Basic Reaction-Equilibrium Relations

This is the same value found in (10.2.63) using the Lewis-Randall standard state. 

Reference-solvent dilute-solution standard state. In this case we do not have values for / pure 1 0 s cannot apply (10.2.61) or (10.2.64). However, we are able to 

Again the value is less than unity. Now (10.2.67) becomes 

This is the same value found in (10.2.63) using the Lewis-RandaU standard state and found in (10.2.66) using the solute-free dilute-solution standard state. 

In general we can say that the reference-solvent dilute-solution standard state is easier to use than the solute-free dilute-solution standard state (except, of course, when Y can be assumed to be unity). This is because is completely independent of composition, while depends on the solute-free mole fractions. But more generally, the lesson is that the three kinds of activity coefficients are simply proportional they are aU embedded with the same information, so they aU give the same value for a fugacity. We use the particular standard state that allows us to take advantage of available data and that simplifies calculations. 

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BASIC REACTION-EQUILIBRIUM RELATIONS 453 10.3.6 Nonstoichiometric Development... 

As discussed in Section 3.10.3, in the gas phase the basicity of simple amines follows the order NMe3 > NHMe2 > NH2Me > NH3 because of the electron donating effect of the methyl (Me) groups. In solution, however, we can define a basicity constant as the equilibrium constant for the reaction shown in Equation 3.4. Note it is important to specify temperature, solvent (usually water) and solution ionic strength, 1 Basicity constants are related to the acid dissociation constants (/Q of the base s conjugate acid via the dissociation constant of water, K = 10 14 at 25 °C. Thus Kbx K = Kw. 

The quantity, represents the chemical potential of the undissociated species M2A in its standard state, and not that of the component. The other standard-state quantities represent the chemical potential of the designated species in their standard states, but, for the present, we cannot separate the two standard chemical potentials from the sums, neither is it important to do so. The standard-state quantities appearing in Equations (8.199)—(8.201) are not all independent, because the three equations are equivalent. If we equate Equations (8.199) and (8.200), we obtain an expression that can be evaluated experimentally for the quantity (/ij + + ma- m2a)- Similarly, we obtain an expression that can be evaluated experimentally for the quantity (2— /i 2a) when we equate Equations (81.99) and (8.201). These last two quantities are related to the equilibrium constants for the chemical reactions. This relation is developed in Chappter 11 and the basic experimental methods are discussed in Chapters 10 and 11. 

If there is no doubt that many Friedel-Crafts halides can undergo selfdissociation in certain solvents (particularly halogenated hydrocarbons), and that this ionc enic reaction is not due to basic impurities in the system, the attainment of the equilibrium relating ions to undissociated molecules often requires a long time This... 

Ammonia, iVffy a colorless gas with a sharp, characteristic smell condensation temperature at normal pressure and about -40°C.It is very soluble in cold water, but completely driven off by boiling. The aqueous solution is weakly basic, because NH3 takes up a proton to form an ammonium ion NH3 -1- HjO NH4 + OH. The reaction equilibrium lies far to the left, so that NH3 can be displaced from ammonium compounds by bases. The toxicity of NH3 is related to the high permeation rate of the nonprotonated form and its tendency to become protonated. 

Scales for bases that are too weak to study in aqueous solution employ other solvents but are related to the equilibrium in aqueous solution. These equilibrium constants provide a measure of thermodynamic basicity, but we also need to have some concept of kinetic basicity. For the reactions in Scheme 5.4, for example, it is important to be able to make generalizations about the rates of competing reactions. 

The last chapter in this introductory part covers the basic physical chemistry that is required for using the rest of the book. The main ideas of this chapter relate to basic thermodynamics and kinetics. The thermodynamic conditions determine whether a reaction will occur spontaneously, and if so whether the reaction releases energy and how much of the products are produced compared to the amount of reactants once the system reaches thermodynamic equilibrium. Kinetics, on the other hand, determine how fast a reaction occurs if it is thermodynamically favorable. In the natural environment, we have systems for which reactions would be thermodynamically favorable, but the kinetics are so slow that the system remains in a state of perpetual disequilibrium. A good example of one such system is our atmosphere, as is also covered later in Chapter 7. As part of the presentation of thermodynamics, a section on oxidation-reduction (redox) is included in this chapter. This is meant primarily as preparation for Chapter 16, but it is important to keep this material in mind for the rest of the book as well, since redox reactions are responsible for many of the elemental transitions in biogeochemical cycles. 

Later we shall see how fundamental quantities such as /i can be estimated from first principles (via a basic knowledge of the molecule such as its molecular weight, rotational constants etc.) and how the equilibrium constant is derived by requiring the chemical potentials of the interacting species to add up to zero as in Eq. (20). The above equations relate kinetics to thermodynamics and enable one to predict the rate constant for a reaction in the forward direction if the rate constant for the reverse reaction as well as thermodynamic data is known. 

The correlation is useful but not exact. This is because basicity is a measure of the position of equilibrium between a substrate and its conjugate acid (see Section 4.4), whereas nucleophilicity relates to a rate of reaction. The above relationship breaks down when one looks at atoms in the same column of the periodic table. As atomic number increases, basicity decreases, whilst nucleophilicity actually increases (Table 6.3). This originates from the size of the atom, so that electrons associated with larger atoms become less localized, consequently forming weaker bonds with protons (see acidity of HX, Section 4.3.2). On... 

Only those components which are gases contribute to powers of RT. More fundamentally, the equilibrium constant should be defined only after standard states are specified, the factors in the equilibrium constant should be ratios of concentrations or pressures to those of the standard states, the equilibrium constant should be dimensionless, and all references to pressures or concentrations should really be references to fiigacities or activities. For reactions involving moderately concentrated ionic species (>1 mM) or moderately large molecules at high pressures ( 1—10 MPa), the activity and fiigacity corrections become important in those instances, kineticists do use the proper relations. In some other situations, eg, reactions on a surface, measures of chemical activity must be introduced. Such cases may often be treated by straightforward modifications of the basic approach covered herein. 

The solubility product, Ksp, for an ionic compound is the equilibrium constant for dissolution of the compound in water. The solubility of the compound and Ksp are related by the equilibrium equation for the dissolution reaction. The solubility of an ionic compound is (1) suppressed by the presence of a common ion in the solution (2) increased by decreasing the pH if the compound contains a basic anion, such as OH-, S2-, or CO32- and (3) increased by the presence of a Lewis base, such as NH3, CN-, or OH-, that can bond to the metal cation to form a complex ion. The stability of a complex ion is measured by its formation constant, Kf. 

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