Standard state hydrogen half-cell

The cell reaction for cells without liquid junction can be written as the sum of an oxidation reaction and a reduction reaction, the so-called half-cell reactions. If there are C oxidation reactions, and therefore C reduction reactions, there are C C — 1) possible cells. Not all such cells could be studied because of irreversible phenomena that would take place within the cell. Still, a large number of cells are possible. It is therefore convenient to consider half-cell reactions and to associate a potential with each such reaction or electrode. Because of Equation (12.88), there would be (C - 1) independent potentials. We can thus assign an arbitrary value to the potential associated with one half-cell reaction or electrode. By convention, and for aqueous solutions, the value of zero has been assigned to the hydrogen half-cell when the hydrogen gas and the hydrogen ion are in their standard states, independent both of the temperature and of the pressure on the solution. 

This is arbitrarily assigned a standard reduction potential Eo= 0.0 V. At the biochemical standard state of pH 7, the hydrogen half-cell has an Eq = —0.421 V. 

Because the standard-state half-cell potential, , is measured relative to the zero potential of the hydrogen half-cell, = El, and the definition of " given by equation 7.27 is substituted into equation 7.22 to give... 

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... 

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. 

Redox half-reactions are often written for brevity as, for example, Li+ + e - Li. with the state symbols omitted. The electrode system represented by the half-reaction may also be written as Li+ /Li. The standard redox potentials for ion-ion redox systems can be determined by setting up the relevant half-cell and measuring the potential at 298 K relative to a standard hydrogen electrode. For example, the standard redox potential for the half-reactions... 

Many half-reactions of interest to biochemists involve protons. As in the definition of AG °, biochemists define the standard state for oxidation-reduction reactions as pH 7 and express reduction potential as E °, the standard reduction potential at pH 7. The standard reduction potentials given in Table 13-7 and used throughout this book are values for E ° and are therefore valid only for systems at neutral pH Each value represents the potential difference when the conjugate redox pair, at 1 m concentrations and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2ET/H2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the standard (pH 0) cell the measured E ° for the 2ET/H2 pair is -0.414 V... 

Because (like AG) refers to a difference in a state property, it can be evaluated in additive fashion along many alternative pathways. For this purpose, it is convenient to assign conventional ° values to each half-cell reaction [e.g., standard oxidation potentials as compiled in W. M. Latimer. Oxidation Potentials, 2nd edn (Prentice-Hall, New York, 1952)], such that the algebraic sum of the two half-reaction potentials equals the overall cell °. Such half-reaction ° values can in turn be obtained by choosing some standard electrode reaction as the conventional zero of the scale [such as the standard hydrogen electrode (SHE) for the l/2H(g) —> H+ aq) + e oxidation reaction, with she = 0]. Sidebar 8.2 illustrates a simple example of this procedure. 

Factors Involved in Galvanic Corrosion. Emf series and practical nobility of metals and metalloids. The emf. series is a list of half-cell potentials proportional to the free energy changes of the corresponding reversible half-cell reactions for standard state of unit activity with respect to the standard hydrogen electrode (SHE). This is also known as Nernst scale of solution potentials since it allows to classification of the metals in order of nobility according to the value of the equilibrium potential of their reaction of dissolution in the standard state (1 g ion/1). This thermodynamic nobility can differ from practical nobility due to the formation of a passive layer and electrochemical kinetics. 

Hydrogen gas at 1 atm pressure is bubbled over a Pt electrode in a 1 mol L 1 H+ solution. En values for other half-cell reactions (with their components in their standard states) may be measured using electrochemical cells in which the hydrogen electrode is linked to an electrode at which the reaction of interest occurs. 

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. 

To determine the potential for the cell, the hydrogen electrode is by convention assigned a half-reaction potential E of 0.00 volt (V) under standard state conditions. This means that if the H ion activity and the H2 gas activity are both 1.00, for the hydrogen electrode is 0.00 volt. This convention can be symbolized as... 

We adopt the following convention the standard-state potential difference of a cell consisting of a hydrogen electrode on the left and any other electrode on the right is called the standard reduction potential of the right half-cell or the right electrode. It is also sometimes called the standard electrode potential. 

