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I've stumbled across an electrochemistry problem that I need help on:

$$\ce{Fe~|~Fe^{+2}~||~Cl^{-},Cl2~|~Pt}$$

I understood that iron is at the anode and is being oxidized, however I do not know what was going on with chlorine, it looks like it is being oxidized while it is supposed to be reduced.

Maybe this is a non-spontaneous reaction? Can this be possible, can chlorine be used as an oxidation reaction? And if so how can we calculate the voltage?

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  • $\begingroup$ Finely divided iron powder spontaneously combusts in chlorine gas, so, assuming aqueous solutions, iron is oxidized at the anode and dissolved chlorine gas is reduced to chloride ions at the cathode. The cell notation is simply incomplete. $\endgroup$
    – Ed V
    May 17 at 16:00
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Here the iron electrode is made by dipping an iron plate or wire in a $\ce{FeCl2}$ or $\ce{FeSO4}$ solution. It is an anode, where $\ce{Fe}$ is oxidized according to $$\ce{Fe -> Fe^{2+} + 2 e-}$$ The chlorine electrode is a cathode made by dipping a platinum electrode in a solution where some chlorine gas $\ce{Cl2}$ is arriving from under the platinum plate, so that a few bubbles remain adsorbed on its surface. This $\ce{Cl2}$ gets reduced to $\ce{Cl-}$ by absorbing the electrons coming from the iron anode according to the equation : $$\ce{Cl2 + 2 e- -> 2 Cl-}$$ The global effect of these reactions is equivalent to the following equation that describes the general change of composition in the cell : $$\ce{Fe + Cl2 -> Fe^{2+} + 2 Cl-}$$

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  • $\begingroup$ I am very surprised this one went 6 years without even a comment! Anyway, (+1) for closing the books on it. $\endgroup$
    – Ed V
    May 17 at 16:12

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