CHEM 1112: Buffers and the Vampire Slayers
By K. Stone
CSU, Stanislaus
Introduction:
Acid-base chemistry is very important in living systems. Biochemical reactions are catalyzed by enzymes that have very narrow ranges of optimum pH values. The pH of a living system is maintained with buffers. In humans, the pH of blood is maintained between 7.35 and 7.45. Typically 5 liters of blood contains enough buffering capacity to absorb 150 mL of 1M acid. The principal buffers in blood are bicarbonate, proteins (including hemoglobin and oxyhemoglobin), and phosphates.
A buffer is a mixture of a weak acid and its conjugate base. Because the solute can either absorb protons or release protons, the buffer solution can maintain the pH at a constant value.
For example, when 10 mL of a 0.1M solution of HCl is added to 1.0 L of a 0.1M Acetic acid solution at pH 4.8, the equilibrium shifts to the left and more acetic acid is made. The added protons are now part of acetic acid. The amount of free hydrogen ions does not change, despite the fact that acid has been added. The new pH is 4.79.
Scheme 1. Acetic acid equilibrium
A buffer solution is at its maximum buffering capacity when the ratio of conjugate base to acid is 1. We will use the familiar equilibrium expression to demonstrate this.
Ka = [H+] [A-]/[HA]
This equation can be rearranged to:
[H+]= Ka [HA]/[A-]
By taking the log of both sides, Eq 2 becomes:
log [H+] = log Ka + log [HA] - log [A-] (3)
Since pH = -log [H+], we can rewrite the above equation into the form known as the Henderson-Hasselbalch equation:
pH = pKa + log [A-]/[HA] (4)
What happens to EQ 4 when the concentration of weak acid equals the concentration of its conjugate base? (When [A-] = [HA]?) What is the log of 1? At this point we see that pH = pKa. In order for a buffer to have maximum buffering capacity, the ratio of acid to conjugate base should be close to 1.0 and the pH of the solution will be close to the pKa of the weak acid.
Overview of the experiment:
There are two parts of this laboratory exercise. In part A you will investigate acid base equilibrium and determine the pKa of a weak acid using two methods. In part B of this exercise, you will prepare a buffer solution and test the buffering capacity of this solution. Each student will prepare their own buffer.
A. Titration of a weak acid with a strong base. You will use a pH meter along with methyl orange and phenolphthalein indicators to follow the titration of an acetic acid solution. You will be asked to compare pH readings and color transitions of the indicators. You will determine the pKa of acetic acid from this titration curve. You will also determine the pKa of acetic acid by measuring the pH of an acetic acid solution.
From the titration data, you will complete a plot of pH vs. mL NaOH titrant. In your conclusions indicate the pH at which you observe color changes of the methyl red and the phenolphthalein. From your pH measurement of the acetic acid solution before any titrant is added, you will determine the Ka of acetic acid. See scheme 1.
|
Ka = |
[H+][CH3CO2-] |
(5) |
|
[CH3CO2H] |
But [H+] = [CH3CO2-], therefore
| Ka = | [H+]2 | = | [H+]2 | (6) |
| [acetic acid] | 0.1- [H+] |
Procedure:
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
Before making any pH measurements, make sure that your pH meter is properly calibrated. Use the pH 7 (yellow) standard solution and adjust the intercept (set) knob until the meter reads 7.0, then use the pH 4 (pink) standard solution and adjust the slope knob until the meter reads 4.0. Return the electrode to the pH 7.0 solution and again adjust the intercept knob, alternate back and forth between pH 7.0 solution and the pH 4.0 solution until no adjustment is necessary.
A. Determining the Ka for acetic acid.
Use both of these methods to determine the pKa of acetic acid. Compare your results from
each method in your conclusion.
Method I.
Place 25 ml of the 0.10 M acetic acid solution in a clean 100 ml beaker. Determine the pH
of the solution using the pH meter. (Before making any pH measurements, make sure that
your pH meter is properly calibrated. Your lab instructor will provide you information on
calibration). Using the measured pH of the acetic acid solution, determine the pKa of
acetic acid. Be sure to use the equations that were given in the introduction. (Note: pH
does not equal pKa, pKa = -log Ka, calculate the Ka from the pH value and the starting
acetic acid concentration. The equilibrium concentration of acetic acid does not equal the
starting concentration.)
Method II.
Add 4-5 drops of methyl red and phenolphthalein indicators to your solution, insert the pH
meter probe, and begin your titration (Did you calibrate the pH meter?, Instructions are
at the top of the procedure section). Add 0.1 M NaOH in small increments of approximately
1 ml at the beginning. Decrease the increments of NaOH when you begin to approach the end
point. Near the end point you will wish to add dropwise (approximately 0.05 ml). Continue
your titration to a pH of 11-12 and plot your data; on your graph show the color
transition range for each of the indicators you have used. Determine the pH at the point
at which the acetate ion concentration equals the acetic acid concentration. (That is,
where the acetic acid is 50% ionized).
