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Extraction Of Lye From Ashes


Kzoppistan

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*Notes regarding a successful attempt to extract and concentrate (possibly) sodium hydroxide.

For an end result, it wasn't too bad. As an experiment, it was a horrible attempt. Since this was a first run, I was lax in any sort of proper procedure in regards to documentation or organization of that procedure. As it is one of the easiest chemicals to produce I didn't think to approach this with the seriousness that could have given me the opportunity to practice the skills of observation and record keeping needed for a scientific discipline.

So the lesson I learned was to approach every experiment with those principles in mind.

The ashes were obtained from burning the hardwood found in the park (incidentally, from my post about cooking beans) and from a friend's bbq grill, which was commercial charcoal. After settling I had about a 1/3rd full sauce pan and some hardwood charcoal chunks.

After letting it soak for a few hours on the stove on medium heat, I went through the process of straining out all the water (by using a ladle and coffee filters stretched over the collection cup) and separating the material (the mostly pure, silt, and chunks) into different tall cups. I continued to soak the material from each of the cups and straining it into the saucepan in a simmer/boil.

Observations:

7-18-11

* Despite the amount of soot present in the slurry, spills were surprisingly easily to clean, leaving behind no residue.

* The fine silt quickly settled to the bottom from the rest of the solution. The fluid resting on top was clear.

* Did leave a waxy residue on the hands. Which I found strange, considering that the fluid was as viscous as water and left no tacky residue when dried.

(Upon further reflection of this fact, I hypothesize two likely explanations, one, that the lye was actually breaking down the oils on my hands, a sort of local saponification, or two, more likely seeing as my hands were constantly run under water and very little oil was present, that it was breaking down the top layer of skin.)

7-18-11

* After 5 days no noticable amount of fluid has evaporated despite being subjected to the summer sun all day. I think this confirms the NaOH's hydrophilic properties.

7-26-11

* After 8 more days, a drop of about 2% of volume has occurred.

* The solution maxed out the pH test strips I obtained, ensuring that it is, at the very least, 8.4 pH.

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Persisting Questions:

I still don't know the exact pH of this solution. The test strips I had were crappy overpriced ones designed for aquarium usage (but, hey, surely I get points for even being able to find something that tests pH in walmart.) I think I'll spring for the real kind, but thinking about this makes me wonder if their is a mathematical formula for taking a diluted solution (getting the pH in the range of the test) and figuring out how much reduction of volume would be needed to increase the concentration....?

Also, I'm not exactly certain that I have sodium hydroxide, for it could be potassium hydroxide, which it is often confused for, or some other mixture. I must find more tests to apply to the solution to determine its exact makeup.

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Easier to steal soap from a religious institution or the local homeless shelter, or a quick half-day job at the local car wash, you're putting too much effort into the easy stuff.

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Congratulations you have made soap.

If you want a more exact measure of the pH of the solution you have made buy a pH meter or use universal indicator.

To test a soap for pH using universal indicator add a small drop to a soap sample and if the drop turns a blue or purple color then this indicates that there is an excess of lye in your soap and it should not be used on human skin.

Soap should have a pH of no greater than 9 although ideally it should be neutral and have a pH of 7.

Your solution maxed out the litmus strips with a maximum reading of 8.4 so I would not recommend that you use it on yourself.

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Dont you have any soapnut trees in your area? too much effort for soap you cant use to bathe with. The lye you made, would more likely than not also be harmful for any clothes you decide to wash with it. Maybe if you have a stable, you could use the lye to wash it out. or use some of it for dipping sheep.

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Some good advice, guys, thanks.

I plan on making a lot of this stuff. While soap will be one of the end products, NaOH has many uses. Though, really, I just wanted my own vial of corrosive liquid. He he he

Edit: Oh, also, lye is used in the soap making process to separate the soap from the glycerol. There shouldn't be any, or very little, actual lye in the end product. Soap works its magic not by the addition or subtraction of hydrogen, like with an acid or base, but by lowering the viscosity of water, which makes it better at removing grime and as an antibacterial agent also permeates the cells which weakens the membranes and causing them to rupture.

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Easier to steal soap from a religious institution or the local homeless shelter, or a quick half-day job at the local car wash, you're putting too much effort into the easy stuff.

Science is fun!

Congratulations you have made soap.

Close, but not quite yet. The addition of rendered fat is required to make soap.

buy a pH meter
I intend to, but the cheapest I found online is about 50 bucks while the most expensive is up in 1,000+ dollar range. Got bills to pay first.

Dont you have any soapnut trees in your area? too much effort for soap you cant use to bathe with.

Wow, ironically, I just learned what soapnuts are yesterday. And, no, we don't have any in the area, though it be cool if we did.

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Having washed my hands in a sodium hydroxide bath (multiple times) I can tell you that indeed the waxy substance was probably from skin.

Lye is used to break triester bonds to make a Na+ O Long hydrocarbon chain

Its used to make "micelles" which wrap itself around nonpolar, insoluble stuff so that it goes into water (with the ionic Na+ bond) effectively making anything nonpolar, polar.

I'd suggest you get something glass that's clean (bottle, cup, whatever you want that's not important), put your mixture in it, heat the crap out of it and see what's left behind. If it's white, it's probably a base. If its slippery, its definitly a base. I'd then suggest trying to recrystallize it by heating up some water, methanol, or ethanol, adding just enough to dissolve it, cool it off, and decant the extra solvent. Then boil off the excess solvent. It might purify it, depending on what you have.

also, http://www.sigmaaldrich.com/catalog/ProductDetail.do?lang=en&N4=221465|SIAL&N5=SEARCH_CONCAT_PNO|BRAND_KEY&F=SPEC

MUCH easier.

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"I still don't know the exact pH of this solution. The test strips I had were crappy overpriced ones designed for aquarium usage (but, hey, surely I get points for even being able to find something that tests pH in walmart.)"

You could get a pH probe, but they aren't cheap. An alternative would be to titrate it with an acid of known concentration. You can check when you get close to neutral using your pH paper, or (much more accurately) using phenolphthalein indicator (only $4 on amazon: http://www.amazon.com/Phenolphthalein-15-ml/dp/B003EE1KRQ/ref=sr_1_1?ie=UTF8&qid=1320122912&sr=8-1). If you know the volume required to titrate (you can get a buret pretty cheaply too) and the concentration of acid then you can calculate how much base it reacted with. If you know the original volume of base then you can get the concentration. The problem here, of course, is that you need an acid of known concentration.

"I think I'll spring for the real kind, but thinking about this makes me wonder if their is a mathematical formula for taking a diluted solution (getting the pH in the range of the test) and figuring out how much reduction of volume would be needed to increase the concentration....?"

I'm not sure what you mean here. Volume and concentration are inversely related; if you halve the volume you double the concentration. That doesn't double the pH though, if that's what you mean. pH should be related to concentration of OH (in a simple, aqueous, arrhenius system like you should have) as concentration of OH = 10^(pH-14).

Well done on a successful extraction though. I've always figured this is probably harder than it sounds.

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