Standard Electrode Half-Cell Potentials at

Table 2.1 Standard aqueous half-cell potentials at 25 °C (also known as standard electrode, redox, or oxidation potentials, and as the standard emf series)(a)...

Equations 2.72 and 2.73 are Nemst half-cell equations. For example, with Eq 2.73, when aMm+= 1, E m Mn,+ = E m+ Hence, E Mm+ is the half-cell potential at unit activity of the ions (i.e., the standard electrode half-cell potential). Values of the standard potentials of many electrode reactions are available in the literature, some of which are given in Table 2.1 (Ref 2, 7, 8). All values are given in sign and magnitude relative to the standard hydrogen electrode as previously discussed. 

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. 

We then construct a standard cell consisting of a standard hydrogen electrode and some other standard electrode (half-cell). Because the defined electrode potential of the SHE contributes exactly 0 volt to the sum, the voltage of the overall cell then lets us determine the standard electrode potential of the other half-cell. This is its potential with respect to the standard hydrogen electrode, measured at 25°C when the concentration of each ion in the solution is 1 M and the pressure of any gas involved is 1 atm. 

The potential difference is closely related to the difference of the electrochemical potential based on the electrochemical affinity. If we could measure A(p directly, we could organize the table of electromotive forces based on the Galvani potential difference. However, A

reference electrode to measure the half cell potential at an electrode. When a certain electrcxle is coupled with a reference electrode, then the electromotive force can be measured. Since we usually use some reference electrodes as standards, the electromotive force is defined as the equilibrium potential of the reaction. The table was made in such a way and the hydrogen reference electrode was used to measure and calculate potentials for the half cell reactions. 

If electron flow between the electrodes is toward the sample half-cell, reduction occurs spontaneously in the sample half-cell, and the reduction potential is said to be positive. If electron flow between the electrodes is away from the sample half-cell and toward the reference cell, the reduction potential is said to be negative because electron loss (oxidation) is occurring in the sample halfcell. Strictly speaking, the standard reduction potential, is the electromotive force generated at 25°C and pH 7.0 by a sample half-cell (containing 1 M concentrations of the oxidized and reduced species) with respect to a reference half-cell. (Note that the reduction potential of the hydrogen half-cell is pH-dependent. The standard reduction potential, 0.0 V, assumes 1 MH. The hydrogen half-cell measured at pH 7.0 has an of —0.421 V.)... 

SHE, standard hydrogen electrode The electrode used as a standard against which aU other half-cell potentials are measured. The following reaction occurs at the platinum electrode when immersed in an acidic solution and cormected to the other half of an electrochemical cell 2H (aq) -H 2e —> H2(g). The half- cell potential of this reaction at 25°C, 1 atm and 1 m concentrations of aU solutes is agreed, by convention, to be OV... 

The reduction-oxidation potential (typically expressed in volts) of a compound or molecular entity measured with an inert metallic electrode under standard conditions against a standard reference half-cell. Any oxidation-reduction reaction, or redox reaction, can be divided into two half-reactions, one in which a chemical species undergoes oxidation and one in which another chemical species undergoes reduction. In biological systems the standard redox potential is defined at pH 7.0 versus the hydrogen electrode and partial pressure of dihydrogen of 1 bar. 

M) half-cell and a standard hydrogen electrode. Electrons flow from the zinc anode to the S.H.E. (cathode). The measured standard cell potential at 25°C is 0.76 V. 

The numerical value of an electrode potential depends on the nature of the particular chemicals, the temperature, and on the concentrations of the various members of the couple. For the purposes of reference, half-cell potentials are taken at the standard states of all chemicals. Standard state is defined as 1 atm pressure of each gas (the difference between 1 bar and 1 atm is insignificant for the purposes of this chapter), the pure substance of each liquid or solid, and 1 molar concentrations for every nongaseous solute appearing in the balanced half-cell reaction. Reference potentials determined with these parameters are called standard electrode potentials and, since they are represented as reduction reactions (Table 19-1), they are more often than not referred to as standard reduction potentials (E°). E° is also used to represent the standard potential, calculated from the standard reduction potentials, for the whole cell. Some values in Table 19-1 may not be in complete agreement with some sources, but are used for the calculations in this book. 

