Sciencemadness Discussion Board - ZnCl2 * nH2O, purificiation

ZnCl2 * nH2O, purificiation

semiconductive - 12-10-2022 at 16:46

Hello, I'm exploring some ionic liquids / deep eutectic solvents.
I have an interest in ZnCl2 based mixtures.

I tried making some ZnCl2 by taking zinc oxide 2g and mixing it with 2.8g of ~32% HCl.
I expected (novice chemistry), the hydrogens of HCl to attack the oxygen on zinc and make water. (Hardware store muriatic acid is contaminated by a small amount of iron, but I ignored that.)

When I tried it ... the result was white - water insoluble crystals and a lot of heat.
I think I made a Sorrel-like cement or Simonkollite.

https://en.wikipedia.org/wiki/Zinc_chloride_hydroxide_monohy...

Adding a slight excess of HCl will dissolve the crystals into a clear mixture.

I tried the reaction a few more times with excess water diluting the HCl.
It seems that the more dillute the solution, the slower the reaction proceeds, and the less of this white crystal is formed.

But I was also noticing that there is some fizzing as the reaction proceeds.
I don't see why that would happen if the conversion ZnO + HCl + HCl -> ZnCl2 + H2O.
When the reaction is very hot, there is a lot of fizzing.

So, I suspect that heat in the reaction can drive off a small amount of HCl? Therefore, not all oxygen on Zinc Oxide gets converted to water.

As a matter of reaction design: I don't want excess HCl in the solution, afterward, if possible.

Q1) Is there a reliable/theoretical formula to accuractely estimate the efficiency of such a reaction so I can predict how much Chlorine loss needs to be compensated for during the reaction? (Obviously I could do it empirically, but I'm trying to understand the theory.)

I suspect that heating the ZnCl2 * nH2O liquor could drive off a certain amount of chlorine/HCl during crystallization. In such a case, it seems likely that ZnCl2 would act similar to MgCl2 -- and form a basic hydroxide.
MgCl2*H2O -> MgOHCl + HCl (gas),
ZnCl2*H2O -> ZnOHCl + HCl (gas).

I've watched a video of this being done starting from metallic zinc + muriatic acid, and then drying the solution to ZnCl2 crystals.

https://www.youtube.com/watch?v=WMODu5IiCIw

Which surprised me. The teacher didn't mention the possible conversion of ZnCl2 to ZnOHCl, during drying. But it seems to me that this is a likely side-reaction.
Is only a small amount of hydroxide formed, and ignorable?

When I try heating my solution to dryness, (unsurprisingly) the color darkens slightly and I can smell chlorine coming off the mixture. I don't get nice crystals, but a sludge on top of white-ish crystalling mass.

I'm wondering if the video only works so well because it has high purity HCl involved, or if there is something about ZincOxide which could produce a slightly different solution than solid zinc would when etched by HCl? (There is going to be a small amount of zincOxide on a metal surface, but it will be a tiny percent.)

With MgCl2, I k now that when I heat the mixture with KCl salt, added, that the amount of hydroxide produced is reduced. KCl seems to stabilize the drying of other chloride crystals under heat. When I check literature on heat storage salts, I find the same comment made -- KCl reduces conversion of MgCl2 to basic hydroxide during dehydration.

So, I am thinking, maybe the same thing will happen with ZnCl2.
if so, could I selectively dissolve out the ZnCl2 after crystallization?

I know, for example, that isopropyl alcohol will not dissolve KCL, and I *think* (memory could be bad) acetone doesn't either. Acetone and alcohol, can be 'salted out' to remove excess water content when azetropes prevent distillation from doing a good job.

So, I'm thinking ... if I add KCL to to the aqueous ZnCl2, would it be reasonable to expect to be able to selectively dissolve the ZnCl2 crystals out again using acetone or alcohol?

eg: Could this be made into a recrystallization process to purify the ZnCl2 and reduce ZnOHCl production?

The only other way I can think of getting rid of hydroxides, like ZnOHCl, is by making a HCL gas generator. Maybe KCl + Sodium Metabisulfite. If I were to pressurize the flask with excess HCl gas during drying, I expect the hydroxides would be suppressed. But, that seems rather complicated.

Thoughts & suggestions, welcome.

