Answer :
E°cell for the process is0.78 V.
The standard cell potential is the obtained value from the subtraction of the known cell of the potential of reduction species to the known cell potential of the oxidation species. It can be shown as [tex]$E_{\text {cell }}^{\circ}=\left(E_{\text {oxi }}^{\circ}\right)-\left(E^{\circ}{ }_{\text {red }}\right)$[/tex].
The reaction is given below.
[tex]$\mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{Fe}(\mathrm{s}) \rightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Fe}^{2+}(\mathrm{aq})$[/tex]
The oxidizing agent is that which shows the addition of electrons in the ion, and the oxidizing agent in this above reaction is [tex]$\mathrm{Cu}^{2+}$[/tex]. The reducing agent is that which presented the loss of an electron from the species, and the reducing agent in this reaction is Fe.
Therefore, the reducing agent is [tex]$\mathrm{Cu}^{2+}$[/tex] and the oxidizing agent is [tex]$\mathrm{Fe}$[/tex].
The cell potential of [tex]\mathrm{Cu}^{2+}$ is $0.34 \mathrm{~V}$[/tex].
The cell potential of [tex]$\mathrm{Fe}$[/tex] is [tex]$-0.44 \mathrm{~V}$[/tex].
The standard cell potential can be calculated is given below.
[tex]$\left(E_{\text {cell }}^{\circ}\right)=$[/tex] Cell potential of copper [tex]$\left(E^{\circ}\left(\mathrm{Cu}^{2+} / \mathrm{Cu}\right)\right)-$[/tex] Cell potential of iron [tex]$\left(E^{\circ}\left(\mathrm{Fe} / \mathrm{Fe}^{2+}\right)\right)$[/tex]
Standard cell potential
Substitute the respective values in the above equation.
[tex]$\begin{aligned}E_{\mathrm{cell}}^{\circ} &=0.34 \mathrm{~V}-(-0.44) \mathrm{V} \\&=0.78 \mathrm{~V}\end{aligned}$[/tex]
Therefore, the value of [tex]$E_{\text {cell }}^{\circ}$[/tex]for the process is [tex]$0.78 \mathrm{~V}$[/tex].
The potential difference between the two electrodes of an electrochemical cell, which develops when electrons are sent via the external circuit of a cell that has not reached equilibrium, is known as the cell potential.
An oxidizing agent, often known as an oxidizer or an oxidant, is a type of chemical that has the tendency to oxidize other substances, increasing their oxidation state by causing them to lose electrons. Halogens (such as chlorine and fluorine), oxygen, and hydrogen peroxide are typical examples of oxidizing agents (H2O2).
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