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Understanding Buffer Solutions in Chemistry

Explore the concept of buffer solutions, Henderson-Hasselbalch equation, preparing buffer solutions, acid-base titrations, and solubility of salts in chemistry.

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Understanding Buffer Solutions in Chemistry

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  1. Chapter 16 Buffers Suroviec Spring 2014

  2. I. Buffer Solutions A. Buffer is a solution that resists a change in pH with the addition of small amounts of acid or base • Examples of buffers: • Phosphate • Acetic

  3. Example • What is the pH of the buffer solution that contains 2.2 grams of NH4Cl in 250mL of 0.12M NH3?

  4. B. General expression for buffer solutions • In the previous example we found the [H3O+] by solving for x in:

  5. B. General expression for buffer solutions • Henderson – Hasselbalch equation • The equation shows that pH is controlled by 2 factors • Strength of acid (Ka or pKa) • The two species are present in the same solution, the ratio of their concentrations is also their mole ratio.

  6. C. Using the H – H equation • Calculate the pH of a 1.0L solution that has an 0.050M acetic acid and 0.075M sodium acetate. The Ka of acetic acid is 1.8 X 10-5

  7. D. Preparing a buffer solution • To be a useful a buffer must have 2 characteristics • Be able to control the pH at the desired value • The buffer should have the capacity to control the pH after the addition of small amounts of acid/base

  8. Example I want a buffer of pH 2.50, which of the following can I use to make the buffer and in what ratio? • NaCl and HCl • CH3CO2H and NaCH3CO2 • H3PO4 and NaH2PO4

  9. Example • Given a 25.00mL of 0.100M NaOH and 35.00mL of 0.125M HC2H3O2 will this make a buffer?

  10. E. Effect of added H3O+ or OH- on Buffer Systems • A buffer solution was prepared by adding 4.95 g of NaCH3CO2H to 250mL of 0.150M acetic acid. • What is the pH of the buffer initially? • What is the new pH when 82mg of NaOH is added to a 100 mL aliquot of the buffer?

  11. II. Acid/Base Titrations • A titration is one of the most powerful and accurate ways we have to determine the quantity of an acid/base in solution.

  12. A. Strong acid – strong base • What happens to the pH as 0.10M HCl is slowly added to 50.0mL of 0.10M?

  13. A. Strong acid – strong base

  14. B. Titration of Weak Acid and Strong Base • There are 3 important points to a titration: • pH before base added is calculated from the weak acid Ka and the acid concentration • pH at the equil. point can be calculated from the conjugate base as the conjugate acid and the strong base has been consumed • pH at the ½ equil. point is equal to the pKa of the weak acid

  15. C. Titration of weak base with strong acid Given a 25.0 mL sample of 0.10M NH3 being titrated with 0.10M HCl • What is the pH of the solution before the titration begins? • What is the pH at the equilivance point? • What is the pH at the midpoint?

  16. III. pH Indicators • An indicator is an organic compound that is itself a weak acid or weak base • In aqueous solution the acid from is in equilibrium with its conjugate base.

  17. IV. Solubility of Salts • Precipitation reactions are reactions in which one of the products are water soluble

  18. A. The Solubility Product Constant, Ksp • Ksp reflects the solubility of a compound and is usually called the solubility product constant

  19. B. Solubility and Ksp • Ksp is measured by experiments in the lab determining concentrations of ions in solution

  20. C. Soluability and the Common Ion • What if we had a test tube of saturated AgNO3 and I added more KNO3?

  21. Example • What is the final molarity of AgSCN (s) if it is placed in • Water • 0.010 M NaSCN

  22. D. Common Ion Concepts • The solubility of a salt will always be reduced by the presence of a common ion • We made the approximation that the amount of common ion added to the solution will also be large compared with the amount of ion coming from the insoluble salt.

  23. E. Precipitation Reactions • How do we know when a precipitate will form? • Look at Q and how is Q related to Ksp

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