[Edited on 13-10-2022 by semiconductive]

j_sum1 - 13-10-2022 at 16:17

I have found that these things tend to work better when you work with more dilute acids.
I would also distil the HCl beforehand to eliminate any contaminants. It is a simple step and worthwhile. If you add approximately the right amount of water beforehand you get azeotropic HCl at 20.2%. For this particular reaction I would probably dilute that back to 10% or lower. I would also use a reasonable excess of acid.

The result is that you will have more controlled evolution of H2 gas, the temperature will be lower, you will have less splashing and loss of product, the possibility for side reactions is lower, and your product is more likely to dissolve. Starting with more pure HCl will also help. (Fe is a transition metal. In the absence of specific knowledge, I never trust transition metals to not catalyse something unexpected.)

In my mind I always think of zinc as being quite reactive. But in my experience, I find it a pain. The oxide tends to stick around longer than I expect and passivates surfaces. Excess acid helps. Time helps. You don't want to be impatient. [edit] On re-reading I realise you are using ZnO and not Zn. If you have access to Zn metal then that may make things easier.

If you use excess HCl, you will have to evaporate it off. That is, fuming. I would recommend popping your solution in a flask and bringing it to the boil to drive off HCl. You can scrub your gas to collect dilute HCl if you wish. If you keep adding water to your flask you will eventually drive off all the excess acid and be left with a nice pure solution of ZnCl2 which you can then evaporate in a beaker or evaporating dish.

Depending on how pure your zinc is and whether you use clean acid, you may or may not want to recrystallise. This will be hard because ZnCl2 is hygroscopic and extremely soluble. Crashing out with ethanol also will not work. If it was me, I would probably just evaporate to dryness and be content with some impurity.

Another approach is to react zinc turnings with gaseous HCl. That could give you an anhydrous product if your HCl is free from water. This will be a more difficult setup and probably only practical at rather small scale. But it is not beyond the realm of possibility.

[edit]
I can't comment knowledgeably on your ideas of adding KCl.
My instinct is always to avoid adding anything extra if I am trying to purify something.
I would try with clean acid first to try to prevent the brown sludge problem. And then see if I could get an acceptable result without adding anything.

[Edited on 14-10-2022 by j_sum1]

teodor - 14-10-2022 at 03:32

You can get ZnCl2 anhydrous "sticks" (it looks like a gray mass which you can break to pieces) by evaporating ZnCl2 solution in HCl (so, you evaporate HCl-water). Those sticks are water-soluble, like normal ZnCl2. You can get a very dense (neutral pH?) liquid by dissolving it in water. Also it have a low melting point, but the fumes are very poisonous (can cause a pulmonary edema).
But it is no possible to get ZnCl2 by evaporating water solution.

[Edited on 14-10-2022 by teodor]

[Edited on 14-10-2022 by teodor]

semiconductive - 2-11-2022 at 22:16

Thanks for the input.

I thought about this, and tried an acetone process first just to experiment.
This isn't a industrial kind of process, but I was just curious if it would work at all.
I expect the losses to be fairly high.

I'm not sure if Fe transition metal should dissolve in acetone or not.
So, I just decided to try it and see....

Here's what I did:
NaCl (1cc's worth), about 3cc's of acetone, and the rest 1cc of ZnO2.
I added HCl 32% (muriatic acid from Ace hardware) drop by drop and shook it until a clear solution was in the test tube.
ZnO2 sticks to everything this took a bit of doing.

Because of the common ion effect, and the acetone, the NaCl falls to the bottom immediately as expected after every shaking.

The theory I was working under is that:
The ZnCl2 being soluble in acetone and water would allow the NaCl to absorb some of the water, eg: since the acetone could displace water around the ZnCl2.
By doing this, I don't have to heat the ZnCl2 solution, and should not form ZnOHCl.

I then decanted the test tube into another test tube allowing it to drip as much as possible; then I popped *just* the salt into the microwave, boiled off the water for 60 seconds, allowed it to cool again; dropped enough acetone into it to wet it again, throwing away excess with any contaminant -- and then dumped the decanted solution back into the test tube and shook it.

I repeated this cycle about ten to twenty times ... and each time the acetone turned slightly more yellow and more viscous. So, it does seem to be able to pull the ZnCl2 out of the water solution (at least partially).

