BUFFER SOLUTIONS LAB
Purpose:
-
To prepare a buffer of a given pH with a reasonable degree of
accuracy
To predict what quantity of acid or base is necessary to
the buffering ability of your solution
To predict what indicator is necessary to indicate when
the buffer is “broken”
“break”
Background Information:
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 or ion 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 appreciably, despite the fact that acid has been
added. The new pH is 4.79.
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] (1)
This equation can be rearranged to:
[H+]= Ka [HA]/[A-] (2)
By taking the log of both sides, Equation 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 HendersonHasselbach equation:
pH = pKa + log [A-]/[HA] (4)
What happens to equation 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. A buffer is “broken” when the amount of acid or base in solution has
been completely neutralized stoichiometrically by a strong acid or base, called the equivalence point, and it can
no longer buffer changes in pH.
Pre lab Questions:
1. Using the Henderson-Hasselbach equation, calculate the theoretical pH for a 1 liter buffer containing
0.100 M propanoic acid, and .100M sodium propanoate.
2. Calculate how the pH will change in problem 1 if 25 mL of a .1M solution of HCl is added.
3. Calculate the pH of a buffer solution made by adding 10.0 g of anhydrous sodium acetate (NaC2H3O2) to
100 mL of 0.100 M acetic acid. Assume there is no change in volume on adding the salt to the acid. pKa
for acetic acid is 4.74.
Materials:
Acid-base pair combinations, burette, 250 mL or 500 mL Erlenmeyer flasks, appropriate indicators, standardized
solution for titration, logger pro, Labpro device, computer, pH probe
Mr. Jauss will assign you a specific pH to target through the creation of your buffer. You will use one of
the three conjugate acid-base pairs listed below to build your buffer. Make sure you choose the right system for
the pH you are shooting for. Note: read labels to determine state of hydration of the salts.
1. Sodium bicarbonate Assume Ka = 4.8 x 10-11 and Sodium carbonate
2. Glacial acetic acid (pure CH3COOH – 17.4M) Ka = 1.8 x 10-5 and Sodium acetate NaC2H3O2
3. Sodium or potassium phosphate (monobasic) NaH2PO4 or KH2PO4 Assume Ka = 1.2 x 10-7 and sodium or
potassium phosphate (dibasic) Na2HPO4 or K2HPO4
4. Concentrated ammonia NH3 (13.4M) Assume Kb = 1.8 x 10-5 and ammonium chloride NH4Cl
Procedure:
Your first task will be to make a buffer solution of defined pH. Part of your grade will be based on how close
your solution comes to the desired pH. Your second task will be to titrate your buffer using 0.5 M NaOH or 0.5 M
HCl until the buffer “breaks”. You will predict what quantity of acid or base is needed to break the buffer (as
judged by a color change of a chosen indicator). Each pair in the class will be responsible for making a buffer
solution of a specific pH.
To create these solutions, you will mix together a certain amount of a particular acid or base with a certain
amount of its conjugate base or acid. The ratio of both will determine the pH that will result. The choice of acid
or base depends on what pH you are assigned. Try to use reasonable quantities of acid or base and conjugate base
or acid - somewhere between one and five grams of each chemical would be great, and an appropriate amount of
liquid. Calibrate the pH meter with the standard reference buffers provided to make sure that the meter has not
drifted in its readings.
Once you have made your particular solution, you will perform the following tasks in order to obtain a
signature from Mr. Jauss. Make sure you attach the signature sheet to your lab report!
1. Test the solution with a pH meter to see how close you have come to your desired pH. You must demonstrate
this to Mr. Jauss.
2. Figure out what would be an appropriate indicator to test when your buffer has been “broken”, and what color
change you expect. Report this indicator to Mr. Jauss.
3. Predict the quantity of acid or base needed to “break” the buffer, and the pH this will occur. Report
this quantity and this pH to Mr. Jauss.
4. Perform a titration, using a standard concentration of HCl or NaOH, recording the volume needed to “break”
the buffer. This will be the solution’s buffering capacity, or the point where there is no acid or base left in
solution. Use the file marked “Buffer Lab” on the K drive.
5. When conducting the titration, make sure you use a standardized solution, wash your burette several times
with the solution, and make sure your pH probe is properly calibrated. Use experiment 19 in logger pro called
“Buffer” to run your titration. Make sure that you use half drops of standardized solution slowly through the
equivalence point, as the pH will change very rapidly through this area.
6. Make sure you save your titration graph, title it, and print it for reference.
Data:
Create an appropriate data table that illustrates all collected and calculated data during the experiment.
Calculations/Conclusions:
1. Show all calculations performed to determine the concentrations of weak acid or base, and the corresponding
salt, that were necessary to create your buffer solution at the assigned pH.
2. Write net ionic equations that demonstrate how your buffer system resists changes in its pH when strong acid
(H+1) or strong base (OH-1) is added.
3. Show all calculations performed to determine the amount of HCl or NaOH that was required to reach the
equivalence point and “break” or titrate the buffer.
4. Show all calculations performed that allowed you to determine the pH of your equivalence point, when no acid
or base was left, based on the salt that then underwent hydrolysis.
5. What indicator did you select to indicate the “breakage” of the buffer?
6. Calculate the pH of your solution if 5 mL of standardized acid or base solution had been added to your buffer
solution. Then, calculate the pH of your solution if 5 mL of standardized acid or base solution had been added
and there was no corresponding analogous salt added initially to your acid or base (meaning, you just had an
acid or base – not a buffer containing the analogous salt).
7. Using number 6 as a guide, what is the advantage to using a corresponding salt to a weak acid or base, versus
having the weak acid or base as a lone buffer?
VERIFICATION TABLE FOR BUFFER LAB
DESIRED PH OF
BUFFER:
DESIRED
QUANTITY OF
ACID OR BASE TO
BREAK OR
TITRATE THE
BUFFER:
FINAL PREDICTED
PH AFTER BUFFER
IS TITRATED:
ACTUAL PH OF
BUFFER:
ACTUAL
QUANTITY OF
ACID OR BASE TO
BREAK OR
TITRATE THE
BUFFER:
ACTUAL PH AFTER
BUFFER IS
TITRATED:
SIGNATURE:
SIGNATURE:
SIGNATURE:
DESIRED
INDICATOR AND
COLOR CHANGE:
SIGNATURE: