1 / 35

Chemistry 232

Chemistry 232. Reaction Equilibria in Nonideal Systems. The Equilibrium Constant. For a nonideal system, the nonstandard Gibbs energy of reaction is written. The Equilibrium Condition. If we apply the equilibrium conditions to the above equation. The Autoionization of Water.

zelda-carroll
Download Presentation

Chemistry 232

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemistry 232 Reaction Equilibria in Nonideal Systems

  2. The Equilibrium Constant • For a nonideal system, the nonstandard Gibbs energy of reaction is written

  3. The Equilibrium Condition • If we apply the equilibrium conditions to the above equation

  4. The Autoionization of Water • Water autoionizes (self-dissociates) to a small extent 2H2O(l) ⇌ H3O+(aq) + OH-(aq) H2O(l) ⇌ H+(aq) + OH-(aq) • These are both equivalent definitions of the autoionization reaction. • Water is amphoteric.

  5. The Autoionization Equilibrium • From the equilibrium chapter • But we know a(H2O) is 1.00!

  6. The Defination of Kw Kw = a(H+) a(OH-) Ion product constant for water, Kw, is the product of the activities of the H+ and OH- ions in pure water at a temperature of 298.15 K Kw = a(H+) a(OH-) = 1.0x10-14 at 298.2 K

  7. The pH scale • Attributed to Sørenson in 1909 • We should define the pH of the solution in terms of the hydrogen ion activity in solution pH = -log a(H+) • Single ion activities and activity coefficients can’t be measured

  8. Determination of pH • What are we really measuring when we measure the pH? pH = -log a(H+) • a (H+) is the best approximation to the hydrogen ion activity in solution. • How do we measure a(H+)?

  9. For the dissociation of HCl in water HCl (aq)  Cl-(aq) + H+(aq) • We measure the mean activity of the acid a(HCl) = a(H+) a(Cl-) a(H+) a(Cl-) = (a(HCl))2

  10. Under the assumption a(H+) = a(Cl-) • We obtain a´(H+) = (a(HCl))1/2 = a(HCl)

  11. Equilibria in Aqueous Solutions of Weak Acids/ Weak Bases • By definition, a weak acid or a weak base does not ionize completely in water ( <<100%). • How would we calculate the pH of a solution of a weak acid or a weak base in water?

  12. Equilibria of Weak Acids in Water: The Ka Value • Define the acid dissociation constant Ka • For a general weak acid reaction HA (aq) ⇌ H+ (aq) + A- (aq)

  13. Equilibria of Weak Acids in Water • For the dissolution of HF(aq) in water. HF (aq)  H+ (aq) + F- (aq)

  14. The Nonelectrolyte Activity HF (aq) ⇌ H+ (aq) + F- (aq) • The undissociated HF is a nonelectrolyte  a(HF) = (HF) m[HF]  m[HF] (HF)  1

  15. Equilibria of Weak Bases in Water • Calculate the percentage dissociation of a weak base in water (and the pH of the solutions) CH3NH2 (aq) + H2O ⇌ CH3NH3+(aq)+ OH- (aq)

  16. The Kb Value • Define the base dissociation constant Kb • For a general weak base reaction with water B (aq) + H2O (aq) ⇌ B+ (aq) + OH- (aq)

  17. Calculating the pH of Solutions of Strong Acids • For the dissolution of HCl, HI, or any of the other seven strong acids in water HCl (aq)  H+ (aq) + Cl- (aq) • The pH of these solutions can be estimated from the molality and the mean activity coefficient of the dissolved acid pH = -log ( (acid) m[H+])

  18. Calculating the pH of Solution of Strong Bases • For the dissolution of NaOH, Ba(OH)2, or any of the other strong bases in water NaOH (aq)  Na+ (aq) + OH- (aq) pOH = -log ( (base) m[OH-])

  19. Calculating the pH of a Weak Acid Solution • The pH of a weak acid solution is obtained via an iterative procedure. • We begin by making the assumption that the mean activity coefficient of the dissociated acid is 1.00. • We ‘correct’ the value of (H+) by calculating the mean activity coefficient of the dissociated acid. • Repeat the procedure until (H+) converges.

  20. The Definition of a Buffer • Buffer  a reasonably concentrated solution of a weak acid and its conjugate base • Buffers resist pH changes when an additional amount of strong acid or strong base is added to the solutions.

  21. How would we calculate the pH of a buffer solution?

  22. note pH = -log a(H+) Define pKa = -log (Ka )

  23. The Buffer Equation • Substituting and rearranging

  24. The Generalized Buffer Equation • The pH of the solution determined by the ratio of the weak acid to the conjugate base. • Henderson-Hasselbalch equation often used for buffer calculations!

  25. Buffer  CH3COONa (aq) and CH3COOH (aq)) CH3COOH (aq) ⇄ CH3COO- (aq) + H+ (aq) The Equilibrium Data Table

  26. The pH of the solution will be almost entirely due to the original molalities of acid and base!!

  27. Solubility Equilibria • Examine the following systems AgCl (s) ⇌ Ag+ (aq) + Cl- (aq) BaF2 (s) ⇌ Ba2+ (aq) + 2 F- (aq) • Using the principles of chemical equilibrium, we write the equilibrium constant expressions as follows

  28. The Common Ion Effect • What about the solubility of AgCl in solution containing NaCl (aq)? AgCl (s) ⇌ Ag+ (aq) + Cl- (aq) NaCl (aq) Na+ (aq) + Cl- (aq) AgCl (s) ⇌ Ag+ (aq) + Cl- (aq) Equilibrium is displaced to the left by LeChatelier’s principle (an example of the common ion effect).

  29. Solubility in the Presence of an Inert Electrolyte • What happens when we try to dissolve a solid like AgCl in solutions of an inert electrolyte (e.g., KNO3 (aq))? • We must now take into account of the effect of the ionic strength on the mean activity coefficient!

  30. The Salting-In Effect AgCl (s) ⇌ Ag+ (aq) + Cl- (aq). • Designate the solubility of the salt in the absence of the inert electrolyte as so = m(Ag+) = m(Cl-) at equilibrium.

  31. For a dilute solution • Designate s as the solubility of the salt in the presence of varying concentrations of inert electrolyte.

  32. Reaction Equilibria in Nonideal Gaseous Systems • For a nonideal system gaseous, the nonstandard Gibbs energy of reaction is written

  33. The Equilibrium Condition • Calculate the equilibrium composition from the fugacity coefficients from compression factor data

  34. Temperature and Pressure Dependence of Ko • As a function of temperature • As a function of pressure

More Related