounts of sulfuric acid in hardware stores.

[Edited on 15-10-2014 by Awesomeness]

xfusion44 - 15-10-2014 at 12:30

@Awesomeness

No, I live in Europe. Actually it says 30-32% on acid bottle, so I said 31

May I ask, why NaCl?

And thanks for your suggestions

PS: I can get acid for lead battery, but I need to concentrate it then... I know it's not worth spending so much conc. H2SO4 just to make conc. HCl, but I just want to know, how can be done, just to have a sample of 38% HCl, maybe it comes handy some time.

[Edited on 15-10-2014 by xfusion44]

woelen - 17-10-2014 at 00:28

If you have 30% HCl, then you can indeed increase its concentration fairly easily.

Take two portions of 30% HCl of equal volume. Chill one of the two portions (put the beaker with 30% HCl in a tank, filled with ice and/or snow).
To the other portion, add calcium chloride (you can purchase this from a hardware store as drying material) and heat gently. Copious amounts of HCl will be released. Lead these through the chilled HCl and in this way you can easily increase the concentration to 37%. Be careful with this reaction, the liquid may start foaming and then you may get contaminated liquid in your 37% HCl. Especially if the CaCl2 used in not really pure, then it may foam easily. I have seen CaCl2 which was pale yellow/brown instead of pure white and this is terrible. Also avoid the perfumed drying agent, just buy the cheapest pure white flakes.

The portion to which the CaCl2 is added can also be recovered, by distilling off most of the liquid. What remains is a concentrated solution of CaCl2 with a small amount of HCl left. This can be flushed down the drain with a lot of water.

Instead of CaCl2 you can also use conc. H2SO4, but using valuable conc. H2SO4 for such a thing is not what I like. But if you go for ease of reaction, and do not mind using conc. H2SO4 for this purpose, then you can use that as well. Simply carefully drip the H2SO4 in the 30% HCl and lead the produced gas through the chilled 30% HCl.

Even simpler, using no additional chemicals at all, is heating the 30% HCl and leading the produced gas though the other (chilled) portion. In this way, however, you will have more difficulty producing HCl gas, soon the 30% HCl hardly produces any free gas and stronger heating is required.

Be careful and prevent suckback of your chilled HCl into the HCl gas generating bottle! Use of an small intermediate bottle reduces the risk of suckback.

[Edited on 17-10-14 by woelen]

xfusion44 - 17-10-2014 at 13:19

@woelen

Thank you!

I'm sure, I will try it with one of the methods you mentioned.

I have another question: can I also use CaCl2/ice bath, to reach temperatures of -40°C? I always wanted to buy CaCl2 to try this out. That way I could probably freeze HCl in one beaker

Do you think It's possible to make liquid chlorine with CaCl2/ice bath?

Thanks again!

PS: sorry for offtopic with CaCl2/ice bath.

[Edited on 17-10-2014 by xfusion44]

subsecret - 17-10-2014 at 14:02

Quote: Originally posted by xfusion44

@Awesomeness

No, I live in Europe. Actually it says 30-32% on acid bottle, so I said 31

...

[Edited on 15-10-2014 by xfusion44]

Pardon my insensitivity.

xfusion44 - 17-10-2014 at 17:37

Quote: Originally posted by Awesomeness

Quote: Originally posted by xfusion44

@Awesomeness

No, I live in Europe. Actually it says 30-32% on acid bottle, so I said 31

...

[Edited on 15-10-2014 by xfusion44]

Pardon my insensitivity.

No, no, everything's fine

It's obvious, that you couldn't know for sure, where I live, just by knowing HCl concentration

DrMario - 18-10-2014 at 07:14

Quote: Originally posted by woelen

If you have 30% HCl, then you can indeed increase its concentration fairly easily.

Take two portions of 30% HCl of equal volume. Chill one of the two portions (put the beaker with 30% HCl in a tank, filled with ice and/or snow).
To the other portion, add calcium chloride (you can purchase this from a hardware store as drying material) and heat gently. Copious amounts of HCl will be released. Lead these through the chilled HCl and in this way you can easily increase the concentration to 37%. Be careful with this reaction, the liquid may start foaming and then you may get contaminated liquid in your 37% HCl. Especially if the CaCl2 used in not really pure, then it may foam easily. I have seen CaCl2 which was pale yellow/brown instead of pure white and this is terrible. Also avoid the perfumed drying agent, just buy the cheapest pure white flakes.

