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blogfast25
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Quote: Originally posted by Fulmen  | | There is one question that hasn't been discussed yet, how do one establish a primary standard? I'm thinking spectroscopy is the only way, chemical
analysis can only compare to other standards or detect individual contaminates. And since there is virtually an infinite number of possible
contaminates, you can't hope to eliminate all. |
No matter how you look at it, it's a kind of chicken-or-egg situation. In practical terms I would simply buy ONCE, a Primary Na2CO3 or KHP from a
reputable supplier and compare it to your home made standard in a real life acid/base titration.
I would caution also against over-emphasising the importance of super-dooper, near-absolute purity of PMs. A 99,99999 % standard is useless in the
hands of a poor analyst, while a decent analyst will get decent results with a 99.9 % standard. Technique and knowledge of the method matter a lot
too.
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Fulmen
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That last question was intended in general, not applicable to this project. I just started wondering about this chicken/egg-problem as you so aptly
put it.
A 99.9% standard would be more than enough for any of my needs, even 99% would do. But why set low goals? Question is, how far can one get with simple
techniques like crystallization? I know, how long is a rubber band? Chicken and egg all over.
I did find a bottle of Puris K2CO3 (min 98%) in the cupboard, I guess I could use that as a comparison. Not sure how pure it will actually be though.
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Dr.Bob
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The amount of water in a chemical can change the composition much more than a trace of other salts. If you have K2CO3, it may have traces of
Na2CO3, KHCO3, KCl, K2SO4, KOH, and other salts in it, but most will be basic still, so the effect on a titration may not be huge. But many common
chemicals can absorb enough water to change the weight by a good bit, so you wan to choose chemicals which do not easily for hydrates if possible.
So solid NaOH or KOH are poor choices, as they absorb water, CO2, and other acids from the atmosphere. An old bottle of NaOH may contain 10% or more
Na2CO3, at least that is what my analytical teacher used to teach. But K2CO3 is less hygroscopic than NaOH, so it is a good start.
Benzoic acid is quite stable, but not a strong acid. So you want to pick something that is easy to get and keeps well, like maybe sulfuric acid, if
you can get a fresh bottle of that, it will not change much with time if kept sealed. Or if you can get a 1M H2SO4 stock solution, that will keep
very well, as it is already in water, so it will not absorb much from the atmosphere. The good news if that if you learn how to do a titration with
mediocre standards, your technique would be the same with better ones and you will get better results. But for synthetic work, titrations are not
that common; whereas for analytical ones, there are ways to get good results with certain standards, but most are dependent on the assay. For home
or amateur work, 99% is usually fine.
Bob
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blogfast25
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Again, how long is a nice crystal? 
Only certain techniques for a given compound can give the answer to that.
I use the ones I use because I know they are used and because thermal recrystallization is easy to do. But STRICTLY speaking I don't know the purity
of my PM(s).
| Quote: Originally posted by Dr.Bob | | So you want to pick something that is easy to get and keeps well, like maybe sulfuric acid, if you can get a fresh bottle of that, it will not change
much with time if kept sealed. |
Try and weigh conc. H2SO4 to a mg or better: it's hygroscopic! Bad example. At a minimum PMs have to be neither hygroscopic nor deliquiescent.
H2SO4 titrant solution are an example of why standardisation is necessary because they can't be prepared accurately enough.
[Edited on 19-2-2015 by blogfast25]
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unionised
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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.
These days sulphamic acid is probably a good choice but, if you can't find it you can use the azeotrope of HCl in water (obviously, you need to be
happy distilling acid to do that).
There are tables of %HCl w/w vs atmospheric pressure.
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Fulmen
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Azeotropic HCl isn't a bad idea, I kinda dismissed it offhand but I now realize it's worth considering. I'll do a bit of reading and see what I can
come up with.
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blogfast25
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| Quote: Originally posted by Fulmen | | Azeotropic HCl isn't a bad idea, I kinda dismissed it offhand but I now realize it's worth considering. I'll do a bit of reading and see what I can
come up with. |
It does have the problem of high volatility though. Not what you routinely want from a PM...
HCl titrant solutions are mostly standardised because it's difficult to prepare a solution of very precisely known HCl concentration.
[Edited on 20-2-2015 by blogfast25]
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Dr.Bob
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I'm not stupid enough to weigh mgs of acid, but you can make a 1 or 2 Normal solution by the liter and it will keep stable as a solution for months.
It is just an example of an acid that is easier to get than many others. 1N HCl could be made from fresh acid as well, but HCl goes down in
concentration quickly as well. But using K2CO3 or Na2CO3, you could then titrate a 1N acid solution, albeit as a secondary standard, but it will
stay stable for a long time. For people with only access to OTC chems, that is likely sufficient enough.
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Fulmen
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Blogfast: It can't be THAT hard? The azeotrope should be both stable and have an accurate composition, so weighing it into a volumetric flask should
produce an equally accurate solution if freshly prepared. This would then be used for assaying the carbonate as a secondary standard.
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Chemosynthesis
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Blogfast's suggestion of buying a standard of 'known' purity and comparing is definitely a good one if you don't find what you consider trustworthy
purification procedure with enough detail to replicate and determine % purity.
One reason for probably not needing a 99.99*% pure standard is equipment. Even if you have good enough technique from your experience in labs that are
equipped with analytical equipment, do you really think your home equipment is as good? Perhaps it is, and perhaps it is better maintained. (Other
than new labs, analytical, and pharmaceutical chem, I question that in most cases.)
