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Problematic Synthesis of Sodium Acetate

  1. Jan 15, 2006 #1
    Problematic "Synthesis" of Sodium Acetate

    Greetings all, was hoping someone more knowledgable than me (not a tall order, by any means) could point out some errors in my theory or method here.
    Fascinated by pictures of supersaturated NaCH3OOH solutions crystallizing at the drop of a dime (or a seed crystal, at least) and thinking a synthesis should be easy (vinegar and baking soda, right?) I decided to give it a try.
    I started with 250mL of White Distilled Vinegar, which I used in lieu of diluted acetic acid (which I don't have). The labeling claimed the contents had been adjusted to 5% acidity, which, after a little web searching, I understood to be 5% acetic acid by volume (could this be the error?).
    So my calculations were:
    .05 * 250 = 12.5mL of acetic acid.
    12.5mL = 12.5cm^3;
    12.5cm^3 / 1.049g/cm^3 density of CH3OOH = .198436... mol (of course taking that many decimal places is absurd with the inaccuracy of my measurements but w/e.)
    .198436 mol * 84g/mol (molar mass of baking soda) = about 16 2/3g NaHCO3.
    So I weighed 16 grams of sodium bicarb on my fairly inaccurate gram scale and stirred it bit by bit into my 250mL distilled vinegar.
    A few more calculations were made here.
    Using the density of Sodium Acetate (which I supposed I probably had a good bit of, though probably contaminated with excess sodium bicarb/acetate/unlisted anti-caking agent from baking soda/impurities in vinegar) I figured about 21mL of water would be the saturation point.
    I boiled the liquid down, using the microwave as I wanted to avoid boiling any possible residual acid in a metal container on the stove, to a little below this point (can't tell precisely; was below the accuracy of the container) and it abruptly turned into a bunch (about 16g) of chunky white powder.
    I had been hoping for something far more crystalline, but still had hope because within the chunks, shimmering bits of salt-like crystal were visible.
    Seeing the decomposition point for sodium acetate (and baking soda, etc, for that matter) was pretty high, I decided to heat this, since it might be sodium acetate tryhydrate, which can be returned to a liquid state at 54C. At this point the texture changed utterly and it seemed two disparate substances were present: a still mostly white, VERY chunky material, and a browned, extremely fine grain powder.
    The white chunks are definitely not sodium bicarb, and they are not very soluble. The brown powder leaves a glassy coating if I dissolve it in water let dry, but I have so little of this left after doing a few careless tests that I am not able to test if it is mostly sodium acetate.
    Any thoughts? Should I retry the whole bit reading the 5% acidity labelling as acetic acid by weight? (with my inaccurate equipment and the parity of h2o's and ch3ooh's densities this may not even make much of a difference.) Are the unexpected characteristics of my result due to impurities in the vinegar?
    lates,
    cotarded.
     
  2. jcsd
  3. Jan 15, 2006 #2
    Update:
    Definitely found some rectangular prism type crystals all over a dried solution of the brown powder. By letting the slurry run slowly (as it dries it becomes very viscous) I was able to purify crudely the crystals as whatever is lending the brown color is the last to come out of solution. Anyone know a better way of taking advantage of this property to separate whatever it is from my crystals (which I hope are sodium acetate)?
     
    Last edited: Jan 15, 2006
  4. Jan 15, 2006 #3

    mrjeffy321

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    Not that it matters a great deal in the calculations but, I am pretty sure the 5% vinegar means 5% acetic acid by weight, and not volume.

    Also, if your boiling down your solution trying to percipitate out the solute, you will end up with mostly a powder or very small crystals as you have found. If you want to make larger crystals, you need to slow down the process and keep the solution undisturbed, otherwise it will never reach the super saturated point.

    I would re try the experiement on a slightly larger scale to give you more product to test.
    You might try different ways of crystalizing out the Sodium Acetate, set some of the solution aside in a warm place where it will not be disturned and let it sit there and evaporate for a while. Take some more of the solution and boil it down dry, then take this and try to saturate another sample of the solution.
     
  5. Jan 15, 2006 #4

    GCT

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    There are flaws in the experiment and careless assumptions, if I find an actual online source for a related experiment I'll post it here and the chemistry webpages subforums.