The portion to which the CaCl2 is added can also be recovered, by distilling off most of the liquid. What remains is a concentrated solution of CaCl2 with a small amount of HCl left. This can be flushed down the drain with a lot of water.

[Edited on 17-10-14 by woelen]

Wouldn't this method (adding CaCl2) work even with the azeotropic solution of HCl? Adding CaCl2 to obtain a 40% solution in water, increases the boiling point to 120C, which is above the 20.2% HCl azeotrope's boiling point of 110C.

pneumatician - 21-10-2014 at 13:52

with all this tech at last waht MAXIMUM concentra. of hcl people can obtain???

45%? 55%?...

what is the MAXIMUM concentration of HCL is obtained by the oficial science stablisment???

what can be done with hcl at 80%? what industrial applications, what science aplications? if anyone get hcl at 80% or more win the nobel?

DrMario - 21-10-2014 at 14:14

Quote: Originally posted by pneumatician

with all this tech at last waht MAXIMUM concentra. of hcl people can obtain???

45%? 55%?...

what is the MAXIMUM concentration of HCL is obtained by the oficial science stablisment???

what can be done with hcl at 80%? what industrial applications, what science aplications? if anyone get hcl at 80% or more win the nobel?

Normally 37 wt% which corresponds to a 12M solution, and this is the concentration most labs have in storage - then they can dilute it with DI water to the desired molarity.

[Edited on 21-10-2014 by DrMario]

Famousroger - 21-10-2014 at 14:19

Hcl is a gas; concentrating above this;
Solubility in water
823 g/L (0 °C)
720 g/L (20 °C)
561 g/L (60 °C)

Will result in a water mixture that will fume hydrogen chloride... 31% "muriatic" will fume, but 20% largely will not...if you need hcl at higher concentrations than ~36%, one likely ought use hydrogen chloride, and not aqeous hcl...(as indicated in above, solubility is inversely effected by increase in temp.). You could super saturate the hcl/water mixture,but this wouldn't be worthy of a Nobel prize...

pneumatician - 21-10-2014 at 14:22

DrMario -> Normally 37 wt%, and this is the concentration most labs have in storage - then they can dilute it with DI water to the desired molarity.

oh yes! this is the maximum % in catalogs of companys selling chemicals for labs, BUT I are asking for a bit more... i want if someone have at hand references about hcl at very big concentration! (of course not pure gas) (of course very fuming!!!)

Famousroger - 21-10-2014 at 14:22

NaCl and h2SO4 (~95%) will produce hydrogen chloride gas. Using this gas requires apropos experience and equipment...

Famousroger - 21-10-2014 at 14:30

You would have to essentially pressurize the water in order for it to hold more hcl. It's a gas/liquid solution, so concentration will drop whenever muriatic is at higher concentration and in an environment where the hcl can escape - you would have to create a container capable of withstanding the pressure required to keep the gas solvated in the water...you could generate pressurized hydrogen chloride that feeds into an airtight flask of concentrated muriatic...but the success of this will be...likely lackluster...depending on pressure and purity of hydrogen chloride...if it reaches high pressure and purity and if fed into say 37% HCl, you could super saturate it...it would be a serious danger though if you were successful. Opening such an enclosure would shoot hcl gas likely quickly and dangerously.

pneumatician - 21-10-2014 at 14:30

Famousroger -> You could super saturate the hcl/water mixture,but this wouldn't be worthy of a Nobel prize...

I already seemed to me !!

[Edited on 21-10-2014 by pneumatician]

pneumatician - 21-10-2014 at 14:44

Famousroger -> You would have to essentially pressurize the water in order for it to hold more hcl. It's a gas/liquid solution, so concentration will drop whenever muriatic is at higher concentration and in an environment where...

ok this is all obviously. so the maximum handable in a normal lab is around 37% maybe with a maximum of 60%

so people in lab are playing with far more dangerous chems. I think in the correct lab and hands the hcl can be more strong, no?

so can this hcl at for exemple 80% disolve 24k gold?

DrMario - 22-10-2014 at 08:09

Quote: Originally posted by pneumatician

Famousroger -> You would have to essentially pressurize the water in order for it to hold more hcl. It's a gas/liquid solution, so concentration will drop whenever muriatic is at higher concentration and in an environment where...

ok this is all obviously. so the maximum handable in a normal lab is around 37% maybe with a maximum of 60%

so people in lab are playing with far more dangerous chems. I think in the correct lab and hands the hcl can be more strong, no?

so can this hcl at for exemple 80% disolve 24k gold?

