Standard equilibrium half-cell reactions

The cathodic reaction has been generalized in the form Xx+ + xc —> X. Representative specific cathodic reactions are classified in Table 1.2 along with the standard equilibrium half-cell potentials, E°, relative to the standard hydrogen electrode (SHE), where E° H+ = 0. The variables that must be set to correct the standard poterftials. E°, to values... 

Examples of cathodic reactions Standard equilibrium half-cell potentials(a), E° (mV vs. SHE) Variables required for correction of E° to E ... 

A pure metal can be anodic only if its equilibrium half-cell potential, E M, is less than the half-cell potential of some cathodic reaction, E x, such that the total cell potential (Ex - E" ) causes current to flow as in Fig. 1.6, that is, current away from the anode area as ions in the solution. A few representative anodic reactions are listed in Table 1.3 along with their standard equilibrium half-cell potentials. 

Hydrogen gas is bubbled over a platinum surface that is coated with platinum black, an electrolytically deposited coating of colloidal platinum, which is an excellent catalyst for the above equilibrium. The hydrogen electrode has been selected as the standard against which the potentials of other electrodes are measured. Equations of the type of reaction (I) are called half-cell reactions, because they include electrons. Reaction I is a reduction half-cell reaction. 

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. 

Because the half-cell reactions are equilibrium reactions, the standard potential of each is the same, yet opposite in sign, as that of the corresponding oxidation i.e., from Table 2.4, the potential for the first reaction is —(—0.44) V, whereas for the second reaction, E° is 0 V. 

Reversible cell. One in which the half-cell reactions are reversed by reversing the current flow such a cell is said to be in thermodynamic equilibrium. Standard Hydrogen Electrode SHE). This consists of a platinum electrode coated with platinum black to catalyse the electrode reaction and over the... 

The reduction potential is an electrochemical concept. Consider a substance that can exist in an oxidized form X and a reduced form X . Such a pair is called a redox couple. The reduction potential of this couple can be determined by measuring the electromotive force generated by a sample half-cell connected to a standard reference half-cell (Figure 18.6). The sample half-cell consists of an electrode immersed in a solution of 1 M oxidant (X) and 1 M reductant (X ). The standard reference half-cell consists of an electrode immersed in a 1 M H+ solution that is in equilibrium with H2 gas at 1 atmosphere pressure. The electrodes are connected to a voltmeter, and an agar bridge establishes electrical continuity between the half-cells. Electrons then flow from one half-cell to the other. If the reaction proceeds in the direction... 

The foregoing example illustrates how equilibrium constants for overall cell reactions can be determined electrochemically. Although the example dealt with redox equilibrium, related procedures can be used to measure the solubility product constants of sparingly soluble ionic compounds or the ionization constants of weak acids and bases. Suppose that the solubility product constant of AgCl is to be determined by means of an electrochemical cell. One half-cell contains solid AgCl and Ag metal in equilibrium with a known concentration of CP (aq) (established with 0.00100 M NaCl, for example) so that an unknown but definite concentration of Kg aq) is present. A silver electrode is used so that the half-cell reaction involved is either the reduction of Ag (aq) or the oxidation of Ag. This is, in effect, an Ag" Ag half-cell whose potential is to be determined. The second half-cell can be any whose potential is accurately known, and its choice is a matter of convenience. In the following example, the second half-cell is a standard H30" H2 half-cell. 

Each electrode reaction, anode and cathode, or half-cell reaction has an associated energy level or electrical potential (volts) associated with it. Values of the standard equilibrium electrode reduction potentials E° at unit activity and 25°C may be obtained from the literature (de Bethune and Swendeman Loud, Encyclopedia of Electrochemistry, Van Nostrand Reinhold, 1964). The overall electrochemical cell equilibrium potential either can be obtained from AG values or is equal to the cathode half-cell potential minus the anode half-cell potential, as shown above. 

F is the Faraday constant, K is the equilibrium constant of the reaction, R is the gas constant, and T is the thermodynamic temperature. However, E jj is not the standard potential of the electrode reaction (or sometimes called half-cell reaction), which is tabulated in the tables mentioned. It is the standard potential of the reaction in a chemical cell which is equal to the standard potential of an electrode reaction (abbreviated as standard electrode potential), E when the reaction involves the oxidation of molecular hydrogen to solvated protons... 

As is standard practice, the above reactions are written as equilibrium reduction reactions and the total chemical change in an electrochemical cell is found by adding the individual anode and cathode reactions (the so-called half-cell reactions). Thus, the overall reaction for the Cu/Zn cell in Figure 26.1 is... 

We have seen in Section 26.2.1 that thermodynamics (i.e., equilibrium half-cell potentials) can be used to determine which of two half-cell reactions proceeds spontaneously in the anodic or cathodic direction when the two reactions occur on the same piece of metal or on two metal samples that are in electrical contact with one another. The half-cell reaction with the higher equilibrium potential will always be at the cathode. Thus, under standard conditions any metal dissolution (corrosion) reaction with an E° less than 0.0 V vs. SHE will be driven by proton reduction while metal dissolntion reactions with an E° less than -e1.23 V vs. SHE will be driven by dissolved... 

The initial consideration in analyzing an existing or proposed metal/environment combination for possible corrosion is determination of the stability of the system. According to Eq 1.18, the criterion is whether the equilibrium half-cell potential for an assumed cathodic reaction, E x, is greater than the equilibrium half-cell potential for the anodic reaction, E M. A convenient representation of relative positions of equilibrium half-cell potentials of several common metals and selected possible corrodent species is given in Fig. 1.7. To the left is the scale of potentials in millivolts relative to the standard hydrogen electrode (SHE). The solid vertical lines identified by the name of the metal give... 

