Lytzu
Harmless
Posts: 11
Registered: 8-1-2020
Member Is Offline
Mood: curious
|
|
Acetate Salts From Vinegar For Synthesis of Acetic Acid
I would like to prepare glacial or near glacial acetic acid for use as a reagent. I think I've read all the relevant threads at least half a dozen
times and am starting to feel confident that it is something I can accomplish. However, none of these threads seem appropriate for this post as, for
now, I will be focusing specifically on acetate salt production. If this would be better merged to a previous thread, the most recent may be applicable.
Anyway, from what I've read the simplest way for the home chemist to produce acetic acid is to "salt out" an acetate from vinegar. This can be
achieved by the addition of sodium carbonate, sodium bicarbonate, calcium carbonate, etc, which reacts with the acetic acid in the vinegar to form an
acetate salt. Excess water can then be boiled off. A second reaction of the acetate salt and sulfuric acid turns the salt back into acetic acid.
I have a very small amount (<=10g) of sodium carbonate on hand from an old chemistry kit I got as a kid and never really used. Eventually I would
like to try this with other bases but for now I'll work with what I have.
My plan is to follow the chemical equation 2CH3COOH (acetic acid) + Na2CO3 (sodium carbonate) = 2NaCH3COO
(sodium acetate) + H2O + CO2
at 3/100ths of a mole. From the molar masses I calculated this to be 72g 5% vinegar + 8.58g sodium carbonate will yield about 4.92g sodium acetate.
The only reason I am doing the experiment on such a small scale is because of the limiting factor of the amount of sodium carbonate I have. After I
try this I will obtain either more sodium carbonate or sodium bicarbonate.
The scale I bought is coming in early next week and in the meantime I would appreciate the help determining if this is a viable plan. I will be
updating this thread as I progress beyond this one experiment, which is part of why I thought I'd make my own post. Further experiments of this
subject will likely include producing acetates from 30% vinegar if I can find it (the depot lied, said it was there and it wasn't) and 85-90%
photography indicator stop bath solution. As I mentioned before I would also like to try with sodium bicarbonate, and perhaps calcium carbonate and
others. For now I am focusing on this acetate salt step of production but will eventually progress to reverting the acetate back to acetic acid via
sulfuric acid.
Thanks in advance. If anything needs correcting or I missed a step or something PLEASE let me know! I am excited for my first chemistry experiment
outside of school or work.
[Edited on 10-1-2020 by Lytzu]
|
|
|
vibbzlab
Hazard to Others
Posts: 241
Registered: 6-11-2019
Member Is Offline
Mood: Always curious
|
|
This is actually a viable experiment.I had tried this experiment once by just crystallising the salt from vinegar. Your yield would not be that great
amounts though as vinegar has a low concentration of acetic acid in it. But yes definetly you can get glacial acetic acid by adding concentrated
Sulfuric into the salt. I would say but glacial acetic acid since that's not that difficult to get . I wouldn't use up my concentrated Sulfuric acid
for acetic acid. But yes if you have no other way or if you want to try it out,then yeah,go ahead. All the best.
Amateur chemist. Doctor by profession
Have a small cute home chemistry lab.
Please do check out my lab in YouTube link below
This is my YouTube channel |
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
You have a very small amount of sodium carbonate (Na2CO3)? You can buy it dirt cheap as washing soda (although it's probably either the heptahydrate
or decahydrate). Or you can buy baking soda (sodium hydrogen carbonate- NaHCO3).
Your sodium carbonate may also be hydrated, and the sodium acetate that you get will probably be the trihydrate.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
vibbzlab
Hazard to Others
Posts: 241
Registered: 6-11-2019
Member Is Offline
Mood: Always curious
|
|
I think by strong heating them u can get anhydrous form right? Correct me if I am wrong
Amateur chemist. Doctor by profession
Have a small cute home chemistry lab.
Please do check out my lab in YouTube link below
This is my YouTube channel |
|
|
Lytzu
Harmless
Posts: 11
Registered: 8-1-2020
Member Is Offline
Mood: curious
|
|
Na2CO3, ah, did I get that formula wrong? Good catch thanks Draconic, fixed.
