I made sodium acetate using the following reaction:

$$\ce{NaHCO3 + CH3COOH ->CH3COONa + H2O + CO2}$$

The resulting solution of sodium acetate isn't pure, it has impurities. I want to crystallize the sodium acetate to extract the pure crystals. I crystallized some if it in the fridge, sodium acetate has a higher freezing point than water but I would like to know the technical side of it. Do temperature and saturation affect the crystallization and how can I crystallize sodium acetate reliably?


2 Answers 2


The Purification of Laboratory Chemicals by Perrin and Armarego states about sodium acetate this procedure:

Cryst[alise]d from acetic acid and pumped under vacuum for 10h at 120$^\circ$. Alternatively, cryst[alise]d from aqueous EtOH, as the trihydrate. This material can be converted to the anhydrous salt by heating slowly in a porcelain, nickel or iron dish, so that the salt liquefies. Steam is evolved and the mass again solidifies. Heating is now increased so that the salt melts again. (NB: if it is heated too strongly, the salt chars.) After several minutes, the salt is allowed to solidify and cooled to a convenient temperature before being powdered and bottled (water content should now less than 0.02%).

  • $\begingroup$ This doesn't really help. I was asking about crystallization and I don't have a vacuum pump or a way to melt salt. $\endgroup$
    – Aeolus
    Jun 12, 2017 at 23:20
  • $\begingroup$ The quote from Perrin is offering several options. 1) Crystallisation of sodium acetate from acetic acid; subsequently the crystals obtained are dried -- which, because it is hygroscopic and in order to do it faster, is done at higher temperature and under reduced pressure. Or 2) a crystallisation from aqueous ethanol, yielding crystals which incorporate three molecules of water per molecule of acetate (like the one at shown at wikipedia). These (by way 2) obtained crystals could be molten to eventually remove the water incorporated, but it will destroy the crystals and yield a powder. $\endgroup$
    – Buttonwood
    Jun 13, 2017 at 12:51
  • $\begingroup$ @TheTesseract'sShadow.. the book cited is the go-to place for finding out how to purify a compound, you should maybe edit your post with the equipment/resources you have available, but in general you're going to struggle to do a decent job without even basic lab resources. $\endgroup$
    – NotEvans.
    Jun 13, 2017 at 17:47

Sodium Acetate is particularly hard to recrystallize due to its high tendency to super cool. What usually happens is the crystals crash out entombing the impurities. I am assuming your making it from vinegar and baking soda. I would recommend concentrating your solution down to saturation and then boiling off 25mL of water at a time and allowing the solution to cool fully and collecting the resulting crystals each time. These WILL probably still be fairly impure. Because these will be the trihydrate salt they will melt with heat (anhydrous will not). You can melt these crystals in a hot water bath (boiling) or cough a microwave cough. Melt most of these crystals leaving a very small amount on the bottom of your beaker to act as seed crystals and allow the sodium acetate to cool very slowly. Once about half of it has crystallized add a small amount of water to the sodium acetate, mix well and pour it onto a towel and pat the remaining crystals dry.


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