Electrode potentials standard reduction half-reaction

The electrochemical series [Ch ter 10, Table 10.13] or, in fact, a list of the standard electrode potentials for reduction half-reactions can be very useful to quickly calculate the standard value of the potential. An example of the electrochemical series just for a few electrochemical half-reactions taken from [Chapter 10, Table 10.13] is given in Table 4.1. 

If you select any two half-reactions from the chart of standard electrode potentials, the half-reaction higher on the list will proceed as a reduction, and the one lower on the list will proceed in the reverse direction, as an oxidation. Beware Some references give standard electrode potentials for oxidation half-reactions, so you have to switch higher and lower in the rule stated in the preceding sentence, though this is not common. 

The standard electrode potential for a half-reaction refers exclusively to a reduction reaction that is, it is a relative reduction potential. 

The emf of a cell can be calculated from the standard electrode potentials of the half-reactions. In order to find the emf, we have to look at the two halfreactions involved in the reaction. Then, set up the two half-reactions so that when they are added we will get the net reaction. Once we have set the equations properly and assigned the prpper potentials to those half-reactions, we can add the standard electrode potentials. A common mistake that students make is that they forget the fact that the standard electrode potentials are given in terms of reduction reactions. Redox reactions involve both oxidation and reduction. If one half-reaction is reduction, the other should be oxidation. So we must be careful about the signs of the half-reaction potentials, before we add the two half-reaction potentials to get the emf value. Do the next example. 

As we have seen, different concentrations of ions in a half-cell result in different half-cell potentials. We can use this idea to construct a concentration cell, in which both halfcells are composed of the same species but in different ion concentrations. Suppose we set up such a cell using the Cu +/Cu half-cell that we introduced in Section 21-9. We put copper electrodes into two aqueous solutions, one that is 0.10 M CUSO4 and another that is 1.00 M CUSO4. To complete the cell construction, we connect the two electrodes with a wire and join the two solutions with a salt bridge as usual (Figure 21-15). Now the relevant standard reduction half-reaction in either half-cell is... 

From a table of electrode potentials, write the two reduction half-reactions and standard electrode potentials for the cell. The cell notation assumes that the anode (the oxidation half-cell) is on the left. Change the direction of this half-reaction and the sign of its electrode potential. (Assuming that the cell notation was written correctly, you change the direction of the half-reaction corresponding to the smaller, or more negative, electrode potential.) Multiply the half-reactions (but not the electrode potentials) by factors so that when the half-reactions are added, the electrons cancel. The sum of the half-reactions is the cell reaction. Add the electrode potentials to get the cell emf. 

Standard potentials are also called standard electrode potentials. Because they are always written for reduction half-reactions, they are also sometimes called standard reduction potentials. 

STRATEGY Find the standard potentials of the two reduction half-reactions in Appendix 2B. The couple with the more positive potential will act as an oxidizing agent (and be the site of reduction). That couple will be the right-hand electrode in the cell diagram corresponding to the spontaneous cell reaction. To calculate the standard emf of the cell, subtract the standard potential of the oxidation half-reaction (the one with the less-positive standard potential) from that of the reduction half-reaction. To write the cell reaction, follow the procedure in Toolbox 12.2. 

The standard reduction potential for this half-reaction (from Table 14.1) is +0.22233 V. The potential is dependent only on the [CT], as was the potential of the SCE, and once again [CT] is constant because the solution is saturated. Thus this electrode is also appropriate for use as a reference electrode. 

This means that the Ni electrode is the anode and must be involved in oxidation, so its reduction half-reaction must be reversed, changing the sign of the standard half-cell potential, and added to the silver half-reaction. Note that the silver half-reaction must be multiplied by two to equalize electron loss and gain, but the half-cell potential remains the same ... 

Step 3 The standard electrode potential for the reduction half-reaction is Rred = 1.066 V. Changing the sign of the potential for the oxidation half-reaction gives... 

In this section, you learned that you can calculate cell potentials by using tables of half-cell potentials. The half-cell potential for a reduction half-reaction is called a reduction potential. The half-cell potential for an oxidation half-reaction is called an oxidation potential. Standard half-cell potentials are written as reduction potentials. The values of standard reduction potentials for half-reactions are relative to the reduction potential of the standard hydrogen electrode. You used standard reduction potentials to calculate standard cell potentials for galvanic cells. You learned two methods of calculating standard cell potentials. One method is to subtract the standard reduction potential of the anode from the standard reduction potential of the cathode. The other method is to add the standard reduction potential of the cathode and the standard oxidation potential of the anode. In the next section, you will learn about a different type of cell, called an electrolytic cell. 

