Standard reduction potential equilibrium constant

Preparation and chemistry of chromium compounds can be found ia several standard reference books and advanced texts (7,11,12,14). Standard reduction potentials for select chromium species are given ia Table 2 whereas Table 3 is a summary of hydrolysis, complex formation, or other equilibrium constants for oxidation states II, III, and VI. 

In addition to defined standard conditions and a reference potential, tabulated half-reactions have a defined reference direction. As the double arrow in the previous equation indicates, E ° values for half-reactions refer to electrode equilibria. Just as the value of an equilibrium constant depends on the direction in which the equilibrium reaction is written, the values of S ° depend on whether electrons are reactants or products. For half-reactions, the conventional reference direction is reduction, with electrons always appearing as reactants. Thus, each tabulated E ° value for a half-reaction is a standard reduction potential. 

Equations and provide a method for calculating equilibrium constants from tables of standard reduction potentials. Example illustrates the technique. 

This is a quantitative calculation, so it is appropriate to use the seven-step problem-solving strategy. We are asked to determine an equilibrium constant from standard reduction potentials. Visualizing the problem involves breaking the redox reaction into its two half-reactions ... 

Given the above standard reduction potentials, estimate the approximate value of the equilibrium constant for the following reaction ... 

These conventions are arbitrary, but you must use them if you wish to use tabulated values of equilibrium constants, standard reduction potentials, and free energies. 

Use the standard reduction potentials in Table 18.1 to calculate the equilibrium constant at 25°C for the reaction... 

Calculate E° for the reaction from standard reduction potentials, as in Worked Example 18.5. Then use the equation log K = nE°/0.0592 V to determine the equilibrium constant. 

Ethanol Production in Yeast When grown anaerobically on glucose, yeast (S. cerevisiae) converts pyruvate to acetaldehyde, then reduces acetaldehyde to ethanol using electrons from NADH. Write the equation for the second reaction, and calculate its equilibrium constant at 25 °C, given the standard reduction potentials in Table 13-7. 

The third largest class of enzymes is the oxidoreductases, which transfer electrons. Oxidoreductase reactions are different from other reactions in that they can be divided into two or more half reactions. Usually there are only two half reactions, but the methane monooxygenase reaction can be divided into three "half reactions." Each chemical half reaction makes an independent contribution to the equilibrium constant E for a chemical redox reaction. For chemical reactions the standard reduction potentials ° can be determined for half reactions by using electrochemical cells, and these measurements have provided most of the information on standard chemical thermodynamic properties of ions. This research has been restricted to rather simple reactions for which electrode reactions are reversible on platinized platinum or other metal electrodes. 

The standard reduction potentials of the most relevant half-reactions involved in the four- and two-electron reduction of dioxygen in acid and alkaline aqueous media are listed in Table 3.1. It follows from these values, that, under full thermodynamic control, the equilibrium concentration of peroxide at the reversible potential for the four-electron reduction of oxygen in acid media, that is, 1.23 V, is of the order of 10 18 M. Hence, a stepwise reduction of dioxygen to yield currents of about 1A cm 2 will require values for the standard heterogeneous rate constants for... 

However, based strictly on the standard reduction potentials of the two redox couples involved, the equilibrium constant of the reaction as written would be exceedingly small. To further complicate matters, the rates of disappearance of 02 in the presence of Fe(II)TMPyP, as monitored by RRDE techniques, and the rates of disappearance of 02 in the presence of Fe(III)TMPyP, as measured by spectrophotometric methods [52], yielded very similar values. The most likely resolution of this seeming quandary invokes formation of a macrocycle-dioxygen adduct as a short-lived, albeit yet to be detected, intermediate. 

The thermodynamic and kinetic data are summarized in Table 4. The equilibrium constants as listed in Table 4 are calculated from log K = A °/0.059 and A °= — (°M) + (°0i) and involve the standard reduction potentials of the melal couple and of oxygen at [02] = 1 M. This standard state is more suitable for kinetic calculations than the more widely used convention p0l = 1 atm. The calculation of thermodynamic equilibrium constants for the surface complexes if more difficult (Sposito, 1983) and requires further work. Figure 8 displays the resulting LFER. A theoretical line of slope one (its the data over a broad range of 13 log k units. This analysis supports an outer-sphere mechanism for the... 

Use the standard reduction potentials to find the equilibrium constant for each of the following reactions at 25°C ... 

Further we looked at galvanic cells where it was possible to extract electrical energy from chemical reactions. We looked into cell potentials and standard reduction potentials which are both central and necessary for the electrochemical calculations. We also looked at concentration dependence of cell potentials and introduced the Nemst-equation stating the combination of the reaction fraction and cell potentials. The use of the Nemst equation was presented through examples where er also saw how the equation may be used to determine equilibrium constants. 

First, calculate the standard emf of the cell from the standard reduction potentials in Table 19.1 of the text. Then, calculate the equilibrium constant from the standard emf using Equation (19.5) of the text. 

The reactions described above all involve the displacement of a carbanion from the cobalt atom of methylcobalamin. These reactions occur under aerobic conditions with rate constants in the order of milliseconds. It is apparent that metals which react by electrophilic attack on the Co-C bond (SE2 mechanism) occur with the more oxidized state of the metal, i.e., Pb, Tl, Hg, Pd, which have standard reduction potentials greater than 0.8 volts. Because of the "base on - base off" equilibrium these reactions are pH dependent. 

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