[organosurf]

Traditionally, Sodium bisulfite ( NaHSO₃ ) is used to destroy excess Hypochlorous acid ( HOCl ) in bleach solution :

HOCl  +  NaHSO₃    NaCl  +  H₂SO₄

Can Hydrogen peroxide ( H₂O₂ ) be used instead ?
If yes, it is a much “greener” reagent and leaves no by products in solution to be separated :

2HOCl  +   H₂O₂      2H₂O  +  Cl_2(g)  +  O_2(g)

I tried the above using household bleach as a source of HOCl at room temperature. The reaction was immediate but of short duration with much effervescence and a strong smell of chlorine. On cessation, excess H₂O₂ was added, but no further effervescence occurred, possibly suggesting that ALL the HOCl was destroyed. There was no smell of chlorine. On adding MnO₂ black powder ( catalyst ), caused the formation of a stream of bubbles, no smell of chlorine, indicating that the excess H₂O₂ was dissociating into H₂O and O₂ :

2H₂O₂     2H₂O  +  O_2(g)

The above is not mentioned in any related examples on the internet.

[discodermolide]

This is not green, you produce chlorine gas which is rather nasty.
I advise you to remain with the bisulfite method

[orgopete]

You should try it again, but you must modify it. The bisulfate reaction is with the sodium salt. Hypochlorite is a product of chlorine and water. I presume the equilibrium favors chlorine at low pH and appears to be what was observed. If this same reaction were to occur at high pH, the chlorine would react with water to give hypochlorite plus sodium chloride, and not chlorine.

The problem is whether the same reaction will occur at high(er) pH's?

[curiouscat]

I presume the equilibrium favors chlorine at low pH and appears to be what was observed. If this same reaction were to occur at high pH, the chlorine would react with water to give hypochlorite plus sodium chloride, and not chlorine.

If he replaces Sodium bisulfite  with H2O2 where will the Na in NaCl come from?

Also, how does one raise pH without facing the same or similar by-product separation problem?

[orgopete]

I presume the equilibrium favors chlorine at low pH and appears to be what was observed. If this same reaction were to occur at high pH, the chlorine would react with water to give hypochlorite plus sodium chloride, and not chlorine.

If he replaces Sodium bisulfite  with H2O2 where will the Na in NaCl come from?

Also, how does one raise pH without facing the same or similar by-product separation problem?

I was thinking NaOH could raise the pH and be a source of sodium for NaCl.

[curiouscat]

I presume the equilibrium favors chlorine at low pH and appears to be what was observed. If this same reaction were to occur at high pH, the chlorine would react with water to give hypochlorite plus sodium chloride, and not chlorine.

If he replaces Sodium bisulfite  with H2O2 where will the Na in NaCl come from?

Also, how does one raise pH without facing the same or similar by-product separation problem?

I was thinking NaOH could raise the pH and be a source of sodium for NaCl.

Gotcha. Thanks.

[discodermolide]

I thought the whole point was to destroy the hypochlorite, so why re-generate it by using higher pH? Furthermore peroxide is not stable at high pH, as we all know.
Am I missing something here?
So what's wrong with the bisulfite?

[organosurf]

Dear All,
Thanks for your input. The aim is to destroy the NaOCl or HOCl by using H2O2 instead of sodium bisulfite.
Even acidifying the hypochlorite destroys it.
I found the earlier reaction 2HOCl  +   H2[/size]O2[/size]    [/size]  2H2O  +  Cl2(g)  +  O2(g) on the net, which is in contradiction ( or Correct ? ) as that one given in wiki reference below. Dont know which to believe. So discodermolide was not wrong to say that H2O2 is not that "green" as it generates Cl2(g), but nevertheless I'am still inclined to using it as I later found out in wiki.
These myriad of co-reactions relating to hypochlorites have always got me muddled up, believe you me,  , as you will see from the list of reactions (and references cited ) that possibly occur simultaneously. Just to recap once again :

1) Hypochlorites are stable at alkaline ( high ) pH, NaOH is added in commercial bleach sol for this purpose.
    Conversely, H2O2 decomposes at alkaline ( high ) pH.
2) Hypochlorites decompose at acidic ( low ) pH and conversely, H2O2 is stable at acidic ( low ) pH
3) The Na comes from the sodium hypochlorite, NaOCl and the NaOH, both present in bleach sol.

http://en.wikipedia.org/wiki/Sodium_hypochlorite

It reacts with other acids, such as acetic acid, to release hypochlorous acid:
NaClO + CH3COOH → HClO + CH3COONa
It decomposes when heated to form sodium chlorate and sodium chloride:
3 NaClO → NaClO3 + 2 NaCl
In reaction with hydrogen peroxide it gives off molecular oxygen:
NaClO + H2O2 → H2O + NaCl + O2↑
When dissolved in water it will slowly decompose, releasing chlorine, oxygen and sodium and hydroxide ions.
4 NaClO + 2 H2O → 4 Na+ + 4 OH- + 2 Cl2 + O2
Addition of chlorine to water gives both hydrochloric acid (HCl) and hypochlorous acid:[8]
Cl2 + H2O   HClO + HCl
When acids are added to aqueous salts of hypochlorous acid (such as sodium hypochlorite in commercial bleach solution), the resultant reaction is driven to the left, and chlorine gas is evolved. Thus, the formation of stable hypochlorite bleaches is facilitated by dissolving chlorine gas into basic water solutions, such as sodium hydroxide.
In aqueous solution, hypochlorous acid partially dissociates into the anion hypochlorite OCl−:
HClO  OCl− + H+
Salts of hypochlorous acid are called hypochlorites. One of the best-known hypochlorites is NaClO, the active ingredient in bleach.
HClO is a stronger oxidant than chlorine under standard conditions.
2 HClO(aq) + 2 H+ + 2 e−   Cl2(g) + 2 H2O E = +1.63 V
HClO reacts with HCl to form chlorine gas:
HClO + HCl → H2O + Cl2

http://en.wikipedia.org/wiki/File:HOCl_%2B_H2O2.png
HClO + H2O2 → HCl + O2 +  H2O

With reference to all the above, would the use of H2O2 to destroy HOCl / NaOCl be "green" ? 

[discodermolide]

No, no chemistry is green, there is no such thing.

[curiouscat]

I thought the whole point was to destroy the hypochlorite, so why re-generate it by using higher pH? Furthermore peroxide is not stable at high pH, as we all know.
Am I missing something here?
So what's wrong with the bisulfite?

Agreed. I'm puzzled too.

[orgopete]

In reaction with hydrogen peroxide it gives off molecular oxygen:
NaClO + H2O2 → H2O + NaCl + O2↑

Addition of chlorine to water gives both hydrochloric acid (HCl) and hypochlorous acid:[8]
Cl2 + H2O HClO + HCl

HClO reacts with HCl to form chlorine gas:
HClO + HCl → H2O + Cl2

Regenerate? Check the reactions posted by organosurf. The chlorine or hypochlorite is due to pH and stoichiometry. It appears that at higher pH, NaCl, water, and oxygen are produced.

I do not wish to engage in any green chemistry debate, but I could understand that hydrogen peroxide could have some advantage depending on ones point of view. For the purpose of this forum, I'd consider that a useful contribution. Personally, I'd be concerned with gas evolution becoming a potential problem. It may not be, but if we were considering the utility of hydrogen peroxide, then it should be a factor as well.