Sciencemadness Discussion Board

Extraction of ASA from Pills

hacker - 23-12-2018 at 22:37

I extracted acetylsalicylic acid from aspirin with acetone. Each bottle contained 500 pills with 325 milligrams of ASA each.

I made a huge mess with the first two bottles. I ground the pills to a fine powder in a coffee grinder. Then I put them into a beaker with 600 milliliters of acetone. I brought the acetone to a boil with stirring then filtered. Some of the binders went through the filter paper. The ASA condensed in the filter funnel and clogged it several times. After everything filtered, I put the mixture into a beaker and washed everything with acetone. I put the beaker into the freezer. After it became ice cold, I filtered the crystals and put them into a big dish.

The filtrate had some white powder in it. I don't know what it was. My guesses are microcrystalline cellulose, starch, and titanium dioxide. I decanted the filtrate and left the powder behind. Then I topped off the acetone to 600 milliliters and extracted one bottle of pills. I let most of the binders settle then filtered. The filtration was much easier the second time. I washed the beaker and filter funnel with a little bit of acetone. Then I used the acetone to wash the product. I cleaned the beaker, and put the filtrate into it. A small amount of white powder remained in the flask.

More crystals formed. I repeated the process until all of the pills had been extracted. The filtrate looked nasty. I put about 500 milliliters of water into it. No crystals formed. I discarded it.

I dried the aspirin crystals for a couple of days. I obtained 679.7 grams of dry ASA crystals for an extraction efficiency of 84%. I saved a small sample and hydrolized the rest to salicylic acid. The ASA is mostly pure but makes a slightly cloudy solution.

84% is a pretty bad extraction efficiency. It could have been higher if I had not discarded the final filtrate.

The purity is not great, but I think it is not bad. Almost all of the product dissolves in boiling water. A tiny bit of white powder does not dissolve.

I think I should try extracting the pills with methanol and precipitate the ASA with water.

Tsjerk - 24-12-2018 at 03:56

When I did this I didn't powder the pills, this way, when the ASA is soaked out the cellulose sort of stays coarse and the acetone is easily decanted off. Then rinse the cellulose with a bit more acetone and you should be able to obtain a higher yield.

[Edited on 24-12-2018 by Tsjerk]

arkoma - 24-12-2018 at 08:33

91% isopropyl alcohol is my go to for ASA.

CharlieA - 24-12-2018 at 17:57

You can get an idea of the purity of your ASA by titration. Then your estimate of the purity would have more credibility than just saying it seems to be pretty pure. Characterizing your product is behaving more like a chemist than a chef. This is not meant to belittle or criticize you, but it seems to me to be a prevalent practice on SM to mix A and B and say I got C, just because there is a procedure that says so, or I can write (a hopefully balanced) chemical equation that says A + B gives C. Unfortunately, characterizing products and analyzing them for purity is one of the tedious parts of chemistry, especially because as hobbyists we have little, if any, access to instrumentation that would make the characterization easier. Wet chemistry is tedious and requires a good selection of reagents.
Again, please don't take this as an attack on you. I promise this will be my last rant of the year.
Merry Christmas,
Charlie

Abromination - 24-12-2018 at 18:04

When I do this, I crush up my tablets, add the powder to water, decant off the filler junk from the top and recrystalize in isopropanol. I distill off the isopropanol and generally end with a very high yeild of decently pure ASA.

hacker - 24-12-2018 at 21:04

Quote: Originally posted by CharlieA  
You can get an idea of the purity of your ASA by titration. Then your estimate of the purity would have more credibility than just saying it seems to be pretty pure. Characterizing your product is behaving more like a chemist than a chef. This is not meant to belittle or criticize you, but it seems to me to be a prevalent practice on SM to mix A and B and say I got C, just because there is a procedure that says so, or I can write (a hopefully balanced) chemical equation that says A + B gives C. Unfortunately, characterizing products and analyzing them for purity is one of the tedious parts of chemistry, especially because as hobbyists we have little, if any, access to instrumentation that would make the characterization easier. Wet chemistry is tedious and requires a good selection of reagents.
Again, please don't take this as an attack on you. I promise this will be my last rant of the year.
Merry Christmas,
Charlie


Merry Christmas.

I agree with most of this. I would not use titration to check the purity. ASA can hydrolize to acetic acid. Titration would cause the purity to appear higher than it really is. I would dry to constant weight and check the melting point to verify purity.

Checking turbidity of a solution is a perfectly valid means of quick characterization.

I know the ASA is not completely pure. I do not care if it is completely pure. I only care if it is pure enough for the next step. I already hydrolized almost all of it to salicylic acid. The process went smoothly. I can say the remaining ASA is decently pure.

