Cell potential standard values

The half cell potential (s° values) corresponding to the given half cell reactions are sometimes referred to as standard reduction potentials and may be found in tables. Typically one of the half cell reactions should be turned around as each redox reaction contains one reduction reaction and one oxidation reaction. This means that one of the signs from the table must be turned around. Further may the half cell reaction from the table typically be multiplied with an integer in order to make sure that the number of electrons fits the number of consumed electrons from the overall cell reaction. These integer multiplications are associate with the determination of the free energy of Gibb s for a half cell reaction ... 

It is possible to obtain values for ii°ceU experimentally, although it is usual in the laboratory to work with solutions of concentrations measure values of Fceii (rather than standard cell potentials). Such values are dependent on solution concentration (strictly, activity), and Fceii and 7s°ceii are related by the Nernst equation (see equation 7.21). ... 

Find the standard cell potential, using values from Table 21-1 in your textbook. 

The potential of the reaction is given as = (cathodic — anodic reaction) = 0.337 — (—0.440) = +0.777 V. The positive value of the standard cell potential indicates that the reaction is spontaneous as written (see Electrochemical processing). In other words, at thermodynamic equihbrium the concentration of copper ion in the solution is very small. The standard cell potentials are, of course, only guides to be used in practice, as rarely are conditions sufftciendy controlled to be called standard. Other factors may alter the driving force of the reaction, eg, cementation using aluminum metal is usually quite anomalous. Aluminum tends to form a relatively inert oxide coating that can reduce actual cell potential. 

To calculate the open circuit voltage of the lead—acid battery, an accurate value for the standard cell potential, which is consistent with the activity coefficients of sulfuric acid, must also be known. The standard cell potential for the double sulfate reaction is 2.048 V at 25 °C. This value is calculated from the standard electrode potentials for the (Pt)H2 H2S04(yw) PbS04 Pb02(Pt) electrode 1.690 V (14), for the Pb(Hg) PbS04 H2S04(yw) H2(Pt) electrode 0.3526 V (19), and for the Pb Pb2+ Pb(Hg) 0.0057 V (21). 

Thus, because the standard cell potential for reaction 15 is positive, the reaction proceeds spontaneously as written. Consequendy, to produce chlorine and hydrogen gas, a potential must be appHed to the cell that is greater than the open-circuit value. This then becomes an example of an electrolytic process. 

Figure 21.2a shows a sample/reference half-cell pair for measurement of the standard reduction potential of the acetaldehyde/ethanol couple. Because electrons flow toward the reference half-cell and away from the sample half-cell, the standard reduction potential is negative, specifically —0.197 V. In contrast, the fumarate/succinate couple and the Fe /Fe couple both cause electrons to flow from the reference half-cell to the sample half-cell that is, reduction occurs spontaneously in each system, and the reduction potentials of both are thus positive. The standard reduction potential for the Fe /Fe half-cell is much larger than that for the fumarate/ succinate half-cell, with values of + 0.771 V and +0.031 V, respectively. For each half-cell, a half-cell reaction describes the reaction taking place. For the fumarate/succinate half-cell coupled to a H Hg reference half-cell, the reaction occurring is indeed a reduction of fumarate. 

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... 

A table giving the cell potentials of all possible redox reactions would be immense. Instead, chemists use the fact that any redox reaction can be broken into two distinct half-reactions, an oxidation and a reduction. They assign a potential to every half-reaction and tabulate E ° values for all half-reactions. The standard cell potential for any redox reaction can then be obtained by combining the potentials for its two half-reactions. 

V) from the more positive value (-0.447 V) to obtain the standard cell potential ... 

The linkage between free energy and cell potentials can be made quantitative. The more negative the value of A G ° for a reaction, the more positive its standard cell potential, as the following two examples illustrate ... 

Equation expresses an important link between two standard quantities. The equation lets us calculate standard electrical potentials from tabulated values for standard free energies. Equally important, accurate potential measurements on galvanic cells yield experimental values for standard potentials that can be used to calculate standard free energy changes for reactions. 