    Here are some things to consider

    I'm not familiar with the composition of vinegar, nevertheless, a solution of acetic acid in its buffer region, will have comparable concentrations of acetic acid and acetate (equilibrium of a weak acid). This means that significant concentrations of acetate are already present in your vinegar.

    The best crystals require some time and more lengthy and tedious processes. But to simplify, you will first need to dissolve the vinegar and baking soda completely, you've used too little water, so nothing may have dissolved in the first place. Then you need to filter out most of the excess fillers.

    You used a microwave, I'm not quite sure if this will be suitable with your compounds of interest, you want a more continuous process such as heating the solution instead, rather than with bombarding it with a microwave. The brown substance you've found might be due to impurities in the microwave or for anything else for that matter (such as sugar, or iodoform, compounds that have been oxidized).
     
  6. Jan 15, 2006 #5
    EDIT: Started writing before GCT's post, will address it at bottom.

    Thanks for the reply. Back from work now so I'll probably mix another, larger batch in a couple minutes and give alternate methods a try.

    (") But you're not getting off that easy! (" he says, powerless to control others) I have a couple more questions.

    First, I still have no idea why heating the original chunky powder caused the product to divide into two seemingly disparate substances: Large hard white relatively insoluble chunks and a slightly brown very fine powder. What happened? Are these two different chemicals, or what? Did I decompose some of the sodium acetate, or are these impurities that are being educed?
    Second, as far as the more instant crystals goes, I'd planned on heating some water to boiling temperature, letting it settle for a second, then using it to barely dissolve a bit of Sodium Acetate. Then I would let that solution cool and I'd have a supersaturated solution, I could drop a seed in, right? (neglecting to mention the thirty tries it'd take to get a batch that didn't auto-crystallize)

    As far as the by volume thing goes: Doh! Thanks for the correction.

    EDIT: Addressing GCT
    I'd apologize for my ignorance, but I've still got more coming.
    I understand that weak acids, like acetic, don't dissociate completely into ions; but doesn't the sodium bicarb-acetic acid reaction drive the ph towards 7 and cause more of the undissociated acetic acid molecules to dissociate? And what does "this means that significant concentrations of acetate are already present" mean? Of course there are; if there weren't my reaction wouldn't take place. You're talking anions, right?

    Vinegar is 95% water - there is plenty in 250mL to dissolve the 16g of baking soda and already dissolved acetic acid.
    Remember, this is white DISTILLED vinegar. There aren't any visible fillers, if that's what you're thinking. Now, there may be impurities; those, I don't know how to handle, and I'd be willing to bet they're behind the brown.


    Thanks.

    lates,
    cotarded
     
    Last edited: Jan 15, 2006
  7. Jan 16, 2006 #6
    Success

    Mrjeffy, it seems your advice did the trick. I redid my calculations based on weight and after boiling the water to my guesstimate of the saturation point (I figured I'd just get the powder again and hope there wasn't as much brown) I went to the bathroom and came back to see crystals blooming. I now have a large bowl frozen solid with beautiful crystals. Was wicked to watch the needles bloom radially, each needle eventually bursting to yield more.

    Now I'll see if I can melt these and repeat the spectacle.
    thanks all.
    lates,
    cotarded
     
    Last edited: Jan 16, 2006
  8. Jan 16, 2006 #7

    mrjeffy321

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    well congradulations. You dont happen to have any pictures do you?

    To try to recreate your results from the original post, I broke out some of my own chemicals and gave it a shot. Using a 20% vinegar solution, I reacted enough Sodium Bicarbonate to neutralize the solution, leaving a clear liquid. i then took this is boiled away the water leaving behind some white power which I could blacken and melt with my butane lighter. This black stuff was interesting to watch, while a liquid, it would move around in the powder (but not mix), then when it cooled, it would turn dark brown and form snake-like crystals as it hardened.
     
  9. Jan 16, 2006 #8

    GCT

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    I've read about this experiment just yesterday, and it seems that a stoichiometric reactions is important for the purity of the product, otherwise you'll have bicarbonate or acetic acid in the final solid. The bicarbonate is used to "neutralize (not quantitatively as with a strong base)" the acetic acid, and in my post I was referring to the original concentration of acetic acid and acetate that you'll need to take into account if you were to perform an exact experiment (the initial concentration of acetic acid to neutralize); but heating the solution ensures a relatively higher concentration of acetate nevertheless. But perhaps this procedure is much simpler than I thought.

