Sciencemadness Discussion Board - Standardisation from scratch?

Of course non-analytical grades might contain impurities that can interfere with different reactions, but is it possible to get reasonable standards with a home setup?

Materials I use as Primary Standards in a home set up:

Thrice recrystallized sodium carbonate, dehydrated. Acid/base standard.

Potassium hydrogenphthalate ('KHP'):

Not very expensive, easy to recrystallize. Acid/base standard.

Copper conductor:

Highly pure copper for primary standard in iodometry.

Ebay 99.99 % Zinc:

Primary standard for EDTA titrations.

Home synthesised Mohr's Salt, recrystallized:

Fe(II) redox primary standard.

[Edited on 18-2-2015 by blogfast25]

Fulmen - 18-2-2015 at 08:36

I checked out the sulfamic acid, but it's a bit too complex just yet.

Working in labs I've always taken simple chems and titrations for granted, now I'm starting to miss the possibility of doing accurate chemistry. I need to acquire analytical chems and equipment, acid/base titrations seems like a good place to start. Where do I start? Buying reagent grade would be the most sensible, but they are expensive and hard to get. Pharmaceutical grades are easier to get hold of, and I assume they would be pure enough for my uses, but over the years I've become increasingly paranoid about flagging my chemistry unless I really have to.

I think I will attempt to assay sodium carbonate. It should be simple enough to recrystallize, and could perhaps be tested by precipitation with calcium. Any better suggestions?

blogfast25 - 18-2-2015 at 08:44

I checked out the sulfamic acid, but it's a bit too complex just yet.

It should be simple enough to recrystallize, and could perhaps be tested by precipitation with calcium. Any better suggestions?

Problem is that CaCO3 precipitates suffer from occlusion. Not a great thing for gravimetry.

Preparing pure Na2CO3 from a 'good' grade must lead to high purity by recrystallization(s).

[Edited on 18-2-2015 by blogfast25]

DJF90 - 18-2-2015 at 08:48

I checked out the sulfamic acid, but it's a bit too complex just yet.

How do you mean? It is readily available OTC as a toilet cleaner. The purification IIRC is just several recrystallisations.

Also, potassium dichromate is a suitable primary standard for redox titrations, as is thiourea. Availability may vary on location.

[Edited on 18-2-2015 by DJF90]

blogfast25 - 18-2-2015 at 08:52

I've never heard of sulfamic acid as a PM, but I don't mean anything by that.

Fulmen - 18-2-2015 at 08:53

Never seen OTC sulfamic acid around here, but I must confess I've never looked for it either. I'll keep an eye out for it.

Sulaiman - 18-2-2015 at 11:38

In my teens I enjoyed chemistry as a hobby,
40+ years later I've just re-started less than a year ago,
and my first 'standard' was Sodium Carbonate.
Stated as 99.3% pure by the vendor.
The sodium Carbonate lost 0.56% weight on de-hydrating.
eBay scales, calibration weights, pH meter, and used volumetric flask and burette.
The pH meter is great if you leisurely want to draw a titration curve for one acid-base titration
(e.g. Sodium Carbonate + Sulphuric Acid)
but I found it so slow that I gave up cross-titrations (to check my errors).
I'll do my titrations again with Phenolphthalein or similar.
So, in my opinion, it's not too difficult to get results with less than 1% error
with a small investment in equipment that will be generally useful,
but for small-scale acid/base titrations (hobby, cheap) don't bother with a pH meter,
acid/base indicators are good enough I believe.

For daily qualitative chemistry,
it generally seems adequate to use the concentration and purity stated by the supplier
and dilute or use accordingly.

That's my impression so far, others more experienced may differ !

[Edited on 18-2-2015 by Sulaiman]

Fulmen - 18-2-2015 at 12:57

Zyklon: Missed your post until now, looks like I have some reading to do. Thank you, I love them old books.

papaya - 18-2-2015 at 13:05

Materials I use as Primary Standards in a home set up:

Thrice recrystallized sodium carbonate, dehydrated. Acid/base standard.

[Edited on 18-2-2015 by blogfast25]

I'm sorry, but Na2CO3 can form different hydrates if I remember properly, isn't it? How do you use it as a standard then?

EDIt: Only now I noticed "dehydrated", ok then how do you dehydrate it completely, how high the temperature must be? Also what indicator is best with it if trying to titrate strong acids?

[Edited on 18-2-2015 by papaya]

S.C. Wack - 18-2-2015 at 13:22

http://pac.iupac.org/publications/pac/18/3/0443/pdf

papaya - 18-2-2015 at 14:02

thanks!

Fulmen - 18-2-2015 at 15:34

I've looked at both, good finds. The IUPAC-paper used reagent grade bicarbonate as a source, so that doesn't help much. It does mention a silver standard for hydrochloric acid, but again based on purity graded silver. This can of course not regress in infinity, question is whether or not it ends at something useful to a home lab.

Looks like the only practical route is to purify something like sodium carbonate by recrystallizing a few times and do qualitative tests for "the usual suspects" of contaminates listed in "Chemical Reagents".

blogfast25 - 18-2-2015 at 15:57

Quote: Originally posted by papaya

Dehydrate at 200 C until constant weight is achieved.

blogfast25 - 18-2-2015 at 16:07

Quote: Originally posted by Sulaiman

In my teens I enjoyed chemistry as a hobby,
40+ years later I've just re-started less than a year ago,
and my first 'standard' was Sodium Carbonate.
Stated as 99.3% pure by the vendor.
The sodium Carbonate lost 0.56% weight on de-hydrating.
eBay scales, calibration weights, pH meter, and used volumetric flask and burette.
The pH meter is great if you leisurely want to draw a titration curve for one acid-base titration
(e.g. Sodium Carbonate + Sulphuric Acid)
but I found it so slow that I gave up cross-titrations (to check my errors).
I'll do my titrations again with Phenolphthalein or similar.
So, in my opinion, it's not too difficult to get results with less than 1% error
with a small investment in equipment that will be generally useful,
but for small-scale acid/base titrations (hobby, cheap) don't bother with a pH meter,
acid/base indicators are good enough I believe.