The Nernst equation is used to calculate electrode potentials or cell potentials when the concentrations and partial pressures are other than standard state values. The Nernst equation using both base 10 and natural logarithms is given by ... 

The standard cell potentials do not tell the whole story, for we might think that the reaction with the more positive value of E°ell would occur. In truth, Al(s) is coated with a thin layer of tightly adhering Al203(s), which protects the metal from attack by water. Na(s) has no such protective coating. 

In the discussion of the Daniell cell, we indicated that this cell produces a voltage of 1.10 V. This voltage is really the difference in potential between the two half-cells. The cell potential (really the half-cell potentials) is dependent upon concentration and temperature, but initially we ll simply look at the half-cell potentials at the standard state of 298 K (25°C) and all components in their standard states (1M concentration of all solutions, 1 atm pressure for any gases and pure solid electrodes). Half-cell potentials appear in tables as the reduction potentials, that is, the potentials associated with the reduction reaction. We define the hydrogen half-reaction (2H+(aq) + 2e - H2(g)) as the standard and has been given a value of exactly 0.00 V. We measure all the other half-reactions relative to it some are positive and some are negative. Find the table of standard reduction potentials in your textbook. 

The more the two half-reactions are separated in the table, the greater is the tendency for the net reaction to occur. This tendency for an overall redox reaction to occur, whether by direct contact or in an electrochemical cell, is determined from the standard reduction potentials, E° values, of the half-reactions involved, and the value of this potential are indications of the tendency of the overall redox reaction to occur. We will now present a scheme for determining this potential, which is symbolized E"d. ... 

The standard cell potential for a galvanic cell is a positive value, E° > 0. 

The numerical values of cell potentials and half-cell potentials depend on various conditions, so tables of standard reduction potentials are true when ions and molecules are in their standard states. These standard states are the same as for tables of standard enthalpy changes. Aqueous molecules and ions have a standard concentration of 1 mol/L. Gases have a standard pressure of 101.3 kPa or 1 atm. The standard temperature... 

Because you can measure potential differences, but not individual reduction potentials, all values in the table are relative. Each half-cell reduction potential is given relative to the reduction potential of the standard hydrogen electrode, which has been assigned a value of zero. The design of this electrode is shown in Figure 11.12. 

This calculation of the standard cell potential for the Daniell cell used the mathematical concept that the subtraction of a negative number is equivalent to the addition of its positive value. You saw that 0.342 V - (-0.762 V) = 0.342 V + 0.762 V... 

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. 

IfflD For which half-cell are the values of the standard reduction potential and the standard oxidation potential equal ... 

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. 

Cell potential and reaction free energy are related by Eq. 2 (AGr = -nFE) and their standard values by Eq. 3 (AGr° = —nFE°). The magnitude of the cell potential is independent of how the chemical equation is written. 

A problem with this procedure is that we know only the overall cell potential, not the contribution of each individual electrode. A voltmeter placed between the two electrodes of a galvanic cell measures the difference of their potentials, not the individual values. To get around this difficulty, we arbitrarily set the standard potential of one particular electrode, the hydrogen electrode, equal to 0 at all temperatures ... 

The value quoted here, like all the values in this text, is for 25°C. Because, according to convention, the hydrogen electrode contributes 0 to the standard cell potential, the standard potential of the cell, 0.76 V, can be attributed entirely to the zinc electrode. Moreover, because the zinc electrode is known by experiment to be the anode, its standard potential is subtracted from that of hydrogen ... 

Strategy First, write the balanced equation and the corresponding expression for Q and note the value of n. Then determine E° from the standard potentials in Table 12.1 (or Appendix 2B). Determine the value of Q for the stated conditions. Calculate the cell potential by substituting these values into the Nernst equation, Eq. 7. At 25°C, RT/F = 0.025 693 V. 

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. 

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