Half-cells/reactions standard reduction potentials listed

Some typical half-cell reactions and their respective standard reduction potentials are listed in Table 21.1. Whenever reactions of this type are tabulated, they are uniformly written as reduction reactions, regardless of what occurs in the given half-cell. The sign of the standard reduction potential indicates which reaction really occurs when the given half-cell is combined with the reference hydrogen half-cell. Redox couples that have large positive reduction potentials... 

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... 

In the discussion of the Daniell cell we indicated that this cell produces 1.10 volts. This voltage is really the difference in potential between the two half-cells. There are half-cell potentials associated with all half-cells A list of all possible combinations of half-cells would be tremendously long. Therefore, a way of combining desired half-cells has been developed. The cell potential (really the half-cell potentials) depends on concentration and temperature, but initially we ll simply look at the half-cell potentials at the standard temperature of 298 K (25°C) and all components in their standard states (1 M concentration of all solutions, 1 atmosphere pressure for any gases, and pure solid electrodes). All the half-cell potentials are tabulated as the reduction potentials, that is, the potentials associated with the reduction reaction. The hydrogen half-reaction has been defined as the standard and has been given a value of exactly 0.00 V. All the other half-reactions have been measured relative to it, some positive and some negative. The table of standard reduction potentials provided on the AP exam is shown in Table 16.1 and in the back of this book. 

As shown in Table 7.1, values are usually listed for half-cell reactions written as reductions. Thus, the standard half-cell reduction potential for the Zn /Zn couple is -0.76 V. Half-cell reductions that are strong enough to spontaneously oxidize H2(g) have a positive E. Conversely, half-cell reductions that have a negative E are not strong enough to oxidize H2(g). Instead, they spontaneously proceed as oxidations, causing the reduction of H (aq). 

The table below lists the cell potentials for the 10 possible galvanic cells assembled from the metals A, B, C, D, and E, and their respective 1.00 M 2+ ions in solution. Using the data in the table, establish a standard reduction potential table similar to Table 11.1 in the text. Assign a reduction potential of 0.00 V to the half-reaction that falls in the middle of the series. You should get two different tables. Explain why, and discuss what you could do to determine which table is correct. 

Therefore for a given reaction to take place, the cell potential must be positive. The cell potential is taken as the difference between the two half-cell reactions, the one at the cathode minus the one at the anode. The half-cell potential exists because of the difference in the neutral state compared to the oxidized state, such as Fe/Fe + or, at the cathode, the difference between the neutral state and the reduced state, as in These reduction-oxidation (redox) potentials are measured relative to a standard half-cell potential. The chart shown in Table 2 lists potentials relative to the which is set as zero. 

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. 

Over the years, chemists have measured and recorded the standard reduction potentials, abbreviated of many different half-cells. Table 21-1 lists some common half-cell reactions in order of increasing reduction potential. The values in the table are based on using the half-cell reaction that is being measured as the cathode and the standard hydrogen electrode as the anode. All of the half-reactions in Table 21-1 are written as reductions. However, in any voltaic cell, which always contains two halfreactions, the half-reaction with the lower reduction potential will proceed in the opposite direction and will be an oxidation reaction. In other words, the half-reaction that is more positive will proceed as a reduction and the half-reaction that is more negative will proceed as an oxidation. 

Appendix E summarizes the standard reduction potentials for a large number of half-reactions. The table lists the reactions in order of decreasing reduction potentials—that is, with the most positive at the top and the most negative at the bottom. In any galvanic cell, the half-cell that is listed higher in the table will act as the cathode (if both half-cells are in the standard state). 

Table 19.1 lists standard reduction potentials for a number of half-cell reactions. By definition, the SHE has an E° value of 0.00 V. Above the SHE the negative standard reduction potentials increase, and below it the positive standard reduction potentials increase. It is important to know the following points about the table ... 

In tables of standard potentials, all of the half-cell reactions are expressed as reductions. The sign is reversed if the reaction is reversed to become an oxidation. In a spontaneous reaction, when both half-cells are written as reductions, the half-cell with the more negative potential will be the one that oxidizes. The negative sign in Eq. (15.9) reverses one of the reduction processes to an oxidation. Some standard reduction potentials for common half-cells are given in Appendix 15.1. These can be used to calculate for other electrochemical cell combinations as we have done for the Zn/Cu cell. More complete lists of half-cell potentials can be found in references such as Bard et al., listed in the bibliography. 

Equation (5) or (11) can be applied directly to half-cell reactions such as (6) and (7) and the resulting potentials obtained will be identical to those obtained from the overall reactions (9) and (10) because of the definition of the SHE as the universal standard. A selection of standard potentials of half-cell reactions is shown in Table 1 [5]. By international convention, electrode reactions in thermodynamic tables are always written as reduction reactions, so the more noble metals have a positive standard potential. Lists such as that in Table 1 are also called electromotive force series or tables of standard reduction potentials. 

The standard reduction potentials for other half-reactions can be determined in a fashion analogous to that used for the Zx " jZvr half-reaction. TABLE 20.1 lists some standard reduction potentials a more complete list is found in Appendix E. These standard reduction potentials, often called half-cell potentials, can be combined to calculate values for a large variety of voltaic cells. 

Table 19.1 lists standard reduction potentials for a number of half-cell reactions, nie activity senes in Rgure 4.1 e is based... 

