CHM 152LL: Exploring Buffer Solutions
Pre-lab: In addition to the typical pre-lab assignment, you should also complete the calculations at
the top of page 5.
Introduction
Buffer solutions resist changes in pH upon addition of small amounts of acid or base. For example, blood
acts as a buffer to absorb small amounts of acid or base resulting from biochemical reactions while
maintaining a pH close to 7.4 because cells can only survive in a narrow pH range around 7.4. Thus,
buffers are essential to life itself.
A buffer solution is able to resist pH changes because it contains a conjugate acid-base pair that can
neutralize small amounts of added acid or base. Typically, a buffer consists of a weak acid and a salt
containing its conjugate base, or a weak base and a salt containing its conjugate acid. Let’s say we have a
mixture of hydrofluoric acid and potassium fluoride salt. Potassium fluoride is a soluble salt so it dissolves
completely in water:
KF(s)  K+(aq) + F-(aq)
Hydrofluoric acid is a weak acid that only partially ionizes in water: HF(aq)  H+(aq) + F-(aq)
Mixed together the acid (HF) and its conjugate base (F-) set up the above equilibrium (potassium ions are
just spectators). Note that a weak acid and its conjugate base will not react with each other but instead
form an equilibrium mixture that contains both the reactants (HF) and products (F-).
However, when a different acid is added to the equilibrium mixture, the fluoride ions (basic ions) neutralize
it. When a different base is added to this mixture, the hydrofluoric acid neutralizes it. Yet for both these
situations the pH remains basically constant. Let’s explore this more.
The pH of a solution depends on the concentration of the hydronium ions (H3O+) and hydroxide ions (OH–)
present, so a buffer solution must be able to neutralize both these ions to maintain the pH. To neutralize
OH– from a base, an acid must be present. For example, if NaOH is added to the HF-KF buffer system
described above, the OH– (from the added NaOH) will be neutralized by the weak acid, HF(aq), as shown
by the following neutralization reaction equations:
molecular: HF(aq) + NaOH (aq) → H2O(l) + NaF(aq)
.
+
+
total ionic: HF(aq) + Na (aq) + OH (aq) → H2O(l) + Na (aq) + F-(aq) .
net ionic: HF(aq) + OH–(aq) → H2O(l) + F–(aq)
Thus, a buffer solution containing a weak acid will be able to neutralize any OH– added to it as long as the
amount of OH– added does not exceed the amount of the weak acid present. If a small amount of OH– is
added to a buffer the pH doesn’t change much since the ratio of [F–]/[HF ] only changes a little.
Similarly, to neutralize H3O+ from an acid, a buffer solution must contain a weak base. For example,
fluoride (F–), the conjugate base of HF(aq), will react with H3O+. The net ionic equation for the
neutralization reaction is as follows:
F–(aq) + H3O+(aq)
→
HF(aq) + H2O(l)
Thus, a buffer solution containing the conjugate base of a weak acid will be able to neutralize any
hydronium ion (H3O+) added to it as long as the amount of H3O+ added does not exceed the amount of base
present. Again, the ratio of [F–]/[HF ] doesn’t change much when a small amount of H3O+ is added so the
pH change is small.
Note that the conjugate base of a strong acid is actually a neutral ion—e.g. NO3– from HNO3(aq)—and thus
cannot neutralize H3O+. So a buffer solution cannot contain a strong acid! Because a strong acid always
ionizes completely, the conjugate ion of a strong acid never combines with H3O+ to produce the acid.
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When H3O+ is added to a solution containing NO3–, the ions simply remain in solution:
NO3–(aq) + H3O+(aq) → No Reaction
Likewise, a buffer solution cannot contain a strong base. Thus, to maintain pH, buffer solutions must
consist of a large concentration of a weak acid and its conjugate base—e.g. HF(aq) and KF—or a weak
base and its conjugate acid—e.g. NH3 and NH4Cl, not strong acids or strong bases. The weak acid must be
present to neutralize any OH– added to the buffer solution while the conjugate base must be present to
neutralize any H3O+ added to the buffer solution.
In this experiment, you will explore the nature of buffer solutions by preparing a buffer solution and
determining the amount of acid and base the buffer solution can absorb while maintaining its pH. You will
also compare the response of the buffer solution to water upon addition of small and large amounts of acid
and base, and you will calculate the pH after acid and base has been added to these systems.
Note: The solutions tested, including water, will generally be more acidic than you may expect
because CO2(g) dissolves in water to form H2CO3(aq), a weak acid.
Tap water is used instead of deionized water in parts V and VI because the pH readings reported
with the sensor for deionized water are erratic due to a low concencentration of ions.