I'm not sure what the color ought to be; but I suspect that iron was in it.
But the test tube didn't go dark like it did when I boiled off the water from ZnCl solution directly.
I figured if I left it in air for a few days, it would absorb water and evaporate the acetone.
Then I could check the color.

Progressively, the solution turned darker until it looked no different than if I had heated the Zinc Chloride + water solution directly to a boil and driven off most of the water.

It seems, whatever contaminant is in there ... is just as soluble in acetone as in water.

I do have zinc metal, so I can do that route as well;
and I can distill the HCl (which I would prefer not to bother with, but I can...)

@teodor, evaporating HCl water. I've tried driving off 32 percent HCl from a salt mixture, just to see what would happen if I heated it gently. I think the HCl boils out faster than the water does, leaving me with quite wet salt.

Is there an optimal condition for evaporating HCl off, so as to remove as much water as possible? eg: Vacuum, dessicant, or something? I have CaCl2, but I'm not sure how I would use it; maybe with a fishtank pump to re-cycle the dried air + HCl gasses .... but it would likely destroy the rubber diaphragm.

I can easily make a teflon pump body (I have teflon powder, and my heating mantle gets up to 400C, which is enough to sinter teflon), but I'm not sure how to get a chemically resistant membrane/diapragm to make a recycling pump.

Any more thoughts?
Thanks!

semiconductive - 2-11-2022 at 22:30

Azetropic HCl ought to boil off uniformly by itself, and come back with the same concentration after distilling. That seems like a nice way to be able to make a constant concentration of distilled HCl.

Just curious: Would either of you expect NaCl or ZnCl's presence to change the azetropic characteristic of HCl that is boiled off?

kmno4 - 2-11-2022 at 23:04

Quote:

But it is no possible to get ZnCl2 by evaporating water solution.

Of course it is not true. See (rather simple) preparation of ZnCl2 and ZnCl2x1,5 H2O in "Pure Chemical Substances" by Yu. V. Karyakin and I. I. Angelov (1974 year), pages 406-407 (in Russian, available online).
You will not find these informations in Brauer

teodor - 3-11-2022 at 04:16

semiconductive, I don't quite well understand your troubles. Just boil it off in an ordinary distillation setup at atmospheric pressure.
Of course, it is unable to evaporate HCl and then water because they have a great affinity to each other. You will have as much residual water as much as HCl. There is a difference also between "perfectly dry" and "anhydrous".

semiconductive - 3-11-2022 at 22:52

I realize there is a difference between dry and anhydrous.
I'm just noticing with regular salt, that the smell of chlorine goes away before the wetness clumping of the salt.

So, it seems to me that when there is competition about the affinity for water and a common ion (chloride), that HCl may boil off with less water in it than azetropic.

Doing a search I came across this:
https://en.wikipedia.org/wiki/Salt-effect_distillation

That seems to agree with what I noticed. (I'm still a novice). That's all.
:cool:

DraconicAcid - 3-11-2022 at 22:59

Quote: Originally posted by teodor

You can get a very dense (neutral pH?) liquid by dissolving it in water.

I seriously doubt that zinc chloride solution is going to be neutral. I expect it to be quite acidic (and getting a bit of hydrated zinc chloride on my hand once hurt like hell, for the record).

teodor - 4-11-2022 at 04:03

Quote: Originally posted by semiconductive

I realize there is a difference between dry and anhydrous.
I'm just noticing with regular salt, that the smell of chlorine goes away before the wetness clumping of the salt.

So, it seems to me that when there is competition about the affinity for water and a common ion (chloride), that HCl may boil off with less water in it than azetropic.

Doing a search I came across this:
https://en.wikipedia.org/wiki/Salt-effect_distillation

That seems to agree with what I noticed. (I'm still a novice). That's all.
:cool:

Well, what matters here is the relative strengths of bonds between ZnCl2 and H2O and ZnCl2 and HCl. The same way it keeps water (coordination?) it (potentially) can keep HCl. We have to look the properties of this particular compound because I think it more defines the behavior than a common salt-effect.
Also, you brought an interesting question. Is H2O + HCl an azeotrope similar to, let say, benzene + water.