I will be picking up more washing soda and/or baking soda soon, I just figured I'd use up the tiny bit I have as a test.
I believe the sodium carbonate I have is at least partially hydrated. It may not have come that way but after sitting in a basement for years most of
it is in a large relatively hard clump.
If the sodium acetate I obtain is trihydrated how would that affect things? Will the acetic acid created from that be less concentrated because of the
water trapped in the sodium acetate crystal? Does this change the sodium acetate itself in any relevant way?
vibbzlab, I had heard that other acids might be able to be used but I'm not certain. So far I've mostly ignored this step in favor of fully
understanding the first part of the process (the acetate formation).
|
|
|
XeonTheMGPony
International Hazard
Posts: 1636
Registered: 5-1-2016
Member Is Offline
Mood: No Mood
|
|
to purify: Recrystallize in methanol.
all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure) distill out the
ethyl acetate then reflux with sodium hydroxide.
to dry the sodium acetate, you can micro wave it till it no longer heats up, be care full to check often, as once all the water is boiled off it will
be very hard on the micro wave! Or you fuse it (Heat it to melting point)
[Edited on 10-1-2020 by XeonTheMGPony]
|
|
|
Ubya
International Hazard
Posts: 1232
Registered: 23-11-2017
Location: Rome-Italy
Member Is Offline
Mood: I'm a maddo scientisto!!!
|
|
i did this experiment a few years ago.
i used "vinegar of alcohol", vinegar used for cleaning, not made from grapes so it shouldn't have sugars (white vinegar has sugars). i neutralized 2
liters of vinegar (6% acetic acid if i remember correctly) with sodium bicarbonate. i then boiled down the solution until only a powder remain. the
trihydrate melts at 58°C, you want notice when all the water is gone since you'll still have a liquid. i heated until i got a syrup, and then i
heated it more until i got the anhydrous form.
i added sulphuric acid and performed a distillation. i got really concentrated acetic acid (by the smell) but it is not glacial as it doesn't solidify
in the fridge, i haven't titrated it yet as i don't have a burette.
---------------------------------------------------------------------
feel free to correct my grammar, or any mistakes i make
---------------------------------------------------------------------
|
|
|
betamethyl
Harmless
Posts: 2
Registered: 25-12-2018
Member Is Offline
|
|
I did this last year too, using a slightly different method. I have my notes written somewhere, but here is a rough outline:
I wanted to produce glacial acetic acid from vinegar, so i planned to: react the vinegar with sodium bicarbonate to produce sodium acetate in my first
step; and, to distill with 98% sulfuric acid in the second step. I went to coles and bought around 12 litres of home-brand vinegar.
I didn't know what concentration of acetic acid was in my vinegar. I didn't really have any pH paper, so i would have to estimate the quantity of
sodium bicarbonate to add. But if i added too much sodium bicarbonate then i would be left with sodium bicarbonate/carbonate in my acetate solution.
This would contaminate my Na acetate with Na bicarbonate, and would create problems during the distillation step with 98% sulfuric acid.
To prevent this, i bought 5 Kg of "Garden lime" from bunnings. This stuff was only around 20% Calcium carbonate/hydroxide as far as i can remember;
the rest was mainly sand.
My plan was to react the vinegar with excess calcium carbonate to form calcium acetate, then to react this calcium acetate with sodium carbonate and
filter off the calcium carbonate byproduct to get a solution of sodium acetate. Because i couldn't test pH and i didn't know how much acetic acid was
in my vinegar, i planned to do this extra step so that i could weigh the calcium acetate product (with minor calcium carbonate impurity due to its low
solubility in water) and determine the concentration of acetate in my vinegar more accurately. By producing the calcium salt from the vinegar first,
the possibility of having any bicarbonate or carbonate in my RBF when i go to add 98% sulfuric acid is removed (as i had accidentally done this in an
earlier experiment). By converting the calcium acetate to sodium acetate before distilling with sulfuric acid, a cleaner, simpler distillation of
acetic acid is afforded without calcium sulfate byproduct which can coat the calcium acetate solid in the flask when added to sulfuric acid. When this
happens, it kind of ruins the reaction because the mixture is just too viscous to stir magnetically, so you have to pull things apart and poke around
with a glass rod while the sulfuric acid is hot and glacial acetic acid vapour is coming out of the flask which will make your eyes water for sure.
