Lead, standard electrode potential

Metals having an electronegative potential show a tendency to oxidise and thus to corrode in aqueous media, if the conditions allow. This tendency increases as the potential becomes more electronegative. It is well known that magnesium (standard electrode potential, —2380 mV) degrades much more under the effect of moisture than lead (standard electrode potential, — 126 mV). 

To calculate the open circuit voltage of the lead—acid battery, an accurate value for the standard cell potential, which is consistent with the activity coefficients of sulfuric acid, must also be known. The standard cell potential for the double sulfate reaction is 2.048 V at 25 °C. This value is calculated from the standard electrode potentials for the (Pt)H2 H2S04(yw) PbS04 Pb02(Pt) electrode 1.690 V (14), for the Pb(Hg) PbS04 H2S04(yw) H2(Pt) electrode 0.3526 V (19), and for the Pb Pb2+ Pb(Hg) 0.0057 V (21). 

When metals are arranged in the order of their standard electrode potentials, the so-called electrochemical series of the metals is obtained. The greater the negative value of the potential, the greater is the tendency of the metal to pass into the ionic state. A metal will normally displace any other metal below it in the series from solutions of its salts. Thus magnesium, aluminium, zinc, or iron will displace copper from solutions of its salts lead will displace copper, mercury, or silver copper will displace silver. 

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. 

Consider a cell made up of two half cells, where one contains the Fe , Fe couple and the other the Cu, Cu couple. By looking up the respective values of the standard electrode potentials E given in Appendix 3, deduce the spontaneous cell reaction that would occur if the leads connecting the two half cells were allowed to short by touching. 

Electrons are transferred singly to any species in solution and not in pairs. Organic electrochemical reactions therefore involve radical intermediates. Electron transfer between the electrode and a n-system, leads to the formation of a radical-ion. Arenes, for example are oxidised to a radical-cation and reduced to a radical-anion and in both of these intermediates the free electron is delocalised along the 7t system. Under some conditions, where the intermediate has sufficient lifetime, these electron transfer steps are reversible and a standard electrode potential for the process can be measured. The final products from an electrochemical reaction result from a cascade of chemical and electron transfer steps. 

Identify conditions that lead to corrosion and ways to prevent it. Calculate cell voltage from a table of standard electrode potentials. 

We will use standard electrode potentials throughout the rest of this text to calculate cell potentials and equilibrium constants for redox reactions as well as to calculate data for redox titration curves. You should be aware that such calculations sometimes lead to results that are significantly different from those you would obtain in the laboratory. There are two main sources of these differences (1) the necessity of using concentrations in place of activities in the Nernst equation and (2) failure to take into account other equilibria such as dissociation, association, complex formation, and solvolysis. Measurement of electrode potentials can allow us to investigate these equilibria and determine their equilibrium constants, however. 

The application of standard electrode potential data to many systems of interest in analytical chemistry is further complicated by association, dissociation, complex formation, and solvolysis equilibria involving the species that appear in the Nemst equation. These phenomena can be taken into account only if their existence is known and appropriate equilibrium constants are available. More often than not, neither of these requirements is met and significant discrepancies arise as a consequence. For example, the presence of 1 M hydrochloric acid in the iron(Il)/iron(llI) mixture we have just discussed leads to a measured potential of + 0.70 V in 1 M sulfuric acid, a potential of -I- 0.68 V is observed and in 2 M phosphoric acid, the potential is + 0.46 V. In each of these cases, the iron(II)/iron(III) activity ratio is larger because the complexes of iron(III) with chloride, sulfate, and phosphate ions are more stable than those of iron(II) thus, the ratio of the species concentrations, [Fe ]/[Fe ], in the Nemst equation is greater than unity and the measured potential is less than the standard potential. If fomnation constants for these complexes were available, it would be possible to make appropriate corrections. Unfortunately, such data are often not available, or, if they are, they are not very reliable. 

When the follow-up reaction leads to an electroactive product, the effects of the redox properties of the product (Z, above) on the current-potential curves also must be considered. For more than one electroactive product, the nomenclature for the model system changes from and to A and B, below, each compound having its own E (standard electrode potential) value ... 

