Sciencemadness Discussion Board - Write-up on making pure BaCl2.2H2O

Write-up on making pure BaCl2.2H2O

woelen - 5-1-2010 at 12:22

Most of you may know of the very cheap barium carbonate available from potteries for just a few bucks per kilo. This material may seem attractive for making other chemicals, but due to impurities it is not that easy to use it for synthesis purposes. I have made pure barium chloride (snow-white material which gives clear solutions and has a pure green flame color) from this potteries grade barium carbonate. I have made a write-up which reflects my experiences:

http://woelen.homescience.net/science/chem/exps/BaCl2_2H2O/i...

Barium chloride is an interesting chemical on its own, but it also is a nice starting point for making other barium compounds.

watson.fawkes - 5-1-2010 at 16:00

Quote: Originally posted by woelen

I have made a write-up which reflects my experiences

In the appendix, when you discussing purifying 30% HCl solution, where the acid concentration above the azeotrope level, you recommend diluting to below the azeotrope level. It seems that you could also put water into the receiving flask, boil the undiluted acid, and dissolve HCl vapor. This would take less energy to boil, since HCl has much higher vapor pressures than does H2O at the same temperature. And that would translate into time savings. Is there a practical reason why this wouldn't work, say, because of gas diffusion limitations?

smuv - 5-1-2010 at 17:06

Having purified Conc. hydrochloric acid essentially by woelen's method, I can say that it is very convenient from a practical perspective. There is absolutely no fuming, or smells associated and you know pretty accurately what the concentration of your collected acid is without having to perform a titration. The method you outlined would involve the production of HCl fumes and would be much less practical in my opinion.

[Edited on 1-6-2010 by smuv]

JohnWW - 5-1-2010 at 17:09

Do you have any idea as to the level of Ca and Sr impurities in the pottery-grade BaCO3 used; and, in the synthesis of BaCl2.2H2O, have you analyzed the product for them? It would have to be virtually analytical-grade for use to gravimetrically analyze for sulfate as BaSO4.

12AX7 - 5-1-2010 at 18:43

My BaCl2 seems to be fairly pure. I'm not sure about turbidity, but then again my tapwater is so hard you'd never be able to tell anyway.

I don't know if it's the BaCO3 or the acid, but the solution tends to take on the yellowish tint of iron contamination. This can be avoided by using an excess of BaCO3, resulting in Fe(OH)3 precipitating in it. A series of filtering, recrystallizing and washing seems to yield very pure material.

A good filter is definitely valuable. I use glass wool, which adsorbs colloidial particles, turning them against the liquid, resulting in an even tighter filter. It can go quite slowly, but the liquid dripping out is crystal clear.

Tim

woelen - 6-1-2010 at 00:27

I tested my BaCl2 for iron with thiocyanate which is a very sensitive test on iron and it is negative on that.
The flame color of the material is pure green, so I'm quite sure that the Sr-content can be neglected (Sr-flame coloring is strong and deep red). Ca-concent is harder to test, because Ca-flame color is not that strong and the green of the barium could easily overwhelm the flame color of the calcium. Is there is a test which is specific for calcium in the presence of a lot of barium?

The reason why I dilute 30% HCl to around 20% is indeed that the distilled liquid then also is 20% (near-azeotropic boiling at 108 C). With 10% acid as starting point there is more uncertainty in the strength of the final acid (it can be anywhere between 15% and 20%).

JohnWW - 6-1-2010 at 00:45

Quote: Originally posted by woelen

(cut) The flame color of the material is pure green, so I'm quite sure that the Sr-content can be neglected (Sr-flame coloring is strong and deep red). Ca-concent is harder to test, because Ca-flame color is not that strong and the green of the barium could easily overwhelm the flame color of the calcium. Is there is a test which is specific for calcium in the presence of a lot of barium? (cut)

Atomic absorption spectroscopy, which is fairly standard for analyzing large numbers of samples for all metals in aqueous solution. If you have the right filters, you may be also able to use flame photometry, which can be used to analyze for alkali and alkaline earth metals (except Be and possibly Mg, I think).

unionised - 6-1-2010 at 01:18

Just in case Santa didn't bring you an AA machine for Christmas (I guess you were naughty); the classical separation depends on the fact that calcium chloride is a lot more soluble in a 1:1 mixture of ethanol and ether than barium or strontuim chlorides. IIRC you can do a similar separation of Ca from Sr as the nitrates.