The reduction potential is an electrochemical concept. Consider a substance that can exist in an oxidized form X and a reduced form X . Such a pair is called a redox couple. The reduction potential of this couple can be determined by measuring the electromotive force generated by a sample half-cell connected to a standard reference half-cell (Figure 18.6). The sample half-cell consists of an electrode immersed in a solution of 1 M oxidant (X) and 1 M reductant (X ). The standard reference half-cell consists of an electrode immersed in a 1 M H+ solution that is in equilibrium with H2 gas at 1 atmosphere pressure. The electrodes are connected to a voltmeter, and an agar bridge establishes electrical continuity between the half-cells. Electrons then flow from one half-cell to the other. If the reaction proceeds in the direction... 

Each electrode reaction, anode and cathode, or half-cell reaction has an associated energy level or electrical potential (volts) associated with it. Values of the standard equilibrium electrode reduction potentials E° at unit activity and 25°C may be obtained from the literature (de Bethune and Swendeman Loud, Encyclopedia of Electrochemistry, Van Nostrand Reinhold, 1964). The overall electrochemical cell equilibrium potential either can be obtained from AG values or is equal to the cathode half-cell potential minus the anode half-cell potential, as shown above. 

The overall cell potential can be calculated by applying the Nernst equation to the overall cell reaction. We must first find E, the standard cell potential at standard concentrations because the same electrode and the same type of ions are involved in both half-cells, this is always zero. Thus,... 

Thus, electrons flow from the sample half-cell to the standard reference half-cell, and the sample-cell electrode is taken to be negative with respect to the standard-cell electrode. The reduction potential of the X X couple b the observed voltap e at the start of the experiment (when X, X, and are 1 M with 1 atm of H2). The reduction potential of the couple is defined... 

An electrode potential is a measure of the thermodynamics of a redox reaction. It may be expressed as the difference between two half-cell potentials, which by convention are measured against a hydrogen electrode. Tabulated values refer to standard conditions (ions at unit activity). 

Equations (26.20) and (26.27) are combined when two hydrogen electrodes are connected in an electrochemical cell. The left electrode is the standard hydrogen electrode with a half-cell potential of 0.0 V. The right electrode is a hydrogen electrode immersed in a solution at a particular pH (where an+ 1). The resulting equilibrium cell potential, in terms of the right electrode compartment pH, is... 

E Mm+ is called the single electrode or half-cell potential of the M,Mm+ electrode on the standard hydrogen scale. It should be recalled that in this text, E denotes the potential in the general case, E the potential at equilibrium, and E° the potential at equilibrium under standard... 

Erhe andELHE are equilibrium half-cell potentials, or electrode potentials, which depend in sign on the definitions of positive and negative electricity and assignment of Eu + = 0 at standard condi-... 

This silver/silver-chloride electrode is sufficiently important as a reference electrode to find it tabulated in tables of half cells. In these tables, the standard half-cell value is given, which is the potential when the ion functioning as the variable controlling the potential is at unit activity. In the present example, the Cl- ion is the variable and the standard half-cell potential is that which results for aa = 1.0 ... 

The cathode consists of platinum which is an inert conductor in contact with the 1 M ions surrounded by hydrogen gas at 1 atm. Such an electrode is called a standard hydrogen electrode which per definition has a half cell potential (symbolised at 298 K by the symbol of s°) of 0.00 volt. The figure below shows the principle in the build up of the standard hydrogen electrode. 

The standard hydrogen electrode has per definition a half cell potential of 0.0 Volt at a Tf concentration of 1,0 M. 

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