Procedure:
I. Calibrating the pH sensor:
1. Obtain two burets, and the basket with the specialized equipment. Obtain and assemble the notebook
computer and LabPro system. Attach the sensor to the CH 1 outlet on the LabPro unit.
pH sensor handling instructions:
 Keep the sensor in a solution except when you are transferring it between solutions; the sensor
needs to be kept moist to function properly.
 Keep a waste beaker at your station for collecting the water when you rinse off the sensor.
 Don’t pour any other solution, even water, into the sensor storage container; this solution is a
special solution that maintains and preserves the sensor.
2. Double-click on the “Shortcut to GCC Chm Labpro” folder on the desktop. Double-click to open the
“HCl titration curve” file. When Logger Pro launches, you should see a reading for the pH sensor
displayed on the screen.
3. Rinse the pH sensor with deionized water from a wash bottle and gently dry it with a Kimwipe. Place
the sensor in the pH 7 buffer solution. If the pH is between 6.9-7.1, skip the calibration and go to step 5.
If the pH is not in this range, calibrate using the instructions in step 4.
4. From the “Experiment” menu select “Calibrate” ► CH1: pH. When the calibration window opens,
click on “Calibrate Now”.
First calibration point
1. Rinse the sensor with deionized water, and place it in the pH 4 buffer.
2. Type “4” in the box below “Reading 1”.
3. Swirl the sensor rapidly, wait until the voltage for Reading 1 stabilizes, and then click “Keep”.
Second Calibration point
A. Rinse the sensor with deionized water, and place it in the pH 7 buffer.
B. Type “7” in the box below “Reading 2”.
C. Swirl the sensor rapidly, and wait until the voltage for Reading 2 stabilizes. Click “Keep”, and then
“Done”.
II. Making the buffer solution:
5. Use a clean, dry 250 mL beaker to measure the mass of solid NaCH3COO calculated in question 1 on
page 5. Your actual weighed amount should be within 0.02 grams of the calculated mass. Record the
exact mass you measured on the balance in your lab notebook.
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6. Obtain about 10 mL of 2.0 M CH3COOH(aq) in a 30 mL beaker. Use the 5.00 mL pipet to transfer the
volume of CH3COOH(aq) calculated in question 2 on page 5 into a 100.00 mL volumetric flask. Use
the plastic squeeze bottle containing DI H2O to add enough water to fill the flask to the 100.00 mL
mark. Stopper the flask, and mix the solution thoroughly by inverting it several times.
7. Slowly pour the solution from the 100.00 mL volumetric flask into the beaker containing the solid
NaCH3COO sample. Use a stirring rod to mix the buffer solution in the beaker.
III. Titrating buffer with HCl:
8. Obtain a magnetic stirrer, and attach a buret clamp to the vertical pole on the stirrer. This clamp will
hold both 25 mL burets.
9. Obtain approximately 40 mL of 0.50 M HCl(aq) in a labeled 150 mL beaker. Clean and rinse one buret
with a few mL of the HCl, and secure the buret on the left side of the buret clamp. Fill the buret a little
above the 0.00-mL mark with the HCl solution. Drain a small amount of the HCl solution so it fills the
buret tip and leaves the HCl at or below the 0.00-mL level of the buret. (Save the rest of the HCl
solution in the beaker to refill the buret later.)
10. Obtain approximately 40 mL of 0.50 M NaOH(aq) in a labeled 150 mL beaker. Repeat step 9 to
condition and fill the second buret with NaOH(aq) solution. Attach this buret to the right side of the
buret clamp.
11. Use a 100 mL graduated cylinder to measure 30.0 mL of the prepared acetic acid/sodium acetate buffer
solution and add it to a labeled 150 mL beaker for the HCl titration. Repeat this process to measure a
second 30.0 mL sample of the buffer solution into a separate labeled beaker for the NaOH titration.
(Keep the rest of your buffer solution in case you make a mistake and need to repeat one of the parts.)
12. Carefully drop the magnetic stir bar into the first beaker containing 30.0 mL of the buffer solution.
Place the beaker on the stirrer. Turn on the magnetic stirrer and adjust the rate to slowly stir the
solution.
13. Rinse the pH sensor with deionized water and gently pat dry. Clamp the pH sensor above the buffer
solution with a small 3-prong clamp. Slowly lower the pH sensor into the buffer solution and secure it
so that it will not be struck by the stir bar and so that the buret will not drip on it. The pH sensor has
fragile electronic parts and must not come into contact with the stir bar.
14. Check the set up on the screen. The vertical axis of the plot should have pH scaled from 0 to 14 pH
units. The horizontal axis should have volume scaled from 0 to 15 mL. (If the screen does not look like
this, make sure the pH sensor is connected to LabPro, and the LabPro unit is plugged into a power strip
that is turned on, then close and reopen the experiment file.) Double click on the graph. Rename the
plot by entering “Buffer + HCl”. Click on the axis options tab. Click on the 15 for the x-axis, then
type in 8.5 and press the ENTER key.