However, if you have removes the (relatively similar) Sr from your Ba then you have probably removed most of the (relatively different) Ca.

[Edited on 6-1-10 by unionised]

kmno4 - 6-1-2010 at 02:21

The most sensitive chemical test (I know) for Ca is reaction with disodium salt of dihydroxytartaric acid phenyl osazone.
Yellow colour appaers in the presence of traces of Ca.
Sr or Ba does not give such reaction.
BTW. I do not like this woelen's procedure because of many reasons. Besides, purification of HCl(aq) for dissolving BaCO3 technical(or even worse) grade makes no sense to me.

[Edited on 6-1-2010 by kmno4]

woelen - 6-1-2010 at 03:06

Could you please explain what it is which you do not like in this procedure? If the procedure can be improved then that would be nice and any suggestions in this direction are welcome. I personally think that the only real issue is that the process takes a long time (almost one week between starting the process and the time of having the dry white material). The process can be done with a bare minimum of equipment and chemicals and that is what I like in the process. The final result is a snow-white material which gives a clear solution in distilled water. If you have a suitable centrifuge, then of course the process can be sped up considerably, but I don't have one, so I just had to wait.

Of course you can do a final recrystallization where you do not allow the liquid to evaporate to dryness. Doing so most likely will remove any calcium contamination (if that is present), but of course you also loose some BaCl2 in the process. Because I intend to use the BaCl2 for making Ba(BrO3)2 I did not recrystallize because Ca(BrO3)2 is very soluble and Ba(BrO3)2 is much less soluble and hence some CaCl2 in the BaCl2 does not really matter.

Making pure HCl(aq) does make sense. The HCl(aq), available at many hardware stores has a bright yellow/green color (at least where I live) and this definitely cannot be used to make a snow-white and pure product unless you are willing to do multiple recrystallizations from distilled water. The impurity in the HCl is mainly iron (I tested once with thiocyanate and it gives a deep red color with drops of the acid added to a solution of KSCN) but I can imagine that it also may contain other metals and colored organic stuff. You don't want that in the BaCl2. It might be that the impurities are as soluble as the BaCl2 and then they cannot be easily separated, so it is better to avoid having to deal with such impurities in the first place.

Unfortunately that test for Ca you mention is not something I can do. The compound dihydroxytartaric acid phenyl osazone does not sound like a common chemical, available for the average home chemist.

blogfast25 - 6-1-2010 at 13:13

Quote: Originally posted by 12AX7

A good filter is definitely valuable. I use glass wool, which adsorbs colloidial particles, turning them against the liquid, resulting in an even tighter filter. It can go quite slowly, but the liquid dripping out is crystal clear.

Tim

Tim, what kind of glass wool do you use? Where do you get it and how do you use it? Simply put on a funnel?

Do you think it would be usable for 'low precision' gravimetry?

Taoiseach - 6-1-2010 at 23:48

This is a very useful method. I have prepared BaCl2 and Ba(NO3)2 from crude barium carbonate and the crystals give a turbid solution from which slightly off-white impurities settle after a while. I definetly need to purify them and this is where woelen's procedure comes in handy.

Maybe precipating BaCl2 from a solution with dry HCl gas gives an even better product? This is a known procedure to make pure NaCl. Dry HCl gas is easy to make from NaCl and NaHSO4, the latter being sold as a pH regulator for swimming pools.

How would you go about making crystals of BaCl2*2H2O that look as nice as the comercial stuff, i.e. they all have about the same size and shape. The selfmade stuff always consists of chunky crystals of different size.

Btw does barium give basic compounds such as Ba(OH)Cl or Ba(OH)NO3 when the acid is neutralized with BaCO3? I once made Ba(NO3)2 by neutralizing HNO3 until no more CO2 was evolved, then filtrated and evaporated the solution. A mixture of the obtained crystals with sulfur cannot be ignited with a match.

[Edited on 7-1-2010 by Taoiseach]

Satan - 7-1-2010 at 00:13

Quote: Originally posted by Taoiseach

Btw does barium give basic compounds such as Ba(OH)Cl or Ba(OH)NO3 when the acid is neutralized with BaCO3? I once made Ba(NO3)2 by neutralizing HNO3 until no more CO2 was evolved, then filtrated and evaporated the solution. A mixture of the obtained crystals with sulfur cannot be ignited with a match.