In the end, i thought that calcium acetate would be okay to distill with sulfuric acid to form GAA. I found out the hard way that this ended up
causing problems.
Calcium acetate:
1. 12 litres of vinegar (approx 5%) was poured into a 40 litre plastic bucket.
2. "Garden lime" was spooned in to the vinegar with a ladle to prevent splashing. The mixture stirred manually. The production of carbon dioxide was
mild - probably due to all the sand taking up so much room and 'diluting' things.
3. After leaving the 5 Kg lime in the vinegar overnight, a pinch of bicarb soda was added to the mixture; the white powder bubbled as it dissolved.
This indicated that unreacted vinegar was present, so the mixture was poured into a 19L stainless steel metal pot (now in the garbage bin), combined
with as much crushed chalk and calcium carbonate-containing items i could find, and boiled over a campfire stove. The mixture was periodically stirred
to avoid bumping, but the sand acted like a heat sink and just prevented the liquid from boiling vigorously.
4. After leaving the metal pot on top of the campfire stove all night, a pinch of bicarb soda was added to the liquid the next morning and the mixture
did not bubble.
5. Filtering such a large mass of solid would require a series of filtrations and the better half of a day. Instead, i siphoned the supernatant liquid
off from the sand and filtered once through kitchen paper towel. I spooned the sand into some old large socks and squeezed the liquid out - this
liquid only needed a simple gravity filtration through kitchen paper to clear up.
If the liquids were at all cloudy at this point, it didn't much matter at this stage.
6. The large volume of liquid needed to be removed of all water. Because of the large volumes, i found that heating all 10-15 litres of liquid in one
pot was very inefficient as my heater struggled to bring everything to a boil. Instead, i found that using multiple 1-2L pots and using all 4 heating
elements on my kitchen stove boiled the water off fastest. The calcium acetate solution (which was brown/yellow) was periodically ladled into these
pots until almost all the water was removed. When this occurred, the calcium acetate paste was spooned all into one pot and some water was used to
wash the other pots and transfer remaining calcium acetate into the main pot. This main pot was heated until a large amount of solid started
precipitating. Soon after, all the water had been removed and the steam adopted a vinegar-like odor. The mixture became homogenous and it appeared
that the calcium acetate had melted. I just took the pot off the heat and covered it with a piece of slate to prevent moisture getting in as it
cooled. If i was to repeat this, i would recommend pouring the molten calcium acetate onto a metal tray and allowing to cool (preventing moisture),
because getting the product out of that metal pot required a hammer and chisel. The yield was around a kilogram i think? maybe 600-800g? i know that i
ended up with a yield of acetic acid which worked out to a concentration of 4% acetic acid in my original vinegar.
Make sure you store calcium acetate in a ziplock bag or a screw-lid jar and use soon after manufacture because it was very hygroscopic. I believe
sodium acetate behaves much the same.
I also have made sodium acetate. In my first experiment, i combined sodium bicarbonate and vinegar in a pot and boiled to remove the water. It made
horrible smelling fumes and went brown. This commonly happened to me over more experiments with sodium bicarbonate, and i never really knew why. I
could never remove the organics easily, or prevent degradation. This is why i used calcium carbonate in excess for my experiment with 12 litres of
vinegar - i suspected that i was adding too much sodium bicarbonate to the vinegar and it was leading to problems when i heated it to remove the water
in later steps.
For distilling, i would recommend one of these 2 methods:
1: Combining sodium acetate (anhydrous) and concentrated sulfuric acid and distilling.
2: If your acetic acid doesn't need to be glacial, you can make lower-concentration solutions by combining an aqueous solution of calcium acetate with
dilute sulfuric acid and filtering off the calcium sulfate precipitate.
