Standard electrode potentials determination

Chemistry Video Consortium, Practical Laboratory Chemistry, Educational Media Film and Video Ltd, Harrow, Essex, UK - Electrochemical techniques (using galvanic cells, using conductometric cells, determining standard electrode potentials, determining solubility products, thermodynamic characteristics of cells, conductometric titrations and using an automatic titrator). 

The equilibrium potentials and E, can be calculated from the standard electrode potentials of the H /Hj and M/M " " equilibria taking into account the pH and although the pH may be determined an arbitrary value must be used for the activity of metal ions, and 0 1 = 1 is not unreasonable when the metal is corroding actively, since it is the activity in the diffusion layer rather than that in the bulk solution that is significant. From these data it is possible to construct an Evans diagram for the corrosion of a single metal in an acid solution, and a similar approach may be adopted when dissolved O2 or another oxidant is the cathode reactant. 

It should be noted that the simple Nernst equation cannot be used since the standard electrode potential is markedly temperature dependent. By means of irreversible thermodynamics equations have been computed to calculate these potentials and are in good agreement with experimentally determined results. 

Equations 20.176 and 20.179 emphasise the essentially thermodynamic nature of the standard equilibrium e.m.f. of a cell or the standard equilibrium potential of a half-reaction E, which may be evaluated directly from e.m.f. meeisurements of a reversible cell or indirectly from AG , which in turn must be evaluated from the enthalpy of the reaction and the entropies of the species involved (see equation 20.147). Thus for the equilibrium Cu -)-2e Cu, the standard electrode potential u2+/cu> hence can be determined by an e.m.f. method by harnessing the reaction... 

It must be emphasised that standard electrode potential values relate to an equilibrium condition between the metal electrode and the solution. Potentials determined under, or calculated for, such conditions are often referred to as reversible electrode potentials , and it must be remembered that the Nernst equation is only strictly applicable under such conditions. 

During the determination of standard electrode potentials an electrochemical equilibrium must always exist at the phase boundaries, e.g. that of the elec-trode/electrolyte. From a macroscopic viewpoint no external current flows and no reaction takes place. From a microscopic viewpoint or a molecular scale, a continuous exchange of charges occurs at the phase boundaries. In this context Fig. 6 demonstrates this fact at the anode of the Daniell element. 

In the example just studied, the electrolysis of HC1 solution, the ions that transport the current (H+ and Cl-) are also the ones that are discharged at the electrodes. In other cases, however, the main ionic transporters of current may not be of the same species as the ions that are discharged. An excellent example is the electrolysis of CuS04 solution between platinum electrodes. A one molal CuS04 solution is quite acid so that the positive current transporters are both Cu2+ and H+ ions. The main negative transporter is the S04 ions. The solution contains, however, a small concentration of OH- ions. In order to determine which ions will be discharged at the electrodes, it is necessary to consider standard electrode potentials of the concerned species ... 

The relationships of the type (3.1.54) and (3.1.57) imply that the standard electrode potentials can be derived directly from the thermodynamic data (and vice versa). The values of the standard chemical potentials are identified with the values of the standard Gibbs energies of formation, tabulated, for example, by the US National Bureau of Standards. On the other hand, the experimental approach to the determination of standard electrode potentials is based on the cells of the type (3.1.41) whose EMFs are extrapolated to zero ionic strength. 

Several descriptions of electrode reaction rates discussed on the preceding pages and the difficulty to standardize electrode potential scales with respect to different temperatures imply several definitions of activation energies of electrode reactions. The easiest way to determine this quantity, for example, for an irreversible cathodic process, employs Eqs (5.2.9), (5.2.10) and (5.2.12) at a constant electrode potential,... 

QB For this cell because the electrodes are identical, the standard electrode potentials are numerically equal and subtracting one from the other leads to the value c°dl = 0.000 V. However, because the ion concentrations differ, there is a potential difference between the two half cells (non-zero nonstandard voltage for the cell). [Pb2+] = 0.100 M in the cathode compartment. The anode compartment contains a saturated solution of Pbl2. We use the Nemst equation (with n = 2) to determine [Pb2+] in the saturated solution. 

Then, knowing F H2, it was relatively easy to determine values of electrode potentials for any other couple. With this methodology, they devised the standard electrode potentials ° scale (often called the E nought scale , or the hydrogen scale ). 

