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Topic: Redox Titrations  (Read 2393 times)

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Offline Kate01

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Redox Titrations
« on: August 14, 2014, 01:12:09 AM »
What is it about the structure/properties of sodium hypochlorite in bleach that makes it unstable and hence attributes to its decomposing state?

Offline rwiew

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Re: Redox Titrations
« Reply #1 on: August 14, 2014, 09:49:55 AM »
What oxidation state is chlorine in the hypochlorite? Could it transform into other oxidation states spontaneously?

Offline Kate01

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Re: Redox Titrations
« Reply #2 on: August 14, 2014, 08:33:23 PM »
Okay, so the oxidation state of chlorine in the hypochlorite is +1 which is unusually high oxidation state for chlorine as it is normally -1. Wouldn't that mean that chlorine is a strong reducing agent, meaning that it easily loses electrons? So this would account for the change in oxidation state, but how does this transformation into other oxidation states occur spontaneously and how can this occur without a catalyst as bleach decomposes readily without a catalyst?

 

Offline rwiew

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Re: Redox Titrations
« Reply #3 on: August 14, 2014, 11:58:11 PM »
Ok, so I agree with chlorine "normally" having the -1 oxidation state, by which I mean it is the most stable one. However, with +1 being unusually stable I won't agree with - the highest oxidation state of chlorine is +7 (perchloric acid and its salts). Now, there's differences in the relative energies and hence stabilities of oxidation states and hence it's possible for an oxidation state to transform to two others if this is energetically favourably - are you familiar with disproportionation? What could Cl(I) disproportionate into? This however only happens when you heat bleach.

At room temperature, when standing in aqueous solution a different reaction can happen - even though I didn't agree with +1 being "unusually" high, we can say it's high enough for it to be wanting to go into -1 - ClO- can also transform into Cl- and O2 (can you write the reaction equation for this reaction?), avoiding creation of a higher oxidation state as in the disproportionation.

Now back to what you said - all this means Cl(I) is a strong oxidizer - it want to get more electrons so it can go down to -1 (it needs 2 electrons / atom to do it). Spontaneity comes from the Gibbs free energy difference between the form it is in and the forms it could be in after the decomposition - are you familiar with Gibbs free energy? Also a catalyst is not necessarily needed for such reactions, if the reaction is fast in itself it does not need help (compare to decomposition of hydrogen peroxides, which needs catalysts for decomposition, as its decomposition reactions are quite slow). Does this make sense?

Offline Kate01

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Re: Redox Titrations
« Reply #4 on: August 16, 2014, 02:31:05 AM »
Oh okay, I understand. So by using Gibbs free energy rule I can explain the spontaneous decomposition of hypochlorite in bleach. Thank you very much! I have another question relating to the decomposition. The decomposition of hypochlorite forms both chloride and chlorate and oxygen and chloride. Through this decomposition, the concentration of active chlorine in bleach is decreased. What happens to these products of the decomposition and what is it about these products that make it inactive?

Offline rwiew

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Re: Redox Titrations
« Reply #5 on: August 16, 2014, 10:34:38 PM »
Right, so what happens to the products of decomposition - chlorate either just sits there in solution, or from looking at standard potentials I can tell you it could further disproportionate to chloride and perchlorate, I might not be happening if that reaction was slow though, and finally it can act as another oxidant in the bleach (it's a better oxidant than hypochlorite); chloride just sits there in solution, oxygen bubbles away. I read somewhere that the oxygen is produced in singlet form though (I don't have much time to confirm this though now, have a look around some reliable sources if you want) - that is very reactive and could be an additional oxidant in the system.

Now what makes them inactive - well, not being the active component does. But it is possible that chlorate and oxygen actually help in the oxidation, chloride is definitely not doing anything. Hopefully you'll be familiar with how bleach works - it oxidizes conjugated systems in dye molecules, killing the colour and hence whitening- i.e. you need the oxidizing power.

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