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Author: Admin | 2025-04-28
The standard electrode potential for that half-reaction. In this example, the standard reduction potential for \(\ce{Zn^{2+}(aq) + 2e^{−} → Zn(s)}\) is −0.76 V, which means that the standard electrode potential for the reaction that occurs at the anode, the oxidation of \(\ce{Zn}\) to \(\ce{Zn^{2+}}\), often called the \(\ce{Zn/Zn^{2+}}\) redox couple, or the \(\ce{Zn/Zn^{2+}}\) couple, is −(−0.76 V) = 0.76 V.We must therefore subtract \(E^o_{anode}\) from \(E^o_{cathode}\) to obtain \(E°_{cell} = 0.76\, V\). This is the origin of the negative sign in Equation \ref{20.4.2}.Because electrical potential is the energy needed to move a charged particle in an electric field, standard electrode potentials for half-reactions are intensive properties and do not depend on the amount of substance involved. Consequently, E° values are independent of the stoichiometric coefficients for the half-reaction, and, most important, the coefficients used to produce a balanced overall reaction do not affect the value of the cell potential.E° values do NOT depend on the stoichiometric coefficients for a half-reaction, because it is an intensive property.Standard Electrode PotentialsTo measure the potential of the Cu/Cu2+ couple, we can construct a galvanic cell analogous to the one shown in Figure \(\PageIndex{3}\) but containing a Cu/Cu2+ couple in the sample compartment instead of Zn/Zn2+. When we close the circuit this time, the measured potential for the cell is negative (−0.34 V) rather than positive. The negative value of \(E°_{cell}\) indicates that the direction of spontaneous electron flow is the opposite of that for the Zn/Zn2+ couple. Hence the reactions that occur spontaneously, indicated by a positive \(E°_{cell}\), are the reduction of Cu2+ to Cu at the copper electrode. The copper electrode gains mass as the reaction proceeds, and H2 is oxidized to H+ at the platinum electrode. In this cell, the copper strip is the cathode, and the hydrogen electrode is the anode. The cell diagram therefore is written with the SHE on the left and the Cu2+/Cu couple on the right:\[\ce{Pt(s) | H2(g,\, 1\, atm) | H^{+}(aq,\, 1\, M)|| Cu^{2+}(aq,\, 1\, M) |Cu(s)} \nonumber\]The half-cell reactions and potentials of the spontaneous reaction are as follows: Cathode: \[\ce{Cu^{2+}(aq) + 2e^{−} -> Cu(g)} \quad\quad E°_{cathode} = 0.34\; V \nonumber\] Anode: \[\ce{H2(g) -> 2H^{+}(aq) + 2e^{−}} \quad\quad E°_{anode} = 0\; V \nonumber\] Overall: \[\ce{H2(g) + Cu^{2+}(aq) -> 2H^{+}(aq) + Cu(s)} \nonumber\]Then via Equation \ref{20.4.2} we can calculate the standard cell potential\[\begin{align*} E°_{cell} &= E°_{cathode}− E°_{anode} \\[4pt] &= 0.34\; V \end{align*}\]Thus the standard electrode potential for the Cu2+/Cu couple is 0.34 V.Calculating Standard Cell PotentialsThe standard cell potential for a redox reaction (E°cell) is a measure of the tendency of reactants in their standard states to form products in their standard states; consequently, it is a measure of the driving force for the reaction, which earlier
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