In article , Elaine T
wrote:

Dr Engelbert Buxbaum wrote:
Rocco Moretti wrote:


Dr Engelbert Buxbaum wrote:


They usually contain a small amount of sodium thiosulphate, which
destroys chlorine catalytically. Thus a very small amount of dechlor can
destroy a huge amount of chlorine, but the less you use the longer it
takes.

How does that work then? If it works catalytically (isn't used up),

The reactions I've seen on the web show the reaction as
Cl2 + H2O = HOCl + H+ + Cl-


Yes, it must start there. Presumably the hypochloric acid gets
destroyed, something like 2 HClO - 2 HCl + O2, but I have no idea about
the exact sequence of events, may be some kine of peroxy-acid
intermediate with the thiosulfate?

The other thing I have never quite understood with these chlorine
destroyers is how they stabilise the small amount of thiosulfate in
there. In the lab such a dilute solution would be prepared fresh daily.

Boy is that an interesting question. Many dechlors are in opaque or
dark bottles, which would help somewhat. I suspect they are also
considerably stronger than necessary, to compensate for thiosulfate
degradation. I use AmQuel anyway - wonder if that degrades similarly.

I would guess that the solutions are stronger than those used in
analysis. The one I use requires only 3 drops per gallon!! Vogel
notes that a 0.01N solution is stable in the dark, provided the solution
is prepared with 18Mohm water; dissolved CO2 seems to be the main
culprit with regards to instability. Since the one I use has a blue
colour, it's possible the manufacturer adds something like methylene
blue, which would function as an antioxidant for the thiosulphate.

I looked in a couple of books here, and none gave much clue as to the
chemistry, other than the disproportionation of the hypochloride would
be (Cotton & Wilkinson):

3HOCl == 2HCl + HClO3

which is noted to be highly favourable, but slow at or below room
temperature. Other alternatives are described as "unfavourable"!