    Try dissolving your purified solid into water, in fact try out a litmus test with both your separate vinegar and bicarbonate solution (acetic acid should have an acid pH, while bicarbonate should have a basic pH). The pH of your dissolved purified solid, if it was truly acetate, would be significantly above 7. You may also want to test the pH before boiling the water off.

    When you react the bicarbonate and acetic acid, do you see any bubbling? I'm not trying to make this experiment more difficult, but it's essential that you check your assumptions of what the final product is with verifiable evidence, at every angle possible. That way, you can truly be satisfied that it is the product you desire, or if not, perhaps a bit of disappointment will ensue, but you'll learn a lot more in the long run. Obtaining a pure product is never easy, and usually the boiling off of a solution always produces some type of crystal (digestion promotes crystallization).
     
  10. Jan 18, 2006 #9
    :rofl: Let's see. I just mixed thirty something grams of baking soda into a half liter of vinegar. Bubbling?
    Yes, I saw so much bubbling - frothing or vigorously foaming might be a better description - that I had to take about 2 minutes to dissolve the 34g of NaHCO3, dropping it in bit by bit, to avoid the loss of my solution over the sides of my quite large vessel (a deep dish with probably >1 liter capacity and 90 degree walls).

    I took a few pictures, will be posting them soon.

    As far as product verification, I'm confident that the verdict is in. The stuff behaves exactly as Sodium Acetate should; I replicated the instant crystals demonstration that inspired this in the first place with a successful crystallization from a super sat solution so rapid that it formed a column on top of a seed crystal as I poured the former from a cup. This stuff is incredible. I dipped my hand in a cup of it and it came back looking like some grotesque chandelier.

    lates,
    cotarded.
     
  11. Nov 10, 2006 #10
    formula

    Sounds like lots of fun...

    I'll have to try it with the kids. Can't wait for the pictures...

    Thanks!

    Tom
     
  12. Nov 21, 2006 #11

    chemisttree

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    Some thoughts about sodium acetate (my "Ice-9" moment)

    While employed at a research institute years ago, a client came in to have us analyze some impurities in their sodium acetate handwarmers. These handwarmers were sold in many of the big box stores in the US. Large crystals that resisted dissolution during reactivation were slowly growing in the bags. While these crystals were not a significant problem when the reactivated bags remained undisturbed, any handling or transporting caused some of these contaminated bags to spontaneously crystallize. Definitely a problem.

    The crystals could not be easily isolated since any tampering with the bag caused the thing to crystallize into mush but we eventually did get the crystals out by holding the temperature slightly above the exotherm temperature of the handwarmer.

    Guess what we found? All of the analyses I performed indicated pure sodium acetate! A list of the tests:

    FTIR (Fourier Transform Infrared)
    HPLC/MS (High Performance Liquid Chromatography - Mass Spectrometry)
    Proton and Carbon-13 NMR (Nuclear Magnetic Resonance)

    The client was going nuts looking for impurities at the plant, gathering Certificate of Analysis from the sodium acetate vendor, looking at water purity, studying extractable materials from the plastic bag and the steel clicker button! To make matters worse, this product was the ONLY product the company made!

    What was the impurity?

    ANHYDROUS SODIUM ACETATE!

    The normal mode of activation for the supersaturated (aqueous) sodium acetate produced the trihydrate and some heat. The manufacturer supplied sodium actate as the trihydrate. It is analyzed but the Certificate of Analysis does not analyze for various hydrated forms of the stuff. There were a few tiny anhydride seed crystals in the product.

    How did I finally discover this anhydride? I was inspired by Kurt Vonnegut, of course! (Cats Cradle) If you don't understand how this could help explain this other form of sodium acetate, read this (http://en.wikipedia.org/wiki/Ice-9) This other form of sodium acetate melted at a much higher temperature than the trihydrate.

    Trihydrate MP = 58C
    Anhydride MP = 382.2C

    All those fancy (expensive) tests gave no joy but my lowly Meltemp melting point apparatus nailed it!

