Standard half-cell potentials, equilibrium

The equilibrium constant for the chemical reaction expressed by Equation 2.15 is related to the difference of the standard half-cell potentials by the relation ... 

Standard half-cell potentials can be used to compute standard cell potentials, standard Gibbs free energy changes, and equilibrium constants for oxidation-reduction reactions. 

The quantity AG° is the free energy of the half-reaction when the activities of the reactant and product have values of unity and is directly proportional to the standard half-cell potential for the reaction as written. It also is a measure of the equilibrium constant for the half-reaction assuming the activity of electrons is unity ... 

This equation allows one to compute the chemical equilibrium constant from measured standard-state electrochemical cell potentials (usually referred to as standard cell potentials). Some standard half-cell potentials are given in Table 14.6-1. The standard potential of an electrochemical cell is obtained by combining the two relevant half-cell potentials. 

Calculation of the Equilibrium Constant from Standard Half-Cell Potentials... 

Using Eq. (17.50), we can calculate the equilibrium constant for any reaction from the standard cell potential which, in turn, can be obtained from the tabulated values of the standard half-cell potentials. The following method and examples illustrate a procedure that will ensure obtaining the with both a correct sign and magnitude. 

Since the values of equilibrium constants are obtained from the standard half-cell potentials, the method of obtaining the S° of a half-cell has great importance. Suppose we wish to determine the of the silver-silver ion electrode. Then we set up a cell that includes this electrode and another electrode the potential of which is known for simplicity we choose the SHE as the other electrode. Then the cell is... 

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. 

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. 

Equations (26.20) and (26.27) are combined when two hydrogen electrodes are connected in an electrochemical cell. The left electrode is the standard hydrogen electrode with a half-cell potential of 0.0 V. The right electrode is a hydrogen electrode immersed in a solution at a particular pH (where an+ 1). The resulting equilibrium cell potential, in terms of the right electrode compartment pH, 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 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. 

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... 

E Mm+ is called the single electrode or half-cell potential of the M,Mm+ electrode on the standard hydrogen scale. It should be recalled that in this text, E denotes the potential in the general case, E the potential at equilibrium, and E° the potential at equilibrium under standard... 

Erhe andELHE are equilibrium half-cell potentials, or electrode potentials, which depend in sign on the definitions of positive and negative electricity and assignment of Eu + = 0 at standard condi-... 

By measuring the cell potential at various concentrations of Cu ", we can determine < cu2+/cu = 0cu2+/cu- This standard potential is tabulated along with the standard potentials of other half-cells in Table 17.1. Such a table of half-cell potentials, or electrode potentials, is equivalent to a table of standard Gibbs energies f rom which we can calculate values of equilibrium constants for chemical reactions in solution. Note that the standard potential is the potential of the electrode when all of the reactive species are present with unit activity, a = 1. 

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. 

Redox reactions are usually described by a disguised form of equilibrium constant called the standard cell potential. It is directly proportional to the log of the Kl. A system of hypothetical half-cell potentials is based on the standard hydrogen electrode. The Nernst equation relates cell potentials to the log of activities involved and to the K°q. For a reaction to which a net change in oxidation states n applies. 

Appendix A Common Mathematical Operations in Chemistry A-1 Appendix B Standard Thermodynamic Values for Selected Substances A-5 Appendix C Equilibrium Constants for Selected Substances A-8 Appendix D Standard Electrode (Half-Cell) Potentials A-14 Appendix E Answers to Selected Problems A-15 Glossary G-1 Credits C-1 Index 1-1... 

In Chapter 8 (pages 119-123) we saw that the position of an equilibrium reaction is affected by changes in concentration, temperature and pressure. Redox equilibria are no different. When we compare the voltage of a standard half-cell, X, with a standard hydrogen electrode, we are measuring E for the half-cell X. If we change the concentration or temperature of half-cell X, the electrode potential also changes. Under these nonstandard conditions we use the symbol E for the electrode potential. 

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