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In this textbook example question, how come the chlorine reaction, which has the greatest reducing power will not react? If it has the greatest E of 1.36 V that means it has most reducing stength and can "pull" electrons from both the Br and I reactions. In my mind I believed it would do so, and since the I-reaction has an even lower E than the Br-reaction the extent of oxidation would be greater for I. But in the answer they say the direct opposite, that Cl does not react and Br will only oxidize I.

Why does Cl which is the strongest oxidizing agent not do anything in this case?

  • $\begingroup$ You are mixing up $\ce{Cl2}$ and $\ce{Cl-}$. $\ce{Cl2}$ is more reducing than $\ce{Br2}$, but you don't have any $\ce{Cl2}$, you only have $\ce{Cl-}$ from the dissolved $\ce{NaCl}$ $\endgroup$ – Tyberius Jan 30 at 18:51
  • $\begingroup$ Oxidation power decreases in range $\ce{Cl2} \gt \ce{Br2} \gt \ce{I2}$, reduction power decreases in range $\ce{I-} \gt \ce{Br-} \gt \ce{Cl-}$ $\endgroup$ – Poutnik Jan 31 at 9:41

Because in the original question there is no elemental chlorine present. It says that you start from a solution of chloride and iodide, so both ions that can be oxidized to the corresponding halogen. And the rest of the answer is already given in your textbook example. Chlorine is a much stronger oxidizing agent. This means that it oxidizes others very well. But what you've got is bromine, and bromine is just between chlorine and iodine, so it can oxidize iodide to iodine but not chloride to chlorine. I think you got confused with the direction and the redox-terms here a bit.

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  • $\begingroup$ I think I understand what they are saying now thanks to your help. The first thing I needed to understand was that in the solution we have iodide and chloride. Next was that since chlorine has the strongest reducing power it "holds strongest" to its electrons, more so than bromine and iodine which have lower E values. So bromine can't pull away those electrons from chloride. However iodine has a lower E than bromine so bromine can pull away the electrons from iodide, causing formation of iodine. Am i correct? $\endgroup$ – Johan Jan 30 at 12:58
  • $\begingroup$ Basically you are right but elemental chlorine is a strong oxidizing agent, not a reducing agent in your case. It oxidizes others for example bromide to bromine while itself it's being reduced in this process. In cases of halogens I would argument with their oxidation strenght. You could say that bromide is a stronger reducing agent than elemental chlorine, hence it reduces chlorine to chloride. $\endgroup$ – Justanotherchemist Jan 30 at 13:19

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