When the standard electrode potentials are listed in decreasing value as shown in Table 14.6, an electromotive series is created which has the hydrogen half-cell reaction listed at a potential of zero. These values are reduction potentials at 25°C referred to the standard hydrogen electrode (SHE). Metals listed at the top of the series are noble or less reactive. Metals listed below the hydrogen reaction are reactive, that is, they corrode more readily. A metal listed below another metal will displace it from a solution containing the higher metal s ions. Iron, for example, wUl have copper metal plated on it when placed in a copper sulfate solution. This is an iron corrosion reaction. Metals listed below hydrogen will displace the H" " ions from acid solutions. 

Table 19.1 lists standard reduction potentials, in order of decreasing reduction potential, for a number of haif-ceU reactions. To avoid ambiguity, all half-cell reaaions are shown as reductions. A galvanic cell is composed of two half-cells, and therefore two half-cell reactions. Whether dr not a particular half-cell reaction occurs as a reduction in a galvanic cell depends on how its reduction potential compares to that of the other half-cell reaction. If it has the greater (or more positive) reduction potential of the two, it wUl occur as a reduction. If it has the smaller (or more negative) reduction potential of the two, it will occur in the reverse direction, as an oxidation. 

Table 21-1 on page 667 of your textbook lists the standard reduction potentials ( ° values) of common half-cell reactions. These values apply to the standard conditions of a IM solution, 25°C, and 1 atm pressure. In any voltaic cell, the half-reaction with the lower reduction potential will proceed in the opposite direction, as an oxidation half-reaction. The half-reaction with the higher reduction potential will proceed as a reduction. The electrical potential of a voltaic cell, also called the cell potential, is found by subtracting the standard reduction potential of the oxidation halfreaction from the standard reduction potential of the reduction half-reaction. 

Use of reduction potential tables—It can be determined which species is reduced or oxidized in a redox couple from standard tables, as exempHfied by the brief listing given in Table 19. The cathode and anode half-cell reactions are noted and added algebraically as noted for Pb and Sn. Note that the algebraic sign of the Sn electrode potential is also reversed (i.e., from —0.136 to +0.136) ... 

In summary, the potential of the standard hydrogen electrode is set at exactly 0 V. Any electrode at which a reduction half-cell reaction shows a greater tendency to occur than does the reduction of H" " (1 M) to H2 (g, 1 bar) has a positive value for its standard electrode potential, E°. Any electrode at which a reduction half-cell reaction shows a lesser tendency to occur than does the reduction of H" "(l M) to H2 (g, 1 bar) has a negative value for its standard reduction potential, E°. Comparisons of the standard copper and zinc electrodes to the standard hydrogen electrode are illustrated in Figure 19-6. Table 19.1 on page 875 lists some cormnon reduction half-cell reactions and their standard electrode potentials at 25 °C. 

Standard half-cell voltages are ordinarily obtained from a list of standard potentials such as those in Table 18.1 (page 487). The potentials listed are the standard voltages for reduction half-reactions, that is,... 

Many reactions that occur in living cells are oxidation-reduction reactions. Appendix IX lists several compounds of biological importance and shows their relative tendencies to gain electrons ai 25°C and pH 7 under standard conditions. The numerical values of Ho reflect the reduction potentials relative to the 2H + 2e" H2 half-reaction which is taken as — 0.414 volt at pH 7. The value for the hydrogen half-reaction at pH 7 was calculated from the arbitrarily assigned value (Ho) of 0.00 volt under true standard-state conditions (1 M H and 1 atm Hs). For those few halfreactions of biological importance that do not involve as a reactant, the Ho and Ho values are essentially identical. 

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. 

Referring to a list of standard electrode potentials, such as in Table 8.3, one speaks of an electrochemical series, and the metals lower down in the se-ries(with positive electrode potentials) are called noble metals. Any combination of half-reactions in an electrochemical cell, which gives a nonzero E value, can be used as a galvanic cell (i.e., a battery). If the reaction is driven by an applied external potential, we speak of an electrolytic cell. Reduction takes place at the cathode and oxidation at the anode. The reduction reactions in Table 8.3 are ordered with increasing potential or pe values. The oxidant in reactions with latter pe (or E°) can oxidize a reductant at a lower pe (or ) and vice versa for example, combining half-reactions we obtain an overall redox reaction ... 

By combining many pairs of half-cells into voltaic cells, we can create a list of reduction half-reactions and arrange them in decreasing order of standard electrode potential (from most positive to most negative). Such a list, called an emf series or a table of standard electrode potentials, appears in Appendix D, with a few examples in Table 21.2 on the next page. 

In general, only the reduction half-reaction potentials are listed in tables, as in Table 9.1. The potential of an oxidation half-reaction is the negative of the value of the reduction half-reaction. Moreover, it is convenient to standardise the concentrations of the components of the cells. If the ceU components are in their standard states, standard electrode potentials, E°, are recorded ... 

Before we discuss standard electrode potential, we will talk about electromotive force (emf). The electromotive force of a cell is the potential difference between the two electrodes. This can be measured using a voltmeter. The maximum voltage of a cell can be calculated using experimentally determined values called standard electrode potentials. By convention, the standard electrode potentials are usually represented in terms of reduction half-reactions for 1 molar solute concentration. The standard electrode potential values are set under ideal and standard-state conditions (latm pressure and 25°C temperature). From the MCAT point of view, you can assume that the conditions are standard, unless stated otherwise. Table 12-1 shows a list of standard electrode potentials (in aqueous solution) at 25°C. 

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