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papaya
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On oxalic acid: can't dihydrate be used directly, or water content is a subject of change depending on age?
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blogfast25
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Quote: Originally posted by papaya  | | On oxalic acid: can't dihydrate be used directly, or water content is a subject of change depending on age? |
Hydrates are rarely very stable, unless they're in equilibrium with the solution from which they originate.
Oxalic acid dihydrate, recrystallized, would not be a bad 'entry' point for a beginning titrator but several of the substances mentioned in the thread
are more accurate (if handled properly).
[Edited on 23-2-2015 by blogfast25]
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DistractionGrating
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| Quote: Originally posted by unionised | IIRC the traditional answer is to use sodium carbonate.
It can be made in very high purity by heating bicarbonate of soda to about 250 C for a few hours.
You need to let it dry in a desiccator.
Bicarbonate of soda is easy to get at very high purity for food use.
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I'm curious how much greater purity can be achieved through recrystallization vs. simply decomposing sodium bicarbonate?
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Fulmen
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Decomposing bicarbonate can't magically remove non-volatile impurities.
We're not banging rocks together here. We know how to put a man back together.
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DistractionGrating
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Fair enough. Although, my question wasn't a qualitative one, but rather, a quantitative one.
One common impurity in sodium bicarbonate or sodium carbonate is sodium sulfate. Considering that sodium sulfate has even less solubility in water
than sodium carbonate at 0C, wouldn't that make recrystallization an ineffective technique for separating sodium carbonate from sodium sulfate? (Not
being argumentative, I'm just trying to learn.)
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blogfast25
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| Quote: Originally posted by DistractionGrating |
One common impurity in sodium bicarbonate or sodium carbonate is sodium sulfate. Considering that sodium sulfate has even less solubility in water
than sodium carbonate at 0C, wouldn't that make recrystallization an ineffective technique for separating sodium carbonate from sodium sulfate? (Not
being argumentative, I'm just trying to learn.) |
Not generally, no. Recrystallisation relies on the impurities being low in concentration and not reaching their solubility limit on crystallisation.
This works kind of MOST of the time. Sodium sulphate in bicar is only a small amount, so it stays in solution, while most of the carbonate
crystallises, hence the purification. That's the general principle anyway.
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DistractionGrating
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I was thinking that the solubility of the sodium sulfate might be very low in the context of a saturated solution of, well, just about anything else,
but sodium carbonate in this case. I tried googling for something to back up my presumption about this, but no luck yet.
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blogfast25
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| Quote: Originally posted by DistractionGrating | | I was thinking that the solubility of the sodium sulfate might be very low in the context of a saturated solution of, well, just about anything else,
but sodium carbonate in this case. I tried googling for something to back up my presumption about this, but no luck yet. |
Well, you can invoke the 'common ion effect' here. The simpler truth is that recrystallizing sodium carbonate works well. The effect is thus likely to
be too small to suppress sodium sulphate solubility that much. Cold saturated sodium carbonate solutions aren't that high when expressed as molarity,
I think...
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DistractionGrating
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So, reading up a bit on the "common ion effect", and I apologize if I'm derailing the thread, even if momentarily, but, does that mean that at 20C, I
can reasonably expect to be able, for instance, to dissolve all of: 39.7g of sodium carbonate, plus 37.2g of potassium chloride, plus 39.7g of
magnesium sulfate in 100g of water (each compound at a saturated solution concentration, and no ions in common)?
EDIT: These may be poor choices, because they may react to form insoluble precipitates. I haven't thought that part through. The essence of my
question is being able to combine multiple, non-reacting substances, with no common ions, each to saturation.
[Edited on 24-2-2015 by DistractionGrating]
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DistractionGrating
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BTW, thank you for this. It is my takeaway.
