BromicAcid
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While adding hot BaCl2(aq) to boiling KBrO3(aq) the solution erupted out of the beaker, what do you think?
Today I was attempting to make Ba(BrO3)2 by a double displacement reaction between BaCl2 and KBrO3.
2KBrO3(aq) + BaCl2(aq) ---> Ba(BrO3)2(s) + 2KCl(aq)
The procedure was taken from preparative inorganic synthesis and scaled down. The prep called for 334g (85g) KBrO3 to be dissolved in 700 ml (175 ml)
boiling water. To this is added a hot solution of 224 g (62 g) BaCl2*2H2O in 400 ml (100 ml) hot H2O. Upon cooling the yield is supposedly almost
quantitative.
My number above are the ones in parenthesis. However upon addition of the BaCl2 solution my main beaker erupted and spattered hot boiling solution
everywhere. There were several variables, 1) My main beaker was under the influence of magnetic stirring 2) My BaCl2 solution was beyond hot, but
still had failed to boil or dissolve all the BaCl2, there was still some solid at the bottom. 3) My BaCl2 was made at my house from ceramic grade
BaCO3 and HCl, the supernate discarded and the crystals dried at 90 C, they still smelled slightly of HCl. 4) There was fine particulate in my BaCl2
solution, little black specks that may have fallen off the tree.
What I think happened could have been the organic matter in the BaCl2 and the acid present reacting with the concentrated hot oxidizing solution to
oxidize the organic which caused a gas evolution. However I have other postulations. Regardless, what does everyone here think happened?
On a side note everything on my table was covered in Ba/K/BrO3/Cl(aq) which hardened to a film. Everything was tossed in a bucket and filled about
one half with water and was acidified with H2SO4 as is my normal method of treating barium wastes (although I was wary due to the bromate, however
bromates do not form bromine oxidies upon acidification in aqueous solution). Still though, when I came back to this mix a few hours later it was a
very strongly acidic bromine mix. It attacked everything, it was very interesting to say the least.
[Edited on 9/14/2004 by BromicAcid]
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JohnWW
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The "eruption" problem on mixing may be due to the two solutions, one of which was boiling and the other heated, having different elevations
of their boiling-points, due to having different molar concentrations (expressed as the anhydrous salt contents) of solutes, which determines
elevations of boiling points of all solutions. The boiling KBrO3 (m.wt. 167.011) is present (334 gm/700 ml) as a 2.857 molar solution. The added
heated BaCl2 (m.wt. 208.246 anhydrous, or 244.278 as the dihydrate) is present (224 gm of dihydrate/400 ml, equivalent to 190.959 gm of anhydrous) as
a 2.292 molar solution.
Adding a solution of lesser molarity, as in this case, which brings the total molarity down to the weighted average of 2.652 before any precipitation
of Ba(BrO3)2 occurs, therefore depresses the boiling-point of the resulting mixed solution. But, assuming that both the KBrO3 solution, of greater
volume and also boiling at a higher temperature, and the BaCl2 solution are at their respecitive boiling-points, adding the BaCl2 solution would
almost instantly result in the latter solution being superheated above its boiling-point, in pockets where it is surrounded by the KBrO3 solution
before mixing is accomplished. This would result in explosive emission of steam from these pockets of unmixed BaCl2.
In addition, if the precipitation of Ba(BrO3)2 is exothermic, which would be the case if it is only sparingly soluble (no data in Perry chapter 3, but
the chlorate is quite soluble), and occurs almost immediately, the heat evolved reflecting its lattice energy, this heat production would increase the
temperature of the mixed solution.
Any fine particulate insoluble impurities in the reagents would further increase the ebullition, as these would provide solid nuclei with porous
surfaces on which water vapor bubbles would form.
BTW treating Ba wastes by acidifying them with H2SO4 would not dissolve them, but instead result in a highly insoluble precipitate of BaSO4, which is
in fact the principle of the gravimetric analysis method for sulfate.
John W.
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BromicAcid
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| Quote: |
In addition, if the precipitation of Ba(BrO3)2 is exothermic, which would be the case if it is only sparingly soluble |
Ba(BrO3)2 is fairly insoluble, .44g/100ml @ 10C and 5.39 g/100ml @ 100C so you're on the right track.
| Quote: | | BTW treating Ba wastes by acidifying them with H2SO4 would not dissolve them, but instead result in a highly insoluble precipitate of BaSO4, which is
in fact the principle of the gravimetric analysis method for sulfate. |
Which was my intention, it is one of the only environmentally benign forms of barium and is therefore the only way that I dispose of it. Because
normally my method for disposing of chlorates/bromates is different but I felt obligated to take care of the potentially poisonous cation over the
oxidizing anions first in this case.
Maybe this is the reason why the BaCl2 solution is supposed to be hot and not boiling, to be able to take up some of the extra heat? The preparation
which is step by step and normally good about telling you what to expect does not even hint at such a possibility.
Also, could it be any metal ions floating around in the barium chloride that might be catalytically decomposing the bromate? Although no bromine or
other scented gasses were evolved.
Although I blame the purity of the BaCl2 I used you may be on the right track JohnWW.
Any further thoughts from anyone would be appreciated 
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JohnWW
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Another important point is that the precipitation of the Ba(BrO3)2 would itself substantially reduce the molarity of the 1100 ml mixture when it
occurs, by up to about 0.833 moles of solutes of the 2.615 present on mixing in 1,100 ml. This would result in substantial depression of the boiling
point, which would be substantially lower than that of the boiling KBrO3 solution - and consequently, almost immediate superheating of the mixture
(not just of pockets of unmixed BaCl2), although depending on how fast the Ba(BrO3)2 precipitates.
Besides, the finely divided solid Ba(BrO3)2 precipitated would result in further surfaces on which water vapor bubbles can form.
John W.
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BromicAcid
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After it erupted there was still some solid that precpitated out. It kept this, and dried this, and today recrystalized it from boiling water.
It appears to be somewhat pure Ba(BrO3)2 it when dillute solutions are made it reacts with H2SO4 to form a precipitate immediately, upon standing the
solution turns yellow then red from decomposition of the HBrO3.
At least I don't have to worry about bromine oxides since they all have very low stability at STP. Concentrated H2SO4 can be added to bromates
without the formation of bromine oxides, just bromine, water, sulfates, and some oxygen.
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Magpie
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I think this is an important observation you have made Bromic and feel John W's hypotheses are good. It is commendable that you are trying to
learn what caused this sudden and potentially dangerous flash boiling. Adding a significant amount of heat quickly to a liquid at or near its boiling
point, such as a hotter material (which you may have done) and/or causing an exothermic reaction (which you likely did) do seem to provide a logical
explanation. Importantly, this provides a warning for future experimentor's to help prevent a scalding that could scar them for life.
The single most important condition for a successful synthesis is good mixing - Nicodem
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