JnPS
Hazard to Self
Posts: 90
Registered: 29-7-2016
Location: PA, USA
Member Is Offline
Mood: Umpolung
|
|
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.
|
|
Lithium
Hazard to Others
Posts: 103
Registered: 25-2-2012
Location: Australia
Member Is Offline
Mood: Thinking!
|
|
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.
Li
|
|
Metacelsus
International Hazard
Posts: 2539
Registered: 26-12-2012
Location: Boston, MA
Member Is Offline
Mood: Double, double, toil and trouble
|
|
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.
|
|
j_sum1
Administrator
Posts: 6292
Registered: 4-10-2014
Location: Unmoved
Member Is Offline
Mood: Organised
|
|
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.
|
|
CHEMxpl
Harmless
Posts: 21
Registered: 14-11-2016
Member Is Offline
Mood: No Mood
|
|
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
|
|
JnPS
Hazard to Self
Posts: 90
Registered: 29-7-2016
Location: PA, USA
Member Is Offline
Mood: Umpolung
|
|
It's a pool chlorinator not bleach
Thanks for the link!
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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)
|
|