how to neutralize high-pH chemicals for disposal

Name: Vicki


Ph paper 100 strips

Ph Paper, 100 Strips

100 strips with 1-14 acid/base chart.


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Jacquard soda ash dye fixer 1 lb.

Jacquard Soda Ash Dye Fixer

Soda ash is a mild alkali that promotes the chemical reaction between fiber reactive dye and cellulose fiber. Soda ash is also known as sodium carbonate, washing soda, or sal soda. One pound of soda ash is required to activate and "set" Procion dye for approximately 15 T-shirts.



ARM & hammer  super washing soda

Arm & Hammer Super Washing Soda

Super Washing Soda works as an all natural detergent booster for cleaning your laundry and can also be used throughout your home as a household cleaner. Washing soda crystals contain sodium carbonate and water.


Message: I am embarrassed to admit that I do not know how to neutralize my chemicals before discarding down the drain. I do have good pH-indicator strips for pH 0-14 from Earthues (the colorpHast/Merck German brand) which I use to get my 11 pH for dyeing cotton. I assume I use these and add chemicals to neutralize. Might you supply the details of how to do each chemical?

Do you have a septic system? Septic systems are pickier about the disposal of high- or low-pH chemicals than municipal sewer systems.  If you have a septic system, it's a good idea to at least partially neutralize any large volumes of very high-pH or low-pH mixtures, and to be careful to dilute small amounts. It is important to the functioning of a septic system to be careful of the health of the bacteria that keep the system running smoothly; it would be bad for these bacteria to dump a lot of one extreme or the other. 

You don't necessarily have to neutralize moderately high-ph leftover solutions, especially if you are using a municipal sewer system. The small amounts you will dump down your drain are diluted by the thousands of gallons of water that go through the rest of a multi-user system. (Obviously, things are very different for high volumes of industrial waste, because of their quantity.) Since I'm using a municipal sewer system, small quantities, and pHs that rarely go above 11, I do not usually neutralize my solutions at all. For a large indigo dye vat and a septic system, I would certainly want to neutralize it before disposal.

It's kind of fun to neutralize the pH of your solutions. You can use baking soda or soda ash to neutralize acid; when you dump baking soda or soda ash into acid, a huge amount of carbon dioxide gas bubbles up. The same thing happens when you add vinegar to soda ash. Watch carefully to be sure your container does not overflow; placing the container in the sink before neutralizing works well. I'd estimate that your starting container should be at least four times the volume of the contents; either that, or add the neutralizing chemical very, very slowly, so that the first bubbles that are given off have time to disperse before you cause more to be formed. In industrial situations, the carbon dioxide can be produced in such quantities as to create a danger of asphyxiation, as carbon dioxide is heavier than oxygen and will fill up a room from the floor on up, but this is unlikely in small-scale neutralizations. Open a window and be sure that your room ventilation is adequate. 

To neutralize a high-pH solution, you will have to add an acid. Vinegar is the usual choice, though in the lab one might use a few drops of hydrochloric acid. To figure out when you've succeeded in neutralizing your chemicals, the pH strips are great, aiming for somewhere closer to a neutral pH of 7, except of course for the dyebaths whose colors make it impossible for the colors of the indicator strips to be read. For colored dyebaths, I'd recommend basing the amount of vinegar to use to neutralize soda ash, or vice versa, on the amount you used to begin with. With weak bases such as soda ash, or weak acids such as acetic acid, neither of which will ionize completely in solution, you cannot calculate the amount to use to neutralize them simply by looking at the starting pH. 

A good rule of thumb would be to use roughly half as many molecules of soda ash as you do of vinegar, because one molecule of soda ash, Na2CO3, reacts with two molecules of vinegar molecules. A mole is a standard number of molecules, used by chemists to simplify calculation. One mole of anhydrous soda ash weighs 106 grams (equals 41 milliliters, or about three tablespoonfuls), while one mole of the decahydrate that is washing soda weighs 286 grams (equals 196 milliliters, or about 3/4 cup). These amounts would neutralize two moles of acetic acid, which would weigh 120 grams; that's the amount of acetic acid in 2400 milliliters of ordinary kitchen vinegar (which is 5% acetic acid, by weight), which is very close to two and a half quarts. So, two and a half quarts of vinegar should neutralize about three tablespoons of soda ash; one cup of soda ash should be neutralized by almost 14 quarts of distilled white vinegar. Not that it's necessary to bring the pH all the way to a perfectly neutral 7.0 pH; 8 or 9 should be okay, too, since it will be diluted by other water that's put down the drain.

One molecule of baking soda, which is sodium bicarbonate or NaHCO3, reacts with each molecule of acetic acid, so it requires half as much vinegar as soda ash does. One mole of sodium bicarbonate weighs 84 grams and occupies about 39 ml of volume, which is about two and a half tablespoons; it will neutralize 1200 milliliters of vinegar. It will also, by the way, produce one mole of water, which is a little over one tablespoon, and one mole of carbon dioxide gas, which occupies about 22 liters of volume. One cup of baking soda should neutralize about 7 quarts of vinegar, producing over a hundred quarts of CO2 gas.

If you're using lime, one mole of calcium oxide, CaO, weighs 56 grams (which works out to 16 milliliters, or about one tablespoon), and, like soda ash, reacts with two molecules of acetic acid, so one tablespoon of lime reacts with 2400 milliliters of vinegar. One cup of lime would require about 35 quarts of vinegar to neutralize it. The reaction would quickly produce enough carbon dioxide bubbles to potentially cause a dangerous spill or splash hazard.

Lye, also known as sodium hydroxide, or NaOH, is a much stronger base, which means that it all ionizes in water at once, leading to a higher pH than a similar number of moles of sodium carbonate in water. (A lot of the sodium carbonate in a solution remains unionized, which is why the pH doesn't go up above 12 no matter how much soda ash you put in, but the pH can go above 14 if you add enough lye.) One mole of sodium hydroxide in the form of lye weighs 40 grams, which is about 19 milliliters of the lye pellets that contain 99% sodium hydroxide, a little more than one tablespoon. Since one mole of sodium hydroxide reacts with one mole of acetic acid, this 19 ml of lye should react with 1200 ml or five cups of vinegar; one cup of lye pellets would be neutralized by about 15 quarts of vinegar. This is based on the assumption that the lye was already diluted in a large quantity of water when the dyebath was prepared. The reaction of strong lye solutions with vinegar could be extremely hazardous to be around, since the lye could be splashed about by the formation of bubbles of carbon dioxide. 

Note that lye spills on the skin should not be neutralized in the same way as lye solutions that are being prepared to pour down the drain. In the event of an accidental skin exposure to strong solutions of lye, it is important to immediately wash the lye from the skin by flushing continuously with large amounts of water. There is a popular misconception that neutralizing lye with vinegar will instantly heal a burn, which is impossible; in fact, it merely corrects the pH that can cause further burns. Nobody should use lye without a thorough understanding of emergency procedures; read an MSDS carefully before starting to use it. If it is used carelessly by those who do not understand how to use it safely, lye can cause permanent blindness or severe burns that can lead to permanent disability or death. It can be quite safe for use by those who understand proper laboratory safety procedures, however.

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Posted: Wednesday - February 17, 2010 at 01:15 PM          

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