Thermodynamic cycles involving standard electrode potentials obtained by cyclic voltammetry have also been used to provide thermochemical information on organometallic compounds. This so-called electrochemical method leads to Gibbs energies of reaction in solution, from which bond dissociation enthalpies may be derived using a number of auxiliary data that are often estimated. For example, the derivation of a metal-hydrogen bond dissociation enthalpy in an L MH species requires (i) an estimate of the reduction potential of in the same solvent where the experiments were carried out (ii) an estimate of the solvation entropies of L MH, L M, and H and (iii) the knowledge of the pK of... 

The standard electrode potential, Epb2+/pb = -0.126 V , shows that lead is thermodynamically unstable in acid solutions but stable in neutral solutions. The exchange current for the hydrogen evolution reaction on lead is very small (-10" - I0 " Acm ), but control of corrosion is usually due to mechanical passivation of the local anodes of the corrosion cells as the majority of lead salts are insoluble and frequently form protective films or coatings. 

The reduction of rare earth metal trihalides, RX3, is in principle possible with all kinds of reducing agents as long as they have standard electrode potentials E° that can overcome that of the respective potentials of E° (R + R +). This is discussed below in more detail. Therefore, the classical reducing agents, nonmetals such as hydrogen or carbon, or like metals (comproportionation route) and unlike metals (metallothermic reduction) are all possible but (may) lead to different products. Cathodic reduction of appropriate melts is also an option. 

The extremely small concentration of free cations leads to a strong negative shift of the standard electrode potential as will be demonstrated for the Au/CN system. The concentration of free Au" ions [Au is given by the dissociation equilibrium and may be introduced into the Nemst equation... 

Keller, Burnett, Carlson, and Nestor [22] have made detailed predictions of the chemical properties of elements 113 and 114. The group IV elements show increasing stability in oxidation state ii relative to the iv state as one goes to higher atomic numbers. Carbon and silicon have very stable tetrapositive oxidation states, and germanium shows a very unstable dipositive oxidation state in addition to a stable tetrapositive state. In tin, both the dipositive and tetrapositive oxidation states are important, and lead is most stable in oxidation state ii. Thus, element 114 should be weakly, if at all, tetrapositive and the most stable oxidation state is expected to be the ii state. The standard electrode potential for the reaction... 

For reasons similar to thdse that lead to the expectation of a stable ii oxidation state in element 114, element 113-a member of group III of the periodic system-is expected to have a preferred oxidation state of i a standard electrode potential of + 0.6 V is predicted. This means that eka-thallium is expected to be more noble than thallium, being more like silver in this respect. In fact, element 113 should have a chemistry somewhere between the chemistries of Tl and Ag". The ion 113 is expected to be much more easily complexed in solution than is the case for T1 . For example, the solubility of TlCl in water is not increased much by adding excess HCl [24] or NH3, whereas AgCl dissolves. Element 113 is expected to be more like silver in this respect. We thus see that the relativistic effects on the 7pi 12 electrons cause a diagonal relationship to be introduced into the periodic table around element 114, causing 114 to be somewhere between Hg or Cd and Pb " in its chemistry and 113 to act more like Ag than Tl" in its chemistry. 

Knowledge of the standard electrode potential of a cell allows us to establish, thermodynamically, the direction of the corresponding cell reaction. Suppose that analysis of the cell using the protocol in Section 1.13 leads to the following formal cell reaction. 

Since electrode potentials are quoted on the hydrogen scale, it might be expected that only metals with positive standard electrode potentials could be deposited from acid solutions, and that hydrogen would be discharged in preference to other cations, such as lead (E = -0.126 V) or nickel (E =-0.25 V). The failure of this prediction is due to the high activation overpotential (q.v.) of hydrogen on many metals, values obtained at a current density of 1 mA cm varying from 0.01 V on a platinised platinum cathode to as much as 0.67 V on lead and 1.04 V on mercury. 

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