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Author: Subject: Hypochlorite decomposition
JnPS
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[*] posted on 15-11-2016 at 18:39
Hypochlorite decomposition


I couldn't find the short questions thread for general questions so I'd figure I'd just post this in Beginnings.

I plan on performing a haloform reaction to make chloroform using acetone and sodium hypochlorite. The hypochlorite solution is labeled as 30% NaClO but I bought these gallon jugs back in July when they were on sale for $3 each. The solution is still colored which I believe indicates a usable concentration of hypochlorite. But I don't have any standards for running a titration to determine the actual concentration, I'm not confident in my scale accuracy for density measurements as I await a standard weight for calibration to arrive, and I know that keeping the hypochlorite in excess is crucial for this reaction to avoid the chloroform-acetone azeotrope.

Would I be safe in assuming that it reduced to half its concentration in calculating my reagent ratios? How fast would the hypochlorite decompose? They were stored away from light in the original plastic containers.
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[*] posted on 15-11-2016 at 19:40


If you have a graduated cylinder you can roughly determine the concentration of hypochlorite in the solution. By decomposing the hypochlorite into chloride and oxygen with a suitable catalyst, and collecting the formed gas in a graduated cylinder (with the aid of tubing and inverting the cylinder in a beaker or tub filled with water) you can find the hypochlorite concentration with a few calculations.





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[*] posted on 15-11-2016 at 21:17


If you want more precision, you could use iodometric titration. Add the bleach to an excess of iodide, and then titrate the iodine that forms using sodium thiosulfate solution of known concentration.



As below, so above.
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[*] posted on 15-11-2016 at 21:51


In any case it is preferable to have the hypochlorite in excess for the haloform reaction. Separating out excess acetone at the end is problematic.
If you bought it for that cheap then a little wastage won't kill you. If you find a rough lower estimate of the concentration and work your stoichiometry from that then you should be fine.

If I was to hazard a guess, unless you have had your bottles open an in a really warm place, your assessment of the amount of decomposition is a large overestimate. I doubt the concentration has halved in that time.




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[*] posted on 19-11-2016 at 00:06


there's no way a bleach would contain 30% NaCLO look carfully maybe it's 30° which meanes it's around 8_10% NaCLO
if you wanna know more about how to make chloroform go to yhis link: https://www.youtube.com/watch?v=j-PrAczOGb0

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[*] posted on 19-11-2016 at 06:13


It's a pool chlorinator not bleach :)

Thanks for the link!
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[*] posted on 20-11-2016 at 15:22


Add an excess of dilute H2O2 to say 10 ml of your aqueous NaOCl. Record the volume of O2 released. Every 22.4 liters of oxygen created equals close to 1 mole of NaOCl.

My experience with this reaction is that it is very fast. As such, one could, for example, place 10 ml of NaOCl in a test tube which is placed upright in a large vessel which contains of an excess of dilute H2O2. Seal with a cover that has a pin whole. Quickly invert the vessel allowing a mixing of the H2O2 and NaOCl producing a stream of O2 that can be captured by an inverted large clear bowl in a water bath. Mark off the level of water displacement observed.

Reaction:

NaOCl + H2O2 = NaCl + H2O + O2(g)
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