Introduction: Crystallization of Homemade Sodium Acetate

Picture of Crystallization of Homemade Sodium Acetate

(Updated - see end, below)

If you've ever wanted to play around with sodium acetate, but you're too much of a nerd to simply go online and buy some from a chemical supply house - noooo, that'd be too easy, you want to make it, from scratch - then this instructable is for you.

Inside, I show the whole process, from baking soda and vinegar, through concentration and filtering, to final crystallization.

(Please note that you, not I, explicitly assume all risk associated with playing with chemicals, fire, or hot things. Use common sense. If you're not an adult, enlist the help of a parent. If you're an idiot, close your browser now before you burn yourself. And regardless, by reading any of the suggestions contained herein, you implicitly assume full responsibility for any and all accidents, burns, lacerations, ruptured spleens, loss of consciousness, death, shin splints, hangovers, spurned advances, or insolvency that may result. Seriously... use your brain.)


Update: When I first wrote this Instructable, I somehow got it in my mind that acetic acid had a boiling point that was slightly lower than water - this is incorrect! Acetic acid (ethanoic acid) has a boiling point of 118.1 C. This may change some steps and I may amend the instructions after I've had a chance to play with a few things; for now, though, when I talk about acetic acid boiling off in some of the steps, take it with a grain of salt (or at least sodium bicarbonate). :-)

Step 1: Preparation

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We'll be making approximately 330 mL (or 11 fl oz) of supersaturated solution.

You will need the following items:

- One 16 oz box of baking soda (sodium bicarbonate)
- One gallon jug of distilled white vinegar (acetic acid)
- Clean pot for boiling (5.5 Qt or larger)
- Another clean pot for filtering (4 Qt or larger)
- Pack of coffee filters (basket style, not funnel shaped)
- Small wire mesh strainer (big enough to hold a coffee filter)
- One cup measuring cup for pouring hot solution through filter
- Large clean cooking spoon for removing samples while boiling
- Small clean dark dish (e.g. custard cup) for holding samples
- Clean jar to hold final solution
- Distilled water (in case you over-boil the final solution - see step 14)

Optional Items - this is for a purification step that I found I needed (see step 6 for an explanation). If you decide to do this too, you'll need:

- Two cups granulated activated charcoal (from drug store or pet supply)
- Lab stand with burette clamp (if you have this)
- Funnel (if you have a lab stand; should be big enough to hold folded coffee filters - see step 10)

Step 2: Combining Baking Soda With Vinegar

Picture of Combining Baking Soda With Vinegar

The first step is the fun part!

Start by pouring all but one cup of the white vinegar into the boiling container (5.5 Qt or larger). (See note about acids below...)

Then, carefully add baking soda to the vinegar, small amounts at a time (no more than a tablespoon). Sprinkle it over the vinegar - don't just dump it in. If you add too much baking soda too quickly, the foam from the reaction may overflow your container.

Stir gently after each addition of baking soda to ensure no unreacted bicarbonate remains.

After adding about a half box of baking soda, you should notice the reaction starting to slow down. At this point you may want to reduce the amount of soda you're adding each time to a teaspoon. When a teaspoon (or less) of baking soda sprinkled over the solution no longer bubbles instantly but bubbles very sluggishly as it sinks into the liquid, it's time to stop. In my experience, it's taken roughly 12 oz of baking soda to get to this point (about 3/4 of a 16 oz box).

Finally, add the retained cup of vinegar to the liquid. Since judging the stopping point can be difficult to eyeball, I've found it helpful to withhold a small amount of vinegar and add it after the baking soda reaction almost stops. This way, we err on the side of excess unreacted acetic acid (which will largely boil off) rather than excess unreacted sodium bicarbonate (which will not, and can interfere with crystallization).

(Note on acids: yes yes, I know, we're all taught in science class that you never add X to acid, you add acid to X... Good! Brownie points for remembering that! This stuff, however, is a 5% solution of acetic acid, and personally, I'm not too concerned about burning my skin off. If you're too uncoordinated to keep it out of your eyes, wear goggles.)

Step 3: What's Going on Here?