Tabulated values of standard reduction potentials, E°, refer to single electrodes. For example, for the half-cell reaction 7.6, the value of °cT+/cu = +0.34 V. However, it is impossible to measure the potential of an individual electrode and the universal practice is to express all such potentials relative to that of the standard hydrogen electrode. The latter consists of a platinum wire immersed in a solution of ions at a concentration of Imoldm (strictly, unit activity) in equilibrium with H2 at 1 bar pressure (equation 7.12). This electrode is taken to have a standard reduction potential E° = QY at all temperatures. 

The splitting of redox reactions into two half cell reactions by introducing the symbol" e , which according to Eq.(II.28) is related to the standard electrode potential, is arbitrary, but useful (this e notation does not in any way refer to solvated electrons). When calculating the equihbrium composition of a chemical system, both e , and can be chosen as components and they can be treated numerically in a similar way equilibrium constants, mass balance, etc. may be defined for both. However, while represents the hydrated proton in aqueous solution, the above equations use only the activity of e , and never the concentration of e . Concentration to activity conversions (or activity coefficients) are never needed for the electron cf. Appendix B, Example B.3). 

The oxidation-reduction potentials for half reactions such as Fe" —> Fe + e are measured by putting a piece of platinum or other inert metal into a solution containing ferrous and ferric ions in standard concentrations, and combining this half cell with the standard hydrogen half cell. Again, the platinum serves to conduct electrons and to catalyze the equilibrium between ferrous and ferric ions. The electromotive force of the cell... 

The potential difference is closely related to the difference of the electrochemical potential based on the electrochemical affinity. If we could measure A(p directly, we could organize the table of electromotive forces based on the Galvani potential difference. However, A

reference electrode to measure the half cell potential at an electrode. When a certain electrcxle is coupled with a reference electrode, then the electromotive force can be measured. Since we usually use some reference electrodes as standards, the electromotive force is defined as the equilibrium potential of the reaction. The table was made in such a way and the hydrogen reference electrode was used to measure and calculate potentials for the half cell reactions. 

The calomel reference electrode has long been a standard reference electrode used in the laboratory. It consists of mercury in equilibrium with Hg2", the activity of which is determined by the solubility of Hg2Cl2 (mercurous chloride, or calomel). The half-cell reaction is... 

Also known as the standard hydrogen electrode (SHE), it is a redox reference electrode which forms the basis of the thermodynamic scale of oxidation-reduction potentials. The potential of the NHE is defined as zero and based oti equilibrium of the following redox half-cell reaction, typically on a Pt surface 2H+(aq) + 2e H2(g). The activities of both the reduced form and the oxidized form are maintained at unity. That implies that the pressure of hydrogen gas is 1 atm and the concentration of hydrogen ions in the solution is 1 M. 

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... 

The most widely used reference electrode, due to its ease of preparation and constancy of potential, is the calomel electrode. A calomel half-cell is one in which mercury and calomel [mercury(I) chloride] are covered with potassium chloride solution of definite concentration this may be 0.1 M, 1M, or saturated. These electrodes are referred to as the decimolar, the molar and the saturated calomel electrode (S.C.E.) and have the potentials, relative to the standard hydrogen electrode at 25 °C, of 0.3358,0.2824 and 0.2444 volt. Of these electrodes the S.C.E. is most commonly used, largely because of the suppressive effect of saturated potassium chloride solution on liquid junction potentials. However, this electrode suffers from the drawback that its potential varies rapidly with alteration in temperature owing to changes in the solubility of potassium chloride, and restoration of a stable potential may be slow owing to the disturbance of the calomel-potassium chloride equilibrium. The potentials of the decimolar and molar electrodes are less affected by change in temperature and are to be preferred in cases where accurate values of electrode potentials are required. The electrode reaction is... 

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. 

E is the standard equilibrium potential, i. e. the potential corresponding to unit activity and RTF. The dissolution reaction leads to the development of an electrical double layer at the iron-solution interface. The potential difference of the Fe/Fe " half cell cannot be measured directly, but if the iron electrode is coupled with a reference electrode (usually the standard hydrogen electrode, SHE), a relative potential difference, E, can be measured. This potential is termed the single potential of the Fe/Fe electrode on the scale of the standard hydrogen couple H2/H, the standard potential of which is taken as zero. The value of the equilibrium potential of an electrochemical cell depends upon the concentrations of the species involved. 

The left half-cell, connected to the negative terminal of the potentiometer, is called the standard hydrogen electrode (S.H.E.). It consists of a catalytic Pt surface in contact with an acidic solution in which JAH, = 1. A stream of H2(g) bubbled through the electrode saturates the solution with H2(aq). The activity of H2(g) is unity if the pressure of H2(g) is 1 bar. The reaction that comes to equilibrium at the surface of the Pt electrode is... 

All species are aqueous unless otherwise indicated. The reference state for amalgams is an infinitely dilute solution of the element in Hg. The temperature coefficient, dE°/dT, allows us to calculate the standard potential, E°(T), at temperature T E°(T) — Ec + (dE°/dT)AT. where A T is T — 298.15 K. Note the units mVIK for dE°ldT. Once you know E° for a net cell reaction at temperature T, you can find the equilibrium constant, K, for the reaction from the formula K — lOnFE°,RTln w, where n is the number of electrons in each half-reaction, F is the Faraday constant, and R is the gas constant. 

Because we can calculate E° from standard potentials, we can now also calculate equilibrium constants for any reaction that can be expressed in terms of two half-reactions. Toolbox 12.2 summarizes the steps involved, and Example 12.7 shows the steps in action. Equation 6 also shows that the magnitude of E° for a cell reaction is an indication of the equilibrium composition. It follows from the equation that a reaction with a large positive E° has a very large K. A reaction with a large negative E° has a K much less than 1. 

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