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Fulmen
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Regarding the bicarbonate method: One source (http://www.sciencemadness.org/talk/viewthread.php?tid=61594&...) specifies sodium carbonate made from high purity sodium bicarbonate, I
wouldn't be surprised if something like this was the source. But I don't see how this process could produce a purer product than the raw materials.
Perhaps the bicarbonate is easier to purify?
We're not banging rocks together here. We know how to put a man back together.
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unionised
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To meet the pharmaceutical specification, sodium bicarbonate has to give an assay value that is equivalent to 99 to 100.5 % NaHCO3.
Food grade stuff is probably comparable.
How much purer do you need?
incidentally, the unsymmetrical limits suggests that the major impurity may be water so heating it will give a better purity in the product than you
started with.
Recrystallising the bicarbonate is a pain, because it decomposes in boiling water.
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blogfast25
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It ain't necessarily so. One food grade bicar I tested went brown on heating. I suspect rice powder as an anti-caking agent as the causal agent.
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Praxichys
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USP "food grade" OTC items are usually quite pure following some simple purification steps.
You can get lots of food grade/USP things like 99% isopropanol, glycerol, NaHCO3, KI, NaNO2, propylene glycol, sulfur powder,
citric/malic/ascorbic/acetic acids, NaOH, sodium metabisulfite, ammnium carbonate... just look for the chemical you need in Wikipedia's List of food additives and see if the chemical you want is on there. If it is, chances are you can find it USP or food-grade in a specialty
grocery store somewhere, and definitely on eBay.
[Edited on 24-2-2015 by Praxichys]
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blogfast25
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| Quote: Originally posted by DistractionGrating | So, reading up a bit on the "common ion effect", and I apologize if I'm derailing the thread, even if momentarily, but, does that mean that at 20C, I
can reasonably expect to be able, for instance, to dissolve all of: 39.7g of sodium carbonate, plus 37.2g of potassium chloride, plus 39.7g of
magnesium sulfate in 100g of water (each compound at a saturated solution concentration, and no ions in common)?
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That's the general principle, yes (insoluble combinations notwithstanding).
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unionised
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| Quote: Originally posted by blogfast25 |
It ain't necessarily so. One food grade bicar I tested went brown on heating. I suspect rice powder as an anti-caking agent as the causal agent.
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I wonder what % starch (or whatever) you would need to add to the pure stuff to get it to go brown on heating.
I wonder if the stuff might still meet the USP or BP spec.
Anyway, there's another interesting route which has the interesting property that you can get two confirmations of purity.
Sodium oxalate is relatively easy to purify.
You can titrate it directly if you like, but it's also possible to roast it to get the carbonate Then you can titrate the carbonate (from a known
mass of oxalate) to get an indication of the purity.
You can also make a solution of oxalic acid and titrate that against a solution of permanganate to confirm the purity of the oxalic acid.
OK, to do that, you need a known concentration of permanganate.
You can get that by titrating the sodium oxalate.
And, of course, you can titrate the oxalic acid as an acid against the sodium carbonate.
As long as the two different titrations of the oxalic acid (against MNO4- and against CO3-- ) give the same result you can be pretty sure that the
oxalate and the carbonate are both pure.
If nothing else, you will end up well practised in titration.
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blogfast25
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| Quote: Originally posted by unionised | I wonder what % starch (or whatever) you would need to add to the pure stuff to get it to go brown on heating.
I wonder if the stuff might still meet the USP or BP spec.
You can titrate it directly if you like, but it's also possible to roast it to get the carbonate. .
If nothing else, you will end up well practised in titration. |
Let us not misunderstand what Food Grade means: purity is not implied. It's possible to market mixtures of food grade products as a food grade
product. Ferricyanide for instance is an E number and a frequent adjuvant in various food products.
Pyrolysis of sodium oxalate may be a text book thingy but purity of the resulting sodium carbonate is far from guaranteed, IMO. I think you're
over-thinking the PM problem.
Agreed of course with the last point. 
[Edited on 25-2-2015 by blogfast25]
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