Quote: Originally posted by DraconicAcid

I seriously doubt that zinc chloride solution is going to be neutral. I expect it to be quite acidic (and getting a bit of hydrated zinc chloride on my hand once hurt like hell, for the record).

Well, I put the question mark because I just never checked the pH of ZnCl2 water solution. What I mean, after drying I expect it to be HCl-free and the remaining acidity due to properties of ZnCl2 itself.

[Edited on 4-11-2022 by teodor]

DraconicAcid - 4-11-2022 at 09:36

According to Merck, zinc chloride gives an aqueous solution with a pH around 4.

teodor - 10-11-2022 at 15:53

So, the ZnCl2 * HCl * 2H2O adduct loses HCl together with water in a vacuum @ 100C leaving anhydrous ZnCl2 (Gmelin, Zink, 1924).
The adduct separates as crystals by passing HCl in the aqueous ZnCl2 solution.

semiconductive - 16-11-2022 at 13:00

Quote: Originally posted by teodor

I got a bump trap today, and I'm going to try distilling concentrated HCL by using a boiling flask and computer controlled mantle, followed by a bump trap that can retain 150ml, followed by a drip spreader & 600mm Vigereux column.

If the cooling liquid can hold more HCL than the azetrope, I'll be curious to find out what concentration I get in the bump trap.

Barely bubbles at 92C, 31.19in Hg.

At least this will remove the iron in the Hardware Store muriatic acid.

Later; I can add KCl to the boiling flask for another batch, and see how/if it affects the rate of distillation and the final concentration.

Nice article.

My German is a bit weak -- it's been 26 years; Could you explain why they go to the trouble of using Phosphorous Pentoxide in the first part of the disccusion ( a dessicant ); if they are forming an aqueous solution?

I see two temperatures involved 25C and possibly 0C.
It seems the crystals form at 25C, so I'm not sure what the other temperature is for.
But, is this a three step process where they run HCl gas into a water solution (slowly) and dissolving Zinc metal? After dissolution, they crystalize slowly as water evaporates (25C) or lower?

I ask, because
I'm curious about the process by which ZnCl2, when driven to dryness by heat, *could* form ZnOHCl.

eg: At what step could hydrogen get released from water to make the hydroxide?

And is the author attempting to avoid that by crystalizing first at 25C *before* raising the temperature to dry the solution?

eg: How can I estimate how much ZnOHCl a process will be contaminated with when I finish drying. It doesn't seem, based on the article, that the process is particularly temperature dependent if they can heat it to 100C.

Zinc doesn't have significant electron affinity, which is one of the reasons it's a weird element. But, I'm not sure if that could be involved in making hydroxides form or not.

[Edited on 16-11-2022 by semiconductive]

clearly_not_atara - 16-11-2022 at 16:00

Quote: Originally posted by teodor

But it is no possible to get ZnCl2 by evaporating water solution.

It is possible to obtain anhydrous ZnCl2 at atmospheric pressure by forming the ammonia adduct ZnCl2*2NH3 and heating this. The required quantity of ammonia may prove unpleasant.

It is also possible to obtain anhydrous FeCl2 by reacting Fe with methanolic HCl. This can react with Zn to give ZnCl2. Preparation of anhydrous methanolic HCl likely requires gassing.

semiconductive - 17-11-2022 at 11:01

Quote:

Preparation of anhydrous methanolic HCl likely requires gassing.

This goes back to something I was trying a couple years ago. Fish tank pumps can be used to re-circulate gasses and force them through a dryer. The problem is ... most plastics are destroyed by HCl (hot gas). So, sintered teflon or glass methods to make an amateur level recirculation pump needs to be invented.

I have the PFTE powder, and my mantle is capable of reaching the sintering temperture of PFTE. I also have access to 3D printer to make molds with, and a high temperature water soluble salt (takes yellow heat from a torch, doesn't break down. So I ought to be able to make a corrosive gas recirculator pump out of either PFTE, Glass, or some combination of them.

I am waiting on 3D printer parts, which should be here within a month; and glass dust making equipment; eg: a ball mill and a few other odds and ends before trying to sinter a pump again. Should be able to give it a go in December.