The midpoint potential of a half-reaction E, is the value when the concentrations of oxidized and reduced species are equal, [Aox] = [Aredl- In biological systems the standard redox potential of a compound is the reduction/oxidation potential measured under standard conditions, defined at pH = 7.0 versus the hydrogen electrode. On this scale, the potential of 02/water is +815 mV, and the potential of water/H2 is 414 mV. A characteristic of redox reactions involving hydrogen transfer is that the redox potential changes with pH. The oxidation of hydrogen H2 = 2H + 2e is an m = 2 reaction, for which the potential is —414 mV at pH 7, changing by 59.2 mV per pH unit at 30°C. 

Table 7.1 Standard Electrode Potentials and Equilibrium Constants for Some Reduction Half-Reactions. ...

The measured voltage of the Daniel cell is 1.10 V. This is the overall voltage of the cell consisting of two half-reactions, namely the oxidation of zinc and the reduction of cupric ion. The individual potentials of the half-reactions cannot be measured. The potentials of the half-reactions can be obtained relative to a standard. The standard is a hydrogen electrode. This consists of a platinum electrode immersed in 1M HC1 with hydrogen gas bubbling through at 1 atmosphere pressure. The reaction is ... 

N electrode potentials potentials, E, of half-reactions as reductions-versus the standard hydrogen electrode. 

Standard reduction potential the potential of a half-reaction under standard state conditions, as measured against the potential of the standard hydrogen electrode. (11.2)... 

Table 8.3 lists a few representative standard electrode potentials (or reduction potentials). Figure 8.6 exemplifies the principle of an electrochemical cell. The hydrogen electrode is made up of a B-electrode (which does not participate directly in the reaction), which is covered by H2(g), which acts as a redox partner [H2(g) = 2H +2e ]. Pt acts as a catalyst for the reaction between H and H2(g) and acquires a potential characteristic of this reaction. The salt bridge between the two cells contains a concentrated solution of salt (such as KCl) and allows ionic species to diffuse into and out of the half-cells this permits each half-cell to remain electrically neutral. 

The standard electrode potential is sometimes called the standard reduction potential because it is listed by the reduction half-reactions. However, a voltmeter allows no current in the cell during the measurement. Therefore, the conditions are neither galvanic nor electrolytic—the cell is at equilibrium. As a result, the half-reactions listed in the table are shown as reversible. If the reaction occurs in the opposite direction, as an oxidation half-reaction, E° will have the opposite sign. 

By convention, the SHE is defined as having a standard electrode potential, EjJ, of zero volts (V) at standard conditions (25 °C and 1 atm). The standard electrode potentials of other couples are similarly determined as reduction half-reactions at unit activity versus the SHE. If the EjJ for a given half-reaction is > 0, that couple has the potential to oxidize the SHE. A negative EjJ indicates a couple that can reduce the SHE. Tables of redox half-reactions and the corresponding EjJ values can be found in Stumm and Morgan (1996). Table 3.7 gives EjJ values and related parameters from these sources for a dozen environmentally important redox reactions. 

Since in a redox reaction electrons are transferred, and since electrons have charge, there is an electric potential E associated with any redox reaction. The potentials for the oxidation component and reduction component of a reaction can be approximated separately based upon a standard hydrogen electrode (SHE) discussed later in this lecture. Each component is called a hall reaction. Of course, no half reaction will occur by itself any reduction half reaction must be accompanied by an oxidation half reaction. There is only one possible potential for any given half reaction. Since tire reverse of a reduction half reaction is an oxidation half reaction, it would be redundant to list potentials for both the oxidation and reduction half reactions. Therefore, half reaction potentials are usually listed as reduction potentials To find the oxidation potential for the reverse half reaction, the sign of the reduction potential is reversed. Below is a list of some common reduction potentials. 

E = electrode potential of the half-ceU E° - standard electrode potential when aRed/ ox = 1 n = number of electrons involved in the reduction reaction N= RxTxhi 10) fF (the Nernst factor if n = 1)... 

By international convention, the standard potentials of electrodes are tabulated for reduction half-reactions. These indicate the tendencies of the electrodes to behave as cathodes toward the SHE. Electrodes with positive values for reduction half-reactions act as cathodes versus the SEfE. Those with negative values for reduction half-reactions act as anodes versus the SEfE. 

Table 1.4. Reduction Half-Reactions and Their Standard Electrode Potentials (Ej) at 25 C...

Standard electrode potential By convention, the potential ( ) of a half-reaction as a reduction relative to the standard hydrogen electrode, when all species are present at unit activity. 

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