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When you mix baking soda with vinegar, what you're really doing is reacting sodium hydrogen carbonate (more commonly, sodium bicarbonate) with acetic acid. Those components react, producing sodium acetate and carbonic acid; the carbonic acid immediately decomposes into carbon dioxide gas and water.

Step 4: Boiling Down the Solution

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The reaction is finished, and we're left with sodium acetate dissolved in water (with a little excess acetic acid). Now, theoretically, all we have to do is boil off the excess water until we get to a metastable supersaturated state! Ideally, anyway. I've had difficulty with this part, but let's continue...

Put the solution over medium-high heat and...

Yes, watch it boil. This isn't terribly fun.

You don't want a violent boil, but you are trying to reduce the volume of the solution, so a happy medium rolling boil should be fine.

Take note of the color of the solution when you start your boil... it should be colorless (more on that later).

The observant reader will correctly infer from the pictures that I'm making a double batch.

Step 5: Boiling, Continued...

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By the time you've reduced the solution to about half of its original volume, you will have noticed a color change. The solution will turn a distinct straw color, and will deepen to gold and amber the more it's reduced.

I've never found a satisfactory explanation for the color change - almost certainly it's due to organic impurities in the vinegar, owing to the natural source of the acetic acid.

My early attempts at getting the solution to a metastable supersaturated point all failed... I would get crystals if I let the solution cool overnight (with or without crystals having formed immediately on the surface of the liquid), but it seemed as though I could never get a supersaturated solution that, when cooled, would immediately crystallize when a seed crystal was introduced.

The next few steps attempt to remove as much of the impurities as possible, and it is up to you whether you want to follow them or continue boiling. (If you decide to skip them, go to step 11.)

(Note about the liquid volume in the picture - recall that I'm making a double batch. There's about a gallon of liquid in that picture, which is half of what I started with.)

Step 6: Purification With Activated Charcoal

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In this next series of steps, we'll use granulated activated charcoal to clean up any large molecules that could interfere with our desired crystallization later.

Stop boiling the solution and let it cool to at least close to room temperature. If you're impatient, you could cover the pot with a lid and set it in a sink full of cold water and ice (or if it's cold enough outside, cover it and set it outside for a while).

Step 7: Adding Activated Charcoal

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After the solution has cooled to near room temperature (or lower), add two cups of activated charcoal directly to the solution and stir gently for several minutes. Let it sit for a while, maybe ten minutes or so, before moving to the next step.

(In the first picture, I'm using a 1/4 c. measure - while this should be more than enough for the liquid in the pot, after I took the picture I decided to trade charcoal for speed, and ended up adding two cups, rather than setting up a funnel with a coffee filter to hold the charcoal, through which the solution could be poured.)

Step 8: Removing Activated Charcoal Granules

Picture of Removing Activated Charcoal Granules

We'll remove the charcoal granules and residual carbon particles in several steps.

First, pour the solution through a small wire strainer into your second pot. This will remove the granules and large particles. If you don't wish to save the charcoal for another project, you could skip this step (the next filtering step will remove them just as well, and it won't significantly slow down that step).

Step 9: First Filtering to Remove Most Carbon Particles

Picture of First Filtering to Remove Most Carbon Particles

Now place one coffee filter in the strainer and begin pouring your solution into it. Be careful not to overfill the filter - if some of the solution overflows the filter, you'll have that much more residual carbon to remove later. (It's not terribly important for this step, but will become extremely important in the next two filtering steps.)

After two or three fills of the filter, you'll notice the solution taking longer and longer to flow through the filter. When it slows noticeably, stop adding solution, pour the solution remaining in the filter back into the unfiltered container if you wish, and discard and replace the filter.

Continue until all solution has been filtered once.

Step 10: Second Filtering to Remove (most) Remaining Carbon Particles

Picture of Second Filtering to Remove (most) Remaining Carbon Particles

From this step forward, cleanliness will be of utmost importance to avoid getting dust or other contaminants into the solution you're working so hard to filter.

First, heat up the solution. It doesn't have to be boiling, but a hot solution will give a higher flow rate through the filter. While it's heating, wash out the empty container you're going to use to receive the filtered solution; if it's the one you poured the solution out of in the previous step, don't just rinse it - wash it with a little soap and a cloth or rag, to make sure all carbon particles are removed. When you finish rinsing it out, do not dry it with a towel or anything else! Shake it out, but leave it wet.