Calculate the Ka of acetic acid from your results and compare this to the value obtained from Method I. What is the known Ka for acetic acid? In your conclusion, restate all of the values that you determined and the literature value. Calculate the relative percent error for each value that you determined. Be sure to discuss the sources of error and how each error would affect your experimental value. For example if the pH meter was not calibrated correctly, and the actual pH was lower than the measured pH, would the calculated Ka be higher or lower than the actual Ka value? This is only one example, there are several others.
| % relative error = | I experimental value- literature value I | x 100% | (7) |
| literature value |
Part B. Preparation of a buffer, and testing its buffer capacity.
Choose an appropriate weak acid from the list below to make 0.5 liter of a 0.1M buffer solution that has a pH of 7.4. You are required to do your own calculations. There are several methods that can be used to make buffers, two of these methods are described below.
| Weak acid | Ka |
| acetic acid | 1.8 x 10-5 |
| phthalic acid | 1.3 x 10-3 |
| dihydrogen phosphate (monobasic) | 6.2 x 10-8 |
| monohydrogen phosphate (dibasic) | 4.8 x 10-13 |
| carbonic acid | 4.6 x 10-7 |
| citrate | 8.4 x 10-4 |
| dihydrogen citrate(monobasic) | 1.8 x 10-5 |
monohydrogen citrate (dibasic) |
4.0 x 10-6 |
In your notebook, be sure to show all calculations for determining the amounts of the weak acid and the conjugate base. Be sure to carefully describe the procedure that you used in the method section of your lab report.
Combining volumes method:
Since the total molarity of the solution must be 0.1 M, you can make up a 0.1 M solution
of the weak acid and a 0.1M solution of the conjugate base, then combine the appropriate
volumes of the two solutions to get the correct ratio of base/weak acid. (Use the
Henderson-Hasselbalch equation to calculate the ratio, and remember that the total volume
is 500 mL) A modification of this method that is used frequently, is to start with the
appropriate volume of either solution (you must first calculate the appropriate volume)
and add the other solution until the correct pH is obtained. For this you will need to
monitor the pH of the solution as you are adding the second solution. You will need a pH
meter, a stirring plate and a stir bar. You may end up with more or less than 500 mL, do
not add water to make up the volume.
Mass method:
Calculate the mass of the weak acid and the mass of the conjugate base that are required
to make a pH 7.4 solution that has a combined molarity of 0.1M.
Weigh out each substance and add distilled water to make a 0.5L solution.
Testing your Buffer.
Place 25.00 ml of your buffer in a clean 100 ml beaker. Determine the pH of the solution using the pH meter. Before making any pH measurements, make sure that your pH meter is properly calibrated.
Add 4-5 drops of phenolphthalein indicator to your buffer solution, insert the pH meter probe, and begin your titration. Add standardized NaOH in small increments of approximately 0.5 ml at the beginning. Decrease the increments of NaOH when you begin to approach the end point. Near the end point you will wish to add dropwise (approximately 0.05 ml). Continue your titration to a pH of 11-12 and plot your data; on your graph show the color transition range for the indicator you used. From the end point you will determine the buffering capacity of your buffer solution. How many moles of NaOH did you add to use up all of the weak acid. Calculate the theoretical amount of base that you should have been able to add. (Hint: how much of the acid form was present in the solution before the base was added? When the acid is gone, there is no more buffering capacity.) In your conclusion, compare the actual buffering capacity of your buffer with this theoretical buffering capacity. If these values differ by 10%, be sure to discuss sources of error.
Note: you do not need to use a pH meter for the titration of your buffer. If a pH meter is not available, just do the titration and look for the very pale pink end point. You can get the pH of your buffer when a pH meter becomes available.
Results, calculations:
Ka of Acetic acid
 | Calculate the Ka for acetic acid using the pH of an acetic acid solution. |
| Show the graph of pH vs mL NaOH added. |
| Show how you used this graph to determine the Ka |
Buffers
| Clearly show your calculations for making a buffer. |
| Include a graph of pH vs mL of NaOH added to your buffer |
| Report the initial pH of your buffer. |
| Calculate the actual buffering capacity of your buffer solution. |
| Calculate the theoretical buffering capacity of your buffer solution. |
Conclusion:
Using well organized paragraphs, restate the results of your experiment and answer the following questions.
| How does the experimentally determined Ka for acetic acid relate to the literature value? |
| Which method gave better results? |
| What weak acid did you choose to make your buffer, why did you select this one? |
| What was the initial pH of your buffer? |
| How does this compare with the solution that you were trying to make? |
| What was the buffering capacity of your buffer solution? |
| How does this compare to the theoretical buffering capacity? |
Additionally, address these questions in your
discussion.
You had to use the standard NaOH solution from the stockroom to determine the buffering
capacity of your buffer. Why can't you use your NaOH solution from last week? (Hint there
is CO2 in the air above the NaOH solution, see EQ 8 below.)
CO2 + H2O <----> H2CO3
<-----> H+ + HCO3- (8)What weak acid would you select if you were going to do experiments with intestinal lumen cells?
Last edited by K.Stone with assistance from John Burt on 09/04/00.