The Titration of of an Amino Acid

Glycine is an example of an amino acid.  The protonated from of glycine has the structure                          

or

or H2A+

 Its reaction with OH- occurs in two steps:

   (H2A+)     

       (HA)   

       (A-)

  The dissociation of of the two acid forms occur according to the following equilibria:

H2A+(aq) + H2O(l)  H3O+(aq) + HA(aq)

Ka1

HA(aq) + H2O(l)  H3O+(aq) + A-(aq)

Ka2

The titration curve of a diprotic acid such as glycine, is shown below.

The first equivalence point, corresponding to an amount of added base equal to Vep1, occurs when all of the acid H2A+ is converted to HA according to the reaction:

OH- + H2A+ HA + H2O

The species HA has the structure 

 

It is a  neutral molecule, and the pH where its concentration is a maximum is the isoelectric point.  This is labeled pHIP , on the graph.

Half-way to the first equivalence point, Vbase = Vep1/2, and (H2A+) = (HA).  The equilibrium constant expression yields:

 

and  pH = pKa1.  

At the second equivalence point, at Vep2 , all of the species HA is converted to A- according to the reaction:

OH- + HA A- + H2O

   Half-way between the two equivalence points, Vbase = (Vep2-Vep1)/2 , (HA) = (A-).  The equilibrium constant expression yields:

and pH = pKa2.

The purpose of this experiment is to measure Ka1 and Ka2 for the protonated form of glycine. 

Special Equipment: Things to use and return on the same day.

Procedure

  1. Wastes and Spills: all of the chemicals in this experiment can be put in the trash or down the drain. Be sure the water is running. Use wet paper towels to wipe up small spills. Thoroughly rinse the pipet, buret and other glassware with water before returning them to the stockroom.  You will want to clean all of your glassware before beginning this experiment. Be sure to triple rinse with distilled water and then triple rinse with the solution you are using before filling burets, beakers or pipets.

  2. Calibration of pH electrode:  Turn on the computer and the labworks interface box.Click on  the Labworks icon on the desktop.  Click on Calibrate, followed by pH.  Follow the instructions on the screen.  A reference solution of known pH for the calibration will be provided by the Stockroom.  The pH electrode is very fragile and very expensive.  It must be kept immersed in water at all times.  Rinse it with deionized water from your wash bottle when moving it from one solution to another.
  3. Place 25 mL of the 0.1 M glycine solution into a clean 100 mL beaker.  Place a stir bar in the beaker, and place the beaker on the stirrer.  Carefully place the calibrated pH electrode in the acid solution.  Position it so that it does not touch the rotating stir bar.  Fill your buret with 0.100 M NaOH solution.  Be sure to expel any air bubbles from the tip, and adjust the volume of base so that the meniscus is at zero on the buret scale.  Use a buret clamp to position the buret above the beaker of acid.  The buret tip should be below the top of the beaker.  

  4. On the computer screen, click on Design, and under File options, click on open an existing file, followed by OK.  Double click  on the folder "shortcut to experiments" and select the experiment "titration3.exp", and open it.  Click on Aquire, and start.  The prompt "enter total volume added" will appearClick OK.   The next prompt asks you to "enter a number ". A zero should appear in the box to the right of the prompt.  Click OK. A pH value will now be displayed on the screen.  This is the pH of the glycine solution before any base is added.

  5. The next step is to obtain data for the pH versus mL plot needed to determine the Ka1 and Ka2  a) In response to the prompt "add more base", click OK.  The prompt "enter total volume added" will reappear.  Click OK.  b) When the prompt "enter a number " appears, add 0.50 mL of base to the acid, and enter "0.50" in the box to the right of the prompt.  Click OK.  The new pH and volume of base will appear on the screen. Repeat this process for each addition of base.  After each addition of base, enter the total volume of base added at that point.  Continue adding the base in 0.50 mL increments until the pH is about 5. At this point, reduce the increment to 0.1 mL or less. Near the equivalence point, a pH of about 6, reduce the volume increment to a single drop. Once the first equivalence point is passed you can increase the volume increment to 0.5 mL, but reduce it again to one drop as the second equivalence point, at about pH = 11,  is approached. Continue the titration to a pH of 11 to 12. 

  6. Stop the data collection by selecting cancel in response to the prompt "add more base" and yes in response to "Do you want to stop the program now?

  7. Save your data in a file in the "Personal" folder in the desktop.  

  8. Click on Analyze in the main Labworks menu.  This will open a spreadsheet program which you can   use to prepare your graphs.  The pH and volume data will be in columns A and B, respectively.  To To calculate the slope of the pH curve versus mL, click in column C, click on column setup on the menu bar, and enter the function DERIV(A,B).  This will calculate D A/D B for each pair of data points.  To plot the pH and the slope versus mL, select Graph setup on the menu bar.  Choose B for the X axis, A for the Y1 axis and C for the Y2 axis.  Click OK, and the desired graphs will appear.

  9. Save your spreadsheet and graph to a diskette.  To obtain a hard copy of your spreadsheet and graph. It will necessary to import the file into Excel and use a computer with access to a printer.

  10. Repeat the titration.  Based on the location of the equivalence points obtained in your first trial, try to improve the  accuracy of your measurements by using drop-sized volume increments in the vicinity of each equivalence point.

Calculations

  1. From your titration graph determine the values of Vep1, Vep2 Ka1 and Ka2. Clearly label these quantities on a full- page-size graph of your titration curve.  Also label the isoelectric point.

  2. Use your value of Ka1 to calculate the initial pH (before any base is added) of the 0.100 M glycine solution.


Conclusion

  1. Report your values of Ka1 and Ka2, and compare them with values from the literature. What are the percentage errors?

  2. What is the pH at the isoelectric point of glycine?

  3. Compare your calculated value of the initial pH of the glycine solution with the measured value.  What is the percent error?