    Thanks Mr. Vonnegut...
     
  13. Nov 21, 2006 #12
    Is this Freezing? Why not?

    Okay, so is this a physical or chemical change? This may sound like a novice question, but it's more subtle I think.

    What's the real difference between freezing and crystallizing out of solution? Or vice versa, what's the difference between melting and dissolving back into solution? It's still hydrogen bonding, yes? So how is this stuff not "freezing?"
     
  14. Nov 21, 2006 #13

    mrjeffy321

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    What is the "this" to which you refer?

    Freezing is the process where a liquid turns into a solid. When something precipitates out of solution, a substance which was originally in solution comes out of solution into the solid phase. A common type of solution is a water / aqueous solution, and in the case of ionic compounds, the solute is disassociated into ions in solution. So, for example, if Sodium Acetate goes from the aqueous state into the solid state by means of precipitation, Na+(aq) and C2H3O2-(aq) ions form NaC2H3O2.
    When melting, a solid turns into a liquid, but it is not necessarily in solution.
    There are other types of solutions beyond just solids dissolved in liquids.
    Hydrogen bonding is a special type of an intermolecular force of attraction which occurs when Hydrogen is bonded with one of a set of specific other elements which causes a partial dipole attraction with other molecules. Not all substances have Hydrogen bonding, but Hydrogen bonding can play a large part in the temperatures a substance will freeze/melt at (example water compared to other, similar types of molecules).
     
  15. Nov 21, 2006 #14
    Okay, so melting doesn't require a solute. But looking at the molecular change itself, is there a difference in what the molecules do? Melting a salt breaks its ions apart, no? Same with dissolving. And the reverse seems to hold too. So what's the molecular difference, besides having a solute to help the process?
     
  16. Nov 22, 2006 #15

    chemisttree

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    The sodium acetate supersaturated solution is just that... a solution. The "supersaturated" nomenclature merely indicates that the solute is present in greater concentration than is seen at equilibrium at the conditions that exist at the time (pressure, temperature).

    If you were referring to the growth of an anhydrous phase in an aqueous environment (my previous post), it is not unheard of. I only mentioned it because it is unusual to my typical experience. If you are asking if a chemical change is occurring when a supersaturated solution precipitates either a trihydrate crystal or an anhydride, I think that you will find that precipitation events are not usually classified as chemical transformations even though in this case different materials are produced (but not different by much in a chemical sense).

    The solution in the handwarmers is definitely a solution. The sodium acetate is present in greater than 50.4 grams per 100 mL of water but it is definitely not pure sodium acetate trihydrate.

    Regarding what the difference between dissolution and melting, the answer is that in dissolution a solvent is present. It is always the case that the presence of this solvent will dissolve (it looks like melting to some) the solute to produce a saturated solution at equilibrium, the concentration of which is a function of temperature and pressure. Supersaturated solutions are made by carefully changing the conditions (lowering the temperature) so that the solution is no longer at equilibrium. In the case of supersaturated sodium acetate, sometimes called supercooled, the crystallization gives off heat (exothermic) until equilibrium conditions are reestablished. At room temperature these handwarmers still have an aqueous solution of sodium acetate contained in the capillary spaces between the crystals of the trihydrate. Reactivation by heating might seem to melt the crystals but in fact it is actually dissolution.

    Regarding the observation that the melting of a salt amounts to the breaking of ionic bonds thereby mobilizing the salt into a fluid state which is essentially what happens in a dissolution process, I agree with the following caveat... For a given substance, melting happens at a unique temperature at a given pressure. It is a fundamental property of that substance whereas dissolution can occur over a very wide temperature/pressure regime. Dissolution in general does not require hydrogen bonding but in this case it is definitely important. Likewise, dissolution does not require ionization of a salt as is the case in this example. In this case, a significant amount of sodium acetate is present in solution in its unionized form. This can be determined experimentally by plotting electrical conductivity against concentration. The resulting graph is nonlinear in the more concentrated regions.
     
