Half-cell potential standard

When a net reaction proceeds in an electrochemical cell, oxidation occurs at one electrode (the anode) and reduction takes place at the other electrode (the cathode.) We can think of the cell as consisting of two half-cells joined together by an external circuit through which electrons flow and an internal pathway that allows ions to migrate between them so as to preserve electroneutrality. 

Reduction potentials Each half-cell has associated with it a potential difference whose magnitude depends on the nature of the particular electrode reaction and on the concentrations of the dissolved electroactive species. The sign of this potential difference depends on the 

In the cell Zn(s) Zn2+(ag) I I Cu2+(aq) I Cu(.s) the zinc appears on the left side, indicating that it is being oxidized, not reduced. For this reason, the potential difference contributed by the left half-cell has the opposite sign to its conventional half-cell potential. More generally, we can define the cell potential or cell EMF as 

Reference half-cells The fact that individual half-cell potentials are not directly measurable does not prevent us from defining and working with them. Although we cannot determine the absolute value of a half-cell potential, we can still measure its value in relation to the potentials of other half cells. In particular, if we adopt a reference half-cell whose potential is arbitrarily defined as zero, and measure the potentials of various other electrode systems against this reference cell, we are in effect measuring the half-cell potentials on a scale that is relative to the potential of the reference cell. 

The reference cell we use for this purpose is the hydrogen half-cell 

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A problem that has fascinated surface chemists is whether, through suitable measurements, one can determine absolute half-cell potentials. If some one standard half-cell potential can be determined on an absolute basis, then all others are known through the table of standard potentials. Thus, if we know E for... 

E° values have been measured for many reactions and tabulated as standard half-cell potentials. Table 9.3 summarizes half-cell potentials as standard reduction potentials for a select set of reactions.aa In the tabulations, E° for... 

In Table 7-1 the relative tendencies of certain elements to react were listed qualitatively. We can give a quantitative measure of relative tendency to react, called standard reduction potential, as shown in Table 14-2. In this table, the standard half-cell potential for each half-reaction, as a reduction, is tabulated in order with the highest potential first. If we turn these half-reactions around, we change the signs of the potentials and we get oxidation potentials. We thus have half-reactions including both elementary metals and elementary nonmetals in the same table, as well as many half-reactions that do... 

Table 8.1 Standard half-cell potentials of some oxidant species in aqueous solution at 25 °C.

We separate the given equation into its two half-equations. One of them is the reduction of nitrate ion in acidic solution, whose standard half-cell potential we retrieve from Table 21-1 and use to solve the problem. 

Thus far, we have based all of our calculations on the standard cell potential or standard half-cell potentials—that is, standard state conditions. However, many times the cell is not at standard conditions—commonly the concentrations are not 1 M. We may calculate the actual cell potential, Ecell, by using the Nemst equation ... 

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 ... 

In section 11.1, you learned that a cell potential is the difference between the potential energies at the anode and the cathode of a cell. In other words, a cell potential is the difference between the potentials of two half-cells. You cannot measure the potential of one half-cell, because a single half-reaction cannot occur alone. However, you can use measured cell potentials to construct tables of half-cell potentials. A table of standard half-cell potentials allows you to calculate cell potentials, rather than building the cells and measuring their potentials. Table 11.1 includes a few standard half-cell potentials. A larger table of standard half-cell potentials is given in Appendix E. 

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 of this standard cell is +0.76 V. By international convention, the half-cell potential of the hydrogen reduction is assigned a value of exactly OV. Thus, the half-cell potential of the zinc oxidation is equal to K.n (i.e., +0.76 V). This voltage is called the standard half-cell potential and is represented by the symbol 1, to indicate that it was determined against a standard hydrogen electrode. 

The standard potential of any galvanic cell is the sum of the standard half-cell potentials for oxidation at the anode and reduction at the cathode ... 

If we could determine E° values for individual half-reactions, we could combine those values to obtain E° values for a host of cell reactions. Unfortunately, it s not possible to measure the potential of a single electrode we can measure only a potential difference by placing a voltmeter between two electrodes. Nevertheless, we can develop a set of standard half-cell potentials by choosing an arbitrary standard half-cell as a reference point, assigning it an arbitrary potential, and then expressing the potential of all other half-cells relative to the reference half-cell. Recall that this same approach was used in Section 8.10 for determining standard enthalpies of formation, A H°f. 

The standard half-cell potential for reduction of quinone to hydroquinone is 0.699 V ... 

The equilibrium constant for the chemical reaction expressed by Equation 2.15 is related to the difference of the standard half-cell potentials by the relation ... 

A third possible convention is due to Gibbs. Here the emphasis is on the cell itself, rather than the depiction of the cell. The emf of the cell is always positive and the difference between the standard half-cell potentials is taken to yield positive values. Then... 

Throughout this discussion we have considered cells in which the electrolytic solution is an aqueous solution. The same methods can be used to define standard half-cell potentials in any solvent system. However, it is important to remember that when the reference state is defined as the infinitely dilute solution of a solute in a particular solvent, the standard state depends upon that solvent. The values so obtained are not interchangeable between the different solvent systems. Only if the standard states could all be defined independently of the solvent would the values be applicable to all solvent systems. 

Table VH shows that hydrazine is the only important variable for technetium sorption on each of the geologic solids. Hydrazine causes technetium to be removed from solution either by sorption or by precipitation of the reduced technetium species. Hydrazine is a powerful reducing agent and should reduce TcO/ to technetium(IV) according to standard half-cell potentials. No Tc02 was observed however, since the technetium passed through...

Sorption/precipitation of plutonium is greatly affected by the presence of hydrazine. Since hydrazine increases sorption, it appears that at least some of the plutonium is present initially as plutonium(V) or plutonium(VI) and is reduced to plutonium(IV) by hydrazine. According to standard half-cell potentials, both plutonium(VI) and plutonium(V) should be reduced to plutonium(IV) or plutonium(HI) under the conditions of the experiments, assuming that the hydroxyl complexes are important at the pHs of the experiments. 

Galvanic Cells and Standard Half-Cell Potentials... 

Standard half-cell potentials are usually tabulated as reduction potentials, e.g.,... 

Standard half-cell potentials can be used to compute standard cell potentials, standard Gibbs free energy changes, and equilibrium constants for oxidation-reduction reactions. 

Standard half-cell potentials Combinations of couples... 

From the table of the standard half-cell potentials ... 

Standard half-cell potentials in aqueous solutions... 

Standard half-cell potentials in aqueous solutions (T = 298 K, Standard state = 1 molal)... 

The quantity AG° is the free energy of the half-reaction when the activities of the reactant and product have values of unity and is directly proportional to the standard half-cell potential for the reaction as written. It also is a measure of the equilibrium constant for the half-reaction assuming the activity of electrons is unity ... 

In the same way, a standard half-cell potential is a measure of the drop in the free energy of the electron when it falls from its source level to the energy of the hydrogen ion at unit effective concentration. By virtue of the defined value of E° = 0 for the H+/H2 couple, the latter level can also be regarded as the zero free energy level of the electron. 

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