This filtering step will use six coffee filters at a time. If you're lucky enough to have a lab stand with clamps, you can use it to hold a funnel over your receiving container as shown in the picture (get 'em cheap at American Science & Surplus). If not, no worries - you can use your strainer instead. If you use the strainer, just be certain to wash it carefully beforehand (again, wash - not just rinse), so that no carbon particles remain that could contaminate the filtered solution.

Place six filters, nested together, into your clean strainer and set it in position on your receiving
container; or, if you're using a funnel, fold the stack of filters the way you were taught in science class - flatten them into a circle (all together, not individually), fold the circle in half, fold that in half again, and open one "pocket", forming the folded filters into a conical shape that fits into the funnel.

Using a one cup measuring cup (or similar small container with an appropriate lip for pouring), fill the filter with hot solution and let it drain. Keep transferring hot solution into the filter until all solution has been filtered. If flow through the filter slows down too much, you can replace the filters. Be careful not to overfill the filter.

While you're waiting for the solution to filter, blanch one head of cabbage... (oh wait, nevermind, that was for dinner...)

Step 11: Final Filtering Step

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If you've gone through the purification steps, this will be the final filtering step (if not, this is the only filtering step). The solution should be hot (if you skipped purification, turn down the heat so the solution isn't actively boiling but is just being kept hot, and go read step 10 which talks about setting up filtering with either a strainer or a funnel).

Clean the strainer and the receiving container and place the strainer in position (or, if using a funnel and lab stand, prepare the funnel). This filtering step will use ten coffee filters at a time. Place ten coffee filters, nested together, into the strainer (or, folded properly, into the funnel).

Using a one cup measuring cup (or similar small container with an appropriate lip for pouring), fill the filter with hot solution and let it drain. Keep transferring hot solution into the filter until all solution has been filtered, taking care not to overfill the filter.

Step 12: Final Boiling - Evaporating Excess Water

Picture of Final Boiling - Evaporating Excess Water

We're almost done! We've got sodium acetate dissolved in excess water; now we just have to evaporate off excess water until we get to the right ratio of sodium acetate to water.

This part can be tricky, and it takes a fair amount of patience.

Start by bringing the solution back up to a gentle boil. Take note of how the solution "sounds" while it's boiling - specifically, how the bubbles sound. When you start to reach saturation, you'll notice that the sound of the bubbles has changed - crinkly and crackly is the best way I can describe it.

By the time you've reduced your solution by about half (see pictures), it will have turned a golden color, and you'll see a spray of fine white powder all over the inside of your pot (and, unfortunately, on your stovetop around the pot - fear not, this stuff is extremely easy to clean up).

Step 13: Final Boiling - When to Stop

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When your solution has been reduced by a little less than half (less than that if you're not sure), we'll begin removing very small samples to see if they crystallize.

Place several ice cubes or a quarter cup or so of crushed ice in a small bowl.

Using a clean cooking spoon (or small ladle), transfer a very small amount of solution (one to two teaspoons) into a clean custard dish or other small bowl; to watch for crystal formation, it's best to use a dark bowl.

Cool the custard dish containing the sample by setting it on (or holding it on) the ice (first picture). After a minute or so of cooling, blow gently on the solution to see if you can trigger crystallization. If nothing happens, you can either discard the solution or return it to the boiling solution. If you're not absolutely certain it hasn't been contaminated with dust or anything else, it's better to discard it.

Rinse the custard dish, shake excess water off (don't dry it with a towel or cloth), and take another sample after 5-10 minutes.

Eventually, you should be rewarded by seeing long, thin crystals bloom across your sample as shown in pictures 2 and 3, below (taken about 2 seconds apart).

As soon as you see this happen, remove the solution from heat and put a lid on it!

Set the crystallized sample aside and cover it with another bowl or with plastic wrap for the moment - you'll want to save these crystals to use them as "seed" crystals later, to trigger crystallization in the cool solution when you're done.

