Half cell conventions

Also, by convention, potentiometric electrochemical cells are defined such that the indicator electrode is the cathode (right half-cell) and the reference electrode is the anode (left half-cell). 

Potentiometric electrochemical cells are constructed such that one of the half-cells provides a known reference potential, and the potential of the other half-cell indicates the analyte s concentration. By convention, the reference electrode is taken to be the anode thus, the shorthand notation for a potentiometric electrochemical cell is... 

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 convention is adopted of writing all half-cell reactions as reductions M"+ + ne - M... 

The potential of an electrode measured relative to a standard, usually the SHE. It is a measure of the driving force of the electrode reaction and is temperature and activity dependent (p. 230). By convention, the half-cell reaction must be written as a reduction and the potential designated positive if the reduction proceeds spontaneously with respect to the SHE, otherwise it is negative. If the sign of the potential is reversed, it must be referred to as an oxidation potential. 

By convention, in the representation of the cell, the anode is represented on the left and the cathode on the right. The anode is the electrode at which oxidation occurs (AN OX), and the cathode is the electrode at which reduction takes place (RED CAT). The single vertical lines (I) indicate contact between the electrode and solution. The double vertical lines (II) represent the porous partition, or salt bridge, between the two solutions in the two half-cells. The ion concentration or pressures of a gas are enclosed in parentheses. 

Most common reference electrodes are silver-silver chloride (SSC), and saturated calomel electrode (SSC, which contains mercury). The reference electrode should be placed near the working electrode so that the W-potential is accurately referred to the reference electrode. These reference electrodes contain concentrated NaCl or KC1 solution as the inner electrolyte to maintain a constant composition. Errors in electrode potentials are due to the loss of electrolytes or the plugging of the porous junction at the tip of the reference electrode. Most problems in practical voltammetry arise from poor reference electrodes. To work with non-aqueous solvents such as acetonitrile, dimethylsulfoxide, propylene carbonate, etc., the half-cell, Ag (s)/AgC104 (0.1M) in solvent//, is used. There are situations where a conventional reference electrode is not usable, then a silver wire can be used as a pseudo-reference electrode. 

To understand potentiometric methods, those that measure electrical potentials and determine analyte concentrations from these potentials, it is necessary that numerical values for these tendencies be known under conventional standard modes and conditions. What are these modes and conditions First, all halfreactions must be written as either reductions or oxidations. Scientists have decided to write them as reductions. Second, the tendencies for half-reactions to proceed depend on the temperature, the concentrations of the chemical species involved, and, if gases are involved, the pressure in the half-cell. Scientists have defined standard conditions to be a temperature of 25°C, a concentration of exactly 1 M for all dissolved chemical species involved, and a pressure of exactly 1 atm. Third, because every cell consists of two half-cells, it is not possible to measure the value directly. However, if we were to assign the tendency of a certain half-reaction to be zero, then the tendencies of all other half-reactions can be determined relative to this reference half-reaction. 

Electrode potential, E The energy, expressed as a voltage, of a redox couple at equilibrium. E is the potential of the electrode when measured relative to a standard (ultimately the SHE). E depends on temperature, activity and solvent. By convention, the half cell must first be written as a reduction, and the potential is then designated as positive if the reaction proceeds spontaneously with respect to the SHE. Otherwise, E is negative. 

The of this standard cell is +0.76 V. By international convention, the half-cell potential of the hydrogen reduction is assigned a value of exactly OV. Thus, the half-cell potential of the zinc oxidation is equal to K.n (i.e., +0.76 V). This voltage is called the standard half-cell potential and is represented by the symbol 1, to indicate that it was determined against a standard hydrogen electrode. 

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... 

E is the standard redox potential in Volts. By convention, the half cell reactions are always written as reduction reactions. 

If we adopt the usual convention that the half-cell potential for the cell in Eq. (4) is taken to be zero, then we also need to adopt a relative convention in which AGj (H+(aq)) is zero. Consequently, values of AG , on the relative scale are... 

Since in experiments such as the one we have just discussed, it is only possible to determine potential differences between two electrodes (and not the absolute potential of each half cell), it is now useful to choose a reference system to which all measured potential differences may be related. In accord with the IUPAC 1953 Stockholm convention, the standard hydrogen electrode (SHE) is commonly selected as the reference electrode to which we arbitrarily assign a zero value of electrical potential. This is equivalent to assigning (arbitrarily) a standard free energy change, ArG°, of zero at all temperatures to the half reaction ... 

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. 

The electrode potential is defined as the potential difference between the terminals of a cell constructed of the half-cell in question and a standard hydrogen electrode (or its equivalent) and assuming that the terminal of the latter is at zero volts. Note therefore that the electrode potential is an observable physical quantity and is unaffected by the conventions used for writing cells. The statement. . . the electrode potential of zinc is —0.76 volts. . . implies only that a voltmeter placed across the terminals of a cell consisting of standard hydrogen electrode and the zinc electrode would show this value of potential difference, with the zinc terminal negative with respect to that of the hydrogen electrode. An electrode potential is never a metal/solution potential difference , not even on some arbitrary scale. 

It is not appropriate in this chapter to tabulate quantities of electrochemical data since that required may be obtained from texts on electrodeposition.1 2 However, a brief mention of sign conventions must be made since, particularly in the early literature, confusion can arise. Two conventions have been used the European and the American .3 It is sometimes erroneously stated that the conventions differ only in sign however, the real difference lies in the distinction between the potential of an actual electrode and the EMF of a half cell reaction. 

The American convention would assign a positive value to E° for the Zn Zn2+(aq) half cell written as an oxidation, but a negative sign if written as a reduction. It is seen that the European convention refers to the invariant electrostatic potential of the electrode with respect to the SHE, whereas the American convention relates to the thermodynamic Gibbs free energy which is sensitive to the direction of the cell reaction. 

IUPAC recommends that electrode potential be reserved for the European convention, whereas the EMF of a half cell is dealt with via the American convention. 

From experiments of the sort just described, hundreds of half-cell potentials have been determined. A short list is presented in Table 18.1, and a more complete tabulation is given in Appendix D. The following conventions are observed when constructing a table of half-cell potentials ... 

In the past, different sign conventions were used in electrochemistry, which led to difficulty in interpretation of experiments and results. Consequently the electrochemical literature requires an understanding of this problem to avoid confusion. The approach followed in this book is summarised in this section. As pointed out in the previous section, all electrochemical cells are regarded as a combination of two half cells, with each of the latter represented by a half reaction written as a reduction ... 

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. 

A third possible convention is due to Gibbs. Here the emphasis is on the cell itself, rather than the depiction of the cell. The emf of the cell is always positive and the difference between the standard half-cell potentials is taken to yield positive values. Then... 

An electrochemical cell reaction, like any oxidation-reduction reaction, can be written as the sum of an oxidation half-reaction and a reduction half-reaction. In the case of a cell, these half-reactions correspond to the reactions at the two electrodes. Since the cell reaction is the sum of the half-cell reactions, it is convenient to think of dividing the cell potential into half-cell potentials. Unfortunately, there is no way of measuring a half-cell potential—we always need two half-cells to make a cell, the potential of which is measurable. By convention, the half-cell reaction,... 

In the cell Zn(s) Zn2+(ag) I I Cu2+(aq) I Cu(.s) the zinc appears on the left side, indicating that it is being oxidized, not reduced. For this reason, the potential difference contributed by the left half-cell has the opposite sign to its conventional half-cell potential. More generally, we can define the cell potential or cell EMF as... 

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