Even though I destroyed a few plastic pumps, three years ago, trying experiments, I did notice that being able to move chlorine gas from a high temperature side to a low temperature side allowed distilling of heavier fatty chain alcohols than could be done in air. I recirculated the gas through Dry-Z-Air calcium Chloride, to try and break the chlorine gas apart from moisture.

It seems to me, that if a method could be made to use a smaller quantity of ammonia, or ammonia chloride, in a loop; that a variety of chemicals might be synthesizable, eg: A tertiary alcohol can probably be either chloronated, or metallized; and since it's volatile enough to reflux (like in the make potassium thread); then I think it worth an experiment to find out if chlorine gas could be used in a distillation like process that destroys water, slowly.

It's just hard to get tertiary alcohols; and I'm not sure methanol would survive since there is hydrogen on the same carbon as the hydroxyl group.

I can't electroplate Iron successfully in methanol for any length of time; I always eventually get gunk. I'm surprised there is a method to get anhydrous FeCl2 using methanol. I would expect a lot of loss to 'gunk'.

Electroplating of iron does work in formic acid, though badly.
So there may be alternatives.

I think (from over 325 experiments) that most organic alcohols may be 'promotors' of reactions; but they eventually get destroyed by side reactions. They aren't very stable 'catalysts'. I was using kerosine lamp oil years ago, with pictures on this site, to keep 'air' out of electrolysis experiments; and I even tried electroplating KOH in it, and a few mixed ethers. I've been having a heyday reading the potassium sticky thread.

Yup, I've come to many of the same conclusions ... though I'm a bit slow as a novice chemist. I try a *Lot* of things that fail.

[Edited on 17-11-2022 by semiconductive]

semiconductive - 17-11-2022 at 13:36

First distill test:
12 hours at 95C.
A few drops distilled over, maybe 2CC out of 100. There has been a *lot* of chlorine smell in the first hour, but if the gas coming off were azetropic; it ought to have condensed rapidly and I should have gotten instant reflux in the Vigereux tube. It didn't. I took over half an hour before the slightest haze showed up on the bump trap. Pulling off the Vigereux tube this morning, I could see fog form directly over the trap. The temperature of the vigereux tube is 34F. (It was cold all last night.)
So, I concluded there is a gas in the trap containing HCl; but it won't condense even at 34F.

This suggests that even without salt, HCl above the azetropic mixture amount can be driven off at below boiling temperatures.

Cranking flask temperature up to 105C. ... let's see if any more distills over; if not ... add another degree C every 3 hours or so. The boiling point at pressure, I should at least be able to find accurately.

But: it seems like if I want concentrated HCl of high purity, it might be easier to fill the drying trap with a marble at bottom, and glass or pure silica sand on top of it; add a small amount of distilled water; that way HCl gas would be forced to bubble through it; and it would drip down slowly each time the marble got lifted, so the most concentrated stuff would fall into the bubble trap immediately. After collecting the strongest stuff, I could then boil over the azetropic mixture.

[Edited on 17-11-2022 by semiconductive]

semiconductive - 18-11-2022 at 14:18

It is indeed *more concentrated* Hcl in the trap than in the bottle below it.
Density is 2% higher than starting liquid.
When I allowed liquid in flask to cool about 5 degrees, swirled it, and no bubbles happened; I was sure it was below the boiling point of what was left in the flask.

Therefore, I dumped the stuff from the trap back into the cooling flask.
(Should be safe if my measurement was off, and the mixture was azetropic. )
Immediately upon hitting the non-boiling acid in the flask, all the liquid from the bump trap burst into foggy vapor.
( Choke, Choke. )

Conclusions: Azetropic mixture isn't the only way that liquid from HCl drying/distillation can be driven off.

These phase changes aren't favorable, because the heat energy is insufficient to overcome the difference in boiling points when the gas rises very slowly. ( The stuff in the bump trap want's to boil at a lower temperature than the stuff in the boiling flask. )

This is probably (I'm guessing) an example of the fermi-energy level of chloride gas, and a very slow 'diffusion' process.

After the first hour of loosing HCl gas (and practically no water), the boiling temperature rose in the flask. (Bubbles stopped. )

I raised the temperature from 95C;
The temperature of the pid loop stabalized at 107C with no boiling with the PID controller set to 108C. Therefore, that was the point where the phase change wants to happen in heated liquid remaining in the flask; eg: somewhere between 107 and 108C. That must be the Azetropic boiling temperature, I think? (At the pressure recorded. )

Hypothesis: I lost enough HCl from the flask to reduce the concentration to Azetropic.