Step 14: Cooling Solution

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Gently pour the still-hot solution into the container you've chosen to hold the solution for cooling and crystallization. You can use a glass jar, or a flask if you have one handy. Make sure the container is absolutely clean before pouring in the solution, and loosely cover the container (with a lid, plastic wrap, or plastic baggie).

Allow the solution to cool to room temperature. You can refrigerate it or put it in an ice water bath if you wish.

There's a small chance that some of the crystals on the side of the boiling pot could have gotten washed into your container when you transferred the solution, in which case your solution may spontaneously crystallize while it's being cooled.

If this happens, recharge the solution while still in the same container (see the next step for instructions).

If you find that the solution still appears to contain crystals after double boiling for 20 minutes or so, it's possible that you've boiled off too much water. In this case, try adding two teaspoons of distilled water to your solution, stirring gently with a clean spoon after adding it, and then continue double boiling for an additional 10-15 minutes. If crystals are still visible, repeat the process - they should dissolve eventually.

Step 15: Try It Out!

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After your final solution has cooled to near or below room temperature (it doesn't have to be cold), it's time to see if you can trigger crystallization.

Gently set your container down on a hard surface and remove the lid or covering. Using tweezers (or clean, dry fingers), pick up one or two small crystals from the test dish you set aside in step 13, and drop them into the solution.

You should see a mesmerizing "bloom" of long, thin crystals growing outward from the crystals you dropped into the solution, growing at maybe 1 cm / sec.

Pick up the container after crystallization finishes - note how warm it is. That's excess energy, effectively stored by the supersaturated solution, and released when the sodium acetate crystallizes.

Note also that the crystals you see in the container are actually sodium acetate trihydrate. Once crystallization begins, one sodium ion and one acetate ion will join together with three water molecules. If you were to gently heat these crystals, as the water is driven off, they'd turn pure white and powdery (instead of milky white and translucent), and you'd be left with anhydrous sodium acetate.

After crystallization, recharge the solution while still in the same container by double boiling it - place the container into a pot, fill the pot with water to or just below the level of the solution in its container, and bring the water to a boil. If your container is sealed (with a lid or stopper), you'll want to break the seal first to allow for air expansion and contraction during heating and cooling.

Here's a video showing about 525 mL of my solution doing its thing:

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.


robot797 (author)2010-04-09

can you eat it?

WhittleMario (author)robot7972017-11-21

Theoretically yes, but you probably shouldn't.

Sodium acetate is used to flavor salt and vinegar potato chips, but rule #1 of chemistry can be summed up in four words; 'never lick the spoon.'

Nidan3 (author)2017-04-24

I've made, but i was failed. Because after i was cooling the solution, it never made an crystal. Why did it happen, ya?

inspecter gadget (author)2016-09-10 ---- I made this 2 years ago & decided to have a play as I was recently making its other close relative Calcium Acetate ;))

Silverfieldwolf (author)2016-08-17

does it need to be that purified? or will it still work without having to worry so much about all of the purification?

This is a very concise & in depth instructable, I have however managed to achieved the same effect without the need for charcoal. It does not auto crystallise either as I can start it off with a seed crystal on most occasions ...

profquatermass (author)2008-10-29

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

KrisK31 (author)profquatermass2016-05-10

Sodium Acetate is the primary ingredient in homemade, reusable pocket warmers. Ideal if you live somewhere chilly.

TylerF5 (author)profquatermass2015-04-19

for fun.

*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. :-)

TrollFaceTheMan (author)2014-04-03

I think the most of the Dark Color your Getting is because of the Formation of Iron Salts with your pot... Aluminium, Glass or Teflon should work without that effect...

It is just a thought but is supported by a claim by a prior comment here, and I quote:

"I just made some 1/9/11 and I boiled over a medium heat and it came out clearer. I did 2 batches the first one I added the vinegar to the baking soda and that batch seemed darker, but the second batch I slowly added the baking soda to vinegar and it was clearer and a different consistency of crystal formation too. Give it a try and see what happens. Good luck. I did boil about 90 percent off too."

I believe he got a better result the second time because he more fully neutralized the acid reducing Iron salt formations...

