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I was under the impression that if I put a semi full bottle of vinegar in the freezer the water would freeze but the vinegar would not until a lower temp. Therefore I could remove the bottle after the water freezes and pour off the vinegar. Apparently the water and vinegar doesn't separate nicely before freezing so this didn't work so well.

I've seen boiling suggested but I'm afraid I'll lose some of the vinegar in the process.

What is the simplest way to concentrate the vinegar at home?

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  • $\begingroup$ I know it says that it is for poison ivy rash, but a product like bayercare.com/domeboro.cfm is aluminum acetate, which would provide you with a stronger "vinegar" that comes in powdered form. $\endgroup$
    – jonsca
    Sep 14, 2013 at 4:09
  • $\begingroup$ Is there a reason you don't want to buy acetic acid and dilute to the desired concentration? $\endgroup$
    – bobthechemist
    Sep 14, 2013 at 11:39
  • $\begingroup$ @jonsca Sorry if I wasn't clear. I'm using it as an herbicide for poison ivy, not for a rash but the suggestion is appreciated. $\endgroup$
    – hortstu
    Sep 14, 2013 at 16:13
  • $\begingroup$ @bobthechemist where can I purchase acetic acid? What is acetic acid derived from? $\endgroup$
    – hortstu
    Sep 14, 2013 at 16:17
  • $\begingroup$ @hortstu No, it's okay, I was trying to clarify that while the label is for ivy rash (just by coincidence), that I understood you were using it against the plant itself. It's still going to give you a source of more concentrated acetate. $\endgroup$
    – jonsca
    Sep 14, 2013 at 16:38

5 Answers 5

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If you have access to a still you can distil the vinegar solution, however this can only go so far. You may find this section of the wikipedia article on azeotropy edifying - to summarise, adding ethyl acetate to a vinegar solution generates a water:ethyl acetate azetrope with a depressed boiling point that makes distillation more effective, however this may be beyond the pale for a home experiment and would obviously render the vinegar non-food-safe.

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  • $\begingroup$ Thanks, I'm not so concerned about getting 100% pure vinegar or having it remain edible. I'm just trying to make a more effective yet safe herbicide for some poison ivy at the back of the yard and in the cracks in the sidewalk. I could improvise a still, but is it the vinegar that gets left behind? It's not volatile at all? I was concerned about carmelizing the vinegar if I tried to boil off the water. $\endgroup$
    – hortstu
    Sep 14, 2013 at 3:43
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Yes,you are right. What you are talking about is a Fractional Freezing. With acetic acid at low concentration you will have ice and a solution of acetic acid more and more concentrated until you reach the eutectic point. The max concentration you will have is about about 58%. If your freezer go to -18°C you can reach only 38%. Take a look to this article.

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  • $\begingroup$ If I had a freezer that went to -18C wouldn't everything be frozen at that point? Is -18 the magic number that separates acetic acid from water? $\endgroup$
    – hortstu
    Sep 16, 2013 at 3:20
  • $\begingroup$ You have to do it in many steps e.g. you freeze it at -4°C, eliminate ice so the new solution is more concentrated and freeze at a lower point, you can do another step to -18°C and so on until you reach the eutectic at -26.77°C. At that point you have ice and acetic acid at 59%, not the best thing for Iced tea! If i have more time I will improve the answer.. $\endgroup$
    – G M
    Sep 16, 2013 at 6:14
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The acetic acid has melting point of 16°C, while water has 0°C. Therefore, when placed into a freezer, the acetic acid will freeze first, not the water. So once the solution freezes, separate the liquid (it contains more water than acetic acid) and keep the ice (which has more concentrated acetic acid).

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  • $\begingroup$ Do you have a source for this? Acetic acid in solution has different freezing points based on concentration. At least that's the impression I'm under but I could be very wrong. $\endgroup$
    – hortstu
    Oct 21, 2017 at 4:00
  • $\begingroup$ Hi Hortstu. I don't have a source for that and I could be indeed wrong (although it seemed "natural" to be like that at the time when I wrote). I also read somewhere that there are different freezing points based on concentration. One could freeze vinegar in the freezer and watch it until a portion of it freezes. Then separate the liquid from ice and then perform titration on them - probably one will be more concentrated than the other. $\endgroup$
    – Nick_F
    Oct 22, 2017 at 5:16
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A more direct way to isolate the acetic acid would be chemically. It requires baking soda (easy to get) and sodium bisulfate (usually found as "spa pH reducer"), and a distillation setup. The idea here is that you can get most of the water out, without also boiling away the vinegar, by first converting it to a salt. Then reform the acid from the salt by mixing it with an acidic salt.

  1. Add baking soda (sodium bicarbonate) to the vinegar until it stops bubbling, or the pH hits neutral. This will take a while as there will be loads of foaming.
  2. Boil the resulting sodium acetate solution until it starts to precipitate. If you want something closer to glacial acetic acid, you'll want to try to get to anhydrous, but "mostly anhydrous" is probably fine. Pour out in a glass dish and let dry, in a 250 C/500 F oven would remove more water.
  3. Mix the acetate very well with 3.5x the weight of sodium bisulfate and put the mix into the distilling flask. Distill over the acetic acid. You'll end up with a brick of sodium sulfate in the boiling flask, takes a long while to get out, but you should also get very concentrated (50% to 90+%) acetic acid in the receiving flask, depending on the level of anhydrousness of the reactants.

More detailed procedures can be found on some other science forums, sometimes split up between 'creating sodium acetate from vinegar', and 'creating acetic acid from sodium acetate and sodium bisulfate'. Here's a video for the latter step.

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Ive tired this via distillation and afterwards it left a black substance so at the very least you need to distill it first to remove any solutes in the viniger.

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    $\begingroup$ Hi Josh, welcome to Chemistry! Take the tour to get familiar with this site. Could you edit your answer for spelling and grammar? It would make the answer better for future users. $\endgroup$
    – pentavalentcarbon
    May 20, 2018 at 17:31

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