I think you misunderstand: 38% is the theoretical maximum wt% HCl that exists - it is the saturated solution and you cannot exceed this under normal conditions. As far as I know, Sigma Aldrich only sells 37%, but that's very close to the absolute maximum concentration. You cannot create a 80% HCl solution under normal conditions (normal pressure and temperature). There do exist some exotic systems with special surfactants etc., in which a higher concentration is possible, but the HCl is not dissociated in water, in those cases. And the water is not pure, of course.

AJKOER - 22-10-2014 at 08:56

Note, there are possible paths to defeating the azeotrope. For example, Wikipedia (see http://en.wikipedia.org/wiki/Azeotrope ) describes several methods including pressure swing distillation, azeotropic distillation involving the introduction of an an entrainer (an additional agent) impacting the volatility of one of the azeotrope constituents, chemical action separation where an entrainer is added also having a strong chemical affinity for one of the constituents, distillation using a dissolved salt where a salt is dissolved in a solvent to alter its relative boiling point, extractive distillation (similar to azeotropic distillation except that the entrainer is less volatile than any of the azeotrope's constituents) to mention a few.

In the case of dilute HCl where the addition of a salt is not problematic for a particular application, one may consider the addition of anhydrous calcium chloride (or, a concentrated solution thereof) as a possibility. The reason relates to the apparent significant increase in the so called 'activity level' upon adding MgCl2 or CaCl2 or NaCl (in declining order of preference). Here is a real world reference relating to practical significance in the field of Hydrometallury where leaching out minerals from ores efficiently and cheaply is a major concern. Source: 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=PA15 .

Note, the author claims there is data confirming that a 2M HCl in 3M MgCl2 or CaCl2 (and also FeCl3) behaves like 7M HCl!

Here is a quote on the matter of discussion in one of the reference sources previously provided by Bfesser on Thermodynamic Activity (see http://en.wikipedia.org/wiki/Thermodynamic_activity ):

"When a 0.1 M hydrochloric acid solution containing methyl green indicator is added to a 5 M solution of magnesium chloride, the color of the indicator changes from green to yellow—indicating increasing acidity—when in fact the acid has been diluted. Although at low ionic strength (<0.1 M) the activity coefficient approaches unity, this coefficient can actually increase with ionic strength in a high ionic strength regime. For hydrochloric acid solutions, the minimum is around 0.4 M.[1]"

[Edited on 22-10-2014 by AJKOER]

xfusion44 - 22-10-2014 at 17:44

Quote: Originally posted by pneumatician

37 or 38% is the point where you can't concentrate HCl solution anymore, because water will simply not dissolve anymore HCl gas. Note, that HCl is a GAS, not a LIQUID. HCl acid is liquid (water with HCl gas dissolved in it). And at 38% wt. water becomes saturated and it cannot accept anymore HCl gas at normal conditions T=20°C, p=100kPa. If you need higher concentration of HCl gas, you must use gas itself, not hydrochloric acid. Gold will not dissolve in HCl, but it will however react with aqua regia (mixture of HCl and HNO3).

pneumatician - 23-10-2014 at 07:03

AJKOER you are near the truth!! in reality (in official teorie) high con. hcl is more easy than anyone can think!! of course very fumming!! but for this bottles have a stopper!!!

hcl at 37% also is fumming!

so the real hcl is around 20% when the "surface tension of water" is not break by the force or pressure of the gas, more than 20% hcl gas in water is gas with water and - is water with gas. (brutal note)

blogfast25 - 23-10-2014 at 07:29

Quote: Originally posted by pneumatician

so the real hcl is around 20% when the "surface tension of water" is not break by the force or pressure of the gas, more than 20% hcl gas in water is gas with water and - is water with gas. (brutal note)

Hmm... do you know that HCl in water dissociates almost 100 % acc.:

HCl(aq) + H2O(l) === > H3O+(aq) + Cl-(aq)

Even saturated hydrochloric acid contains almost no free, undissociated HCl(aq).

fdsailor - 29-11-2015 at 21:13

A curious question - when concentrating HCL by dissolving the dry gas in water, bubbling is usually used as best I can tell in reading the posts, and of course this means you have to plan for/against suckback. On the other hand, I understand HCL to, like ammonia, be VERY soluable in water - so. for preparation of smaller amounts, why can't you simply pass the gas OVER water , hence averting suckback? Is this one of those things that initially works fine, but then as the solution approaches saturation your yield just goes way down?

byko3y - 30-11-2015 at 00:23

fdsailor, I have no idea where you've found the stupid guides with bubbling HCl through water, but the right way is to always have the inlet tube above surface of receiving solution.