Just a theory but it might help...

teachme! (author)2010-01-31

ok im going to sound really stupid but why do the crystals get hot when they crystalise im wondering?

steed1172 (author)teachme!2010-02-10

it has to do with "supercooling"

And Exothermic raection...

MAR! (author)2010-12-28

About at what temp. should i boil this?

ycsmela (author)MAR!2011-01-10

I just made some 1/9/11 and I boiled over a medium heat and it came out clearer. I did 2 batches the first one I added the vinegar to the baking soda and that batch seemed darker, but the second batch I slowly added the baking soda to vinegar and it was clearer and a different consistency of crystal formation too. Give it a try and see what happens. Good luck. I did boil about 90 percent off too.

TrollFaceTheMan (author)ycsmela2014-04-03

I believe because when their is still acid in the Vinegar it make salts with the iron giving it a dark huegh...

MAR! (author)ycsmela2011-01-13

But yours did crystallize like it should, right? Not just dehydrate and turn to mush in the pot? My stove has numbers so to me med high is a little over 3.
Heres how 4 attempts ended in my case, help me please.

nannysaunders01 (author)2013-08-18

hi i just did yet another batch, as my vinegar had sediment in it, i filtered it and it did crystalise afterwards, but still a bit mushy. I recharged it by boiling it down, and have a bottom layer of milky white slimey substance - the brown amber liquid sits on top - what is this?

nannysaunders01 (author)2013-08-18

hi i have asked a few questions on other instructables, i know all experiments are done differently but following nurdrage he uses 1litre vinegar to 3tbls bicarb soda. i have converted yours and you use 1 gallon (3.7l) vinegar to 12 oz ( 20 tbls ) bicarb soda.
Why do you use so much more bicarb soda? and is this why with me using 3 tbls my hot ice is not solidifying but going mushy? thanks

bumsugger (author)2011-12-23

Can anyone tell me the prime use for this chemical,as it appears to have slipped my mind? Thanks.

foxworrior (author)bumsugger2012-10-07

this is used in re useable heat pads, the ones that crystallise and heat up when you snap the little metal disc.

Johenix (author)2010-01-13

Why not use washing soda, sodium carbonate?  It has twice the sodium content.

Wesley666 (author)Johenix2010-11-21

If the chemical formula is different it could change the reaction. It may give the same product in the same quantities, but you would have more waste material, or you would get a bit more product but using alot more vinegar. I don't know the chemical formula for washing soda, but these are some reasons why it might not be used...

jhaas1 (author)Wesley6662012-04-19

I've used washing soda for this.

Checked it, the formula: Sodium Bicarbonate Decahydrate

So it's no problem.

leontom1 (author)2012-01-17

Is This Used To Make Hot Ice????

pc5 (author)2011-11-28

maybe you had trouble seeding "supersaturated solution" because it wasnt supersaturated, if u attempted to seed after using the activated charcoal, which may adsorb your desired product? Just an idea.

I'm skimming, appreciate the cool instructable :)


skulendran (author)2011-09-28

hey, i was just wondering. you said if you heated the crystals, water would be driven off and it would produce a white powder (anhydrous sodium accetate). but how would you go about doing that? could i just put the crystals in a pan on low heat?

julsscott (author)2011-02-04

can this be done as a school project? can you do the prep at home and then have a deminstration of some sort at school? How would that work,,,,,,any suggestions?

SenileFelineS (author)2010-06-05

I thought that once all impurities and contaminates are removed it would turn into a slightly tinted clear solution, not stay amber colored.

At least, that's what pure, lab grade sodium acetate should look like....Is there a way to achieve that?

Wesley666 (author)SenileFelineS2010-11-21

If you look at store bought manufactured hot packs that use Sodium Acetate, they are almost always amber too. Don't know why but they are, so I think its safe to say that it doesn't matter if you can get it clear or if it stays amber colored.

Wesley666 (author)2010-11-21

I was just screwing around and wrote out the vinegar and baking soda reaction equation and realized you could make hot packs with the product. Then I checked Instructables to see if anyone had done it, and they had. Darn, thought I had a really good, original idea too...

