CHEM 215 L Preparation of a Buffer (Revised 11/2004)

 

Supplemental Reading:

Read Chapter 10 and 11 in Quantitative Analysis by Harris (6^th edition) for more details on buffer preparation.

See Appendix G, Harris for various K_a Values.

Manual for pH meter :  http://www.denverinstrumentusa.com/media/pdf/op-man-basic-ph-rev-g.pdf (see link  on CHEM 215 Webpage)

 

Introduction:

Buffers are solutions which contain reasonable amounts of a weak conjugate acid-base pair.   A buffer solution has the ability to resist large changes in pH, because when H₃O⁺ (aq) or OH^- (aq) is added to the solution is converted to the conjugate form of the weak conjugate acid-base pairs by one of the following equations:

H₃O⁺ (aq)  +  A^- (aq) D HA (aq) +  H₂O (l)

OH^- (aq)  +  HA (aq) DA^- (aq)  +  H₂O (l)

The pH of a buffer can be calculated by the Henderson-Hasselbalch equation:

pH = pKa + log

where           

n_B = moles (or mmoles) of the conjugate base

n_A = moles (or mmoles) of the conjugate acid

pKa = -log (Ka)  for the Ka of the conjugate acid

 

The buffer may be prepared (direct method) by mixing the appropriate moles of conjugate acid (n_A) with the appropriate number of moles of conjugate base (n_B).   If the conjugate acid and base are not both available, then the buffer may be prepared (indirect method) by partial conversion of the conjugate acid to the conjugate base (or vice versa) using a strong acid or base as appropriate.  Note:  Dilution of the buffer will cause slight changes in pH because of changes in ionic strength.  The buffer capacity depends on the actual moles of conjugate acid and base that are present. 

Objective:

Students will calculate and prepare buffers at a specific pH’s.  Students will apply theoretical knowledge of the Henderson-Hasselbalch equation and activity coefficients to buffer preparation and to pH calculations.

 

Experimental:

(Prepare any three of the following.)  Show your calculations to the instructor before proceding.  Calculate the number of grams of solid and mL of 1.0 M NaOH or 1.0 M HCl needed to prepare your buffer.

a.  Make a pH = 3.00 buffer using citric acid and sodium hydroxide.

b. Make a pH = 5.00 buffer using sodium acetate and hydrochloric acid.

c. Make a pH = 6.50 buffer using sodium citrate and hydrochloric acid.

d. Make a pH = 9.00 buffer using ammonium chloride and sodium hydroxide.

 

1. Prepare each buffer to be 0.050 M in the conjugate acid, and 100 mL total volume.  (Dilute to volume with deionized water.)  If both of your reagents are available as a solid, you should calculate the mass of each reagent you are going to mix.  If only one reagent is available as a solid, you will have to add 1.00 M NaOH or HCl to create its conjugate form in solution.  Show your calculation of the mass of reagent(s) and, if necessary, the volume of strong acid or base you will mix to create your buffer.  Show your calculations to the lab instructor before you proceed.

 

2. Calibrate your pH meter then measure the pH of your solution with a the pH meter, and record the pH in your notebook.  How close did you come to your target pH?  Can you think of some reasons why the actual pH of your buffer may differ from the theoretical pH?

 

3. Challenge- Calculate the ionic strength of your solution.  See if you can estimate the “true” pH of your solution using Equation 10-18 (page 190) of the Harris Text. 

 

4. The final step in buffer preparation is to bring the solution to the desired pH by the addition of strong acid or base.  Transfer your buffer to a 250 mL beaker and add enough 1.0 M NaOH or HCl from a buret, with magnetic stirring, to bring the pH to the target.  The new volume of your buffer is the original 100 mL plus the volume of acid or base you just added.  Calculate the new theoretical pH by calculating the new weak base to conjugate acid ratio and using the Henderson-Hasselbalch equation.  To get the new ratio, you will need to calculate the moles of strong acid or base you have just added.  Compare the new actual and theoretical pH.