Chlorine removers?

[Scott]

Has anyone ever tested water after it has had a chlorine neutralizer added
to it, to see how well they work. The reason I ask is that I have been using
Kent chlorinex for a year or so. I changed to it after using the more
expensive tetra one, and it seems like its never going to run out. You only
have to add four drops to eight liters of water and its supposed to work
straight away, but I'm not shore how effective they are. And if there not,
couldn't doing regular water changes also be putting chlorine in my aquarium
rather than just taking out impurities?

[Bill Stock]

"Scott" > wrote in message
...
> Has anyone ever tested water after it has had a chlorine neutralizer added
> to it, to see how well they work. The reason I ask is that I have been
> using Kent chlorinex for a year or so. I changed to it after using the
> more expensive tetra one, and it seems like its never going to run out.
> You only have to add four drops to eight liters of water and its supposed
> to work straight away, but I'm not shore how effective they are. And if
> there not, couldn't doing regular water changes also be putting chlorine
> in my aquarium rather than just taking out impurities?

I imagine most dechlorinators are formulated to take out the maximum
expected chlorine content. I believe the US max allowed is something like 4
PPM at the treatment plant, which would equate to much less when it reaches
your house. Mine typically reads .5 PPM or less, so in theory I could get
by without dechlorinator for small water changes, say 10-15%. But I always
add dechlor in the event that my city changes to chloramines or I have an
unexpected blip in the Chlorine content.

I actually mix my own dechlor (Sodium Thiosulphate), so I can mix any
concentration I like. But I prefer not to get down to the drops per litre
strength, as it's too much of a PITA. I typically mix it up to take care of
3 PPM of chlorine at a Tablespoon per 10 gallons. It just makes it easier
not to have to carry an eyedropper around to do water changes. I have a
small pond and I just found it too expensive to keep buying the premixed
dechlor, not to mention the geek factor of making my own.

I don't think you need to worry about adding Chlorine to your tank,
especially for small water changes. It's likely Kent is just using a
stronger formulation than Tetra. Jungle makes fairly inexpensive Chlorine
test strips if you want to be certain.

[Dr Engelbert Buxbaum]

Scott wrote:

> Has anyone ever tested water after it has had a chlorine neutralizer added
> to it, to see how well they work. The reason I ask is that I have been using
> Kent chlorinex for a year or so. I changed to it after using the more
> expensive tetra one, and it seems like its never going to run out. You only
> have to add four drops to eight liters of water and its supposed to work
> straight away, but I'm not shore how effective they are.

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.

[Rocco Moretti]

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), I
would initially imagine that the net rection would be Cl2 => 2 Cl-
But that isn't quite right, as we need to balance charge, so it would
have to be Cl2 + 2e- => 2 Cl-, but that means we need some source of
electrons.

The reactions I've seen on the web show the reaction as
Cl2 + H2O => HOCl + H+ + Cl-
HOCl + 2 S2O3(2-) => Cl- + S4O6(2-) + OH-
------------------------------------------
Cl2 + 2 S2O3(2-) => 2 Cl- + S4O6(2-)

where the thiosulphate acts as the electron donor in the reaction and is
thus used up.

FWIW, I also found similar reactions with sulfite compounds acting as an
electron donor:

Sodium Sulfite: Na2SO3 + Cl2 + H2O => Na2SO4 + 2 HCl
Sodium Metabisulfite: Na2S2O5 + 2Cl2 + 3H2O => 2NaHSO4 + 4HCl

[Dr Engelbert Buxbaum]

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.

[Elaine T]

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.

--
Elaine T __
http://eethomp.com/fish.html <'__><
rec.aquaria.* FAQ http://faq.thekrib.com

[David C. Stone]

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"!