Electrochemistry - Redox & Electroplating 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?
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 $$\ce{Fe -&gt; Fe^{2+} + 2 e-}$$ 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 : $$\ce{Cl2 + 2 e- -&gt; 2 Cl-}$$ The global effect of these reactions is equivalent to the following equation that describes the general change of composition in the cell : $$\ce{Fe + Cl2 -&gt; Fe^{2+} + 2 Cl-}$$