For daily qualitative chemistry,
it generally seems adequate to use the concentration and purity stated by the supplier
and dilute or use accordingly.

That's my impression so far, others more experienced may differ !

I broadly concur with that.

Don't expect miracles at the home lab level and for the most part they aren't needed anyway.

macckone - 18-2-2015 at 16:12

Sodium carbonate is easy to purify by crystallization.
It has a really good solubility curve. It is usually standardized
with potassium hydrogen phthalate for the highest precision.
But it can be purified and dehydrated and used on its own.

S.C. Wack - 18-2-2015 at 17:04

IThe IUPAC-paper used reagent grade bicarbonate as a source, so that doesn't help much.

Actually it wasn't posted *just for you*, and BTW I did not suspect that there was anything I could do to help you in any way. I don't expect that you have a Pt dish either.

It would be an interesting experiment to see if the highest grade of bicarbonate that Sigma-Aldrich sells gives better results in your hands than the cheapest off-brand of USP baking soda straight from the box.

Fulmen - 18-2-2015 at 17:53

You know what, I'm all out of Pt-dishes right now :-)

Interesting experiment, I agree. I wouldn't put money on me telling them apart, I'm not expecting miracles here. It's just about doing it right, so you know instead of believe.
But everything boils down to accuracy, there are no absolutes here.

Zombie - 18-2-2015 at 18:01

I found a PDF version of the 6th edition "Purification of Laboratory Chemicals" here...

http://www.pyrobin.com/files/purification%20of%20laboratory%...

Fulmen - 19-2-2015 at 07:40

Wow, lot of good info coming forth here.

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.

blogfast25 - 19-2-2015 at 08:07

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.

Fulmen - 19-2-2015 at 08:30

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.

Dr.Bob - 19-2-2015 at 10:05

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

blogfast25 - 19-2-2015 at 10:11

Question is, how far can one get with simple techniques like crystallization?

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]

unionised - 19-2-2015 at 12:24

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.

Fulmen - 19-2-2015 at 13:22

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.

blogfast25 - 19-2-2015 at 18:47

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]

Dr.Bob - 19-2-2015 at 19:47

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.

Fulmen - 20-2-2015 at 02:20

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.

Chemosynthesis - 20-2-2015 at 04:01

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.

Fulmen - 20-2-2015 at 04:13

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.

blogfast25 - 20-2-2015 at 05:15

Quote: Originally posted by Dr.Bob

I'm not stupid enough to weigh mgs of acid, [...]

No one is: it's the ACCURACY of 1 mg or better that is needed, not 'weighing a few mg of acid'.

blogfast25 - 20-2-2015 at 05:16

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.

Nope. Trust me, use the anh. Na carbonate as PM. NO ONE uses azeo HCl.

blogfast25 - 20-2-2015 at 05:22

First order of business will be to recrystallize Na2CO3 3-4 times, I think that will be good enough for general labwork.

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]

Fulmen - 20-2-2015 at 05:35

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...

DJF90 - 20-2-2015 at 05:45

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.

blogfast25 - 20-2-2015 at 06:01

Quote: Originally posted by DJF90

Tris (http://en.m.wikipedia.org/wiki/Tris) is a suitable primary standard base if you'd rather use that over sodium carbonate.

I've used TRIS as a buffer, I think it might give 'spongy' end points because of that.

DJF90 - 20-2-2015 at 06:16

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.

blogfast25 - 20-2-2015 at 06:34

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.

Fair enough.

Fulmen - 20-2-2015 at 08:42

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.

macckone - 22-2-2015 at 19:08

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

UC235 - 22-2-2015 at 19:50

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.

Fulmen - 22-2-2015 at 22:44

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.

blogfast25 - 23-2-2015 at 05:14

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]

Fulmen - 23-2-2015 at 06:09

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.

papaya - 23-2-2015 at 06:32

On oxalic acid: can't dihydrate be used directly, or water content is a subject of change depending on age?

blogfast25 - 23-2-2015 at 06:55

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]

DistractionGrating - 23-2-2015 at 13:14

Thrice recrystallized sodium carbonate, dehydrated. Acid/base standard.

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.

I'm curious how much greater purity can be achieved through recrystallization vs. simply decomposing sodium bicarbonate?

Fulmen - 23-2-2015 at 14:54

Decomposing bicarbonate can't magically remove non-volatile impurities.

DistractionGrating - 23-2-2015 at 16:45

Decomposing bicarbonate can't magically remove non-volatile impurities.

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.)

blogfast25 - 23-2-2015 at 17:35

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.

DistractionGrating - 23-2-2015 at 18:21

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.

blogfast25 - 23-2-2015 at 19:12

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...

DistractionGrating - 23-2-2015 at 21:40

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]

DistractionGrating - 23-2-2015 at 21:41

The simpler truth is that recrystallizing sodium carbonate works well.

BTW, thank you for this. It is my takeaway.

Fulmen - 23-2-2015 at 23:41

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?

unionised - 24-2-2015 at 13:31

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.

blogfast25 - 24-2-2015 at 13:53

Food grade stuff is probably comparable.

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.

Praxichys - 24-2-2015 at 13:53

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]

blogfast25 - 24-2-2015 at 13:55

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)?

That's the general principle, yes (insoluble combinations notwithstanding).

unionised - 25-2-2015 at 13:05

Food grade stuff is probably comparable.

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.

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.

blogfast25 - 25-2-2015 at 13:43