I do not want to waste chemicals. I will be more rigorous with the purity of the phenol.

CharlieA - 28-12-2018 at 17:13

Quote: Originally posted by hacker  
Quote: Originally posted by CharlieA  
You can get an idea of the purity of your ASA by titration....
Charlie


Merry Christmas.

I agree with most of this. I would not use titration to check the purity. ASA can hydrolize to acetic acid. Titration would cause the purity to appear higher than it really is. I would dry to constant weight and check the melting point to verify purity.

Checking turbidity of a solution is a perfectly valid means of quick characterization.

I know the ASA is not completely pure. I do not care if it is completely pure. I only care if it is pure enough for the next step. I already hydrolized almost all of it to salicylic acid. The process went smoothly. I can say the remaining ASA is decently pure.

I do not want to waste chemicals. I will be more rigorous with the purity of the phenol.



The titration starts with titrating the ASA to the acetic acid/acetate endpoint with standardized NaOH. The excess standardized NaOH is added to the solution and it is heated to hydrolyze the ASA to salicylic acid. The solution if then titrated with standardized HCl to determine how much of the NaOH was consumed. Each mol of ASA requires 2 mol NaOH. With this method, I analyzed some aspirin tablets with a 3% error, based on the calculated amount of ASA per tablet. The ASA I isolated from pills, had a relative error of 9.19 ppt.

If you search for BackTitration-lab4.pdf you will find a laboratory experiment used at Lasalle University for the analysis of aspirin. If you can't find it, U2U me with your email address and I will email you a copy.
Happy New Year,
Charlie

Murexide - 28-12-2018 at 19:10

Quote: Originally posted by CharlieA  
Quote: Originally posted by hacker  
Quote: Originally posted by CharlieA  
You can get an idea of the purity of your ASA by titration....
Charlie


Merry Christmas.

I agree with most of this. I would not use titration to check the purity. ASA can hydrolize to acetic acid. Titration would cause the purity to appear higher than it really is. I would dry to constant weight and check the melting point to verify purity.

Checking turbidity of a solution is a perfectly valid means of quick characterization.

I know the ASA is not completely pure. I do not care if it is completely pure. I only care if it is pure enough for the next step. I already hydrolized almost all of it to salicylic acid. The process went smoothly. I can say the remaining ASA is decently pure.

I do not want to waste chemicals. I will be more rigorous with the purity of the phenol.



The titration starts with titrating the ASA to the acetic acid/acetate endpoint with standardized NaOH. The excess standardized NaOH is added to the solution and it is heated to hydrolyze the ASA to salicylic acid. The solution if then titrated with standardized HCl to determine how much of the NaOH was consumed. Each mol of ASA requires 2 mol NaOH. With this method, I analyzed some aspirin tablets with a 3% error, based on the calculated amount of ASA per tablet. The ASA I isolated from pills, had a relative error of 9.19 ppt.

If you search for BackTitration-lab4.pdf you will find a laboratory experiment used at Lasalle University for the analysis of aspirin. If you can't find it, U2U me with your email address and I will email you a copy.
Happy New Year,
Charlie


Yes this is certainly required. Do not attempt direct titration (like I did just moments ago, based on this procedure: http://www.bellevuecollege.edu/wp-content/uploads/sites/140/...). The ASA takes very long to dissolve even in 50% ethanol and what you end up is a strong pink colour which fades over the next 5 minutes (repeat with each drop). The dissolving could be causing the reversal of the colour to transparent, or it could be the hydrolysis to acetic acid. Regardless, do not attempt the direct titration.

Now I will attempt the back titration procdeure. I will report on how it ends up.

@Charlie A Did you grind it to a powder as written or titrate each tablet seperately?

hacker - 29-12-2018 at 02:04

Quote: Originally posted by CharlieA  
The titration starts with titrating the ASA to the acetic acid/acetate endpoint with standardized NaOH. The excess standardized NaOH is added to the solution and it is heated to hydrolyze the ASA to salicylic acid. The solution if then titrated with standardized HCl to determine how much of the NaOH was consumed. Each mol of ASA requires 2 mol NaOH. With this method, I analyzed some aspirin tablets with a 3% error, based on the calculated amount of ASA per tablet. The ASA I isolated from pills, had a relative error of 9.19 ppt.

If you search for BackTitration-lab4.pdf you will find a laboratory experiment used at Lasalle University for the analysis of aspirin. If you can't find it, U2U me with your email address and I will email you a copy.
Happy New Year,
Charlie


I own two burettes. I will read the procedure. I may require more indicators.

