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Acid-Base Equilibrium

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    Acid/Base Properties of a Pharmaceutical Compound

    Many pharmaceutically interesting compounds are weak acids or weak bases, a fact of some importance when preparing them for use by the public. For example, if the active ingredient in a nasal spray is too acidic, then the preparation might include an additional ingredient to neutralize some of the acid; after all, no one wants to spray something as acidic as, say, lemon juice up his or her nose!

    Pseudoephedrine is a central nervous system stimulant used in many cold and allergy tablets.  In its molecular form it is a weak base (it is an amine), which, for convenience, we may represent as B. In water, the following equilibrium reaction exists

    \[\ce{B}(aq) + \ce{H2O}(l) ⇌ \ce{OH-}(aq) + \ce{HB+}(aq)\nonumber\]

    The Merck Index reports that a 0.030 M solution of pseudoephedrine has an equilibrium pH of 11.44. What is the value of Kb for this compound?

     

     

     

     

     

     

     

     

     

     

     

    Because its base form is only slightly soluble in water, pseudoephedrine typically is dispensed in its weak acid form, HB+. Pseudoephedrine hydrochloride, therefore, is an ionic compound consisting of protonated pseudoephedrine, HB+, and Cl– as a counter ion. The formula for this often is written as B•HCl; don’t confuse this with the strong acid HCl. A solution of pseudoephedrine hydrochloride, therefore, is acidic due to the presence of HB+. Write the Ka reaction responsible for making the solution acidic and report the value for Ka.

     

     

     

     

    Suppose you dissolve three tablets of Sudafed®, each containing 30.0 mg of pseudoephedrine hydrochloride, in 200.0 mL of water. What is the pH of the resulting solution? The molar mass for pseudoephedrine hydrochloride is 201.7 g/mol. 

     

     

     

     

    Buffers and pH

    1. What is the pH of a buffer that is 0.55 M in formic acid, HCOOH, and 0.63 M in sodium formate, NaHCOO?

     

     

     

     

     

     

     

     

     

     

    1. What is the ratio of hypobromite, BrO–, to hypobromous acid, HBrO, in a buffer with a pH of 7.88?

     

     

     

     

     

     

     

     

     

     

    1. Human blood contains two buffer systems, one based on phosphate species and one on carbonate species. If blood has a normal pH of 7.4, what are the principle phosphate and carbonate species present? What is the ratio between the two phosphate species? At the temperature of human blood, the Ka values for phosphoric acid are 1.3×10–2, 2.3×10–7, and 6×10–12, respectively. The Ka values for carbonic acid are 8×10–7 and 1.6×10–10.

     

     

     

     

     

     

     

     

     

    Problem: Prepare 500 mL of a buffer with a pH of 9.87 such that the total concentration of buffering agents (conj. weak acid/conj. weak base) is 0.200 M. The following solutions are available to you

    6.0 M HCl, 6.0 M H3PO4, 6 M CH3COOH, 6 M NaOH

    and the following solids are available to you

    NaHCO3, Na3PO4•12H2O, NaH2PO4, Na2CO3, Na2HPO4, CH3COONa•3H2O

    Possibly useful information

    • acetic acid: pKa = 4.757
    • carbonic acid: pKa = 6.35
    • bicarbonate: pKa = 10.33 (only suitable choice)
    • phosphoric acid: pKa = 2.15
    • dihydrogen phosphate: pKa = 7.20
    • monohydrogen phosphate: pKa = 12.38

      

     

     

     

     

     

     

     

     

     

     

     

     

     

    Prepare 500 mL of a buffer with a pH of 9.87 such that the total concentration of buffering agents (conj weak acid/base pair) is 0.200 M.

    • The following solutions are available to you:
      • 6.0 M HCl
      • 6.0 M H3PO4
      • 6.0 M CH3COOH
      • 6.0 M NaOH
    • The following solids are available to you:
      • NaHCO3,
      • Na3PO4.12H2O,
      • Na2CO3,
      • Na2HPO4,
      • CH3COONa.3H2O
    • Possibly useful info:
      • acetic acid: pKa = 4.757
      • carbonic acid: pKa = 6.35
      • bicarbonate: pKa = 10.33
      • phosphoric acid: pKa = 2.15
      • dihydrogen phosphate: pKa = 7.20
      • monohydrogen phosphate: pKa = 12.38

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    More Buffer Problems

    1. A biochemist must prepare a solution as a medium in an experiment involving acid-producing bacteria. The pH of the medium must not change by more than 0.05 pH units for every 0.0010 mol H3O+ generated by the organisms per liter of medium. A medium consisting of 0.10 M HA and 0.10 M A– is prepared with a total volume of 1.0 L. Will the buffer capacity be sufficient for this experiment?
    2. A 1.00-L buffer is prepared that is 0.2000 M in the weak acid, HA, and 0.1500 M in the weak base NaA. The buffer has a pH of 3.35. What is the pKa for the weak acid? Is this buffer better at neutralizing strong acid or strong base? What is the buffer’s capacity to neutralize strong acid? What is the buffer’s capacity to neutralize strong base? What is the buffer’s pH of the buffer if 0.0015 mol NaOH is added to 0.5000 L of the buffer?
    3. An environmental chemist needs a carbonate buffer of pH 10.00 to study the effects of the acidification of limestone-rich soils. How many grams of Na2CO3 must she add to 1.5 L of freshly prepared 0.20 M NaHCO3 to prepare this buffer?

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    Characterizing a TRIS Buffer

    When microorganisms reproduce they release waste products that may change the pH of their environment in a manner that prevents their further reproduction. This can present a  problem in biology labs where microorganisms are grown in the closed environment of a culture dish. Culture media, therefore, are usually buffered. One common buffer is made using the weak base tris(hydroxymethyl)aminomethane, which is more commonly called TRIS, and its conjugate weak acid, TRISH+. Relevant information about this buffering system is provided here:

    TRIS is (HOCH2)3CNH2 and has a Kb of 1.19×10–6

    TRISH+ is (HOCH2)3CNH3+ and is available as the salt (HOCH2)3CNH3Cl

    1. What range of pH values are possible for buffers made using TRIS and TRISH+?
    2. At what pH will a TRIS/TRISH+ buffer have its greatest capacity for neutralizing against the addition of strong acid?
    3. Suppose you want to make a TRIS/TRISH+ buffer with a pH of 9.00 and that the equilibrium concentration of TRIS needs to be 0.100 M. What concentration of TRISH+ will you need?
    4. A TRIS/TRISH+ buffer is prepared by dissolving 50.0 g of TRIS and 65.0 g of TRIS•HCl in deionized water and diluting to 2.00 L. What is the pH of this buffer?
    5. What is the pH after adding 0.5 mL of 12.0 M HCl to a 200.0-mL portion of the buffer from Problem 4?
    6. What is the buffer capacity against the addition of strong base for a 200.0-mL portion of the buffer from Problem 4? Express your answer in terms of the maximum volume, in mL, of 6.0 M NaOH that can be added?

     

     

     

     

     

     

     

     

     

     

     

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