Crystallization of homemade sodium acetate

Step 16 Additional purification

As an addendum, I've discovered a simple way of further purifying the crystals, should you wish to do so. I've no pictures, but I wanted to at least add this description for those of you obsessive enough to consider it.

I knew that sodium acetate is not very soluble at all in isopropanol (isopropyl alcohol), and it occurred to me that i could essentially "wash" away most of the impurities from the crystals after crystallization. There's no real chemistry here - it's just the mechanical action of the liquid isopropanol that does the trick.

Crystallize your solution in a beaker or a bowl; after it's completed, use a clean glass rod or a clean fork to break up the crystals as thoroughly as possible, yielding an amber colored mush of small crystals and remaining liquid and impurities.

Then, a couple of tablespoons at a time, "wash" the crystals with isopropanol in a coffee filter in a small strainer over an appropriate collection container. It's best if you can use a wash bottle to hold the alcohol; otherwise, pour it in small amounts at a time. You can agitate the crystals with a glass rod or small spoon. To help with drainage, periodically sweep away the crystals from the center area of the filter (be careful not to rip the filter doing this).

I went through approximately two liters of isopropanol to wash my entire batch of crystals. It's not very cost-effective, but it's worth the fun of learning.

After cleaning the crystals, you'll want to eliminate residual isopropanol. The only practical way to do this is to spread the crystals evenly inside a large baking dish and heat them, driving off both the isopropanol and the water. I did mine in several small batches in my oven, at 230 degrees F or so, holding the oven open slightly by closing it on a heat-proof trivet, which allowed some air circulation (to avoid igniting the isopropanol vapor).

Check the crystals after a half hour or so. You may need to break up still-moist areas with a spoon. As the water is driven off, the trihydrate will turn into plain sodium acetate (anhydrous). When it's finished, it will probably be a light, puffy mass.

Once the crystals are dry, you will need to redissolve them in fresh distilled water. Use about 50% more water than the volume of the solution you started with. (That is, if your supersaturated solution was 200 mLs, add 300 mLs.) Heat the solution to almost boiling, filter it to remove any dust or other contaminants (using 4-5 coffee filters - see step 10), then boil off the excess water again as you did in step 12. You'll need to boil off at least the 50% extra volume of water that was added. A small amount of sodium acetate will have been lost in the washing step, so the resulting volume of supersaturated solution will be somewhat less than what you started with.

Comments

[JCarter718]

Hi! I've been playing around with this reaction for the past few days and then stumbled on your instructable and was surprised to see this section on using isopropyl alcohol. The reason I'm making this solution in the first place is to see if I can lower the boiling point by adding alcohol. But since you're saying it's not soluble, I'm now thinking that won't work. Any thoughts?

[hot ice freak]

is it possible that if you freeze the solution past the recommended point, then boil it so that the water is frozen @ the bottom. when that happens, could u just take out the water?

[indigoandblack]

Short answer: no, doesn't work that way. :-) I'm not sure exactly what would happen if you dropped the temperature on the supersaturated solution. My guess is that you'd get ice with sodium and acetate ions in it (the same as happens when salt water is frozen), and I also imagine that as you then heated the resulting ice mass, the liquid produced from melting would simply be the original water + sodium and acetate ions (as long as nothing triggered crystallization while you were melting the ice). The water won't separate from the ions, in other words.

[aspiring_alchemist]

i would like to see your other chemistry experiments. thank you very much for your help.

[indigoandblack]

Heh... Thanks for the vote of confidence, but this is the only one I've documented! :-)

[Mongo_boy]

You know I used to think I was smart but this surpasses me

[Paper1]

Very fascinating! May I ask what happens if no filtration of the super saturated solution is done? Would the impure solution simply crystallise as soon as it cools?

[indigoandblack]

Which part, exactly? :-) If the purification step isn't done (with the activated charcoal), I've found that I cannot get the classic supersaturated state... Either no crystallization occurs, or slow crystal growth occurs when the solution is cooled. If you do the purification step, without fine filtering I'm not sure what will happen... there would be an awful lot of carbon dust in the solution. Truthfully, even with a pure sodium acetate solution (not homemade), it will tolerate a good deal of visible contamination (dust, cat hair...). The supersaturated solution requires the right kind of "trigger" - you could drop in, say, grains of sugar or salt, and it wouldn't cause crystallization. Drop in a crystal of sodium acetate, however, and that will cause the sodium and acetate ions to come together in just the right configuration. In the hand warmers, a little metal disc is used to trigger crystallization - it's usually embossed and "domed" slightly... what's not clear to me, though, is what exactly happens when it's flexed and it clicks, and triggers crystal growth. I've heard a number of different explanations, but none backed up with verifiable specifics.

[Paper1]

Yes that metal disk is patented I think. The explanation I heard was that it pressurised a small number of the molecules into the right configuration causing all the rest to do the same as a seed crystal would. If no purification of any kind is used, then crystallisation doesn't occur as effectively? Do you know the scientific explanation for this? Would it be that the impurities disrupt the structures?

[killbox]

from one of the patients i looked up years ago, it indicated the groves in the disk, caused small nuclei of crystals to form but they cant spread past just a few atoms thick and only a long the grove. flexing the disk fires off a few small crystals.

[indigoandblack]

I've no specific explanation of why the impurities interfere with the desired crystallization, other than that they effectively lower the concentration of sodium and acetate ions in solution - clearly not enough to keep *any* crystals from forming, since I'd see significant crystal growth overnight from an unfiltered batch, but enough to make the sought-after "instant" crystallization not happen.

[profquatermass]

OK, I'll bite. Just why would I want to make this?

[indigoandblack]

*You*? I've no clue. *Me*? I was talking with some kids about the vinegar and baking soda reaction, and when we looked up the formulas, I realized that the product was sodium acetate - never realized that before. So I decided it would be interesting to see if I could produce it from scratch to a sufficient degree of purity to be able to show the crystallization of a solution. It took more than a month to get it right (not full time, mind you...), and meantime I worried a neighbor was going to call the police on me, thinking I was cooking meth in my kitchen, owing to all the glassware and so forth visible through my window. :-)

[lileva]

Great instructable! I am a high school science teacher interested in having the students do this as a multi-day lab. Can the recipe be reduced by 1/3 to save costs? Can the procedure be done over a course of days (i.e. at the points when the procedure calls for bringing the solution down to room temperature)? In your opinion, do you think the experiment is feasible w/ a class of 25 teenagers working in groups of two? Thanks for your help!

[indigoandblack]

Ah, very nice! :-) You should be able to cut down the proportions more or less arbitrarily. I'd definitely suggest that you try it out first to fine tune things. And breaking up the steps is no problem - especially considering that the smaller solution amount will heat and cool more quickly than mine did. I think it ought to be completely manageable with that class size. Though, again, you should definitely do it once yourself first to note any problems you run into.

[lileva]

Absolutely!! I always give my lab experiments a dry run first. Thank you so much for getting back to me quickly. I'll be in touch in case I encounter a glitch during my attempt.

[Soupraok]

how long to the crystals last for?

[indigoandblack]

They'll last up until the time you wash 'em down your sink with water. :-) If stored either in crystallized or solution form, the only thing you'd want to keep in mind is to avoid evaporative losses by keeping the container closed, air-tight. The trihydrate crystals, if left, say, out in the open, will dehydrate; if the prepared solution loses water to evaporation, it'll eventually spontaneously crystallize.