Worked Example 7.18 Determine a value for the standard electrode potential + Ag with the data below. Assume that y = 1 throughout. 

For an irreversible reduction the half-wave potential is determined not only by the standard electrode potential but also by the polarographic overvoltage. For a simple electrode process the metal ion-solvent interaction is mainly responsible for the polarographic overvoltage and hence E[ j of such irreversible reductions may also be considered as a function of the solvation 119f... 

It is a good idea, when using such simple Nernst plots as an analytical method of determining an activity, to check that the intercept at jc = 0 is indeed the standard electrode potential. (There are many compendia listed in the Bibliography at the end of this book that cite large numbers of values, as does Appendix 3.)... 

The amount of zinc in the soil was determined by immersing a rod of clean, pure zinc in the solution of zinc sulfate plus sulfuric acicf and measuring Ez 2+ z,i as —0.864 V. What is the concentration of the zinc solution formed by digesting in acid and what is the activity of the zinc salt In order to answer there questions, we will need to know the standard electrode potential, n —0.760 V. 

Variations in ionic strength are such an important concern that it is recommended for solutes to be analysed by a potentiometric procedure only if the ionic strength is known and controlled. Furthermore, calibration steps, i.e. to determine the standard electrode potential E should also be performed in a solution of the same, known, ionic strength, e.g. in a solution of perchloric acid of — 1.0 mol dm K Provided that 1 is always much higher than the concentration of the analyte, the latter does not contribute more than a tiny fraction of the overall ionic strength and so fluctuations in the activity coefficient y can be safely ignored. 

Having revised a few basic electrochemical ideas, such as the nature of reference electrodes, the standard hydrogen electrode and the scale based on it, we next looked briefly at thermodynamic parameters such as the electrode potential E, the standard electrode potential f and emf, and then discussed how AG, AH and AS (where the prime indicates a frustrated cell equilibrium ) may be determined. 

Figure 4.2 Plot of the variable (emf — she) against the volume of ceric ion soiution during a potentiometric determination of [Fe ]. The end point is clearly shown by a sharp transition from the standard electrode potential of the analyte couple, p 2+. to that of the titrant couple, 3 (cf. equation (4.1)).

As it has been shown that the Gibbs function for formation of an individual ion has no operational meanings [12], no way exists to determine such a quantity experimentally. However, for the purposes of tabulation and calculation, it is possible to separate AfGm of an electrolyte arbitrarily into two or more parts, which correspond to the number of ions formed, in a way analogous to that used in tables of standard electrode potentials. In both cases, the standard Gibbs function for formation of aqueous H" " is defined to be zero at every temperature ... 

Let us determine whether we can use the displacement deposition technique to deposit Sn on a Cu substrate. The simplest way to determine this is to use the principle presented in Figures 5.10 and 9.1. According to this principle, Sn cannot be deposited by displacement on a Cu substrate since the standard electrode potential of a Cu /Cu couple is more positive than that of an Sn +/Sn couple ... 

Americium and californium have been prepared by the reduction, using noble metals and hydrogen at temperatures greater than 1110 °C, of the oxides MOj 5. A new determination has been made of the heat of dissolution of Am in aqueous HCl, and the standard enthalpies of a series of Am compounds and ions have been reported (Table 1). The standard electrode potential of the Am -Am" couple was +2.06 + 0.01 V, making the metal only slightly more electropositive than Pu. 

Determination of the Standard Electrode Potential ( from Electrochemical Measurements... 

R. Parsons, The Single Electrode Potential Its Significance and Calculation and Standard Electrode Potentials Units, Conventions and Methods of Determination, in Standard Potentials in Aqueous Solution, A. J. Bard, R. Parsons, and J. Jordan, eds. Chs. 1 and 2, Marcel Dekker, New York (1985). 

A. Ardvalo and G. Pastor, Verification of the Nernst Equation and Determination of a Standard Electrode Potential, J. Chem. Ed. 1985,62, 882. 

Marcus has introduced a model for, S N 2 reactions of the ET type based on two interacting states which takes into account the relevant bond energies, standard electrode potentials, solvent contributions, and steric effects.87 The rate constant for intramolecular electron transfer between reduced and oxidized hydrazine units in the radical cation of the tetraazahexacyclotetradecane derivative (43) and its analogues has been determined by simulation of then variable temperature ESR spectra.88 The same researchers also reported then studies of the SET processes of other polycyclic dihydrazine systems.89,90... 

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. 

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