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Author: Subject: Easy to obtain standards for titrations and other analytical methods
j_sum1
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[*] posted on 28-2-2020 at 19:52
Easy to obtain standards for titrations and other analytical methods


Good analysis is part and parcel of good chemistry – figuring out what you have got and how much theer is. But qualitatively measuring concentrations can be painful as an amateur. We are basically limited to 3sf precision in any measuring equipment that we own. We resort to density measurements to get approximate concentrations of many liquids. And that good old method of titration is hamstrung by the lack of good standard solutions. (Besides, it can be tedious, time-consuming and require meticulous practice and patience that I, at least, am often lacking.)

With that dismal introduction I propose this thread as a place to share ideas regarding good, reliable standards that are easily accessible to the amateur. I begin with the following:

Oxalic acid
This can be either the anhydrous or dihydrate or some mix between. But it can be easily dehydrated using toluene and a dean-stark trap or improvised alternative. The acid itself is strong enough that it makes for ok titrations. I have also heard of people oven-drying OA to get the anhydrous form or recrystallising from water to obtain the dihydrate but I am not sure of the accuracy of these methods. (I have done reactions with OTC oxalic acid dihydrate straight out of the box and got quantitative results. Reaction with excess Fe2+ solution to produce and isolate iron oxalate is a great student experiment for confirming stoichiometry.)


Sodium carbonate
This has numerous hydrated forms so is completely unsuitable out of the box. But it is possible to make good quality anhydrous Na2CO3 by heating NaHCO3 and liberating H2O and CO2. Oven drying works: 2h at high temperature with occasional stirring. Heating first in a beaker over a flame will speed up the process and is slightly less messy. (You do get a white film on the inside of the oven.)

Copper sulfate
High purity copper sulfate pentahydrate is easy to get. This is easily decomposed to anhydrous CuSO4 with oven baking: again with stirring. I have heard of this being used as a standard for redox titration but I forget the details.


Any others?


The past couple of days I have been thinking about this and wondered at the possibility of using pharmaceutical products such as aspirin and vitamin C tablets which are presumably manufactured with tight tolerances. But I am not sure how reasonable this is.

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Sulaiman
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[*] posted on 28-2-2020 at 20:08


For REDOX,
potassium permanganate is quite pure, and not too expensive,
and can be bought via eBay as >= 99.4% and >=99.9% purity.

relatively cheap and pure iodine and sodium thiosulphate are available via eBay.
_______________________________________________________________
AFAIK, until recently USP weight measurements had classes of accuracy (e.g. 1% and 0.1%) but now USP only requires manufacturers to work to what they consider suitable accuracy,
which for benign chemicals (e.g. vit.C) could be - not accurate at all.
Qualified pharmaceutical chemists probably know more on this :D

[Edited on 29-2-2020 by Sulaiman]




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Fery
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[*] posted on 28-2-2020 at 23:25


Potassium hydrogenphtalate for acidobasic titrations.
Zn metal for EDTA complexometric titrations (must be dissolved in an acid after weighing it of course).
(NH4)2Fe(SO4)2 . 6 H2O for redox titrations
KBrO3 for standardisation of Na2SO3
NaCl for standardisation of AgNO3 (precipitation titration)




If there is a heaven, it seems not to be materially based. Does chemistry exist there and if yes, how does it look like? Are there good souls well supplied with laboratory equipment, glass, chemicals and information?
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[*] posted on 29-2-2020 at 02:37


Sulfamic acid can be used as a primary standard. See attached pdf for a purification procedure

Attachment: Purification of sodium carbonate and sulfamic acid.pdf (240kB)
This file has been downloaded 119 times

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[*] posted on 29-2-2020 at 04:03


A quick mention for borax and cream of tartar as pH references.
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DavidJR
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[*] posted on 29-2-2020 at 09:35


In order to use sodium carbonate as a primary standard you really have to oven dry it first even if you have nominally anhydrous material.
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clearly_not_atara
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[*] posted on 29-2-2020 at 10:32


Potassium tetraoxalate - KH3(C2O4)2 is easily precipitated from mixtures of almost any soluble potassium salt and oxalic acid. It is used as a pH calibration standard because it has a low solubility of 2.5% w/w or just about 0.2 molar although the more common solution used is half that concentration -- 0.1 molar with a pH of 1.7. I think I have a picture of the prep somewhere; you just combine solutions and filter, couldn't be easier. The dihydrate is stable and non-hygroscopic.

More importantly, because the saturated concentration is low and the protons are all fairly acidic (the last one is lost at about pH 5.5) it can be used to titrate low concentrations of alkali in conjunction with an indicator.

