Standard states electrode potential

I11 this equation pt/rai. ox stands for the standard state electrode potential, when the activities of both ions ilf t and M t equal unity. ) Thus e. g. the electrode potential attaining equilibrium in the following process... 

V°rev = 1.229V is the standard state reversible potential for the water splitting reaction and Vaoc is the anode potential at open circuit conditions. Term Vmeas-Vaoc arises from the fact that Voc represents the contribution of light towards the minimum voltage needed for water splitting potential (1.229V) and that the potential of the anode measured with respect to the reference electrode Vmeas has contributions from the open circuit potential and the bias potential applied by the potentiostat (i.e. Vmeas= Vapp+Vaoc). The term Vmeas-Vaoc makes relation (3.6.16) independent of the electrolyte pH and the type of reference electrode used. Thus the use of V°rev in relation (3.6.16) instead of VV or V°hz as in the case of relation (3.6.15) is justified. 

In the last section it was shown that instead of representing an electrode potential on a relative scale (arbitrarily setting the standard hydrogen electrode potential equal to zero), it is possible to numerically calculate the actual value of the latter, with a reference state of zero energy for the stationary electron at infinity in a vacuum. 

W. M. Latimer, The Oxidation States of the Elements and their Potentials in Aqueous Solutions, 2nd ed., Prentice-Hall, Englewood Cliffs, NJ, 1952. Although a useful source in its time, this book is now outdated. Unfortunately, no comprehensive compilation of electrochemical thermodynamic data has entirely replaced it. However, useful data sources include A. 1. de Be-thune, N. A. S. Loud, Standard Aqueous Electrode Potentials and Temperature Coefficients at 2S°C, C. A. Hampel, Skokie, IL, 1964 A. J. Bard, ed.. Encyclopedia of Electrochemistry of the Elements, Vols. 1-10, Marcel Dekker, New York, 1973-1976. 

The standard electrode potential at the above mentioned standard state conditions is denoted by °. For the MCAT, the values of the standard electrode (reduction) potentials will be given to you if you are required to solve such a question. Do not try to memorize those values. The standard electrode potentials are based on an arbitration with reference to standard hydrogen electrode. The standard hydrogen electrode potential is considered to be 0 volt. 

Both Sn and Sn in the corrosion reaction are at standard state. The potential measured for the Sn Sn electrode is 0.138 V with Sn Sn negative vs. the hydrogen reference electrode. Because the potential of the SHE is defined as zero, the standard potential of the Sn I Sn redox system is e° =—0.138 V, as presented in Table 2.2. Electrodes shown... 

The standard-state cell potential difference of this cell is -0.2223 V, the negative of that of the cell of Figure 8.2. Hydrogen gas at the same pressure is fed into both hydrogen electrodes and the two HCI solutions are at the same concentration. A wire is connected between the two hydrogen electrodes and maintains them at the same electric potential. We can write... 

The back e.m.f. is a voltage that opposes the passage of a current through an electrolytic cell. There are three sources of the back e.m.f. The first is the reversible back e.m.f. due to the cell reaction. For example, in a Daniell cell with unit activities the reversible back e.m.f. is the equilibrium standard-state cell potential of 1.100 V. For activities other than unit activities, the reversible back e.m.f. can be calculated from the Nernst equation. For an infinitesimal electrolytic current, the reversible back e.m.f. is the only contribution to the back e.m.f. For a finite current, the IR drop in the voltage across the electrolyte solution due to its electrical resistance also contributes. In many cases, we will be able to neglect this contribution. The third source of back e.m.f. for a finite current is the overpotential, which is due to the polarization of the electrode. 

Standard Hydrogen Electrode The standard hydrogen electrode (SHE) is rarely used for routine analytical work, but is important because it is the reference electrode used to establish standard-state potentials for other half-reactions. The SHE consists of a Pt electrode immersed in a solution in which the hydrogen ion activity is 1.00 and in which H2 gas is bubbled at a pressure of 1 atm (Figure 11.7). A conventional salt bridge connects the SHE to the indicator half-cell. The shorthand notation for the standard hydrogen electrode is... 

The ladder diagram for this system is shown in Figure 11.24a. Initially the potential of the working electrode remains nearly constant at a level near the standard-state potential for the Fe UFe redox couple. As the concentration of Fe + decreases, however, the potential of the working electrode shifts toward more positive values until another oxidation reaction can provide the necessary current. Thus, in this case the potential eventually increases to a level at which the oxidation of H2O occurs. 