  17. Jan 26, 2007 #16
    chemisttree,

    You spoke of an issue a company was having with the anhydrous form of sodium acetate being in solution with the trihydrate and that the anhydride was always present and could act as a seed crystal. I was wondering what the solution to this problem was. It seems to me that the problem was that there wasn't enough water in the solution because adding more water would hydrate the sodium acetate into the trihydrate form and put it into solution. Were the bags thin enough that they were dehydrating through the membrane or was it a manufactoring issue?

    I'm taking particular interest in this thread (stumbled onto it from a google search). I'm a high school chemistry teacher and am thinking of cooking up a sodium acetate supercooled/super saturated lab.
     
  18. Jan 29, 2007 #17

    chemisttree

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    The bags were plasticized PVC and no leakage through the bag material was assumed. The service life of unaffected bags was greater than 5 years unless the thing was overheated in a microwave.

    I don't know what was done to fix the problem at the plant but your idea would work. I was told that the exotherm temperature was a function of the concentration of the sodium acetate, but I never measured it. I believe that adding more water would change the product's properties unless measures were taken to later remove the extra water.

    Some other ideas that might work are 1) heating for a longer time before initially filling the bag, 2) ultrafiltration of the hot sodium acetate solution, 3) centrifugation of the hot sodium acetate solution. I don't know what would be best in the manufacturing environment.

    One last comment, the anhydrous form was not always present in the material from the manufacturer of the sodium acetate. This might never be a problem for you or your students but you can make it one for instructional purposes. If you do decide to dope the trihydrate with the anhydride, be sure to grind it very finely and add only a trace (don't ask me what I mean by "trace"...). The growth of the anhydride phase may not happen every time but you may get a statistical difference between undoped and doped product. Anoher processes that can crystallize an anhydrous phase from an aqueous environment in nature is seen for gypsum (calcium sulfate). Gypsum is usually crystallized as a dihydrate but can crystallize as the anhydride from a saltwater (NaCl) solution.
     
  19. Jun 4, 2007 #18
    Making Sodium Acetate using 99% pure Acetic Acid

    Hi...

    I am an English Teacher in S. Korea and I have a science class.....I heard about the 'Freezing Ice' experiment using Sodium Acetate.

    My school doesnt have and sodium acetate, so i thought i would try to make some. I came across this thread....

    I can get vinegar in korea of various strengths: 7%, 19%, 40%.....But these are not distilled white vinegar. It is apple vinegar and not white.

    BUT, I can also buy 99% pure white vinegar at the supermarket. So, with what am i better off making the sodium acetate?? The yellow apple vinegar, or the pure acetic acid??

    If i use the 99% pure Acetic Acid and disolve as much Sodium Bicarbonate in it as I can, should that do the trick? Or should i dilute it with water and work from there???

    Thanks for your help
     
  20. Jun 4, 2007 #19

    chemisttree

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    I would use distilled white vinegar. Natural vinegars from grape or apple juice have other compounds in them that may interfere. A neat experiment for your students would be to make sodium acetate both ways and look for the differences in the 'Freezing Ice' experiment.

    I would add excess vinegar to the bicarbonate and evaporate away the excess liquids. The solid sodium acetate can be isolated and recrystallized several times to purify it... another idea for an instructional experiment.

    99 percent pure white vinegar at the supermarket is probably only 5% to 10% acetic acid. Pure in this sense has a different meaning. 99% acetic acid is called glacial acetic acid. If you could get your hands on that stuff it would be best, of course. It is somewhat hazardous material and should be handled with gloves, rubber apron, goggles, ventilation, etc...
     
    Last edited: Jun 4, 2007
  21. Jun 4, 2007 #20
    OK, well.....I tried it. I believe the acid i have is real acid. I wouldnt want to touch it (i used rubber gloves), the fumes are VERY strong.

    Anyways, after adding excess 'vinegar' to the sodium bicarbonate and boiling it away, I do get crystals forming on the insides of the pot. I take it these are what I am after?? If i boil it dry so i am just left with crystals, collect these and then disolve these into water and make the supersaturated solution, I should be in business??

    I tried this, but it seemed like there was a little bit of liquid that wouldnt evaporate off the crystals - they stayed like a 'gel'.

    I also added water and tried to guess when the solution was supersaturated, but this didnt work. I could not get the 'ice' to form after the liquid had cooled.

    Thanks again for your help
     
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