Analytical Chemistry - The Effect of Electrolytes on Chemical Equilibrium | CHEM 321, Study notes of Analytical Chemistry

Download Analytical Chemistry - The Effect of Electrolytes on Chemical Equilibrium | CHEM 321 and more Study notes Analytical Chemistry in PDF only on Docsity! 1 CHEM 321: Analytical Chemistry Chapter 8 Notes I. The Effect of Electrolytes on Chemical Equilibrium A. In general, because of the interaction between electrolytes in solutions (i.e. the electrical double layer) the tendency of two reactant electrolytes to form product will be affected by the concentrations and charge strength of "spectator" or non-reacting electrolytes. In general, the tendency to form precipitates will be decreased with increasing ionic strength due to interactions (but not necessarily reactions) of the reactant species with other charged species, while the tendency to form charged product species from uncharged reactant species increases, due to interactions of the charged product species with the spectator ions. Examples: Ag+ + Cl- <==> AgCl (s) , Ksp increases with increasing ionic strength, so AgCl becomes more soluble. H2O + H2CO3 <===> H3O+ + HCO3-, Ka1 increases with increasing ionic strength, so the solution becomes more acidic. B. Equilibrium Constant We define the experimental or "apparent" equilibrium constant, K, as approaching the activity or thermodynamic equilibrium constant, K', in very dilute solutions. In the remaining chapters we will use concentration-based K' as if they were K-values, you should remember that the values for all thermodynamic equilibrium constants for reactions in solution are dependent on the ionic strength of the solution as well as the solution temperature. For this chapter, we will designate concentration-based equilibrium constants with a ', and examine the relationship between K' and K. See Figure 8-1 and note that Ksp's are more greatly affected by ionic strength than acid dissociation constants or the ion-product of water. II. The Effect of Ionic Strength A. Ionic strength is a measure of the concentration and charge of the ions in solution. It is defined by the equation: ionic strength = 1/2([A]ZA2 + [B]ZB2 + [C]ZC2 + ...) where A, B, and C are charged species in the solution, [A], [B], [C], … represent the species molar concentrations of ions A, B. C, and ZA, ZB, Zc, … are their respective charges. B. Note that the ionic strength is not dependent on the chemical properties of the ions, only their concentrations and charges. Therefore, the tendency of a reaction to occur is not dependent on the types of spectator ions provided the ionic strength of the solution remains the same. Please note this statement is only true for the low ionic strength solution (µ ≤ 0.1 M). C. Activity Coefficients: these are "fudge factors" which adjust the theoretical concentration to equal the observed concentration: 2 ax = [X]γx where ax = is the activity or observed concentration, [X] = the expected or theoretical concentration and γx = the activity coefficient, and for aA + bB <====> cC + dD The ratio of the product to reactant activity coefficients approach unity in very dilute solutions (K ----- > K'). D. Properties of Activity Coefficients 1). As ionic strength increases, its activities coefficient (γ) decreases. At moderate ionic strengths, γ < 1 As the solution approaches infinite dilution, γ ~ 1 and hence α ~ [X] and K’ ~ K At high ionic strengths (µ > 0.1 M), the activity coefficients often increase and may even become larger than unity. 2). In solutions that are not too concentrated (µ ≤ 0.1 M), the activity coefficient for a given species is independent of the types of electrolyte and dependent only upon the ionic strength. 3) For a given ionic strength, the activity coefficient of an ion departs farther from unity as the charge of the species increase. (Fig. 8-2) 4) At a given ionic strength, the activity coefficients of ions of the same charge are approximately the same. There are exceptions. E. The Debye-Hűckel Equation: m)(10 nanometersin Xion hydrated theofdiameter effectiveα solution theofstrength ionic µ X species on the chargeZ X species oft coefficienacitivity γ µ3.3α1 µ0.51Z γlog- 9- x x x x 2 x x = = = = + = Activity coefficients are shown in Table 8-1. Since it is impossible to have an excess of positively- or negatively-charged species in a neutral solution, we really can only calculate the mean activity coefficient for the solution. For example: The mean activity coefficient of the electrolyte AaBb is defined as b B ba A a d D dc C c γ]B[γ[A] γ]D[γ[C] K =