Bleach Oxidation Reactions

  • Thread starter Cesium
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  • #1
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I'm having trouble determining the products of these reactions with bleach. I know for sure that all of these reactions do occur because it's part of a lab in which we find the stoichimetric coefficients for the reactants using calorimetry. Since they don't want us to be able to balance the equation, we are not given the products. Completing the experiment will not tell us what the products are (although it does determine the coefficients which may help figure it out). I've picked these three reactions just because they were the ones given, but I'm open to a larger discussion about the oxidation power of bleach.

I've been using a redox potential list as my guide:
ClO- + H2O + 2e- <-> Cl- + 2OH- 0.81

a) NaClO + KI

I'm pretty sure here that iodine would be oxidized to its elemental form. My guess is...

NaClO + 2KI + H2O --> I2 + NaCl + 2KOH

b) NaClO + Na2S2O3

I've looked online for this one and I only find this reaction which occurs in basic solution. I'm not sure if the thiosulfate ion could be oxidized to the sulfate ion in a neutral solution. Perhaps to S2O6-2?

4ClO- + S2O3-2 + 2OH- --> 2SO4-2 + 4Cl- + H2O

c) NaClO + Na2SO4

This one I have no clue. I would think that no reaction would occur. Oxidation to the peroxydisulfate ion seems difficult:

S2O82- + 2e--> 2SO42- 2.010
 

Answers and Replies

  • #2
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I talked to my teacher about these after we completed the lab and she said that she honestly did not know what the products were. She made the lab up several years ago and had forgotten (I'm not quite sure I believe that). When mixing the bleach with the unknown reagent that I had, I got a gas that vaguely smelled of eggs. In any case, it didn't smell great and I think it was SO2. The experimentally derivied stoichiometric experiments were 4 for NaClO and 1 for the unknown. I am thinking I had sodium thiosulfate, but I don't get any equation that has a 4:1 ratio.

2NaClO + Na2S2O3 + H2O --> 2NaOH + 2NaCl + 2 SO2
 
  • #3
ShawnD
Science Advisor
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Cesium said:
I've been using a redox potential list as my guide:
ClO- + H2O + 2e- <-> Cl- + 2OH- 0.81

a) NaClO + KI

I'm pretty sure here that iodine would be oxidized to its elemental form. My guess is...

NaClO + 2KI + H2O --> I2 + NaCl + 2KOH
The reduction table I have supports this reaction, but try not to include spectator ions. For example, Na+ and K+ do absolutely nothing, so it's a bit confusing if you include them.
ClO- + H2O + 2e- --> Cl- + 2OH- cathode half
2I- --> I2 + 2e- anode half
ClO- + H2O + 2I- --> Cl- + 2OH- + I2 net cell reaction

b) NaClO + Na2S2O3

I've looked online for this one and I only find this reaction which occurs in basic solution. I'm not sure if the thiosulfate ion could be oxidized to the sulfate ion in a neutral solution. Perhaps to S2O6-2?

4ClO- + S2O3-2 + 2OH- --> 2SO4-2 + 4Cl- + H2O
Thiosulfate is a reducing agent which turns into tetrathionate
2S2O32- --> S4O62- + 2e2-
You are on the right track though.

c) NaClO + Na2SO4

This one I have no clue. I would think that no reaction would occur. Oxidation to the peroxydisulfate ion seems difficult:

S2O82- + 2e--> 2SO42- 2.010
I believe you are correct in saying no reaction will occur. Hypochlorite and sulfate are both oxidizers. That peroxydisulfate one would have a voltage of -2.010 if it were the anode reaction, but hypochlorite only has +0.84 on the cathode side; this means the reaction would not be spontaneous (won't happen).
 
  • #4
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Thanks for the response ShawnD!

ShawnD said:
Thiosulfate is a reducing agent which turns into tetrathionate
2S2O32- --> S4O62- + 2e2-
Doh! I did not see this in my table. The standard redox potential for this half reaction is -0.08, indeed sodium thiosulfate is a good reducing agent, used in photography for this very reason I think.

So the complete reaction (with spectator ions, I know) is:

NaClO + H2O + 2Na2S2O3 --> NaCl + 2NaOH + Na2S4O6

ShawnD said:
I believe you are correct in saying no reaction will occur. Hypochlorite and sulfate are both oxidizers. That peroxydisulfate one would have a voltage of -2.010 if it were the anode reaction, but hypochlorite only has +0.84 on the cathode side; this means the reaction would not be spontaneous (won't happen).

This is exactly what I told my teacher, but she said "No, no! That reaction must occur or else we would not have a lab using it!".

In general, is it safe to use the redox table to determine if a reaction occurs? I realize that they give standard redox potentials occuring at 298 K, 1 atm, and with 1 M concentrated for both solutes. If I want to deviate from these values, then I should use the Nernst equation, right? (I am thinking that perhaps a reaction does occur between hypochlorite and sulfate ions if the conditions are correct.)
 
  • #5
ShawnD
Science Advisor
668
1
Yes you use the nernst equation, but the voltage of the cell generally will not switch from negative to positive until the concentration of one of your theoretical products is less than the ksp value.
 

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