CH 150: Introduction to Biochemistry

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Buffers and Amino Acids

Objective: The objective of this exercise is to reinforce your understanding of buffer systems and to learn the use of a pH meter.

Introduction:

        Buffers are chemical systems that are resistant to the change in the pH of a solution.  They are composed of a weak acid (HA) and its conjugate base (A^- ). By converting between the acid and base form of the molecule by accepting or donating a proton from solution, the buffer can help control the pH of the solution.

 HA « H⁺ + A^-

When acid (H⁺) is added to a buffer, the equilibrium shown above is shifted to the left, according to LeChatilier’s Principle), as the available conjugate base (A^-) accepts the added H⁺, forming the acid (HA).  Similarly, when base (OH^-) is added, the H⁺ in the solution is consumed, forcing the equilibrium to the right, dissociating HA to form more H⁺ and A^-.

        The equilibrium expression for the above acid dissociation equation is given as

where [H+], [A-], and [HA] are the molar concentrations of each species in solution, and K is the equilibrium constant for the reaction.  By taking the log of each side of the equation and rearranging the result, the Henderson-Hasselbach equation is derived.

The pKa is the –log (Ka), and the [A-] and [HA] are the molar concentrations of each species in the solution.  The pKa for each acid is a constant value for a given acid at 25 ^oC, and is available in tabulated forms. 

         Keep in mind that the [HA] in a solution is not necessarily equal to the concentration of the material given.  For example, if an acid has a concentration of 0.1 M, that means that the total concentration of the acid is 0.1 M.  Since weak acids do not dissociate 100 %, this value is the sum of the acid present in the dissociated form [A^-] and the undissociated form [HA], or

 X M = [HA] + [A^-]

Since every A^- results from the dissociation of one HA molecule, the [H⁺] is equal to [A^-]

[H⁺] = [A^-]

         A buffer is effective in a pH range equal to ± 1unit from the pKa.  For example, if the pKa of a buffer is 4.76, then the solution will function as a buffer (i.e. maintain a near constant pH) from pH 3.76 to pH 5.76.  In other words, the system can absorb many H⁺ or OH^- ions and maintain nearly the same pH.

         The buffering capacity of a solution is defined as the amount of H⁺ or OH^- a solution can  “absorb” and maintain the pH at a nearly constant level.  This is dependant on the concentration of HA and A^- present in the solution.  Once all the A^-, for example, has been consumed in a reaction by added H⁺, the buffering capacity of the solution has been exceeded. 

         In this experiment, you will create buffers of different buffering capacities by adding various ratios of the acid (HA) form and the conjugate base (A^-) form of either the tris buffering system of the imidazole system.  You will add acid (HCl) or base (NaOH) to each buffer you generate and measure the pH, comparing the measured values to those calculated based on the Henderson –Hasselbach equation.   Lastly, you will titrate a glycine solution by adding NaOH to a solution containing only the undissociated, acid form of glycine.  By plotting the volume (milli-equivalents) of OH^- added against the pH, you will see the effective buffering range(s) of glycine.

 Preparation before Laboratory:

Part 1:

 1.      Determine for your acid-base pair which is the acid component and which is the base component.

2.      Calculate the amount of material (g) that will be required to prepare the 100 mL of each 0.25 M solution.  Fill in amount on data tables.

3.      From theory and literature pKa values, calculate the exact pH for each of the mixtures.  Do on a separate sheet of paper or on a spread sheet and hand in with this report.  Fill in answers in the data tables.

4.      Repeat step 3 for the mixtures plus HCl (or NaOH), the mixtures prepared from the diluted stock solutions, and the diluted stock solutions plus HCl (or NaOH).  Ignore the volume change that results from the addition of 0.1 mL of HCl or NaOH to 10 mL solution.  Fill in answers in data tables.

5.      Calculate the theoretical pH of the 10 mM KCl solution with and without added HCl (or NaOH).  Fill in answers in data tables.

Part 2:

 6.      Draw the three possible ionic forms of glycine and identify which form predominates at pH 1, pH 6, and pH 11.   Which does not exist in water?

7.      What are the two pKas for glycine?

8.      Write the chemical equations showing how glycine can buffer solutions at both pKas by consuming added H⁺ or OH^-.

Part 1:  Buffer Systems

 Experimental Procedure:

a. Preparation of Solutions:

Each laboratory group will prepare 100 mL solutions of one of the acid-base pairs (HA for acid, A- for base) listed below.  The final concentration of the solutions should be 0.25 M in terms of the buffer component.  Thus, each group will have two solutions: 100 ml of 0.25 M Tris[1] and 100 mL 0.25 M Tris-Cl.  Weigh the appropriate amounts of each component and dilute each with 10 mM KCl solution[2] to the final volume.

 Acid - Base Pairs

                        Tris (MW 121.1)                                Tris-HCl (MW 157.6)

                        Imidizole (MW 68.06)                       Imidizole-HCl (MW 104.54)

 b. Experimental Set up:

Prepare mixtures (10 mL final volume for each) of your acid-base pair with the following compositions:

             (a) 100 % HA

            (b) 75 % HA, 25 % A-

            (c) 50 % HA, 50 % A-

            (d) 25 % HA, 75% A-

            (e) 100 % A-

Measure the pH of each of these solutions.

 Add 0.1 mL of 2 M HCl (or 0.1 mL of 2 M NaOH) to each of these solutions.  Mix thoroughly and measure the pH of each.

 Dilute a portion of each 0.25 M stock solution (acid and base) by a factor of 5 using 10 mM KCl as the diluent.  Prepare mixtures of these dilute solutions as above (a-e), measure the pH of each solution, add 0.1 mL of 2 M HCl  (or 0.2 M NaOH), and measure the pH again.

 Measure the pH of 10 mL of 10 mM KCl before and after the addition of 0.1 mL of 2M HCl (or 2 M NaOH).

Data Tables:

Acid (HA) component _________________

Weigh out _____ g  to dissolve in water to make 100 mL of a 0.25 M solution.

Base (A-) component _________________ 

            Weigh out _____ g  to dissolve in water to make 100 mL of a 0.25 M solution.

Table 1: Buffer pH Measurements

Table 2: Buffer Test

Add 0.1 mL of 2 M  _____ (HCl or NaOH)

Table 3: Diluted Buffer pH Measurements

Table 4: Diluted Buffer Test

Add 0.1 mL of 2 M  _____ (HCl or NaOH)

Results and Discussion:

1.      Compare your calculated pH values to those observed and offer explanations for any discrepancies. 

2.      Draw the structures of the acid form and base form of your acid base pair.

3.      Using your data to support your statements, discuss the effectiveness of your buffer solutions in controlling pH. Briefly, extend your conclusions to buffers in general.

Part 2: Titration of Glycine

 Experimental Procedure:

Obtain a 40 mL sample of 1 M glycine.  Read and record the pH of the solution. Set up a buret containing 2.5 M NaOH. Set up the LAb-works pH metera nd computer recording sation. Add 2.5 M NaOH drop-wise, with Lab Works recording the pH after each addition until the pH of the solution reaches 12.

Save your data to a disc.  Open the data in Excel. Turn in the following:

a. Data: # drops NaOH added, vol. NaOH Added, and pH

b. graph: vol NaOH (x-axis) added vs pH (y-axis)

Results and Discussion:

1.    From the graph, determine the pKa values for glycine.

 2.      How well do your measured pKa values compare to the theoretical values? 

3.      Compare the amount of glycine present with the amount of titrant needed to change the solution from the first pKa to the second.

4.      What is the isoelectric point (pI) of glycine?