blogfast25
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pH modification problem
I have a material in fine powder form that contains both water soluble alkali and water insoluble alkali.
The total Alkali Reserve is about 2.5 gram equivalent of NaOH / 100 g of substance (determined on an aqueous leachate by acidometric titration).
The water soluble alkalinity is not caused by carbonates: an aqueous leachate of the material does not test positive for carbonates (CO3(2-) (aq)).
The material also contains at least 30 % of calcium carbonate, the rest I assume is mainly silica or siliceous material (non alkaline).
I want to modify the slurried pH of this material, currently about 12.5, by adding a powdered pH modifier to neutralise the water soluble alkalinity.
This has proved far more difficult than initially anticipated. So far adding (acc. the Alkali Reserve) substances like NaHSO4, CaCl2 (at one point I
thought the Alkali Reserve was caused by K2CO3 but I was wrong), aluminium sulphate octadecahydrate and citric acid monohydrate have not even put a
dent in the slurried pH.
I assume that the reason for these failures is that these solid acids (bar CaCl2) are in fact attacked and neutralised by the ‘sea’ of water
insoluble alkali that the substance contains, i.e. the calcium carbonate.
Does anyone has any smart ideas for powdered pH modifiers that might not attack or not be attacked by fine CaCO3 but will neutralise small amounts of
accompanying water soluble, non-carbonate based alkali, when slurried with water?
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watson.fawkes
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It's behaving like a buffered solution. The first thing I might suggest is to try your pH adjustment on a filtered solution, to see if there's
something acting as a buffer in the solution. It might be in both, but you can at least start the divide-and-conquer process by working separately
with the solution and the filtrate.
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blogfast25
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Quote: Originally posted by watson.fawkes  | | It's behaving like a buffered solution. The first thing I might suggest is to try your pH adjustment on a filtered solution, to see if there's
something acting as a buffer in the solution. |
This is in the pipeline for tomorrow, as well as another similar experiment.
I don't think the filtrate has much buffering power though: I would have noticed that when I was determing the AR by titration of the filtrate with
0.1 M HCl. I titrated against methyl orange because I was convinced the alkalinity was carbonate based. I guess I should now repeat the AR titration
against phenolphtalein or straight against a pH meter.
Another assay showed the material (as such) to contain about 75 % HCl soluble matter. Much effervescense ensued when carrying out this leaching. The
leachate was part neutralised with NaOH, then an excess Na2CO3 was added causing a large amount of white precipitate to drop out. It strongly points
to CaCO3, as the acid soluble matter.
If CaCO3 is causing the problem then modifying the pH of this material will require some nifty chemistry, IMHO...
Another candidate pH modifier, NH4Cl, also failed to lower the pH by even 0.1 points.
I should also make clear that apart from cost considerations, there are two other constraints imposed on the pH regulator:
* must be a stable solid
* no more than about 10 - 15 % can be added to the substance.
Yesterday I was pondering about a hypothetical material that would be a medium strong acid, encapsulated in a neutral but water soluble 'coating'. Dry
mixing the pH regulator into the substance would leave the pH moderator intact but on carrying out the pH test, i.e. upon wetting, the pH regulator
would dissolve and then its acid reserve would react preferentially with the already dissolved alkali reserve of the substance.
[Edited on 2-3-2013 by blogfast25]
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blogfast25
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Two more experiments were performed.
Firstly the pH test was simulated but by filtering the watery phase off. pH was the determined to be 12.7. The the calculated amount of NaHSO4 was
added as a solid and allowed to dissolve by stirring. pH was again determined and found to be to be about 5.0, so a slight overshoot (7 - 9 would be
the target.
It does show that adding the pH modifier in the absence of the solid matrix works (but it’s not a legal way of doing it).
Slightly curious was that some turbidity was observed after addition of the pH modifier, which didn’t go away. That could point to CaSO4. Maybe
there’s a little CaHCO3 in the leachate? No effervescence was observed though.
Secondly , the pH of the unmodified material was determined and found to be 12.7 too. Then the calculated amount of pH modifier NaHSO4 was added,
already dissolved in a minimum of DIW. The pH initially went down, a small amount of effervescence was observed and the pH then stabilised at about
12.2. Clearly the pH modifier preferentially reacts with the presumed CaCO3 matrix, to the leachable alkalinity.
[Edited on 2-3-2013 by blogfast25]
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AJKOER
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If the target is "a hypothetical material that would be a medium strong acid, encapsulated in a neutral but water soluble 'coating'", perhaps
H2C2O4.2H2O?
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blogfast25
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Asssuming you mean oxalic acid dihydrate, then definitely no. Apart from toxicity concerns, there is no 'water soluble coating'. Oxalic acid would
also preferentially attack the massive CaCO3 matrix.
Having looked at all the data (I've spent quite a few hours on this project) it's become clear that much, perhaps all, of the water soluble alkalinity
is Ca(OH)2. Small amounts of CaO wouldbe plausible in the dry product, with water this then forms Ca(OH)2 which is sufficiently water soluble to cause
the high, slurried pH values.
If I'm right about this, then several water soluble sulphates should provide at least a partial fix: MgSO4 and ZnSO4 would react as:
Ca(OH)2 (aq) + MSO4 (aq) === > CaSO4 (s) + M(OH)2 (s)
CuSO4 and FeSO4 should also work but there are other issues with these.
So tonight will be the hour of truth, so to speak...
[Edited on 3-3-2013 by blogfast25]
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