Other Coordination Complexes. Because carbonate and bicarbonate are commonly found under environmental conditions in water, and because carbonate complexes Pu readily in most oxidation states, Pu carbonato complexes have been studied extensively. The reduction potentials vs the standard hydrogen electrode of Pu(VI)/(V) shifts from 0.916 to 0.33 V and the Pu(IV)/(III) potential shifts from 1.48 to -0.50 V in 1 Tf carbonate. These shifts indicate strong carbonate complexation. Electrochemistry, reaction kinetics, and spectroscopy of plutonium carbonates in solution have been reviewed (113). The solubiUty of Pu(IV) in aqueous carbonate solutions has been measured, and the stabiUty constants of hydroxycarbonato complexes have been calculated (Fig. 6b) (90). 

In addition to simple dissolution, ionic dissociation and solvolysis, two further classes of reaction are of pre-eminent importance in aqueous solution chemistry, namely acid-base reactions (p. 48) and oxidation-reduction reactions. In water, the oxygen atom is in its lowest oxidation state (—2). Standard reduction potentials (p. 435) of oxygen in acid and alkaline solution are listed in Table 14.10- and shown diagramatically in the scheme opposite. It is important to remember that if or OH appear in the electrode half-reaction, then the electrode potential will change markedly with the pH. Thus for the first reaction in Table 14.10 O2 -I-4H+ -I- 4e 2H2O, although E° = 1.229 V,... 

The standard electrode potential of magnesium is given, along with the potentials of other metals, in Table 4.17 and the steady-state potentials of magnesium in various solutions are listed in Table 4.18. ... 

The two elements have similar electronegativity. (Note electronegativity is the power of an element to attract electrons to itself when present in a molecule or in an aggregate of unlike atoms it is a different property from the electrode potential, which depends on the free energy difference between an element in its standard state and a compound or ion in solution (see Section 20.1).) In addition a metal of a lower valency tends to dissolve a metal of a higher valency more readily than vice versa. 

Electrode Potential (E) the difference in electrical potential between an electrode and the electrolyte with which it is in contact. It is best given with reference to the standard hydrogen electrode (S.H.E.), when it is equal in magnitude to the e.m.f. of a cell consisting of the electrode and the S.H.E. (with any liquid-junction potential eliminated). When in such a cell the electrode is the cathode, its electrode potential is positive when the electrode is the anode, its electrode potential is negative. When the species undergoing the reaction are in their standard states, E =, the stan-... 

Standard Electrode Potential (E ) the equilibrium potential of an electrode reaction when the components are ail in their standard states. 

When metals are arranged in the order of their standard electrode potentials, the so-called electrochemical series of the metals is obtained. The greater the negative value of the potential, the greater is the tendency of the metal to pass into the ionic state. A metal will normally displace any other metal below it in the series from solutions of its salts. Thus magnesium, aluminium, zinc, or iron will displace copper from solutions of its salts lead will displace copper, mercury, or silver copper will displace silver. 

Similarly to aqueous electrochemistry, potentials in solid state electrochemistry utilizing YSZ are expressed in terms of the potential of a reference metal electrode exposed to P02 = 1 atm at the temperature T of interest. Thus a standard oxygen electrode scale (soe) can be defined. Similarly to equation (7.2) one has ... 

In redox couple notation, E°(HJ"/H2) = 0 at all temperatures. A hydrogen electrode in its standard state, with hydrogen gas at 1 bar and the hydrogen ions present at 1 mol-L 1 (strictly, unit activity), is called a standard hydrogen electrode (SHE). The standard hydrogen electrode is then used to define the standard potentials of all other electrodes ... 

Details are given only if relevant for the stated value of iisHE- l All potentials refer to a standard hydrogen electrode in the used solvent. Exemptions are stated explicitly. 

If no concentration of the educt is given the standard exchange current density y oo is stated. Values of)t are printed in italics values of the apparent rate constant k pp are printed in parentheses in italics. For electrode potentials where the latter rate constant was actually determined the reader is referred to the original literature. 

In the introductory chapter we stated that the formation of chemical compounds with the metal ion in a variety of formal oxidation states is a characteristic of transition metals. We also saw in Chapter 8 how we may quantify the thermodynamic stability of a coordination compound in terms of the stability constant K. It is convenient to be able to assess the relative ease by which a metal is transformed from one oxidation state to another, and you will recall that the standard electrode potential, E , is a convenient measure of this. Remember that the standard free energy change for a reaction, AG , is related both to the equilibrium constant (Eq. 9.1)... 

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