Deathunter88 - 30-11-2015 at 06:23

Quote: Originally posted by byko3y

fdsailor, I have no idea where you've found the stupid guides with bubbling HCl through water, but the right way is to always have the inlet tube above surface of receiving solution.

Um, IDK about you but all the sources I have read suggest bubbling the HCl into the water as it makes the process more efficient...Just be sure to plan for suck back by having a second flask.

annaandherdad - 30-11-2015 at 09:57

Use an inverted funnel, placed just under the surface of the water, to avoid suck-back. This is an old trick and it works very well, also with ammonia.
Design it so that as the water starts to rise in the funnel, it lowers the level of the water outside the funnel until it breaks loose and lets air in. In my experience you get very little of the soluble gas (HCl or NH3) coming out into the room, even when the solution is near saturation.

aga - 30-11-2015 at 12:12

Quote: Originally posted by pneumatician

disolve 24k gold?

Salt, vinegar and bleach does that, as everyone knows.

https://www.sciencemadness.org/whisper/viewthread.php?tid=63...

Aqua regia also works.

Chlorine too.

ave369 - 30-11-2015 at 13:08

Just use a funnel-and-beaker trap and forget about suckback.

aga - 30-11-2015 at 13:48

Suck-back is always something to consider in this type of generator/gas/liquid arrangement.

The inverted funnel is an excellent idea. and works very very well.

crackedbits - 1-3-2018 at 13:20

I found this forum and more specifically this thread, because I'm wanting to concentrate the store bought Muriatic Acid which is in the 30% to 34% by volume range, to the reagent level of 36% to 38% concentration by volume. I'm an amateur chemist and have a small lab setup, but I'm not sure of what other equipment I may need to make this endeavor more efficient and repeatable. I'm not looking to make gallons of the stuff, but a process that yields any where from 500ml to 1L would be nice. I've seen the method by NurdRage on Youtube and a few other examples, nothing of which gives me a comprehensive layout of what equipment and expected yield. Any help would be appreciated, thanks.

ave369 - 1-3-2018 at 13:43

In my experience, hydrochloric acid is easier synthesized from scratch than bought and purified.

To make azeotropic HCl, I use table salt and battery acid (36% sulfuric acid). I mix table salt and battery acid in a flask and start distilling the mixture. Azeotropic HCl comes over, along with a small amount of HCl gas. Since both reagents are very cheap, this method can be used to distill large amounts of 20% HCl.

To make concentrated HCl, I use boiled-down battery acid. I don't "boil the bat" all the way down to 95%, 70% is more than enough for making HCl gas. Boil the battery acid until it starts fuming, then cool it. Put table salt in a flask, pour sulfuric acid in, stopper the flask quickly and start heating it, trapping the resulting HCl gas in azeotropic HCl you made previously using a funnel and beaker trap. This way you can make concentrated HCl all the way to 36%.

j_sum1 - 1-3-2018 at 15:22

@crackedbits
It seems to me that you want to bump up the concentration a little bit. I am not sure why. Fuming of HCl increases greatly as you raise the concentration into the thirties. I am not sure what the solubility limit is but it can't be too far from teh azeotrope.

My approach would be to chill the muriatic acid as low as it will go. Then produce HCl gas via NaCl and concentrated sulfuric acid. Bubble this through your chilled muriatic acid. Keep a thermometer in there so that hopefully you notice when no more gas is dissolving. Pass the excess gas through a suckback trap followed by iced water to avoid rusting everything in your fumehood.

HCL Concentration Setup

crackedbits - 1-3-2018 at 20:35

I was thinking of something along these lines, based on what I've seen and some suggestions.

ave369 - 1-3-2018 at 23:23

Quote: Originally posted by crackedbits

I was thinking of something along these lines, based on what I've seen and some suggestions.

What is the second flask for? Drying HCl gas? Why do you want to dry it, if you dissolve it in water anyway?