Your crystals probably contain water of crystallization.

woelen - 7-1-2010 at 04:43

Quote:

Maybe precipating BaCl2 from a solution with dry HCl gas gives an even better product? This is a known procedure to make pure NaCl. Dry HCl gas is easy to make from NaCl and NaHSO4, the latter being sold as a pH regulator for swimming pools.

I do not think that is a good way to purify BaCl2 which gives turbid solutions. If you precipitate BaCl2 from a liquid which also has other solid particles in it, then the BaCl2 will become mixed with the solid particles. Once you have material which gives completely clear solutions, then it _might_ be a useful method of further purifying, but this needs further testing. This only works if indeed crystals of BaCl2 separate from the liquid when the concentration of HCl increases.

Quote:

How would you go about making crystals of BaCl2*2H2O that look as nice as the comercial stuff, i.e. they all have about the same size and shape. The selfmade stuff always consists of chunky crystals of different size.

The crystals I have are fairly large, some well over 5 mm while others are as small as a grain of salt. But I do not care about that. If I really want a more uniform material then I simply crush the crystals. What matters for me is that the crystals are purely white/colorless and give a clear solution.

Quote:

Btw does barium give basic compounds such as Ba(OH)Cl or Ba(OH)NO3 when the acid is neutralized with BaCO3?

I don't think so. There might be a very weak hydrolysis in its solutions, but hardly noticeable. HCl is a strong acid and Ba(OH)2 is a strong and soluble base.

I also have made Ba(NO3)2 and I also found it hard to ignite a mix of Ba(NO3)2 and S. But the same is true for KNO3/S. If some C is added, then ignition is easier, but you loose the green color of the barium. With a tiny amount of red P added to the mix, ignition is MUCH easier.

Ba(NO3)2 is harder to make in a pure state than BaCl2.2H2O, because Ba(NO3)2 only is sparingly soluble in even moderately concentrated HNO3 and that makes it harder to separate from insoluble material in the BaCO3. You can do this though according to the same procedure as posted here, but you need to use much more water and have to work with more dilute solutions. In this way I have made around 5 grams of Ba(NO3)2.

@Satan: As far as I know, Ba(NO3)2 does not have water of crystallization. Igniting mixes of nitrates and sulphur is not that easy, unless you have a somewhat larger amount of the mix.

kmno4 - 7-1-2010 at 06:34

Woelen - there are many variations of procedure you gave.
In my opinion your procedure is too complicated.
I would do like this:

Take 10g of technical BaCO3, 50g of tap water and add enough technical HCl to dissolve all carbonate. If H2S smell is present - boil it all for 30 minutes to remove H2S. Solution may be not clear, but nevermind.
Next add 1g of carbonate (must not dissolve - but if does, add another portion) and boil this mixture for 10 minutes.

In this way all colloids become more compact sediments, HCO3/CO3/SO4/... ions are removed as Ba salts and most of heavy metals are precipitated as carbonates/hydroxides.
Now solution+precipitate should be filtered and coffe filters are very good for this purpose. BaCO3 present in precipitate acts as collector/additional filter for small particles. Anyway, if filtrate is turbid - filter it using the same filter. If still turbid - filter again.
This is known oldschool way for removing even the smallest particles suspended in sollution.
Of course, vacuum filtering is very welcome.
Clear sollution can be partially evaporated and cooled for BaCl2x2H2O crystallization and separated crystals should be crystallised from water again. For this last crystalization demi or distilled water should be used, because any traces of SO4/CO3 ions present in water cause turbidity of solution.
BaCl2x2H2O can be also precipitated with pure, concentrated HCl(aq) or better by same HCl gas, but it generates additional costs.

Effects of this virtual procedure depend on kind of impurities in starting materials.

BTW - mentioned osazone can be prepared from phenylhydrazine and dihydroxytartaric acid. Dihydroxytartaric acid is prepared (in simple way) by nitration of tartaric acis and subsequent hydrolysis (see: http://pubs.acs.org/doi/abs/10.1021/ja01442a013)

12AX7 - 7-1-2010 at 07:01

Quote: Originally posted by blogfast25

Quote: Originally posted by 12AX7

Tim, what kind of glass wool do you use? Where do you get it and how do you use it? Simply put on a funnel?

