Electrolyte Effects: Activity or Concentration Jess Sproul Chapter 9 Notes Chapter 9 Electrolyte Effects: Activity or Concentration

9A-1: How Do Ionic Charges Affect Equilibria? The magnitude of the effect that the electrolyte has is highly dependent on the charge of the ions participating in the equilibrium. When only neutral species are involved in the equilibrium, the electrolyte has no real effect, no matter how concentrated.

9A-1: How Do Ionic Charges Affect Equilibria? (cont.) In the figure seen at the right, the molar solubility of the doubly charged ions in BaSO4 increases by a factor of 2 in a solution of 0.02 M KNO3 over a solution of pure water. On the other hand, the molar solubility of a solution of Ba(IO3)2 with one singly and one doubly charged ion only increases by a factor of 1.25 in a solution of 0.02 M KNO3 over a solution of pure water. And finally, the molar solubility of AgCl with two singly charged ions only increases by a factor of 1.2 in a solution of 0.02 M

9A-2: What Is the Effect of Ionic Strength on Equilibria? (cont.) The ionic strength of a solution of a strong electrolyte consisting solely of singly charged ions is identical with its total molar salt concentration. But when the ions are no longer singly charged, the ionic strength increases. Type Electrolyte Example Ionic Strength* 1 : 1 NaCl c 1 : 2 Ba(NO3)2, Na2SO4 3c 1 : 3 Al(NO3)3, Na3PO4 6c 2 : 2 MgSO4 4c *c = molarity of the salt

9A-3: The Salt Effect The electrolyte effect results from the electrostatic attractive and repulsive forces that exist between the ions of an electrolyte and the ions involved in an equilibrium. These forces cause each ion from the dissociated reactant to be surrounded by a sheath of solution that contains a slight excess of electrolyte ions of opposite charge.

9A-3: The Salt Effect The electrolyte effect results from the electrostatic attractive and repulsive forces that exist between the ions of an electrolyte and the ions involved in an equilibrium. These forces cause each ion from the dissociated reactant to be surrounded by a sheath of solution that contains a slight excess of electrolyte ions of opposite charge. These charged atmospheres lessen the charges on the ions in equilibrium, the Ba2+ and SO42- in the figure, thus decreasing their overall attraction to each other and increasing solubility.

9B: Activity Coefficients The term activity, a, is used to account for the effects of electrolytes on chemical equilibria. The activity, or effective concentration, of species X depends on the ionic strength of the medium and is defined as ax = γx [X] where ax is the activity of X, [X] is its molar concentration, and γx is a dimensionless quantity called the activity coefficient. The activity coefficient and the activity of X vary with ionic strength.

9B-1: Properties of Activity Coefficients In dilute solutions, the activity coefficient for a given species is independent of the nature of the electrolyte and dependent only on ionic strength. For a given ionic strength, the activity coefficient of an ion departs farther from unity as the charge carried by the species increases. At any given ionic strength, the activity coefficients of ions of the same charge are approximately equal. The activity coefficient of a given ion describes its effective behavior in all equilibria in which it participates.

9B-1: Properties of Activity Coefficients (cont.)

9B-2) The Debye-Hückel Equation The Debye-Hückel equation takes the ionic atmosphere model and derives a theoretical expression that permits the calculation of activity coefficients of ions from their charge and their average size.

9B-3: Equilibrium Calculations with Activity Coefficients Equilibrium calculations with activities yield values that are better in agreement with experimental results than those obtained with molar concentrations.

9B-4: Omitting Activity Coefficients in Equilibrium Calculations We shall ordinarily neglect activity coefficients and simply use molar concentrations in applications of the equilibrium law. This simplifies calculations and decreases necessary data. For most purposes, the error introduced by this method is minimal. Significant discrepancies occur when the ionic strength is larger than 0.01 or when the ions involved have multiple charges. With dilute solutions of nonelectrolytes or of simply charged ions, the use of concentrations in a mass-law calculation often provides reasonably accurate results.