Half-cells hydrogen

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.)... 

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... 

If the said hydrogen half-cell is part of a cell, the change in the electrical double layer will mean a change in the e.m.f. of the cell. 

A half-cdl consisting of a palladium rod dipping into a 1 M Pd(NOj)2 solution is connected with a standard hydrogen half-cell. The cell voltage is 0.99 volt and the platinum dectrode in the hydrogen half-cell is the anode. Determine E° for the reaction... 

The potentials of the metals in their 1 mol U salt solution are all related to the standard or normal hydrogen electrode (NHE). For the measurement, the hydrogen half-cell is combined with another half-cell to form a galvanic cell. The measured voltage is called the normal potential or standard electrode potential, E° of the metal. If the metals are ranked according to their normal potentials, the resulting order is called the electrochemi-... 

Three kinds of equilibrium potentials are distinguishable. A metal-ion potential exists if a metal and its ions are present in balanced phases, e.g., zinc and zinc ions at the anode of the Daniell element. A redox potential can be found if both phases exchange electrons and the electron exchange is in equilibrium for example, the normal hydrogen half-cell with an electron transfer between hydrogen and protons at the platinum electrode. In the case where a couple of different ions are present, of which only one can cross the phase boundary — a situation which may exist at a semiperme-able membrane — one obtains a so called membrane potential. Well-known examples are the sodium/potassium ion pumps in human cells. 

The electrochemical potential in a potentiometric cell is inevitably measured with respect to a standard electrode. Several types of electrodes are commonly used. The standard hydrogen electrode (SHE) is a hydrogen half-cell in which the cell reaction is as follows ... 

The overall free energy of the redox reaction can be calculated using the standard free energies for the half reactions. As AG° for the hydrogen half cell is zero and AG for the electrons cancels out. 

Non-redox equilibria are expressed in terms of equilibrium constants based on activities, whereas Eh is given in volts. To compare and combine redox equilibria with other non-redox equilibria it is often convenient to use another term, pe. pe is the negative logarithm of electron activity based on the hydrogen half cell in which the redox activity is set at unity. Because pe is expressed as mol L , this term enables redox equilibria and other equilibria to be combined and expressed in terms of a single constant. 

Two half-cells in the Daniell cell are each connected to the standard hydrogen half-cell (Rae D6jur)... 

The tendency of a system to accept or donate electrons is measured using an inert electrode (typically platinum). Electrons can pass from the system into this electrode, which is thus a half-cell. The Pt electrode is connected via a potentiomenter to another half-cell of known potential (usually, a saturated calomel electrode). All potentials are referred to the hydrogen half-cell ... 

Equation (6) defines the difference between two E° values. To set the individual values, we need to choose a particular redox couple as a reference. The reference used most commonly by biochemists is the standard hydrogen half-cell, in which protons at pH 0 are reduced to H2 at a... 

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. 

The reference cell we use for this purpose is the hydrogen half-cell... 

Reaction 5.6 is known as a half-cell reaction to move from left to right, there should be an electron donor. This electron donor is the hydrogen half-cell reaction ... 

H2 —The hydrogen half cell is an arbitrary standard, which all other activities are measured against. 

The hydrogen half-cell is not very convenient for routine laboratory usage—indeed, 1 m H+ (corresponding to a pH of 0 ) and 1 atm H2 (explosive) are dangerous. Hence, secondary standards are used, e.g., mercury/mercurous (calomel) or silver/silver chloride electrodes, which have midpoint redox potentials of 0.244 V and 0.222 V, respectively. 

V, then measure the potential of another half-cell, such as the copper halfcell, with that half-cell. The entire potential of this cell can then be assigned to the other (copper) half-cell, because the potential of the hydrogen half-cell is zero. Of course, we can do the same thing for every other half-cell, but we need not do so, since the hydrogen/hydrogen ion half-cell is hard to work with because it involves a gas. We can get an unknown half-cell potential from its cell potential with any half-cell of known potential. For example, once we get the copper half-cell potential, we can use it to calculate the unknown zinc halfcell potential from the Daniell cell potential. A collection of half-cell potentials, all written as reductions, is presented as Table 17.2. 

Now consider any two half-cells, with electrode reactions 1 and 2, combined in turn with a hydrogen half-cell (reaction 0) thus ... 

FIGURE 17.3 Hydrogen is a gas at room conditions, and electrodes cannot be constructed from it directly. In the hydrogen half-cell shown here, a piece of platinum covered with a fine coating of platinum black is dipped into the solution, and a stream of hydrogen is passed over the surface. 

For example, if a Cu (l m) Cu half-cell is connected to the standard H30 (1 m) H2 half-cell, copper is observed to plate out therefore the copper halfcell is the cathode and the hydrogen half-cell is the anode. The observed cell voltage is 0.34 V thus,... 

When a Zn (f M) Zn half-cell is connected to the standard hydrogen half-cell, zinc dissolves its half-cell is the anode because oxidation occurs in it. The measured cell voltage is 0.76 V, so... 

A galvanic cell is constructed using a standard hydrogen half-cell (with platinum electrode) and a half-cell containing silver and silver chloride ... 

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. 

Here 5 = - 0.5. Hence, the contribution to 5 from an electron in a half cell reaction is the same as the contribution of a gas molecule with the stoichiometric coefficient of 0.5. This leads to the same value of 5 as the combination with the hydrogen half cell. 

Now the redox potential, Eh, is conventionally defined in terms of the potential of a cell composed of two half-cells the half-cell of particular interest (given in general terms by reaction 7.15) and the standard hydrogen half-cell (with Eh and pe assigned values of zero). The reaction in the latter case is... 

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... 

Further standard voltages could be measured and placed into an electrochemical sequence with other pairs of half-cells, as can be found in many chemistry teaching books. The extent of the introduction of the hydrogen half-cell as a standard electrode can be determined in each individual lesson. If one consequently describes all redox reactions in relation to the metal sequence, redox or electrochemical sequence with ions and offers the students model drawings (see Figs. 8.3 and 8.4), then the electron transfer and the redox definition, in terms of involved smallest particles, becomes even clearer and the mixing at the language level of substances and that of particles can be effectively suppressed. 

The electrochemical cell with iron and hydrogen half-cell reactions... 

The accepted primary reference electrode is the hydrogen half cell described in association with Fig.2.1 (Ref 5). It consists of platinum (which serves as an inert conductor) in contact with a solution at 25 °C, saturated with hydrogen gas at one atmosphere pressure, and containing hydrogen ions at pH = 0 (aH+ = 1). In practice, the major use of the standard hydrogen electrode (SHE) is for calibration of secondary reference electrodes, which are more convenient to use. Two common reference electrodes are the calomel or mercury/saturated-mercurous-chloride half cell with a potential of +241 mV relative to the SHE and the sil-ver/saturated-silver-chloride half cell with a relative potential of+196 mV. Both of these electrodes are saturated with potassium chloride to maintain a constant chloride and hence metal-ion concentration. 

---