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# 物理代写|电化学代写Electrochemical代考|CHEE4302 Precedence of Electrode Reactions

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## 物理代写|电化学代写Electrochemical代考|Precedence of Electrode Reactions

When the potential of an electrode is moved from its open-circuit value toward more negative potentials, the substance that will be reduced first (assuming all possible electrode reactions have fast kinetics) is the oxidant in the couple with the least negative (or most positive) $E^0$. For example, consider Figure 1.1.8a, which relates to a platinum electrode immersed in an aqueous solution containing $0.01 \mathrm{M}$ each of $\mathrm{Fe}^{3+}, \mathrm{Sn}^{4+}$, and $\mathrm{Ni}^{2+}$ in $1 \mathrm{M} \mathrm{HCl}$. The first substance reduced will be $\mathrm{Fe}^{3+}$, since the $E^0$ of this couple is most positive.

Likewise, when the potential of the electrode is moved from its open-circuit value toward more positive potentials, the substance that will be oxidized first is the reductant in the couple of least positive (or most negative) $E^0$. Thus, for a gold electrode in an aqueous solution containing $0.01 \mathrm{M}$ each of $\mathrm{Sn}^{2+}$ and $\mathrm{Fe}^{2+}$ in $1 \mathrm{M} \mathrm{HI}$ (Figure 1.1.8b), the $\mathrm{Sn}^{2+}$ will be first oxidized, since the $E^0$ of this couple is least positive.

One must remember, however, that these predictions are based on thermodynamic considerations (i.e., reaction energetics). Slow kinetics might prevent an electrode reaction from occurring at a significant rate in a potential region where the standard potential suggests that the reaction is possible. For a mercury electrode immersed in a solution of $0.01 \mathrm{M} \mathrm{each} \mathrm{of}^{3+}$ and $\mathrm{Zn}^{2+}$ in $1 \mathrm{M} \mathrm{HCl}$ (Figure 1.1.8c), the first reduction process predicted is the evolution of $\mathrm{H}_2$ from $\mathrm{H}^{+}$. As discussed earlier, this reaction is very slow on mercury, so the first process actually observed is the reduction of $\mathrm{Cr}^{3+}$.
None of the cases in Figure 1.1.8 involves a redox couple for which both redox forms are present; consequently, there is no equilibrium defining the open-circuit potential of the electrode. In every electrochemical system, the open-circuit potential lies between the easiest oxidation and the easiest reduction. If, apart from species defining the background limits, the solution contains only oxidized forms eligible for reduction (as in Figure 1.1.8a,c), the open-circuit potential must lie between the positive background limit and the standard potential for the first electroreduction. If the solution contains only reduced forms eligible for oxidation (as in Figure 1.1.8b), the open-circuit potential must lie between the negative background limit and the standard potential for the first electrooxidation.

## 物理代写|电化学代写Electrochemical代考|Both Redox Forms of a Couple as Solutes

In some systems, both redox forms of a couple are present as solutes, and the behavior differs notably from the cases we have been discussing. Suppose, for example, that we make three changes in the cell of Figure 1.1.4:

• Employing a silver electrode covered with $\mathrm{AgCl}$, rather than one covered with $\mathrm{AgBr}$.
• Adopting a vessel having two compartments, one for each electrode, separated internally by a frit.
• Placing $1 \mathrm{M} \mathrm{HCl}$ in each compartment, but also including $2 \mathrm{mM} \mathrm{Fe}(\mathrm{II})$ and $4 \mathrm{mM} \mathrm{Fe}(\mathrm{III})$ in the solution at the Pt electrode.
Thus, the cell becomes
$$\mathrm{Pt} / \mathrm{H}^{+}(1 \mathrm{M}), \mathrm{Cl}^{-}(1 \mathrm{M}), \mathrm{Fe}(\mathrm{III})(4 \mathrm{mM}), \mathrm{Fe}(\mathrm{II})(2 \mathrm{mM}) / / \mathrm{H}^{+}(1 \mathrm{M}), \mathrm{Cl}^{-}(1 \mathrm{M}) / \mathrm{AgCl} / \mathrm{Ag}$$
The compartments are employed to keep Fe(III) species away from the silver electrode, where they would otherwise react by oxidizing the metal. The frit prevents bulk mixing of the solutions in the two compartments, but allows them to remain in ionic electrical contact.

In this system, the silver electrode becomes an $\mathrm{Ag} / \mathrm{AgCl}$ reference, ${ }^{17}$ and the cell potential is the potential of the Pt working electrode against that reference, which can be re-expressed on the NHE scale, essentially as was done for Figures $1.1 .6$ and 1.1.7.

The iron species exist in $1 \mathrm{M} \mathrm{HCl}$ as chloro complexes, so we simply write the relevant couple as 18
$$\mathrm{Fe}(\mathrm{III})+e \rightleftharpoons \mathrm{Fe}(\mathrm{II}) \quad E^{0^{\prime}}=+0.70 \mathrm{~V} v \text { v. } \operatorname{NHE}(1 \mathrm{M} \mathrm{HCl})$$

Because the Pt electrode is in contact with both redox forms, the working electrode is poised by the $\mathrm{Fe}(\mathrm{III}) / \mathrm{Fe}$ (II) and shows a true equilibrium potential near $E^{0^{\prime}}$ for (1.1.24). The current-potential curve is depicted in Figure 1.1.9.

## 物理代写|电化学代写电化学代考|一对作为溶质的两种氧化还原形式

• 使用覆盖$\mathrm{AgCl}$的银电极，而不是覆盖$\mathrm{AgBr}$的银电极
• 采用有两个隔间的容器，每个电极一个，内部用熔块隔开。
• 在每个隔间中放置$1 \mathrm{M} \mathrm{HCl}$，但也包括$2 \mathrm{mM} \mathrm{Fe}(\mathrm{II})$和$4 \mathrm{mM} \mathrm{Fe}(\mathrm{III})$在Pt电极的溶液中。因此，细胞变成
$$\mathrm{Pt} / \mathrm{H}^{+}(1 \mathrm{M}), \mathrm{Cl}^{-}(1 \mathrm{M}), \mathrm{Fe}(\mathrm{III})(4 \mathrm{mM}), \mathrm{Fe}(\mathrm{II})(2 \mathrm{mM}) / / \mathrm{H}^{+}(1 \mathrm{M}), \mathrm{Cl}^{-}(1 \mathrm{M}) / \mathrm{AgCl} / \mathrm{Ag}$$
隔间被用来使Fe(III)种远离银电极，否则它们会通过氧化金属发生反应。熔块阻止了溶液在两个隔室中的大量混合，但允许它们保持离子电接触。

$$\mathrm{Fe}(\mathrm{III})+e \rightleftharpoons \mathrm{Fe}(\mathrm{II}) \quad E^{0^{\prime}}=+0.70 \mathrm{~V} v \text { v. } \operatorname{NHE}(1 \mathrm{M} \mathrm{HCl})$$

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