Good 'Ible though! :D

Erebo2005 (author)2010-10-08

Good, clear and simple. Thanks


ryanmuller (author)2010-05-01

All other sodium acetate I've seen dissolves clear in water... is there any way to achieve this? I need it to look like plain ordinary water.

dragon181818 (author)ryanmuller2010-06-23

For a "water" look, you will need either: laboratory grade chemicles or to buy it online. About 500 grams is like $25 If you need help, consult youtube :)

briannac1 (author)2010-02-28

Do we HAVE to have activated charcoal?

JohnJY (author)briannac12010-03-16

No, look at NerdRage's video.

hawk 1sr (author)2010-03-04

i followed the directions, but some how i just have yellowish-gold substance that wont melt all the way and theres chunks of the krystals that i cant get rid of.  Your stuff was clearish, mines not. what did i do wrong?

Pizzapie500 (author)2010-01-27

Cool! But I'm not "too much of a nerd" so I bought the sodium acetate of ebay. For some reason, mine was like 20 times faster than yours, and it would go off w/o me putting a crystal in it. I made this like a year ago and kept some in my fridge. Do you think I can bring it to school to show my teacher? Or will it go off in my backpack?

KonstantinosM (author)2010-01-16

I Just wanna ask if i can use Paper towels instead of coffee filters!!! I'll comment again if I am successful P.S. If i misspelled successful or misspelled its because i'm Greek! 

ryguy428 (author)2009-06-26

For 1 gallon of white vinegar, what's the yield of dry sodium acetate? Assuming vinegar has a density of 1 g/mL, I'm getting a theoretical yield of about 250 g sodium acetate.

corradini (author)ryguy4282009-12-21

Water is 1g/ml; acetic acid is 1.049 g/ml; distilled white vinegar (in the US) is 5% acetic acid - easy math: 1.00245 g/ml (close enough to 1 for government work).

Sodium Acetate is C2H3NaO2, molar mass 82.03g; density 1.528 g ml ^-1
Acetic Acid is C2H4O2, mm is 60.05, density already given above.
Sodium Bicarbonate is CHNaO3, mm 84.01, density 2.173
Water: If you (anyone, not ryguy428) don't know formula or density of water, you're NOT ready to try chemistry. Or graduate from secondary school. Or vote. Molar mass is trickier - it's about 18.0152833 - depending on a few things... ;-)

Rxn: CH3–COOH + Na+[HCO3] → CH3–COO Na+ + H2O + CO2

Do the molar balance, substitute for density, and you'll get theoretical yield.
(Hint: 84g of soda should give about 82g of sodium acetate. If you were using 8% vinegar, you'd need about 750ml. If you can make that calc come out right, do it for 5% and you'll have your answer.)

AndyGondorf (author)corradini2010-01-09

Nice explanation of the sums! Very clear.

A tad harsh about the assumed knowledge though, there are lots of bright and creative people who don't need that info.

What things did you have in mind for the water that would affect the mass within the accuracy limits and significant figures of your other data?

corradini (author)AndyGondorf2010-01-09

A) Thanks!

B) I'm reminded of the story that someone asked Einstein what the quadratic equation was and he (supposedly) said "I know where to look it up - why should I bother memorizing it?"

<inappropriate political rant> I shudder to think, though, that one can get through high school without knowing that water is H2O, and the absolute *basics* of the metric system (i.e., that mass, length, and volume (and, in fact, temperature and energy, by extension) are tied together by the single, simple fact that 1 cc of water weighs (masses, actually) one gram. Now, I was being a wee bit sardonic, but -- heck, the U.S. makes people applying for citizenship learn all kinds of stuff that (generously) 40% of voters couldn't answer. I sure don't want people voting on global warming, stem-cell research, and energy policy, if their scientific knowledge base doesn't include the absolute scrapings from the bottom of the barrel... ;-) </ inappropriate political rant>

More fairly - I liked Heinlein's quote (Google "specialization is for insects"). IMHO, and only my opinion, no adult in the modern world should be considered educated without a grasp of some plurality of some basic knowledge set. Defining that is politically 'difficult' - but I stand by the concept:

Things are made of atoms. The earth is round and goes around the sun, because of gravity. Traits that help survival get passed on. Wood burns, which gives light, warmth, and tasty food. Men and women think and prioritize differently. (Some) germs cause (some) diseases. Too much sunlight is bad for you (as is not enough). Stay away from wild animals. Seeds grow into plants, which need light and water. Don't eat wild mushrooms, or build a fire in an enclosed area. Yadda yadda. And water is H2O.