I have also read things on SciMad about combining oxalic acid with calcium acetate solutions and filtering off the calcium oxalate which is very
insoluble, but i have no experience with this.
|
|
|
annaandherdad
Hazard to Others
Posts: 387
Registered: 17-9-2011
Member Is Offline
Mood: No Mood
|
|
I bought some washing soda as a cheap source of sodium carbonate. When I dissolve it in water it makes some foam which my reagent sodium carbonate
does not, and the solution is not as clear. Therefore I have suspected that my washing soda contains some contaminants.
On the other hand, sodium bicarbonate can be bought in the grocery store and it is fit for human consumption. That does not prove it is pure, but
when I bake it in the oven (300 degrees F for an hour) I get a yield that by weight is within 1% of what I would expect for conversion of pure sodium
bicarbonate into pure sodium carbonate. (This is for Arm and Hammer brand baking soda, I'm in the US.) Also, the resulting product dissolves in
water like my reagent sodium carbonate, making a clear solution without foam.
To summarize, I have suspicions that commercial washing soda is not a good source for high quality sodium carbonate, while baking soda, heated to
convert to sodium carbonate, may be.
Any other SF Bay chemists?
|
|
|
Cou
National Hazard
Posts: 958
Registered: 16-5-2013
Member Is Offline
Mood: Mad Scientist
|
|
Nile red made sodium acetate from 5% food vinegar.
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
What about forming calcium acetate from cheap sodium acetate which then could be used to form acetic acid as calcium acetate reacts with sulfuric acid
to form insoluble calcium sulfate and acetic acid not sure if this could also work or just make chloroacetic acid but sodium acetate could react with
HCl acid to form insoluble salt which NaCl is insoluble in HCl and acetic acid
|
|
|
karlos³
International Hazard
Posts: 1520
Registered: 10-1-2011
Location: yes!
Member Is Offline
Mood: oxazolidinic 8)
|
|
Quote: Originally posted by XeonTheMGPony  | | to purify: a all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure)
distill out the ethyl acetate then reflux with sodium hydroxide. |
You probably meant, reflux with NaOH, then distill out the formed ethanol?
|
|
|
Lytzu
Harmless
Posts: 11
Registered: 8-1-2020
Member Is Offline
Mood: curious
|
|
Sodium Acetate Experimant 1 Complete
Alright so I've just finished my first official home chemistry experiment  . As
far as I can tell it's been a success.
As planned in my original post my procedure was as as follows:
9 grams of sodium carbonate was added to a 250mL Erlenmeyer flask. Much of it had to be crushed before hand because it had
formed a large clump in the jar.
72g of 5% vinegar were added to the Erlenmeyer flask. A white foam was immediately produced. When fizzing completed there was still sodium
carbonate powder on the bottom of the flask. I added an additional 54g(!) of vinegar before it all disappeared. Swirling of the flask was used to aid
in this process.
Because I do not have a hot plate or Bunsen burner yet an improvised hot plate was created by placing an old pan on the stove top. The flask was
set directly on this surface. The burner was set to medium. Bubbles began to rise in the solution very soon after this. It reached a true boil after
maybe 20 minutes. By then it was beginning to take on a yellowish tinge. After about an hour to an hour and a half nearly all the liquid had boiled
off. When the sodium acetate began to crystalize it happened very quickly, almost all at once.
The crystalized solution was allowed to cool slightly before I attempted to scrape it into a small jar. Due to lack of scoopulas a long
screwdriver was used. I was worried about scrapping the glass so I was only able to recover maybe half of the solids. This weighted in at a surprising
4g. It was still very slightly damp so that is probably the majority of why.
Conclusion: Sodium acetate was produced in this process, although likely mixed with impurities.
Questions:
Why did I have to add extra vinegar? Were my calculatons off? Should I allowed the sodium carbonate to stay in the solution
and become part of the resulting powder? Which reactant do I want more off? I would assume I'd rather have extra acetic acid in the solution because
that is ultimately what I am trying to form, although this probably boils off.
The foam was the result of CO2 produced from the reaction correct? And when boiling the bubbles were just water vapor, just like when boiling
water?
There was one spot on the bottom of the flask where more and larger bubbles were produced and in the beginning "sticking" to the bottom. Why is
this? I have seen a similar phenomenon when boiling water.