The standard reduction potential (E°) provides a measure of the stability of a metal in a particular oxidation state. The E° value is the voltage generated in a half-cell coupled with the standard hydrogen electrode (SHE), which itself has a defined half-cell potential of 0.0 V. Put simply, the more positive is E0 the more difficult is it for metal oxidation to a hydrated metal ion to occur. Alternatively, we could express it by saying that the less positive is °, the more stable is the metal in the higher oxidation state of its couple... 

Hydrogen gas is bubbled over the platinum electrode at 1 atm of pressure, in an H" solution of 1 mol L . The Eq value for another half-cell reaction (with its components in their standard states as well) is measured by linking its electrochemical cell with the standard hydrogen cell. 

The half-cell of the hydrogen reference electrode consists of platinum foil, which serves as a conductor. The platinum foil is in contact with a sulfuric acid solution that contains cations of unit activity, in equilibrium with H2 gas, in its standard state of 1 atm [12-17]. The standard hydrogen electrode may be represented as ... 

The value of this electric field depends on the electrochemical potentials of the electrode metals and is related to their position in the electrochemical series. Table l7.3 shows some values for some common electrode metals at romn temperature. Electrochemists have by convention adopted the hydrogen electrode as a standard of reference potmitial and assigned it a value of 0 V. All other metals have a ncmzero potential with respect to it Metal half-cell potentials depend on their electrochemical oxidation state, and they are usually arranged in a table showing their activity relative to others, such as seen in Table 17.3. 

Measured by the cell S.H.E. II half-reaction of interest, where S.H.E. is the standard hydrogen electrode and all reagents in the right half-cell are in their standard state (= 1 M, 1 bar, pure solid, or pure liquid)... 

When the and Zn ions are in their standard states (that is, both have an activity of 1 in the solution), we find that the emf of a Daniell cell is 1.10 V at 25°C (see Figure 13.4). This emf must be related directly to the redox reactions, but how Just as the overall cell reaction can be thought of as the sum of two half-cell reactions, the measured emf of the cell can be treated as the sum of the electric potentials at the Zn and Cu electrodes. Knowing one of these electrode potentials, we could obtain the other by subtraction (from 1.10 V). It is impossible to measure the potential of just a single electrode, but if we arbitrarily set the potential value of a particular electrode at zero, we can use it to determine the relative potentials of otha electrodes. The hydrogen electrode, shown in Figure 13.5, serves as the reference for this purpose. 

Redox reactions are usually described by a disguised form of equilibrium constant called the standard cell potential. It is directly proportional to the log of the Kl. A system of hypothetical half-cell potentials is based on the standard hydrogen electrode. The Nernst equation relates cell potentials to the log of activities involved and to the K°q. For a reaction to which a net change in oxidation states n applies. 

Each half-cell reaction has a specific standard potential reported as the potential of the reduction reaction vs. the normal hydrogen electrode (NHE). In an elecdochemical cell, there is a half-cell corresponding to the working electrode (WE), where the reactions under study take place, and a reference half-cell. Experimentally the cell potential is measured as the difference between the potentials of the WE half-cell and the reference electrode/ref-erence half-cell (see Chapter 4). The archetypal reference electrode is the NHE, also known as the standard hydrogen electrode (SHE) and is defined, by convention, as 0.000 V for any temperature. Although the NHE is not typically encountered due to difficulty of operation, all conventional electrodes are in turn referenced to this standard to define their absolute potential (i.e., the Ag/AgCl, 3 M KCl reference has a potential of 203 mV vs. the NHE). In practice, experimental results are either stated as being obtained vs. a specific reference electrode, or converted to potentials vs. NHE. 

These potential values are those of reduction potentials. Hence, these values are those of zero-current cell potentials in which the hydrogen electrode is located on the left and that under study on the right (see Fig. 2.5 in Chap. 2). Moreover, all the species that participate in the half-reduction equilibria are in their standard states their activities are equal to unity. We have already seen that the hydrogen electrode necessarily plays the part of the anode and the electrode under study that of the cathode. Hence, the studied system suffers the electrodic reaction... 

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