LearnedAmateur - 2-3-2018 at 01:14

When I want to avoid suckback in a distillation, I attach a vacuum adapter to the receiving flask and put on a piece of rubber tubing which I kink using my hand. As the solution rises up the tube, the kink is relaxed and atmospheric pressure is restored. Probably not as suitable for large distillations, but it’s always worked perfectly well for <200mL quantities. The only times suckback has been too rapid is when the solution is allowed to climb to the top of the tube and reaches the horizontal (I use a long ~80 degree distillate receiver to bubble gases), but before this point it maxes out at about 1cm per second and that’s near the end of the reaction when heat is taken off, so plenty of warning to take action.

Sulaiman - 2-3-2018 at 01:34

My HCl is 36% because that's what is commonly available here,
I can't think of aan occasion that I needed 36% concentration, even for aqua-regia.
I've even considered diluting my 36% to Azeotropic to reduce the rate at which my tools rust.

So my question has to be - why do you want highly concentrated hydrochloric acid ?

or, is there any good reason to keep my HCl at 36% instead of diluting to 20% ?

crackedbits - 2-3-2018 at 04:37

@ave369; Yes, I'm not wanting to add anymore water to my end product so I dry it through the second flask with sulfuric.

@LearnedAmateur; thanks for the tip.

@Sulaiman; I can actually buy ACS grade HCL, but it's six times the price of Muriatic, I have equipment, so I figured, I'd save myself some money. I'm also wanting to recycle the HCL I use in other processes. Bottom line, I want to save money, using less chemicals and recycling.

Sulaiman - 2-3-2018 at 05:28

An easy purification of muriatic acid is distillation,
if the starting concentration is greater than azeotropic
then the first hydrochloric acid to distill over will be more concentrated than the starting muriatic acid,
until azeotropic hydrochloric acid distills over.

AJKOER - 3-3-2018 at 06:23

Here an extract from some related comments on increasing the strength of HCl I once did:

Quote: Originally posted by AJKOER

.......
........
In the case of dilute HCl where the addition of a salt is not problematic for a particular application, one may consider the addition of anhydrous calcium chloride (or, a concentrated solution thereof) as a possibility. The reason relates to the apparent significant increase in the so called 'activity level' upon adding MgCl2 or CaCl2 or NaCl (in declining order of preference). Here is a real world reference relating to practical significance in the field of Hydrometallury where leaching out minerals from ores efficiently and cheaply is a major concern. Source: 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=PA15 .

Note, the author claims there is data confirming that a 2M HCl in 3M MgCl2 or CaCl2 (and also FeCl3) behaves like 7M HCl!

Here is a quote on the matter of discussion in one of the reference sources previously provided by Bfesser on Thermodynamic Activity (see http://en.wikipedia.org/wiki/Thermodynamic_activity ):

"When a 0.1 M hydrochloric acid solution containing methyl green indicator is added to a 5 M solution of magnesium chloride, the color of the indicator changes from green to yellow—indicating increasing acidity—when in fact the acid has been diluted. Although at low ionic strength (<0.1 M) the activity coefficient approaches unity, this coefficient can actually increase with ionic strength in a high ionic strength regime. For hydrochloric acid solutions, the minimum is around 0.4 M.[1]"

[Edited on 22-10-2014 by AJKOER]

Link: http://www.sciencemadness.org/talk/viewthread.
---------------------------------------------------------------

Interestingly, depending on the intended use of the HCl, say in organic formations of chlorides, an interesting idea is perhaps not aiming so much at concentrating the HCl itself, but instead focus on the formation of the monoatomic chlorine radical to effect reactions. This could proceed by adding some Cl2 to the HCl along with UV light and warming. Some chemistry on the proposed HCl/H2O/Cl2/UV mix::

Cl2 + H2O = HCl + HOCl

HOCl + heat → HCl + 1/2 O2

Cl2 + Cl-(aq) = Cl3-(aq) (or perhaps, Cl2 + HCl(aq) = HCl3(aq) )

HOCl + hv = OH• + Cl•

OH• + Cl- = OHCl•-

OHCl•- + H+ → H2O + Cl• (see http://pubs.acs.org/doi/abs/10.1021/j100497a003 )

Cl2 + hv = Cl• + Cl• (gas phase reaction is more efficient at higher pressure)