Oh! Almost forgot. It could just be me working in my kitchen (I know, not the best but its the only available space right now) but the sodium
acetate had a very slight sent. Like...biscuits? The fumes from my solution had no scent, unlike betamethyl's | Quote: | | i combined sodium bicarbonate and vinegar in a pot and boiled to remove the water. It made horrible smelling fumes and went brown.
|
though they were using bicarbonate and I was using carbonate. Ideas?
Also, | Quote: | | the trihydrate melts at 58°C, you want notice when all the water is gone since you'll still have a liquid. |
I don't believe this is what happened with mine, unless it went from a syrup to the anhydrous form very quickly. Is Ubya correct
that this is what happens?Additional questions from the thread not yet answered:
If the sodium acetate I obtain is trihydrated how would that affect things? Will the acetic acid created from that be less
concentrated because of the water trapped in the sodium acetate crystal? Does this change the sodium acetate itself in any relevant way? Is Ubya
correct that heating it results in an anhydrous crystal?
I believe vibbzlab also had a question about if heating the sodium carbonate(?) could create anhydrous sodium carbonate. I believe he was
referring to that chemical. Notes:The sodium acetate crystals formed 2 sort of layers on the bottom of the flask. On top was slightly brownish and
still very slightly damp while below this was a harder white substance. I believe this has to do with impurities rising with the water vapor. I think
I saw it mentioned in another thread.
I think that is all. My next experiment will likely be similar, although using sodium bicarbonate instead. After that I may move on to trying this
with 85-90% acetic acid stop bath solution. Thanks for the help so far.
|
|
|
XeonTheMGPony
International Hazard
Posts: 1636
Registered: 5-1-2016
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by karlos³ | | Quote: Originally posted by XeonTheMGPony | | to purify: a all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure)
distill out the ethyl acetate then reflux with sodium hydroxide. |
You probably meant, reflux with NaOH, then distill out the formed ethanol? |
That's what I said lol
But you must distill out the ethyl acetate first as there is a bunch of other garbage mixed in usually, once that is distilled out then you reflux
over Sodium Hydroxide
Once you have re-fluxed for an hour or better, distill off the fresh freed ethanol, you should have a liquid sludge left.
React to slightly acidic with vinegar (As I hope you used excess base) then evaporate to dryness.
Take this dry mass and dissolve in methanol and filter out any that doesn't, wash with a bit of warm methanol to get any out that was retained in the
filter paper.
Recover some methanol by distilling to a gell or just evaporate it all off to recover your pure sodium acetate (Sodium carbonate will remain in the
filter paper as a solid.
|
|
|
draculic acid69
International Hazard
Posts: 1371
Registered: 2-8-2018
Member Is Offline
|
|
| Quote: Originally posted by XeonTheMGPony | to purify: Recrystallize in methanol.
all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure) distill out the
ethyl acetate then reflux with sodium hydroxide.
to dry the sodium acetate, you can micro wave it till it no longer heats up, be care full to check often, as once all the water is boiled off it will
be very hard on the micro wave! Or you fuse it (Heat it to melting point)
[Edited on 10-1-2020 by XeonTheMGPony] |
Does this type of hydrolysis use an excess of Naoh or equimolar amounts?
|
|
|
XeonTheMGPony
International Hazard
Posts: 1636
Registered: 5-1-2016
Member Is Offline
Mood: No Mood
|
|
I used an excess (Read as: I didn't bother to look it up!), as I wanted it to be dry as possible, then I just used evaporated down vinegar to fully
neutralize any remaining hydroxide.
Then I purify it in methanol to get rid of carbonates and gunk, some where on the board I did a solubility amount for it.
After that I got very very good concentrated acetic with minimal crap left in the reaction flask, so it was well worth the effort to recrystallize and
dehydrate thoroughly
https://www.youtube.com/watch?v=rq7a5_tvBPM
https://www.youtube.com/watch?v=BzgWq7Dtn3U
[Edited on 15-1-2020 by XeonTheMGPony]
|
|
|
Lytzu
Harmless
Posts: 11
Registered: 8-1-2020
Member Is Offline
Mood: curious
|
|
Experiment 2: Sodium Bicarbonate Complete
Alright so I've continued with this project but I am wondering if this thread should be moved to beginner as I have so many basic questions. I would
like at least some responses related to these questions. I am especially interested in trihydrated vs anhydrated acetic acid. Thanks in advance.