Do you think it would be usable for 'low precision' gravimetry?

No, you aren't likely to recover all the material (or all the liquid, to weigh by difference) from the wool.

For filtering large solutions (like all that chlorate I went through), I have an 8" screen with glass wool wadded on the bottom.

The glass wool specifically is Kaowool 23. Generic glass/mineral wool may have coatings that don't work well, YMMV.

Tim

woelen - 7-1-2010 at 07:16

Thanks for your response. I like the idea of refiltering with the same filter. That indeed can help to make a clear solution.

I do not agree with using technical grade HCl, as I explained in my previous post. Having a lot of colored crap in the solution makes it more complicated to get rid of things afterwards. Use of tap water might be an option. Sulfate, chloride, and carbonate in the tap water is not an issue as they are precipitated or are target ions, but calcium and magnesium ions can be an issue. Of course, with a final recrystallize you can get rid of these and for some purposes these ions are not an issue at all.

Maybe next weekend I can try this modified procedure with another 10 grams of barium carbonate, with the exception of using good quality HCl and not the technical grade crap. I have a vacuum pump and a buchner funnel, so the filtering can be sped up considerably.

blogfast25 - 7-1-2010 at 08:10

Thanks, Tim.

You wrote:

"No, you aren't likely to recover all the material (or all the liquid, to weigh by difference) from the wool."

I'm just looking for alternative filtration media for simple, indicative gravimetry. As long as the precipitate gets fully collected on the filter and the filter doesn't absorb/lose something at, say gas mark 9, I'd be OK with that.

What's Kaowool 23 normally used for?

woelen - 9-1-2010 at 14:16

I have tried the method of kmno4. It indeed is much faster, in one evening you can do his entire procedure, provided you have a vacuum pump and a suitable funnel for fast filtering.

The only thing which is really disappointing is the bad yield. Two times of partial recrystallisation from water leads to loss of a LOT of BaCl2. I again started with 10.0 grams, but at the end I only had 3.1 grams of pure BaCl2.2H2O. The product is of good quality, purely white and it gives clear solutions, but I hoped to have a somewhat better yield.

I did the following:
- Take 10.0 grams of potteries grade barium carbonate
- Add around 40 ml of distilled water and swirl the beaker, such that the barium carbonate is suspended through the water.
- Take 10% HCl (colorless material, not the yellow/green stuff from the hardware store) and add this to the suspension of barium carbonate while swirling the beaker. Add the acid in small amounts, otherwise there will be too strong foaming. Keep on adding acid, until the liquid has become opalescent and then stop. At this point around 30 ml of 10% HCl was added.
- Now add 0.5 gram of barium carbonate and swirl again. The liquid is white again and completely turbid.
- Boil for 15 minutes or so in order to get rid of all H2S. After this, there is no smell of H2S anymore, but there is another sulphur-like smell. The liquid still is white and turbid. Some more coarse granules are at the bottom of the beaker.
- Filter the liquid. A vacuum pump really is needed here, otherwise the filtering takes a loooong time. After one pass of filtering the liquid still is opalescent and pale yellow.
- Use the same filter again and filter another time. This second filtering pass takes a longer time, even with vacuum pump, but the resulting liquid is totally clear. The liquid is not colorless though, it is light yellow and it also has a peculiar smell, somewhat like the smell of a burnt match.
- Boil down this liquid and keep on boiling until some solid barium chloride is formed on the surface of the boiling liquid.
- Set aside the hot liquid and leave it at room temperature. While this is done, many crystals are formed and a crust of solid barium chloride is formed at the bottom and on the surface of the liquid.
- When the liquid has cooled down to room temperature, set it in an ice bath, made of snow (we have a lot of that at the moment) and some water. Temperature is 0 C. More crystals are formed.
- Decant the ice cold yellow liquid and put the solid on a filter paper, which is placed on a heap of paper tissue. Most of the yellow liquid is absorbed. The solid already looks quite nice, it is off-white and only somewhat humid when it is allowed to stand on the filter paper for a few minutes. The paper tissue absorbs the liquid. Do not place the solid directly on paper tissue, because then it will be hard to separate lateron.
- Clean the beaker with some distilled water and then carefully scrape the solid from the filter paper and put it in the clean beaker.
- Add around 10 ml of water and boil the liquid. All of the solid dissolves again and the resulting solution is almost colorless. Keep on boiling until crystals of solid material appear again.
- Set aside the liquid and let it cool down to room temerature. Crystals are formed in the liquid. Then put it in an ice bath and let more crystals be formed.
- Decant the liquid from the ice cold mess. Only 2 ml of liquid can be decanted.
- Put the solid on a fresh filter paper and a fresh pile of paper tissue and let the adhering liquid be absorbed. When the material looks fairly dry, carefully scrape it from the filter paper and transfer it to a petri dish. Spread the material over the petri dish and put that on a warm dry place and allow to dry.
- When the material is almost dry, crunch the material and allow to dry somewhat longer. Then transfer the still warm and perfectly dry material to a vial which must be tighttly closed.