Cl3-(aq) + hv =?= 2 Cl• + Cl- (as the existence of Cl3- is itself speculative, likely unstable, and may be more effective than the prior gas phase reaction of gaseous Cl2)

Or: HCl3 (aq) + hv =?= 2 Cl• + HCl (speculation, but as the very existence of HCl3 is questioned, it lack of stability on irradiation seems plausible)

Cl• + Cl- = Cl2•- (this radical has increased longevity over Cl• but lower reactivity)

Additionally, one can increase the hydroxyl radical concentration (leading to Cl•) via the use of N2O and UV light:

N2O + H2O + UV --> N2 (g) + •OH + OH-

The hydroxyl radical can also be sourced from electrochemical redox or fenton/fenton-type reactions. See comments in hydrometallurgy reference above.

Example of applications:

CH3OH + hv --> •CH3 + •OH

•CH3 + Cl• --> CH3Cl

CH4 + Cl• --> •CH3Cl + HCl

However, some reactions which may not be desirable depending on the intended use of the HCl/H2O/Cl2/UV mix, introducing a range of products:

OH• + ROH --> H2O + •RO
......

Also from the reaction:

Cl• + HOCl = HCl + ClO• (see, for example, http://pubs.acs.org/doi/abs/10.1021/acs.jpca.5b01273 )
........

where the ClO• may further introduce chloride product impurities.

If your dilute HCl is also impure, a much better idea may be to convert it into chlorine and proceed similarly per above with added UV.

[Edited on 3-3-2018 by AJKOER]

AJKOER - 8-4-2018 at 06:44

Found an example of the application of chlorine radical formation in chlorine water with added NaCl and Fe(lll)/Fulvic acid as a photo-fenton source of hydroxyl radicals, OH•, and resulting Cl• and with Cl-, Cl2•− radicals. To quote: "Formation of Chlorinated Intermediate from Bisphenol A in Surface Saline Water under Simulated Solar Light Irradiation", by Hui Liu, Huimin Zhao, Xie Quan, Yaobin Zhang and Shuo Chen, Environ. Sci. Technol., 2009, 43 (20), pp 7712–7717, DOI: 10.1021/es900811c .

"Synopsis
Photoformation of organochlorine compound in surface saline water is demonstrated as one of its natural sources, and its formation mechanism is explained.

Abstract
Chlorinated organic compounds are generally of great concern, but many uncertainties exist regarding how they are generated. To illustrate the possibility of photochemical formation of organochlorine compounds in natural water, the phototransformation of bisphenol A (BPA) in aqueous saline solution containing Fe(III) and fulvic acid (FA), and in coastal seawater under simulated solar light irradiation was investigated. 2-(3-Chloro-4-hydroxyphenyl)-2-(4-hydroxyphenyl) propane (3-ClBPA) and 2,2-bis(3-chloro-4-hydroxyphenyl) propane (3,3-diClBPA) were the main chlorinated derivatives during the processes. Laser flash photolysis (LFP) and electron spin resonance (ESR) results indicated that the chlorination of BPA was most likely due to the formation of Cl2•− radical as a consequence of Fe(III) irradiation, yielding Cl• and OH• radical species and finally forming Cl2•− radical upon further reaction with chloride. The formation of Fe(III)−FA complex, which is a normal coexistence configuration of Fe(III) and FA in natural water, promoted the BPA chlorination through producing more Cl2•− radical. Moreover, FA had two opposite effects: forming Fe(III)−FA complex to enhance Cl2•− formation and competing radicals with BPA, which resulted in different overall effects at different concentrations: BPA chlorination was enhanced with the increasing of FA concentration ([FA]) when [FA] < 3.2 mg L−1; when the concentration of FA was as high as 10 mg L−1, it slowed down obviously. The described BPA photochlorination process took place from pH 6.3 to 8.5 and increased with the increasing of chloride concentration, indicating it could occur universally in natural saline surface water. These results propose a natural photochemical source for organochlorine compounds."

Link: https://pubs.acs.org/doi/abs/10.1021/es900811c
----------------------------------------------------------------

Here is an old work citing the historic known action of chlorine on benzene in the UV light available at: https://smartech.gatech.edu/bitstream/handle/1853/27667/fowl... .

[Edited on 8-4-2018 by AJKOER]

JJay - 8-4-2018 at 23:06

HCl gas doesn't seem to suckback nearly as strongly against an azeotropic solution as against pure water. The suckback vs. water is pretty ferocious.