As far as experimentation goes I just finished a run using basically the same method as before but with baking soda instead. My calculations seemed
correct this time. I'm still not sure why they were off last time. It doesn't seem to have been the vinegar calculation that was the problem at least,
but in case it is...
I determined the molar mass of 5% (by weight) acetic acid vinegar to be 1,200g. My understanding has been that if 5g acetic acid is in 100g vinegar
solution and the molar mass of acetic acid is 60g/mol then 1,200g vinegar solution would be needed to obtain those 60 grams. This being at 1 mole.
Then I just divided it into whatever fraction I needed.
This time was 1/4 mole. The solution fizzed for a LOT longer than I expected but all the sodium bicarb appeared to react. I stirred it until bubbles
no longer formed. Or at least until my patience ran out. About the same time.
On the stove it also took FOREVER to boil off. Eventually I turned up the heat a notch because although it was boiling it wasn't boiling like water
anymore, if that makes sense. I've read about the solution "bumping" before, I'm not sure if that's what it was or not. It's difficult to explain.
Also my vinegar was a mix of two different brands this time because I was running out. It seemed less yellow/fewer impurities but that also could have
to do with scaling up a couple hundred mL.
The crystallization was also...strange. It began as like a weird crystal foam as it was bubbling. I kept the heat on it long after it crystallized and
it became a lot drier than before, as would be expected. I thought about turning the heat up even higher to try to melt it and get anhydrous but it
had been quite a few hours already and that wasn't in the plan. Maybe later in the week as a quick thing to do. Need to read more into it. Some input
on anhydrous and how all that works would also be appreciated.
Anyway, I also bought some 85-90% indicator stop bath solution for the photo store. Need to do some calculating and reading but that is next on the
list of things to try. I am anticipating MUCH less boiling time. Hopefully?
Thanks again, send advice if you've got it.
[Edited on 29-1-2020 by Lytzu]
|
|
|
XeonTheMGPony
International Hazard
Posts: 1636
Registered: 5-1-2016
Member Is Offline
Mood: No Mood
|
|
I strongly urge recrystallizing, as there will be some junk all ways in there that needs to be removed.
|
|
|
B(a)P
International Hazard
Posts: 1112
Registered: 29-9-2019
Member Is Offline
Mood: Festive
|
|
I am not 100% on what your questions are that you are wanting responses on. To help you understand if your solution was bumping, it is the formation
of a pocket of vapour beneath the surface of your liquid that occurs rapidly enough that it vibrates your apparatus. You can use boiling chips or even
a stir bar to prevent or at least minimise the effect.
|
|
|
B(a)P
International Hazard
Posts: 1112
Registered: 29-9-2019
Member Is Offline
Mood: Festive
|
|
I am not 100% on what your questions are that you are wanting responses on. To help you understand if your solution was bumping, it is the formation
of a pocket of vapour beneath the surface of your liquid that occurs rapidly enough that it vibrates your apparatus. You can use boiling chips or even
a stir bar to prevent or at least minimise the effect.
|
|
|
Lytzu
Harmless
Posts: 11
Registered: 8-1-2020
Member Is Offline
Mood: curious
|
|
From my previous post. I don't believe any of these have been addressed.
| Quote: |
Questions: Why did I have to add extra vinegar? Were my calculatons off? Should I allowed the sodium carbonate to stay in
the solution and become part of the resulting powder? Which reactant do I want more off? I would assume I'd rather have extra acetic acid in the
solution because that is ultimately what I am trying to form, although this probably boils off.
The foam was the result of CO2 produced from the reaction correct? And when boiling the bubbles were just water vapor, just like when boiling
water?
There was one spot on the bottom of the flask where more and larger bubbles were produced and in the beginning "sticking" to the bottom. Why is
this? I have seen a similar phenomenon when boiling water.