But, that means the narrow time where extra HCL was present in the fumes could still be exploited using mechanical pumping. I think non-azetropic distillation of chloride is quite possible by adding mechanical energy as air pressure to the boiling flask. ( Am I wrong? )

eg:
If I recirculated the gas from the top of the vigerux columb back into the boling flask through a vacuum tap, using a corrosive gas pump, I suspect I could distill 32% HCl; albeit slowly. Which suggests that HCl can be separated from water // it isn't a super strong adduct. ( I hope that's the right word. )

I wonder, how dry was the HCl in the vigereux column when I smelled chlorine?
Is there an easy way to test that?
The column was just above freezing, and not even a haze of moisture showed up it; but I still smelled chlorine for the first hour of heating the flask to 95C.

I think this means (theoretically, I'll have to test it), that HCl can be separated from water under mild thermal conditions;
if I add sufficient KCl salt into the boiling flask, I hypothesize (Le Shatlier's principle) that it would reduce the solubility of the chloride ion in the boiling flask and increase the release of HCl gas from the solution.

If that's the case, then Zinc atoms in the chloride salt might be able to oxidize water to hydroxide.
Hmmm....
I'll have to search for TGA of ZnCl * nH2O. Maybe I'll get lucky.

Edit: Interesting find:
According to some very old U.S. propaganda, mixing ZnCl2 aqueous with ammonium chloride will precipitate Zn[NH3]2.Cl2. With no water molecule left, thermal degradation can't leave a hydroxide, so heating to 300C or so -- will drive off ammonia, and this is an simple/efficient way to get anhydrous ZnCl2.

That's exactly the kind of thing I was already thinking of trying. But I doubt I would have raised it all the way to 300C on my own. That's pushing my mantle's capability near the edge. I think it maxes at 450C.

[Edited on 19-11-2022 by semiconductive]

semiconductive - 19-11-2022 at 21:35

My beginner luck.

From 107C yesterday, all the way to 125C this morning ... and it still didn't finish distilling. ( The vigereux was at freezing. )
The flask wasn't even bubbling this morning. ( And yes, the probe is on the glass of the Erlenmeyer, not on the mantle itself. )

HCl does not like to condense.

I finally cranked it up to 150C, and watched HCl start to boil very slowly.
But, it nearly stopped again after an hour. Darn weird stuff.
So I cranked it up to 160C; and it went nearly dry, when I turned it off.
I can see the iron contamination, now.

-------------------------------------------

I will try a variation on clearly_not_atara's idea ; unless someone thinks this won't work.

I'm going to measure out a gram of ZnO2.
Then enough NH4Cl to have two molecules for every zinc atom + slight excess.

I'll add HCL drop by drop, until the ZnO2 is fully dissolved.
Then I'll add amonnia until the PH is around 5 or 6.
I think Iron, copper, and most metals will precipitate at near neutral pH; as long as I don't add too much ammonia and re-dissolve them as hydroxides.
This should allow me to get rid of any small impurities in the ZnO2 powder by precipitation.

Decant,

Add ammonium chloride until all the zinc chloride has precipitated.

Filter out.

At this point I can either cook it at 300C and get anhydrous ZnCl2, or I can see if the difference in solubility is enough to selectively dissolve ZnCl2 into acetone, MEK, or ethyl-acetate. Not a lot of amateurs have 300C equipment, so I think I may try the solvent extraction first; just because it might be of interest to more people.

I don't know how much water I have in the acetone, mek, or ethyl acetate ; so I can't guarantee it's anhydrous; but if I put a piece of cooked CaSO4 into it, I think it will get it pretty close to anhydrous.

I could also distill them. (But, with my luck, that would take a week.)

With that done, I could then cook the ZnCl2 * [NH4Cl]2 adduct to dry it (just at 100C).
And see if I can then solvent extract anhydrous ZnCl2 into an organic solvent, filter out the NH4CL, and recrystallize the ZnCl2.

Does anyone see serious mistakes in my thinking?
Let me know.

I'm optimistic that this might actually work.