1 / 27

Chapter 18

Chapter 18. Acid-base Equilibrium II : Buffers and Indicators. 18.1 Buffers 18.2 Calculations Involving Composition and pH of Buffer Solutions 18.3 Acid-base Indicators 18.4 Acid-base Titrations. 2 types of buffers :.

claral
Télécharger la présentation

Chapter 18

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. Chapter 18 Acid-base Equilibrium II : Buffers and Indicators 18.1 Buffers 18.2 Calculations Involving Composition and pH of Buffer Solutions 18.3 Acid-base Indicators 18.4 Acid-base Titrations

  2. 2 types of buffers: 1. Acidic buffer --- mixing a weak acid and its salt of a strong base. (e.g. CH3COOH & CH3COONa) 18.1 Buffers (SB p.151) Buffers A buffer solution is a solution that tends to resist change in pH when a small amount of acid or base is added to it. 2. Basic buffer --- mixing a weak base and its salt of a strong acid. (e.g. NH3 & NH4Cl)

  3. 18.1 Buffers (SB p.151) Action of Acidic Buffers CH3COOH(aq) \-----==\ CH3COO-(aq) + H+(aq) CH3COO-Na+(aq) --------> CH3COO-(aq) + Na+(aq) Note that the buffer contains a large amount of the weak acid (CH3COOH) and its conjugate base (CH3COO-). If a small amount of acid is added to this system,H+(aq) + CH3COO-(aq) \==------\ CH3COOH(aq) If a small amount of alkali is added to this system,OH-(aq) + CH3COOH(aq) \==------\ CH3COO-(aq) + H2O(l)

  4. 18.1 Buffers (SB p.152) If a small amount of acid is added to this system,H+(aq) + CH3COO-(aq) \==------\ CH3COOH(aq) If a small amount of alkali is added to this system,OH-(aq) + CH3COOH(aq) \==------\ CH3COO-(aq) + H2O(l)

  5. 18.1 Buffers (SB p.153) pH = pKa + log pH = pKa + log The pH of an Acidic Buffer Solution CH3COOH(aq) \-----==\ CH3COO-(aq) + H+(aq) CH3COO-Na+(aq) --------> CH3COO-(aq) + Na+(aq) initial conc’s

  6. 18.1 Buffers (SB p.154) Action of Basic Buffers NH3(aq) + H2O \-----==\ NH4+(aq) + OH-(aq) NH4+Cl-(aq) --------> NH4+(aq) + Cl-(aq) Note that the buffer contains a large amount of the weak base (NH3) and its conjugate acid (NH4+). If a small amount of acid is added to this system,H+(aq) + NH3(aq) \==------\ NH4+(aq) If a small amount of alkali is added to this system,OH-(aq) + NH4+(aq) \==------\ NH3(aq) + H2O(l)

  7. 18.1 Buffers (SB p.154) If a small amount of acid is added to this system,H+(aq) + NH3(aq) \==------\ NH4+(aq) If a small amount of alkali is added to this system,OH-(aq) + NH4+(aq) \==------\ NH3(aq) + H2O(l)

  8. 18.1 Buffers (SB p.155) pOH = pKb + log pOH = pKb + log pH = 14 – pKb - log The pH of an Basic Buffer Solution NH3(aq) + H2O \-----==\ NH4+(aq) + OH-(aq) NH4+Cl-(aq) --------> NH4+(aq) + Cl-(aq) initial conc’s

  9. 18.2 Calculations Involving Composition and pH of Buffer Solutions (SB p.156) Calculations Involving Composition and pH of Buffer Solution Example 18-1 A buffer is made by adding 4.1g of sodium ethanoate to 1dm3 of a 0.01 M solution of ethanoic acid. Calculate the pH of the buffer. (Given: Ka of CH3COOH at 298 K = 1.74 x 10-5 mol dm-3; molar mass of CH3COONa = 82 g mol-1; assume there is no volume change on mixing)

  10. Solution Number of moles of CH3COONa = 4.1 g/ 82 g mol-1 = 0.05 mol [CH3COO-(aq)] = 0.05 mol/1 dm3 = 0.05M [CH3COOH(aq)] = 0.01M pH = pKa + log pH = -log(1.74x10-5) + log 0.05/0.01 = 4.76 + 0.70 = 5.46 18.2 Calculations Involving Composition and pH of Buffer Solutions (SB p.156) Answer

  11. 18.2 Calculations Involving Composition and pH of Buffer Solutions (SB p.159) Example 18-3 How many grams of ammonium chloride would you add to 100 cm3 of 0.1 M NH3(aq) in order to prepare a basic buffer of pH 9.0? (Given: Kb of NH3 at 298K = 1.74 x 10-5 mol dm-3; molar mass of NH4Cl = 53.5 g mol-1) Solution Let x M be the concentration of ammonium chloride in the buffer solution. pH = 14 – pKb – log[salt]/[base] 9 = 14 – [-log(1.74 x 10-5)] – log(x / 0.1) 9 = 9.24 – log(x/0.1) x = 0.174 ∴ [NH4Cl(aq)] = 0.174 M (continued) Answer