Oh! Almost forgot. It could just be me working in my kitchen (I know, not the best but its the only available space right now) but the sodium
acetate had a very slight sent. Like...biscuits? The fumes from my solution had no scent, unlike betamethyl's
| Quote: |
i combined sodium bicarbonate and vinegar in a pot and boiled to remove the water. It made horrible smelling fumes and went brown.
|
though they were using bicarbonate and I was using carbonate. Ideas?
Also,
| Quote: |
the trihydrate melts at 58°C, you want notice when all the water is gone since you'll still have a liquid. |
I don't believe this is what happened with mine, unless it went from a syrup to the anhydrous form very quickly. Is Ubya correct that this is what
happens?
Additional questions from the thread not yet answered:
If the sodium acetate I obtain is trihydrated how would that affect things? Will the acetic acid created from that be less concentrated because
of the water trapped in the sodium acetate crystal? Does this change the sodium acetate itself in any relevant way? Is Ubya correct that heating it
results in an anhydrous crystal?
I believe vibbzlab also had a question about if heating the sodium carbonate(?) could create anhydrous sodium carbonate. I believe he was
referring to that chemical.
|
My assumptions to the answers are, in order of the above list
Something to do with how hydrated the sodium carbonate was
yes
impurity on the glass?
no idea, got a little bit of the same smell on this latest batch, some impurity in the vinegar maybe?
maybe?
different crystal structure in the trihydrated vs anhydrated sodium acetate, less concentrated acetic acid final product
...yes?
|
|
|
B(a)P
International Hazard
Posts: 1112
Registered: 29-9-2019
Member Is Offline
Mood: Festive
|
|
My assumptions to the answers are, in order of the above list
Something to do with how hydrated the sodium carbonate was
yes
impurity on the glass?
no idea, got a little bit of the same smell on this latest batch, some impurity in the vinegar maybe?
maybe?
different crystal structure in the trihydrated vs anhydrated sodium acetate, less concentrated acetic acid final product
...yes?
[/rquote]
My two cents worth, in order of your answers.
Point one, you need to know which hydrate you have before you can get your stoichiometry right. Also it is unlikely the 5% conc is accurate.
Point two, yes about the initial CO2, but CO2 will continue to come out of solution as it gets hotter and more concentrated so not just water vapour.
Point three, could be contamination in the glass or could be an imperfection, scratches make good locations for bubbles of gas to form.
Point 4, I also do not know
Point 5, That certainly is the meeting point, but I have not done the procedure so I couldn't say.
Point 6, if you want to react the substance with something you need to know the what you have before you can work out your stoichiometry. Heating it
will work, also you could make a dehydrater if you have time to let it sit.
Point 7, same as point 6
Hope this helps
|
|
|
Lytzu
Harmless
Posts: 11
Registered: 8-1-2020
Member Is Offline
Mood: curious
|
|
Thanks for all the help B(a)p.
Today I'm going to try using the stop bath solution. I also picked up some washing soda (sodium carbonate) so I will try both, but I've got some more
questions.
How do I determine how hydrated the sodium carbonate is? Assuming it possibly isn't all the same hydrated compound, how do I make it all be the same?
Bake it in the oven? I'd like to get the stoichiometry correct instead of just vaguely guessing at amounts.
In a similar vein, the stop bath solution also isn't an exact concentration of acetic acid. It's the Kodak 85-90% stuff. I won't try to figure it out
today but in the future is there a titration or something I can do to determine more precisely the concentration? I will have to look around the
forum, because there might already be a thread on that.
Oh, before I forget to ask, Xeon why methanol? This is probably a very stupid question but why can't you just dissolve the acetate back into water and
filter that? Is it because the impurities will also dissolve in the water and not be filtered out?
Sorry, I'm very beginner and not understanding half of what you're saying. I've kind of been ignoring it for now until I have a better grasp on
things.
I believe you were referring to this method in your first post with the ethyl acetate and sodium hydroxide Xeon? I am planning on trying that eventually. I'm sure it is a better
method for obtaining large quantities of acetic acid, but both my comprehension and my technical skill are nowhere near where I would be comfortable
trying it yet.
By the way, besides the acetone free nail polish remover M.E.K. Substitute should be a good source of ethyl acetate if you can get it where you are. I
already picked some up while I was getting acetone.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
