Chemistry 104 Study Groups
BUFFERS & TITRATIONS
 

CONCEPTS

 

Buffers

    A buffer is a solution containing a weak acid and its conjugate base (i.e. CH3COOH and CH3COO-) or a weak base and its conjugate acid (i.e. NH3 and NH4+).
    A buffer solution will maintain a relatively constant pH even when acidic or basic solutions are added to it.
    If 1 mL of 1 M HCl is added to 1 L of pure water, the pH will change from 7 to 3 (a change of 4 pH units). (show details)
    If 1 mL of 1 M HCl is added to 1 L of a buffer solution consisting of 500 mL of 0.2 M CH3COOH and 500 mL of 0.2 M CH3COO-, the pH will change almost imperceptable. The initial pH of this system is 4.74. (show details) After the HCl is added, the pH is raised to 4.75 (a change of only 0.01 pH unit). (show details)

Calculations involving buffers

    The initial pH of a buffer is determined by an equilibrium calculation. The concentration of both the weak acid and weak base must be considered.

    CALCULATE pH of a buffer containing 0.1 M NH3 and 0.1 M NH4.

    The pH of a buffer after addition of a strong acid requires consideration of the stoichiometry of the reaction (conversion of weak base to weak acid) and the equilibrium as follows:

    1. Calculate the amount of base which is converted into acid. Every mole of added acid converts one mole of weak base into its conjugate acid. Calculate the new concentrations of weak acid and weak base (include increase in volume if necessary).

      CALCULATE concentration of weak acid and weak base present in 1.0L of a buffer containing 0.1 M NH3 and 0.1 M NH4 after 1.0 mL of 10.0 M HCl has been added (neglect volume change).

    2. Use ICE to calculate the new [H3O+] or [OH-] as appropriate and calculate the pH

      CALCULATE [OH-] and pH of buffer solution after addition of 1.0 mL of 10.0 M HCl.

    The pH of a buffer after addition of a strong base requires consideration of the stoichiometry of the reaction (conversion of weak acid to weak base) and the equilibrium in the same manner as for the addition of a strong acid. In this case, the strong base converts some of the weak acid to the weak base

Titrations

    A titration is a reaction between two solutions. The properties and concentration of one of the solutions are known(titrant) and are used to determine the properties and/or concentration of the other solution. Three different kids of acid/base titrations may be performed:

    1. Strong Acid/Strong Base (either may be titrant)
    2. Weak Acid/Strong Base (strong base is titrant)
    3. Weak Base/Strong Acid (strong acid is titrant)
    Calculating the pH at any point in a strong acid/strong base titration requires consideration of the stoichiometry of the acid/base neutralization reaction.

    H3O+(aq) + OH-(aq) ---> 2H2O(l)

    Calculating the pH before the equivalence point in a weak acid/strong base or weak base/strong acid titration requires two parts:

    1. Consideration of the stoichiometry of the acid/base neutralization reaction.

      HA(aq) + OH-(aq) ---> A-(aq) + H2O(l)

      or

      B(aq) + H3O+(aq) ---> BH+(aq) + H2O(l)

    2. From the stoichiometry, the amount of acid-HA and conjugate base-A- (or base-B and conjugate acid-BH+) present in the system may be determined. The new concentrations of acid-HA and conjugate base-A- (or base-B and conjugate acid-BH+) are used in an ICE calculation:

      HA(aq) + H2O(l) ----> A-(aq) + OH-(aq)

      or

      B(aq) + H2O(l) ----> BH+(aq) + H3O+(aq)

    At the stoichiometric point, all of the acid-HA has been converted to its conjugate base A- (or base-B has been converted to its conjugate acid BH+). The pH can be calculated by first calculating the concentration of the conjugate base (or acid) and then using that value in an ICE calculation:

    A-(aq) + H2O(l) ----> HA(aq) + OH-(aq)

    or

    BH+ + H2O(l) ----> B(aq) + H3O+

    After the stoichiometric point, the pH of the solution depends only on the amount of excess base (or acid) present.
 
 

Exercises

  Consider methylamine (CH3NH2 Kb = 4.4 x 10-4). What is the pH of a 0.20 M solution of methylamine? What is the pH of a solution which is 0.20 M in methylamine and 0.20 M in the methylamonium ion?

 
  Consider a solution 0.20 M solution of methylamine.
  1. What mass of methylamonium chloride would be added to 250.0 mL of 0.20 M solution of methylamine to produce a solution with a pH of 10.0?

  2. What mass of methylamonium chloride would be added to 250.0 mL of 0.20 M solution of methylamine to produce a solution with a pH of 11.3?

  3. Comment on the difference between the mass of methylamonium chloride added in parts a and b.

 
  The solutions in a and b of the previous exercise are buffers. Why are they buffers? Which solution would be a better buffer against added acid? Which solution would be a better buffer against added base?

 
  Calculate the pH of the solution in 2a after 0.1 g NaOH has been added (assume no volume change). Calculate the pH of the solution in 2b after the same amount of strong base has been added.

  What ratio of concentrations of methylamine to methylamonium ion would be necessary to produce a solution that would buffer at pH of 9.9?

  Consider the titration 20.0 mL of a 0.20 M solution of ethylamine with 0.10 M HCl. Calculate:
  1. The pH before any titrant has been added.

  2. The pH after 10.0 mL of titrant has been added.

  3. The pH after 20.0 mL of titrant has been added.

  4. The pH at the stoichiometric point.

  5. The pH when 2.0 mL of titrant has been added, past the stoichiometric point.

  6. What volume corresponds to a position half-way to the stoichiometric point? What is the significance of this point?

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