As far as crystallization goes, you can get a very clean product if you don't mind yield loss. At least some instances of repeated crystallization
give NMR worthy purity at at least 300MHz. I knew a gentleman who did not trust chromatography (ancient man) and required his students to re
crystallize instead. They spared themselves the joy of maintaining hundreds of small test tubes, but had to wait twice as long to get results.
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Fulmen
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Buying a good standard is of course the simplest approach, but not the most challenging or educational one. I'm starting to miss analytical work, so
this will be for my own satisfaction rather than simply to serve a need. I have the experience, but not the equipment needed for 99,99% work, 99-99,9
would be more realistic.
First order of business will be to recrystallize Na2CO3 3-4 times, I think that will be good enough for general labwork. But I think the HCl-route is
interesting enough to warrant a trial, comparing those two should produce some interesting data.
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blogfast25
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No one is: it's the ACCURACY of 1 mg or better that is needed, not 'weighing a few mg of acid'.
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blogfast25
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| Quote: Originally posted by Fulmen | | Blogfast: It can't be THAT hard? The azeotrope should be both stable and have an accurate composition, so weighing it into a volumetric flask should
produce an equally accurate solution if freshly prepared. This would then be used for assaying the carbonate as a secondary standard.
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Nope. Trust me, use the anh. Na carbonate as PM. NO ONE uses azeo HCl.
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blogfast25
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| Quote: Originally posted by Fulmen | First order of business will be to recrystallize Na2CO3 3-4 times, I think that will be good enough for general labwork.
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You'll find that after 2 crystallisations of Na2CO3 you'll have to help it along because the third solution supercools! So a seed crystal from a
previous crystallisation is needed to get the stuff to crystallise out.
That is of course a sign of purity in and of itself!
The idea of using something that practically fumes in air as a PM seems a little absurd to me.
[Edited on 20-2-2015 by blogfast25]
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Fulmen
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No one uses azeo HCl, I get that. Why would you when p.a. reagents are available? But for the sake of science, why wouldn't it work? It seems to be an
accepted method, albeit old. I'm having trouble finding actual methods and accurate data of the azeotrope vs pressures, but I'm sure that will turn up
in time.
As for supercooled solutions, I know exactly what that is like. Ever tried recrystallizing calcium nitrate? Son of a b...
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DJF90
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Every volumetric solution should be made using a primary standard, or standardised using one. Azeotropic HCl isn't satisfactory as slight difference
in pressure affects the concentration of the azeotropic distillate by a reasonable amount (IIRC). 6M is far too strong for a working standard anyway,
so I'd suggest just diluting conc. HCl to working concentration and standardising against standard base. Tris (http://en.m.wikipedia.org/wiki/Tris) is a suitable primary standard base if you'd rather use that over sodium carbonate.
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blogfast25
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I've used TRIS as a buffer, I think it might give 'spongy' end points because of that.
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DJF90
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The analytical dept at a company I used to work at used Tris for ALL acid standardisations. If it is good enough for them, it's good enough for me.
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blogfast25
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| Quote: Originally posted by DJF90 | | The analytical dept at a company I used to work at used Tris for ALL acid standardisations. If it is good enough for them, it's good enough for me.
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Fair enough.
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Fulmen
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I will use sodium carbonate for now, just started recrystallizing a batch. However it would be fun to find some material om azeo HCl as I've found
several references to it's use as a primary standard.
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macckone
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Oxalic acid should make a good primary standard.
It is not too difficult to purify and is easily made from
sugar and nitric acid in the presence of a catalyst.
If forms a dihydrate but that is decomposed by boiling
with carbon tetrachloride.
So with carbon tetrachloride, nitric acid, sugar and distilled
water you can have an extremely pure acid.
https://archive.org/stream/systematicstudyo00bluh/systematic...
http://www.orgsyn.org/demo.aspx?prep=CV1P0421
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UC235
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Consider potassium bitartrate as a primary standard for base standardization. Pure tartaric acid is readily available from wine and beer suppliers to
adjust the pH of grape juice before fermentation. The bitartrate has considerably lower solubility than tartaric acid, dipotassium tartrate, sodium
bitartrate, disodium tartrate, and rochelle salts and can therefore be readily crystallized from water solution in high purity.
While the solubility is never high, it does increase in solubility tenfold between 0C and 100C and so repeated crystallizations from distilled water
can provide a product of very high purity. The resulting white crystals are not a hydrate, are nonhygroscopic and are stable in storage.
Potassium Bitartrate is currently listed as a primary standard for pH by NIST. A saturated solution in water has a pH of 3.557 at 25C, so it can also
be used to calibrate a pH probe.
Unlike KHP (potassium acid phthalate), "KHT" has poor water solubility however, so either a very large volume of titration solution is needed, or good
stirring in the presence of finely ground solid KHT is needed. As with KHP, phenolphthalein should be sufficient to detect the endpoint.
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Fulmen
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Interesting. While the solubility is horrible (6-60g/l) it shouldn't be impossible to purify enough for standard use. And I just realized I actually
have 40g of food grade lying around.
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blogfast25
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| Quote: Originally posted by Fulmen | | Interesting. While the solubility is horrible (6-60g/l) it shouldn't be impossible to purify enough for standard use. |
If 6 g/l is the solubility at RT, that's only 0.03 M. Quite low for a standardising titration. 0.1 M gives a stronger pH jump, so a clearer end-point.
Oxalic acid is very OTC (no need to synthesize it) but I found it hard to dehydrate (the dihydrate) to constant weight, for that reason I abandoned
it.
[Edited on 23-2-2015 by blogfast25]
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Fulmen
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That is a good point, but I think it is worth testing. I'll take a look at oxalic acid as well, but if it's hard to dry out properly it probably isn't
the best choice.
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