Make Your Own Darn Buffer !

In this experiment, you will be required to prepare a buffer solution that has a specific pH value. By definition, a buffer is a solution that is resistant to changes in pH, even when small amounts of a strong acid or base are added to the solution. This ability to maintain a constant pH is a result of a buffer containing both components of a weak acid/base pair. When a strong acid is added, the weak base in the buffer accepts the excess protons. When a strong base is added, the weak acid in the buffer donates protons to neutralize the base. For this experiment, you will prepare a buffer that contains acetic acid and its conjugate base, the acetate ions. The equilibrium equation for the reaction is shown below: + - HC2H3O2(aq) + H2O(l) <=> H (aq) + C2H3O2 (aq)

-5 The equilibrium expression for this reaction, Ka, has a value of 1.8 x 10 at 25ºC.

   [H ][C2H 3O2 ] Ka [C2H 3O2 ] Ka  , w h e r e [ ] r e p r e s e n t s m o l a r i t y . R e  a r r a n g e d ,   [HC 2H 3O 2] [H ] [HC 2H 3O2]    pH Ka [C2 H3O2 ] S i n c e t h e [H ]  10 , pH  [10 ] [HC2 H 3O2 ]

Therefore, the ratio between the molarity of the acetate ions to the molarity of the acetic acid in - assigned pH your buffer must equal the ratio between the Ka value and 10 . This ratio should be - reduced , so that either the [HC2H3O2] or [C2H3O2 ] has a concentration of 0.10 M, and the concentration of the other component must fall within a range from 0.10 M to 1.00 M. Dimensional analysis can be used to determine the quantities of each reactant needed to prepare 100.0 mL of your buffer solution that has these specific concentrations. The success of your efforts will be determined by how close the actual pH value of your buffer matches the desired pH value, and whether or not the pH remains fairly constant when a strong acid or base is added to the buffer solution.

Procedure: 1. Obtain an assigned pH value and list of given reactants for the buffer you are to prepare.

2. Calculate the quantities of reactants needed to prepare 100.0 mL of your buffer solution that has a concentration of each component within the range of 0.10 M to 1.00 M. Use the lowest possible concentrations for the acid and conjugate base in the buffer.

3. Use volumetric pipets, a Mettler balance (+ 0.001 g), and a 100.0 mL volumetric flask to prepare your buffer solution.

4. Label three small beakers as 1, 2 and 3. Place 25.0 mL of your buffer into beaker 1, and 15.0 mL into each of the other two beakers. Add 10.0 mL of 0.10 M HCl(aq) to beaker 2, and 10.0 mL of 0.10 M NaOH(aq) to beaker 3.

5. Have your teacher check the pH of the solutions in all three beakers. Your buffer solution must match your assigned pH value (+ 0.5). The pH of the solutions in the other two beakers must be within a range of + 1.0 of the pH value for your original buffer solution.

Results & Calculations: 1. Explain the reasoning used to select the proportions of the components needed to make your buffer.

2. Show all calculations used to determine the quantities needed to prepare your assigned buffer.

3. Calculate the theoretical pH of the mixture made with 15.0 mL of your buffer and 10.0 mL of the 0.10 M HCl(aq). Determine the percent error based upon the measured pH value.

4. Calculate the theoretical pH of the mixture made with 15.0 mL of your buffer and 10.0 mL of the 0.10 M NaOH(aq). Determine the percent error based upon the measured pH value.

Conclusions: Discuss the accuracy of your experiment and any sources of experimental error.