PRELIMINARY READING: Zumdahl, sections 15.2, 15.3
 

INTRODUCTION

A buffer solution contains a conjugate acid-base pair, i.e., two species whose formulas differ by only one H⁺. This pair could consist of a weak acid and its salt, e.g., acetic acid (HC₂H₃O₂) and sodium acetate (NaC₂H₃O₂), or a weak base and its salt, e.g., ammonia (NH₃) and ammonium chloride (NH₄Cl).

Buffer solutions have the useful property of resisting changes in pH. This is because a buffer solution contains both an acidic and a basic component. If any acid should be added to a buffer solution, the basic component would react with it; if any base should be added to a buffer solution, the acidic component would react with it.

Let us examine the equilibria occurring in a buffer solution containing the generic weak acid HA and its salt NaA. Since the salt is a strong electrolyte it is completely dissociated:

                                         NaA(aq)       Na⁺(aq)   +   A^-(aq)                                                    (1)

The weak acid exists in equilibrium with its ions:

                                        HA(aq)       H⁺(aq)   +   A^-(aq)                                                          (2)

The ionization constant for the acid is given by

                                          K_a  =  [H⁺][A^-] / [HA]                                                                         (3)

In the buffer solution, the principal contribution to the [A^-] concentration is from the salt. Therefore, assuming negligible ionization of the acid,

                                          K_a  =  [H⁺] [salt]_o / [acid]_o                                                                   (4)

where the subscripts indicate initial concentrations. As detailed in experiment 7, the Henderson-Hasselbalch equation may be derived from this expression:

                                          pH  =  pK_a   +   log([salt]_o / [acid]_o)                                                     (5)

This equation becomes more general if the term "conjugate base" is substituted for "salt" (A^- is the conjugate base of HA):

                                         pH  =  pK_a   +   log([conj base]_o / [conj acid]_o]                                      (6)

In this form the equation may also be used for buffer solutions containing a weak base and its salt, where the salt is the acidic component.

In practical buffers the concentrations of the two components are generally similar to one another. Thus, the pH of the solution is close to the pK_a value, but can be varied somewhat by adjusting the ratio [conj base] / [conj acid]. In this experiment you will prepare buffer solutions and test the effects of added acid or base upon the pH.
 

PRE-LAB EXERCISE

  1. Calculate the pH of a solution that is 0.120 M in acetic acid, HC₂H₃O₂ , and 0.210 M in sodium acetate, NaC₂H₃O₂. The K_a of acetic acid is 1.76 x 10^-5.

  2. How many grams of sodium acetate must be added to 100 mL of 0.150 M acetic acid in order to produce a solution with a pH of 4.620? Assume that no volume change occurs.
 

EQUIPMENT: three 150-mL or 250-mL beakers, 50-mL graduated cylinder, pH meter with 400-mL beaker and wash bottle.  Burets to dispense the three weak acids, the strong acid solution, and the strong base solution.
 

EXPERIMENTAL

A. Preparation of the Buffer Solution

You will be assigned a weak acid solution to use as starting material. The instructor will assign you a particular pH of the buffer solution to be prepared. The buffer is to be prepared by treating the acid solution with a NaOH solution to partially neutralize the acid and form the conjugate base:

                                             HA(aq)   +   NaOH(aq)      NaA(aq)   +   H₂O(l)                             (7)

The goal is to produce the proper ratio of [A^-] / [HA] to produce the desired pH. Note that this ratio is also equal to (no. mmols A^- / no. mmols HA). Assume that 20 mL of the NaOH solution is to be added to the proper volume of HA. The number of millimoles of A^- produced by reaction (7) will equal the number of millimoles of NaOH added and will also equal the number of millimoles of HA used up in the reaction.

The volume of HA used must contain the number of millimoles of HA present in the final buffer plus the number of millimoles to be converted to A^-. Knowing the total number of millimoles of HA needed to make the buffer, calculate the volume of acid required. Do your calculations in your laboratory notebook. When you have completed the calculation, mix that volume of acid with 20.0 mL of the NaOH solution and measure the pH.
 

B. Effect of Added Acid

If a strong acid such as HCl is added to the buffer solution, the conjugate base A^- will react with the added acid:

                                             H⁺(aq)   +   A^-(aq)       HA(aq)                                                           (8)

In your notebook, calculate the pH expected if 1.0 mL of the supplied strong acid solution is added to 20.0 mL of your original buffer solution. Prepare this solution and measure and record the pH.
 

C. Effect of Added Base
 

If a strong base such as NaOH is added to the buffer solution, the acid HA will react with the added base:

                                         OH^-(aq)   +   HA(aq)       A^-(aq)   +   H₂O(l)                                          (9)

In your notebook, calculate the pH expected if 1.0 mL of the supplied NaOH solution is added to 20.0 mL of your original buffer solution. Prepare this mixture and measure and record the pH.
 

POST-LAB

Carry out the following calculations.

1. Calculate the pH expected if the strong acid in part B were added to 20.0 mL  of pure water instead of the buffer solution.

2. In a similar manner, calculate the pH expected if the strong base in part C were added to 20.0 mL of pure water instead of the buffer solution.

3. Suppose that you are given a 0.698 M acetic acid solution and a 0.500 M sodium acetate solution. Determine the volumes of each solution required to prepare 160 mL of a buffer with pH 4.65.
 

DISCUSSION

Discuss the relative error in each of your pH measurements and postulate reasons for the error. Discuss the results of parts B and C in light of your results for questions 1 and 2 of the post-lab section.