  12. 18.2 Calculations Involving Composition and pH of Buffer Solutions (SB p.159) Number of moles of NH4Cl used = 0.174 M x 100/1000 dm3 = 0.0174 mol Mass of NH4Cl used = 0.0174 mol x 53.5 g mol-1 = 0.931 g

  13. 18.3 Acid-base Indicators (SB p.160) HIn(aq) + H2O(l) H3O+(aq) + In-(aq) weak acid conjugate base (colour 1) (colour 2) KIn = pH = pKIn + log Acid-base Indicators

  14. HIn(aq) + H2O(l) H3O+(aq) + In-(aq) weak acid conjugate base (colour 1) (colour 2) When  , colour 1 is observed When  10 , colour 2 is observed pH = pKIn + log 18.3 Acid-base Indicators (SB p.161)

  15. 18.3 Acid-base Indicators (SB p.161) Thus the indicator changes from colour 1 to colour 2 over a range of: from  to  10 which corresponds to pH =(pKIn-1) to pH = =(pKIn+1) Generally speaking, the colour change takes place over a range of 2 pH units (pH range of the indicator).

  16. 18.3 Acid-base Indicators (SB p.162) Phenolphthalein HPh(aq) + H2O(l) Ph-(aq) + H3O+(aq) colourless pink KIn = 7 x 10-10 mol dm-3

  17. 18.3 Acid-base Indicators (SB p.162) When ≦ 1/10 pH ≦ pKIn + log 1/10 ≦ pKIn - 1 ≦ 8.15 When ≧ 10 pH ≧ pKIn + log 10 ≧ pKIn + 1 ≧ 10.15 Phenolphthalein When [Ph-(aq)] = [HPh(aq)], pH = pKIn = -log (7 x 10-10) = 9.15

  18. 18.3 Acid-base Indicators (SB p.162) HMe+(aq) + H2O(l) Me(aq) + H3O+(aq) red yellow KIn = 2 x 10-4 mol dm-3 Methyl Orange pH = pKIn + log

  19. 18.3 Acid-base Indicators (SB p.162) When 1/10 pH ≦ pKIn + log 1/10 ≦ pKIn - 1 ≦2.7 When ≧ 10 pH ≧ pKIn + log 10 ≧ pKIn + 1 ≧ 4.7 When [Me(aq)] = [HMe+(aq)], pH = pKIn = -log (2 x 10-4) = 3.7

  20. 18.3 Acid-base Indicators (SB p.163)

  21. 18.4 Acid-base Titrations (SB p.164) Acid-base Titrations Remarks base 1. Titration is the determination of the equivalence point, the point at which equivalent quantities of the acid and base have reacted. 2. There are oftensharp changes in pH near the equivalence points of the titrations. 3. The point at which an indicator change colour during titration is called the end point. 4. A good indicator is one whose end pt. matches with the equivalence pt. acid + indicator

  22. 18.3 Acid-base Indicators (SB p.165) pH Titration Curves Strong Acid-Strong Base Titration Sharp change in pH at the equivalence point: 3-11 Both methyl orange &phenolphthalein can indicate the equivalence point accurately. (Their pH ranges lie within the sharp change in pH.)

  23. 18.3 Acid-base Indicators (SB p.165) Strong Acid-Weak Base Titration Sharp change in pH at the equivalence point: 2-6 Methyl orange can indicate the equivalence point accurately. (Its pH range lies within the sharp change in pH.) Phenolphthalein cannot indicate the equivalence point accurately. (Its pH range does NOT lie within the sharp change in pH.)

  24. 18.3 Acid-base Indicators (SB p.165) Weak Acid-Strong Base Titration Sharp change in pH at the equivalence point: 8-12 Phenolphthalein can indicate the equivalence point accurately. (Its pH range lies within the sharp change in pH.) Methyl orange cannot indicate the equivalence point accurately. (Its pH range does NOT lie within the sharp change in pH.)

  25. 18.3 Acid-base Indicators (SB p.165) Weak Acid-Weak Base Titration NO Sharp change in pH at the equivalence point. Both methyl orange &phenolphthalein CANNOT indicate the equivalence point accurately.

  26. 18.3 Acid-base Indicators (SB p.167) Double Indicator Method Na2CO3 + HCl  …….. Na2CO3 + HCl NaHCO3 + NaCl NaHCO3 + HCl  NaCl + CO2 + H2O

  27. The END

More Related