Oxidation-reduction reactions equilibrium constants

Several types of reactions are commonly used in analytical procedures, either in preparing samples for analysis or during the analysis itself. The most important of these are precipitation reactions, acid-base reactions, complexation reactions, and oxidation-reduction reactions. In this section we review these reactions and their equilibrium constant expressions. 

It is now possible to calculate the equilibrium constants of oxidation-reduction reactions, and thus to determine whether such reactions can find application in quantitative analysis. Consider first the simple reaction ... 

Because of the bulk of comparable material available, it has been possible to use half-wave potentials for some types of linear free energy relationships that have not been used in connection with rate and equilibrium constants. For example, it has been shown (7, 777) that the effects of substituents on quinone rings on their reactivity towards oxidation-reduction reactions, can be approximately expressed by Hammett substituent constants a. The susceptibility of the reactivity of a cyclic system to substitution in various positions can be expressed quantitatively (7). The numbers on formulae XIII—XV give the reaction constants Qn, r for the given position (values in brackets only very approximate) ... 

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

When a biochemical half-reaction involves the production or consumption of hydrogen ions, the electrode potential depends on the pH. When reactants are weak acids or bases, the pH dependence may be complicated, but this dependence can be calculated if the pKs of both the oxidized and reduced reactants are known. Standard apparent reduction potentials E ° have been determined for a number of oxidation-reduction reactions of biochemical interest at various pH values, but the E ° values for many more biochemical reactions can be calculated from ArG ° values of reactants from the measured apparent equilibrium constants K. Some biochemical redox reactions can be studied potentiometrically, but often reversibility cannot be obtained. Therefore a great deal of the information on reduction potentials in this chapter has come from measurements of apparent equilibrium constants. 

The work described in the foregoing sections is of a preliminary nature. Nevertheless, it offers hope that experimental scales of free hydrogen ion concentration (pcn or pmn) in seawater may be feasible. One need not know pmn or pan to derive meaningful equilibrium data, such as acid-base ratios and solubilities, provided that suitable apparent equilibrium constants are chosen (7). In these cases, the unit selected for the acidity scale disappears by cancellation. Nevertheless, the acidity of seawater is a parameter of broader impact. It plays a role, for example, in the kinetics of organic oxidation-reduction reactions and in a variety of quasi-equilibrium processes of a biological nature. The concentration of free hydrogen ions is clearly understood, and its role in these complex interactions is more clearly defined than that of a quantity whose unit purports to involve the concept of a single-ion activity. 

As you know, oxidation-reduction reactions can involve molecules, ions, free atoms, or combinations of all three. To make it easier to discuss redox reactions without constantly specilying the kind of particle involved, chemists use the term species. In chemistry, a species is any kind of chemical unit involved in a process. For example, a solution of sugar in water contains two major species. In the equilibrium equation NH3 + H2O NH/ + OH , there are four species the two molecules NH3 and H2O and the two ions NH/ and OH. ... 

Reduction-oxidation red Box Bjod kTeq, reaction equilibrium constant... 

Before we discuss redox titration curves based on reduction-oxidation potentials, we need to learn how to calculate equilibrium constants for redox reactions from the half-reaction potentials. The reaction equilibrium constant is used in calculating equilibrium concentrations at the equivalence point, in order to calculate the equivalence point potential. Recall from Chapter 12 that since a cell voltage is zero at reaction equilibrium, the difference between the two half-reaction potentials is zero (or the two potentials are equal), and the Nemst equations for the halfreactions can be equated. When the equations are combined, the log term is that of the equilibrium constant expression for the reaction (see Equation 12.20), and a numerical value can be calculated for the equilibrium constant. This is a consequence of the relationship between the free energy and the equilibrium constant of a reaction. Recall from Equation 6.10 that AG° = —RT In K. Since AG° = —nFE° for the reaction, then... 

Equilibrium Constants for Zirconium Oxidation-Reduction Reactions in 1 1 Mole KCl NaCl... 

These equations suggest that in oxidation-reduction reactions, the relationship of chemical reactions can be expressed as standard electrode potentials or equilibrium constants. 

The values E°, pe° and are different forms to express equilibrium constants of individual oxidation-reduction reactions. The first one is measured in volts of electric voltage, and the rest of them are dimensionless values. As a rule, as equilibrium constants of oxidation-reduction... 

This shows that the difference between two irreversible dissolution reactions can be written as a reversible oxidation/reduction reaction. An important implication of this is that the addition of a dopant is, by itself, not sufficient to enhance the conductivity of a metal oxide photoelectrode one also needs to ensure that the equilibrium of reaction (2.17) lies at the right-hand side. The factors that affect the equilibrium position are oxygen partial pressure, dopant concentration, and temperature. The equihbrium constant of (2.17) is given by... 

Like any chemical system, oxidation-reduction reactions will progress toward equilibrium. Because of this fact, the cell potential provides a way to measure equilibrium constants or free energy changes. 

An oxidation-reduction reaction using Sn(s) to remove N20(g) from a reaction vessel has been proposed, and you need to find its equilibrium constant. Ifou cannot find any thermodynamic information on one product of the reaction, NH30H (aq). How could you estimate the equilibrium constant of the reaction ... 

The great simplification introduced by this procedure can be seen by examining Table 11-1. This table contains only 56 entries, which correspond to 56 different electron reactions. By combining any two of these electron reactions the equation for an ordinary oxidation-reduction reaction can be written. There are 1540 (56 x- ) of these oxidation-reduction reactions that can be formed from the 56 electron reactions. The 56 numbers in the table can be combined in such a way as to give the 1540 values of their equilibrium constants accordingly, this small table permits a prediction to be made as to whether any one of these 1540 reactions will tend to go in a forward direction or the reverse direction. 

The oxidation reactions already described have been discussed in terms of equilibria. From the equilibrium values it can be seen that some compounds, such as acetaldehyde, are oxidized by DPN or TPN quantitatively, whereas, at the same pH values, malate, lactate and ethanol react to a very slight extent. The equilibrium constants of oxidation-reduction reactions have been used to evaluate a property of members... 

Calculating the equilibrium constant from cell emf Given standard potentials (or standard emf), calculate the equilibrium constant for an oxidation-reduction reaction. (EXAMPLE20.il)... 

We have seen that a positive standard cell potential corresponds to a spontaneous oxidation-reduction reaction. And we know (from Chapter 17) that the spontaneity of a reaction is determined by the sign of AG°. Therefore, ceii 2nd AG° must be related. We also know from Section 17.9 that AG° for a reaction is related to the equilibrium constant (K) for the reaction. Since and AG° are related, then E n and K must also be related. 

A chemical reaction does not have to be an oxidation-reduction reaction for us to calculate its equilibrium constant. It is only necessary to be able to write the reaction as the sum of an oxidation half-reaction and a reduction half-reaction. 

The thermite reaction is an oxidation-reduction reaction that uses powdered aluminum metal as a reducing agent to reduce a metal oxide, such as Fe203, to the free metal. The thermodynamic equilibrium constant, K, is an equilibrium constant expression based on activities. In dilute solutions activities can be replaced by the numerical values of molarities and in ideal gases, by the numerical values of partial pressures in bar. The activities of pure solids and liquids are 1. 

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