Final yield: 3.1 grams of purely white solid which gives a clear solution.

I think that the final product of this procedure is more pure than the product from my procedure, due to the recrystallisations and not letting the liquid evaporate to dryness. But the losses also are much more. Even in ice cold water, barium chloride dissolves quite well, so the losses are high.

A next time, when I do this, I think that I also will keep the decanted liquid, boil that down further and get another crop of (less pure) crystals. Maybe the yield can be improved somewhat when that is done.

I also HAD to do a recrystallization, because the liquid was yellow after heating and filtering. This time I did not add any H2O2 and the yellow color most likely is due to the presence of some colored sulphur compound (also given the smell of the liquid, which was like the smell of burnt matches). In my procedure, the liquid I obtained was colorless and this probably is because the added H2O2 oxidized most of the soluble sulphur compounds to sulphur and the long standing in contact with air probably did the rest.

Maybe next weekend I'll try a combination of both procedures, trying to combine the speed of kmno4's procedure and the yield of my procedure. Kmno4's procedure is nice for its fast work up and most likely purer end product (less calcium in the final product), but the yield also is much lower.

[Edited on 9-1-10 by woelen]

blogfast25 - 23-9-2010 at 06:53

Quote: Originally posted by kmno4

Woelen - there are many variations of procedure you gave.
In my opinion your procedure is too complicated.
I would do like this:

{snip}

BaCl2x2H2O can be also precipitated with pure, concentrated HCl(aq) or better by same HCl gas, but it generates additional costs.

I have some BaCl2.2H2O but it's quite impure (yellowish, dissolves to turbidity...) which I'd like to purify.

Does kmno4 mean that BaCl2 is less soluble in conc. HCl than in water? That it can be precipitated fairly purely with conc. HCl?

That is also the case for ZrOCl2, which can be precipitated by adding conc. HCl but for others it's the opposite case: PbCl2 and Cu(I)Cl are much better soluble in conc. HCl than in pure H2O...

woelen - 23-9-2010 at 09:14

Yes, BaCl2 is less soluble in conc. HCl. You can precipitate it with conc. HCl, but I expect that also quite some acid will co-precipitate. You'll end up with a solid, which is somewhat acidic. Probably heating the solid for a while helps to remove most of the adhering HCl.

blogfast25 - 23-9-2010 at 09:41

Ok, viability will depend on just how quantitatively the BaCl2 precipitates. Worth a try I guess...

12AX7 - 23-9-2010 at 10:41

I had fine results (excepting the turbidity, which settles out) with plain old recrystallization. Do you have a problem with minor impurities not easily removed, e.g. sodium, calcium?

Tim

blogfast25 - 23-9-2010 at 12:27

Not particularly Tim, but I now have access to relatively pure 32 w% HCl and not very expensive either. Gone are my 'Patio cleaner' days! So I'd like to explore that method a bit.

I'm wondering for instance whether you could recover the supernatant HCl which will be weakened a little and should be saturated with BaCl2. One batch of HCl may thus work for several precipitations. Just a thought...

12AX7 - 23-9-2010 at 20:19

I imagine distillation is necessary. *cough* eww.

blogfast25 - 24-9-2010 at 04:14

Quote: Originally posted by 12AX7

I imagine distillation is necessary. *cough* eww.

For highest purity, yes. But depending on solubilities and concentrations a few recrystallisations should do wonders...