[Edited on 29-2-2020 by clearly_not_atara]




[Edited on 04-20-1969 by clearly_not_atara]
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nezza
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[*] posted on 29-2-2020 at 11:40


You can buy volumetric standard acid (HCl) and base (NaOH) as 0.1M solutions on ebay.



If you're not part of the solution, you're part of the precipitate.
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[*] posted on 29-2-2020 at 15:27


For acids:

Anhydrous Na2CO3

NaHCO3 or KHCO3

Borax

Sodium oxalate

For bases:

Oxalic acid

Potassium hydrogenphtalate

Benzoic acid

For manganometry and cerimetry:

Oxalic acid or sodium oxalate

For any reducing agent:

K2Cr2O7

KIO3

For TiCl3, SnCl2 and CrCl2 also NH4Fe(SO4)2.12H2O

For iodine:

Resumblimation

Sodium thiosulfate solution of known concentration (you can make your own and standardize it).

(As2O3 - but this isn't common and it's for experienced chemists)

For chelatometry:

MgSO4.7H2O

ZnSO4.7H2O

Bismuth or copper metal

For argentometry nad mercurymetry:

NaCl

For K4[Fe(CN)6]:

ZnSO4.7H2O

[Edited on 29-2-2020 by Bedlasky]
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j_sum1
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[*] posted on 29-2-2020 at 16:33


Lots of good ideas here. But rather than just a list, I would love procedures for converting otc products or just using them as is.

I did not know about saturated potassium bitartrate as a pH standard. That is cool. Can it also be used as an acid titration standard? It would be extremely useful for an education project I have if that is the case.
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[*] posted on 29-2-2020 at 19:19


If you use a burette with 0.1 mL divisions, use more than 10 mL of titrant and interpolate between the divisions to the nearest 0.01 mL (using a Mark I eyeball, aided or unaided), you can get 4 SF. And 3 SF is still 1 part per thousand, which would make even a purist happy (most of the time :D). To standardize strong acids/bases, KHP is the best, and making up a liter or so of each reagent that you have to standardize should last a reasonable time.

The first sentence of your original post is so true! I suspect many posters perform some qualitative/quantitative tests to characterize their products, but they often don't report that. I'm presently trying to fit a PID + type T thermocouple to an old Fisher-Johns melting point apparatus so I can dispense with my Thiele tube. Breaking thermometers is annoying and original replacements for the F-J apparatus that I have found are exorbitantly priced.
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[*] posted on 1-3-2020 at 08:06


Quote: Originally posted by j_sum1  
Lots of good ideas here. But rather than just a list, I would love procedures for converting otc products or just using them as is.

I did not know about saturated potassium bitartrate as a pH standard. That is cool. Can it also be used as an acid titration standard? It would be extremely useful for an education project I have if that is the case.


For standardization of acids I often use carbonates/bicarbonates. As indicator I use methyl orange. Titrate your standard with acid (HCl, H2SO4) and when indicator turns in to the orange, then boil solution to get all of carbon dioxide from solution. Solution turns back to the yellow. After that let it cool to the room temperature and titrate it again with acid to the pink colour.

Borax have some advantages - higher molar mass, so there is smaller error, and you don't need boil solution, because boric acid is weaker acid than carbonic and doesn't interfere determination.

For standardization of bases (NaOH, KOH) I use oxalic acid. You can use phenolphtalein or methyl orange as indicators. With phenophtalein it's very simple titration - end point is when indicator turns in to the pink.

For methyl orange exist different technique. You titrate oxalic acid with hydroxide until indicator turns in to the orange. Then you add 10ml of 20% CaCl2 - this caused formation of insoluble calcium oxalate and HCl. Solution turns back in to the pink and you titrate it to the yellow colour.

Interesting alternative can be benzoic acid for standardization of KOH in ethanol which can be used for determination of acidity number of oil or for testing of purity of organic acids. Phenolphtalein is suitable indicator for this.

For standardization of KMnO4 I use oxalic acid. Oxalic acid reacts only slowly with KMnO4, so solution must be hot or instead of heating you can use MnSO4 as catalyst. Indicator is KMnO4 itself.

For standardization of TiCl3

I never did this, but I plan it with ammonium ferric sulfate as standard. There is two possible indications:

Methylene blue: Until solution turns from blue to colourless

KSCN: Until solution turns from red to colourless

For standardization of Na2S2O3

Potassium dichromate or iodate is dissolved in water, then you add sulfuric acid, excess of KI and let it stand for 10 minutes. Then you titrate this solution until it turns pale yellow. After that add some starch solution, which forms with iodine dark blue complex. Then titrate it until solutions turns in to the colourless.

