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I have to decide which one has a higher oxidation strength $\ce{ClO-}$ or $\ce{Cl2}$

$\ce{ClO- + H2O + 2e- -> Cl- + 2OH-}$

$\ce{Cl2 + 2e- -> 2Cl-}$

Will this reaction occur?

$\ce{ClO- + H2O + 2Cl- -> Cl- + 2OH- + Cl2}$

In my opinion not, because otherwise the $\ce{ClO-}$ is reduced to $\ce{Cl-}$. If the reaction works, the $\ce{Cl-}$ could react as reduction agent again. Therefore if I would add $\ce{Br2}$ to $\ce{ClO-}$:

$\ce{2 ClO- + 2 H2O + 4Br- -> \mathbf{2 Cl- } + 4OH- + 2Br2 }$

$\ce{ClO- + H2O + 2Br- + \mathbf{2 Cl- } -> \mathbf{Cl- } + 2OH- + Br2 + \mathbf{Cl2}}$

So therefore chlorine has a higher oxidation strength than $\ce{ClO-}$?

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  • $\begingroup$ Adam, I have used mathjax in combination with our chemical equations typesetting on your post to make it a bit more readable. Please check the syntax so you can do it yourself next time. $\endgroup$ – Michiel Nov 6 '14 at 8:27
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    $\begingroup$ Chlorine dissolves in alkaline solutions forming mixture of chloride and hypochlorite. However, in acidic solutions the reaction is reversed. For this reason acidification of bleach solutions is not advised. The reason is that hypochlorite half-reaction includes hydroxide/proton, so its potential is susceptible to change depending on solution pH $\endgroup$ – permeakra Jun 4 '16 at 13:07
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Look right here ... .

http://hyperphysics.phy-astr.gsu.edu/hbase/tables/electpot.html

Hypochlorite ion (+0.90 V) has a lower reduction potential in basic solution than elemental chlorine (+1.36 V), so the hypochlorite ion is less strongly oxidizing and more stable in basic solution (which is what you write the reaction for).

Household "chlorine" bleach, which really uses hypochlorite, is made alkaline as one way to stabilize the hypochlorite ion and prevent evolution of chlorine gas.

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