1 / 34

Aqueous Ionic Solutions and Equilibrium

Aqueous Ionic Solutions and Equilibrium. Chapter 19. Common Ion Effect. Shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction. What happens to the equilibrium if 0.10 M NaF is added?.

madge
Download Presentation

Aqueous Ionic Solutions and Equilibrium

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. Aqueous Ionic Solutions and Equilibrium Chapter 19

  2. Common Ion Effect • Shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction. What happens to the equilibrium if 0.10 M NaF is added? What happens to the equilibrium if 0.10 M NaCl is added?

  3. Buffers • Resist a change in pH when H+ or OH- is added • Components: conjugate acid-base pairs • acidic component reacts with OH- • basic component reacts with H+ • Example: 1.00 L of 0.500 M CH3COOH + • 0.500 M CH3COONa • CH3COOH/CH3COO-

  4. Reactions with H+ or OH-

  5. Key Points on Buffers • 1. Weak acids and bases containing common ion • 2. Problems involve: • stoichiometry first • equilibrium second

  6. Buffer Characteristics • Contain relatively large amounts of weak acid and corresponding base. • Added H+ reacts to completion with weak base. • Added OH reacts to completion with weak acid. • pH is determined by ratio of concentrations of weak acid and weak base.

  7. Buffer Capacity • Amount of H+ or OH- it can absorb without a significant change in pH. • For HA/A- system, buffer capacity depends on: • [HA] and [A-] (higher = higher) • [HA] ratio (closer to 1 = higher) • [A-]

  8. Calculations with Ka • Calculate the pH of a buffer consisting of 0.50 M HF and 0.45 M F- • (a) before and • (b) after addition of 0.40 g NaOH to • 1.0 L of the buffer. • Ka of HF = 6.8 x 10-4

  9. Calculations • Find pH of a buffer • Buffer preparation • Find equilibrium concentrations • Helpful: Hendeson-Hasselbach equation

  10. Calculations with Ka • Calculate the pH of a buffer consisting of • 0.50 M HF and 0.45 M F- using the Henderson-Hasselbach equation.

  11. Preparing a buffer • 1. Choose and acid-conjugate base pair • 2. Calculate the ratio of the buffer component pairs • 3. Determine the buffer concentration • How would you prepare a benzoic acid/benzoate buffer with pH = 4.25, starting with 5.0 L of 0.050 M sodium benzoate (C6H5COONa) solution and adding the acidic component? • Ka of benzoic acid (C6H5COOH) = 6.3 x 10-5

  12. Titration (pH) Curves • Plot pH of solution vs. amount of titrant added • Equivalence point: • Enough titrant added to react exactly with solution being analyzed.

  13. Titration of Strong Acid with Strong Base

  14. Weak Acid-Strong Base Titration • 1. Stoichiometry • Reaction assumed to run to completion • 2. Equilibrium • Use weak acid equilibrium to find pH

  15. Weak Acid-Strong Base Titration

  16. Differences

  17. Differences and Ka

  18. Titration Calculations • 1. Solution of HA • 2. Solution of HA and added base • 3. Equivalent amounts of HA and added base • 4. Excess base • A chemist titrates 20.00 mL of 0.2000 M HBrO • (Ka = 2.3 x 10-9) with 0.1000 M NaOH. Find the pH: • (a) before any base is added • (b) when 30.00 mL of NaOH is added • (c) at the equivalence point • (d) when the moles of OH- added are twice the moles of HBrO originally present?

  19. Titration of Strong Base with Strong Acid

  20. Weak Base-Strong Acid Titration

  21. Acid-Base Indicator • Indicates endpoint of a titration • Endpoint is not necessarily the equivalence point

  22. Polyprotic Acid, H2SO3

  23. Solubility Constant, Ksp • For solids dissolving to form aqueous solutions. • For slightly soluble salts: • equilibrium between solid and component ions

  24. Ksp • Write the expression for Ksp for • (a) CaSO4 • (b) Cr2CO3 • (c) Mg(OH)2 • (d) As2S3

  25. Solubility Product • Solubility s: • Amount of PbCl2 that dissolves • For: • [PbCl2]dissolved = [Pb2+] = ½[Cl-] • s: varies, especially if common ion is present

  26. Calculations • 1.5 x 10-4 g of CaF2 dissolves in 10.00 mL solution at 18°C. • Write the expression for Ksp. • Find the molar solubility of CaF2. • Find the [Ca2+] and [F-]. • Calculate Ksp. • What is the molar solubility of Mg(OH)2 if the value of Ksp is 6.3 x 10-10?

  27. Complex Ions • Complex Ion: • Charged species - metal ion surrounded by ligands (Lewis bases). • Coordination Number: • No. of ligands attached to a metal ion. • (Common: 6 and 4.) • Formation (Stability) Constants, Kf: • Equilibrium constants for stepwise addition of ligands to metal ions.

  28. Overall Formation Constant

  29. Complex Ions • Write the stepwise formation constants for Cr(NH3)63+, starting from Cr(H2O)63+ and NH3(aq) • What is the coordination number of Cr3+?

  30. Applications • Selective precipitation • Exploit differences in Ksp • Qualitative analysis • 1. Insoluble chlorides (Ag+, Hg22+, Pb2+) • 2. Acid-insoluble sulfides (Cu2+, Cd2+, Hg2+, As3+, Sb3+, Bi3+, Sn2+, Sn4+, Pb2+) • 3. Base-insoluble sulfides and hydroxides (Zn2+,Mn2+, Ni2+, Fe2+, Co2+ as sulfides; Al3+, Cr3+ as hydroxides) • 4. Insoluble phosphates (Mg2+, Ca2+, Ba2+) • 5. Alkali metal and ammonium ions

More Related