C) Nothing - given your stated constraints. Hence my smiley-face. I was winking at those who'd say something about "well, that's liquid phase - it's more like 0.917 as solid - you have to say at s.t.p. (or, more accurately, at the triple-point). And - what about deuterium content? And never mind supercooling, or allotropic forms....

AndyGondorf (author)corradini2010-01-09

A : You're welcome :) It's great to see someone apply some real analysis instead of just spouting second hand partially realised facts.

C : I know, and apologise.

I was baiting you a bit about showing off your (absolutely correct, by the way) knowlege.   :)
The constraints in fact were outlined by your good self in terms of the sig. figs., phase and isotopic accuracy you supplied for the other reagents.  Naughty of both of us, I apologise. ;)

B : (sorry about the chaotic order, this reply is harder to compose since it's more subjective than A or B!)

An interesting one and, as you say, very subject to opinion.

Being a real "science geek" (also with degree level chemistry) at heart myself I hear and understand everything you're saying about basic knowledge.
However, being in the education biz myself tends to steer me away from giving particular importance to my own skill set and knowledge base.

Indeed, to quote your own words "no adult in the modern world should be considered educated without a grasp of some plurality of some basic knowledge set.", which I agree with but for the sake of argument we could go with the knowledge set allowing us to create works of art or music.

Neither of these knowledge sets requires the formula for water, but I don't think we can eliminate some of our great creative individuals from the voting process just yet!

Having said that, if you have a look at you'll see a wonderful teardown of some science "facts" put forward by some our of leading "celebrities"

I certainly grant that most visitors to this site will also veer towards the science geekdom side of things so perhaps your argument stands in many cases, but I still think not all.  Perhaps I'm just playing devil's advocate!

So there! *puts away his own soapbox for the time being*

Apologies to anyone I've bored with my own rant :)

corradini (author)AndyGondorf2010-01-11

Oh, we're having fun now! >;-) Good responses, every one.

"also with degree level chemistry" -- whoops! I never said that. In fact, I've had only 2 chem classes, and that's generous - one was junior high and the other high school, and I didn't excel at either. I'm self-taught over the last 4 years or so, with some help from a PhD friend or two and a lot of journal articles....

I agree fully about not eating too much of one's own dog food. My educational and work background is management (MBA) and IT, so my criteria also include basic computer & Internet skills, balancing a checkbook, a rudimentary grasp of economics ("supply and demand", as Father Guido Sarducci's 5-minute University had it), and how interest works.

People ought to have some clue about atoms, electricity, magnetism, light/EM radiation, the Big Bang, stars, and the universe. Ditto nutrition, what a liver's for, vitamins (esp. folic acid), and first aid. Ditto basic literacy, numeracy, and communication skills. Ditto basic geography, history, political science, philosophy, literature, art, and music.

I'm also a big fan of ideas like: everyone should learn to speak another language, play a musical instrument, (try to) get good at a sport, travel to foreign countries, hike/camp in the wilderness, go sailing out of sight of land, summit a mountain, be able to cook one good meal, etc.

OK - I've now fully digressed into a complete rant about how the entire human race should conduct its affairs. So - I'll see your apology, and raise you. ;-) I realize this is not the forum for my opinions on this -- I'll take it offline henceforth.

AndyGondorf (author)corradini2010-01-16

LOL...I'm calling "stalemate" on this one since we've both gone a touch off-topic, but it's been fun for me too. :)

ryguy428 (author)corradini2010-01-09

Rereading my comment, I think I might have been unclear - I wondered what the ACTUAL yield was so I could figure out the efficiency of the reaction and purification. I knew how to get the theoretical yield, I just didn't have a balance at home to measure the actual mass of the sodium acetate that I made. :-)

But thanks for working out the stoichiometry anyway!

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