Similar procedure is used for standardization of iodine (without using of oxidizing agent).

For standardization of AgNO3

Simple titration of NaCl with AgNO3 solution. As indicator you can use K2CrO4 (which forms with Ag+ reddish brown precipitate) or fluorescein (end point is when fluorescence disappear).

For standardization of Hg(NO3)2

Simple titration of NaCl with Hg(NO3)2 solution. As indicator you can use sodium nitroprusside which forms with excess of Hg nitrate white precipitate.

For standardization of Na2H2EDTA

Simple titration of ZnSO4 or MgSO4. Indicator is eriochrome black T. Titration must be done at pH = 10.

Some procedures are here:

https://www.gitam.edu/departments_cms/assets/uploads/syllabu...

http://egyankosh.ac.in/bitstream/123456789/43339/1/Exp-10.pd...

http://www.titrations.info/
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[*] posted on 2-3-2020 at 00:05


Quote: Originally posted by CharlieA  
If you use a burette with 0.1 mL divisions, use more than 10 mL of titrant and interpolate between the divisions to the nearest 0.01 mL (using a Mark I eyeball, aided or unaided), you can get 4 SF. And 3 SF is still 1 part per thousand, which would make even a purist happy (most of the time :D).

(from a transiently unhappy metrology purist)

When I do (e.g. acid:base) titrations the limits to accuracy are;
1 the accuracy of the molarity of the titrant
2 titration equivalency may not be pH7
3 equivalency pH determination errors
4 the minimum dispensable quantity is one drop ... typically between .03ml and .05ml
5 my 50ml class-A burettes have 0.1ml divisions and are supposed to be accurate to 0.05ml
6 each drop of titrant adhering to the inner wall of the burette is also about 0.05ml error
7 near 20oC the volumetric thermal expansion of water is about 207 ppm/oC so 5oC error = 0.1% volume error

This is why I find it near impossible to achieve 0.1% absolute accuracy
but fairly easy to maintain 1% accuracy




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[*] posted on 2-3-2020 at 17:27


Quote: Originally posted by Sulaiman  
Quote: Originally posted by CharlieA  
If you use a burette with 0.1 mL divisions, use more than 10 mL of titrant and interpolate between the divisions to the nearest 0.01 mL (using a Mark I eyeball, aided or unaided), you can get 4 SF. And 3 SF is still 1 part per thousand, which would make even a purist happy (most of the time :D).

(from a transiently unhappy metrology purist)

When I do (e.g. acid:base) titrations the limits to accuracy are;
1 the accuracy of the molarity of the titrant
2 titration equivalency may not be pH7
3 equivalency pH determination errors
4 the minimum dispensable quantity is one drop ... typically between .03ml and .05ml
5 my 50ml class-A burettes have 0.1ml divisions and are supposed to be accurate to 0.05ml
6 each drop of titrant adhering to the inner wall of the burette is also about 0.05ml error
7 near 20oC the volumetric thermal expansion of water is about 207 ppm/oC so 5oC error = 0.1% volume error

This is why I find it near impossible to achieve 0.1% absolute accuracy
but fairly easy to maintain 1% accuracy


Sulaiman, I am truly impressed with your attention to detail! But since I'm just a retired chemist (I last practiced professionally over 35 years ago, so now I consider myself an amateur chemist), I would be satisfied with 1% accuracy, but that is just me.
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[*] posted on 3-3-2020 at 04:45


sulaiman got a point, purity of the standard is not the only thing.
in uni i did a quantitative analytical chemistry class, we used analytical balances (+0.1mg) to weight our standard compounds, dried for 24 hours and kept always in a desiccator, using 50ml class A burettes with schellbach line for easier identification of the meniscus. we tried to not use approximate equations (they can get pretty long lol) to minimize mathematical errors. it was my first time using a burette so i wasn't an expert, but at the end of the lab i did around 35 titrations, being really carefull, slow (the burettes were degreased but still a film of liquid would form on the walls, a fast dripping rate and a quick look at the level of the solution would give an error, as with time the adhered solution would slowly go down the wall and reach the top of the meniscus, giving you now a bigger volume. if you are fast you think to have used more solution than what you actually used) i would still get 1-2% of error minimum. then you have to think that many solutions are not stable, permanganate for example, yea you can get that pretty pure, but it slowly decomposes in solution.

then there's temperature... and if you don't have air conditioning in your lab to keep everything at a constant temperature, it can be a source of error (even if not huge most of the time)





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