Murexide - 29-12-2018 at 03:48

The procedure worked very well. I standardised HCl to Sodium carbonate and NaOH to HCl. The procedure was followed and the ASA amount was determined to within 0.5%. The heating with water bath was replaced with intermittent heating until bubbles were seen, hotplate turned off, and then back on. NaOH required happened to be 50 mL so 25mL pipettes were used. Final endpoint was a cloudy white as listed with a very slight pink trace.

CharlieA - 30-12-2018 at 17:23

@Murexide: I ground up 4 tablets with a mortar and pestle and then weighed (massed!) ~0.3 g samples, which I calculated should require more than 10 ml of titrant, to give me 4 SF for the volumes; but looking back at my laboratory manual notes, I don't know what I was thinking because the balance I was using at that time only measured to 0.01 g, so the 2 SF of the sample weight was the limiting factor in the calculations.

@hacker: the only indicator that you will need is phenolphthalein. What do you plan to do with the phenol? (Just curious.)
Just a thought about my philosophy of using the purest chemicals that you can get or prepare youself. Let's assume 3 step synthesis: A -> B -> C _> D, where A is the starting material B & C are intermediates that are prepared and isolated, and D is the product.

a) assume A, B, C, and D are all isolated and are 100% pure, but the yield is just 0.9 (()%). The overall yield is then 0.9 x 0.9 x 0.9 = 0.7 or 70%.
b) now assume the same conditions but the yield in each step is 70%; the overall yield will be just 50%.

Now assume that the yields are the same in each scenario, and the purities of A, B, C, and D are 90% and 70 % respectively. In the first case you get a 90% yield of 70% pure product; in the second case you get a 50% yield of 50% pure product.
I think this is an over-simplified example, but I think the point is that one would prefer to get as high a yield of as pure a product as possible, both for reasons of efficiency and economy. Do I make sense or am I just whistling in the wind? (In my defense, it has been 40 years since I practiced chemistry professionally, and in getting back to some amateur chemistry in the last year I am very distressed at how much chemistry I have forgotten, let alone how much new chemistry has been discovered since then.)

P.S. I welcome constructive criticism, but please don't be too hard on an old man!:D


[Edited on 12-31-2018 by CharlieA]

Murexide - 30-12-2018 at 18:37

Quote: Originally posted by CharlieA  
@Murexide: I ground up 4 tablets with a mortar and pestle and then weighed (massed!) ~0.3 g samples, which I calculated should require more than 10 ml of titrant, to give me 4 SF for the volumes; but looking back at my laboratory manual notes, I don't know what I was thinking because the balance I was using at that time only measured to 0.01 g, so the 2 SF of the sample weight was the limiting factor in the calculations.

@hacker: the only indicator that you will need is phenolphthalein. What do you plan to do with the phenol? (Just curious.)
Just a thought about my philosophy of using the purest chemicals that you can get or prepare youself. Let's assume 3 step synthesis: A -> B -> C _> D, where A is the starting material B & C are intermediates that are prepared and isolated, and D is the product.

a) assume A, B, C, and D are all isolated and are 100% pure, but the yield is just 0.9 (()%). The overall yield is then 0.9 x 0.9 x 0.9 = 0.7 or 70%.
b) now assume the same conditions but the yield in each step is 70%; the overall yield will be just 50%.

Now assume that the yields are the same in each scenario, and the purities of A, B, C, and D are 90% and 70 % respectively. In the first case you get a 90% yield of 70% pure product; in the second case you get a 50% yield of 50% pure product.
I think this is an over-simplified example, but I think the point is that one would prefer to get as high a yield of as pure a product as possible, both for reasons of efficiency and economy. Do I make sense or am I just whistling in the wind? (In my defense, it has been 40 years since I practiced chemistry professionally, and in getting back to some amateur chemistry in the last year I am very distressed at how much chemistry I have forgotten, let alone how much new chemistry has been discovered since then.)

P.S. I welcome constructive criticism, but please don't be too hard on an old man!:D


[Edited on 12-31-2018 by CharlieA]


I used tablets by themselves, but they were all consistently around 375mg each. The 3rd sig fig did vary a bit, but that's unavoidable when your scale is uneclosed. (It even happens to my 2dp scales for some reason). It is interesting that for the amateur community masses are often less accurate than volumes, but normally scales are seen as much more reliable!

Either the filler is randomly distributed or on the surface. If it's randomly distributed (which it looks like as the ethanol slowly strips away the surface of the tablet in the same manner, then weighing each individual tablet should be fine. However, if it's an outside coat then it becomes necessary to grind the tablets.

lordcookies24 - 3-1-2019 at 05:57

I would use methanol to dissolve the pills. Last time I dissolved acetylsalicylic acid, the filler junk wouldn't dissolve so I could just filter out the stuff then boil off or distill the solution. I would not boil of the methanol inside tho but you provably know that already