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Buffers

Buffers. A Buffer in Action. Definition of a Buffer. A buffer solution is one which resists changes in pH when small quantities of an acid or a base are added to it. The most important way that the pH of the blood is kept relatively constant is by buffers dissolved in the blood.

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Buffers

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  1. Buffers A Buffer in Action

  2. Definition of a Buffer • A buffer solution is one which resists changes in pH when small quantities of an acid or a base are added to it. • The most important way that the pH of the blood is kept relatively constant is by buffers dissolved in the blood. H2CO3 + H2O HCO3- + OH- HCO3- + H+ H2CO3 HCO3- + 3OH  H2O + H2CO3 • When a substance acts as both an acid and a base such as HCO3- does, it is called amphoteric.

  3. How do buffer solutions work? • A buffer solution contains a weak acid and base, which removes any hydrogen ions or hydroxide ions that are add – keeping the pH constant. (not neutral) • The Bicarbonate Buffer: weak acid weak base H2CO3 + H2O  HCO3- + OH-

  4. Another Buffer solution • Ammonia is a weak base, and the position of this equilibrium will be well to the left: • Any added hydrogen ions will react with the ammonia to make more of the conjugate acid, maintaining the pH.

  5. Another Buffer solutions • Any added OH- ions will react with the weak acid to make the conjugate base and again maintain a constant pH.

  6. Determination of the pH of a Buffer solution • Example: What is the pH of a solution of 0.11 M NaC2H3O2 and 0.090 M HC2H3O2? (ka = 1.8 x 10-5) HC2H3O2 C2H3O2- + H+ Initial 0.09 0.11 0 Change -x +X +x Equil. .09-x 0.11 +X x ka = x (0.11)/.09-x X = 1.47 x 10-5, pH = 4.83 • (buffered solution will have a very small dissociation so + or – x can be ignored)

  7. Preparing Buffer Solutions • The Henderson-Hasselbalch Equation gives the ratio of weak acid to base needed to maintain a constant pH. • pH = pka + log [A-]/[HA] (The pH can be approximated: pH = pka +1

  8. Preparing Buffers • Example. What is the mole ratio of acetic acid to acetate ion needed to prepare a buffer solution at a pH = 5.00? ka = 1.8 x 10-5 ka = [H+] [A-] = [H+] mol C2H3O2- [HA] mol HC2H3O pH = 5 so,[H+] = 1 x 10-5 mol C2H3O2- = [H+] = 1x 10-5 = .56 mol HC2H3O ka 1.8 x 10-5

  9. Preparing Buffers • What is the pH of a solution containing 0.11 M of sodium acetate and 0.090 M acetic acid? Ka = 1.8 x 10-5 • pH = pka + log [A-]/[HA] pH = 4.74 + log (0.11mol A-) (0.090 mol HA) pH = 4.83

  10. Titrations • Titration of a strong acid by a strong base The equivalence point occurs at 25.00 mL added base with a pH of 7.0.

  11. Titrations • Titration of a weak acid by a strong base • This can be divided into four regions • Before the titration begins: this is simple a solution of weak acid • During the titration, but before the equivalence point: the solution is a buffer • At the equivalence point: the solution contains a salt of the weak acid, and hydrolysis can occur • Past the equivalence point: the excess added OH- is used to determine the pH of the solution

  12. Buffered Region The titration curve for the titration of 25.00 mL of 0.200 M acetic acid with 0.200 M sodium hydroxide. Due to hydrolysis, the pH at the equivalence point higher than 7.00.

  13. Titrations • Titration of a weak base by a strong acid • This is similar to the titration of a weak acid by strong base • Again dividing into four regions • Before the titration begins: this is a solution of a weak base in water • During the titration, but before the equivalence point: the solution is a buffer • At the equivalence point: the solution contains the salt of the weak base, and hydrolysis can occur • Past the equivalence point: excess added H+ determines the pH of the solution

  14. Buffered Region Titration curve for the titration of 25.00 mL of 0.200 M NH3 with 0.200 M HCl. The pH at the equivalence point is below 7.00 because of the hydrolysis of NH4+.

  15. Titrations • Titration curves for diprotic acids • The features are similar to those for monoprotic acids, but two equivalence points are reached The titration of the diprotic acid H2A by a strong base. As each equivalence point is reached, the pH rises sharply.

  16. Titrations • A few general comments about indicators can be made • Most dyes that are acid-base indicators are weak acids, which can be represented as HIn • The color change can be represented as:

  17. Titrations • The color change will “appear” to the human eye near the equivalence point of the indicator • At the equivalence point, the concentration of the acid and base form are equal, so that • The best indicators have intense color(s) so only a small amount will produce an intense color change that is “easy” to see and won’t consume too much of the titrant.

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