15. Before adding any HCl, click “Collect” and monitor the pH for 5-10 seconds. Once the pH reading has
stabilized, click “Keep” (top right corner). An “Events with Entry” window opens. In the volume box,
type “0” (for 0 mL added). Click “OK” to store the first data pair.
16. Add approximately 1.00 mL of HCl to the buffer solution. When the pH stabilizes, again click the
“Keep” button. (Even if the pH readings continue to fluctuate; you can click “Keep” about 20-30 secs
after adding HCl.) In the volume box, type the current buret reading, to the nearest 0.01 mL, and click
“OK”. You will also need to record the volume and pH readings in a data table in your lab
notebook after each addition of HCl. You will continue adding HCl solution in 1.00 mL increments,
clicking “Keep” to store the pH and buret readings after each increment of HCl added until you have
added approximately 8.00 mL of HCl.
17. When you have finished collecting data, click the “Stop” button.
18. To print out copies of the HCl titration curve, go to the “File” menu, select “Print”, enter your names
in the footer box, and click “OK”. When the “Print” window opens, click on “Properties” (to the right
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of PS-149), and select the Layout tab at the top. Select the “Landscape” orientation and click “OK”.
Print one copy for your group. Make sure the graph AND data table are printed. (You may want to
sketch the titration plot in your lab notebook for reference in answering questions.)
IV: Titrating the buffer with NaOH:
19. Prepare the computer for the addition of NaOH to the buffer solution. Close and reopen the HCl
titration file. Double-click on the plot and rename the plot: “Buffer + NaOH”. Reset the x-axis for 8.5
mL again.
20. Set the second beaker containing a 30.0 mL portion of buffer solution on the stirrer. Refer to steps 1517 and substitute NaOH for HCl to collect data for the addition of NaOH to the buffer solution. Repeat
step 18 to print out the plot and data table.
21. Wash and dry the 150 mL beakers. Refill your burets with HCl and NaOH and adjust the volume so
that each solution will be at or below the 0.00 mL level on the buret.
V: Titrating tap water with HCl:
22. Add 30.0 mL of tap water (NOT DI water) to two separate labeled 150 mL beakers.
23. Place the pH sensor in the pH 7 buffer and swirl to check its calibration. If the pH is not between 6.97.1 then you should follow the directions in step 4 to recalibrate it.
Note: Placing the probe in the pH 7 buffer before measuring the pH of tap water will give us a
much more accurate initial pH reading for our tap water trials!
24. Prepare the computer for the addition of HCl to the water sample. Close and reopen the titration file.
Double-click on the plot and rename the plot: “HCl + water”. Reset the x-axis for 8.5 mL.
25. Repeat steps 15-17 to collect data for the addition of HCl to the water sample. Repeat step 18 to print
out the plot and data table.
VI: Titrating tap water with NaOH:
26. Prepare the computer for the addition of NaOH to the water sample. Close and reopen the titration file.
Double-click on the plot and rename the plot: “NaOH + water”. Reset the x-axis for 8.5 mL.
27. Repeat step 23 to check the calibration of the probe and to ensure that the pH of the probe is close to
neutral before measuring the pH of tap water. Repeat steps 15-18 to collect data for the addition of
NaOH to water. Repeat step 18 to print the plot and data table.
28. When you finish using the pH sensor, rinse it with deionized water, and return the sensor to the jar
containing the sensor storage solution.
Waste Disposal: Pour the solutions in the sink with running water.
Internet Activity Questions: View the animation and complete the internet activity questions on page
5 before you disassemble and return the computer and LabPro equipment.
The buffers lab report should include the following sections:
1) Title, date, name, partners
2) Prelab calcs and prelab page with instructor’s initials
3) Data/graphs: mass of the solid and the pH/volume data for the 4 trials – lab notebook. Plots – attached to
report for 1 of students in group. (Other students should include a reference for where to find the plots.)
4) Questions and drawings on page 5-8.
5) Calculations on page 8 – lab notebook copy.
6) Conclusion – a paragraph summarizing the findings. Also discuss two possible experimental errors and
how they would affect results.
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CHM 152LL:
Exploring Buffer Solutions
Name: ________________________
Section Number: _______________
Turn in one set of graphs per group with your lab report.
Pre-lab Calculations. Do these calculations in your lab notebook. Your instructor will initial these
during the pre-lab check.
1. Calculate the mass of solid NaCH3COO required to prepare 100.0 mL of a 0.10 M solution.
2. Calculate the volume in mL of 2.0 M CH3COOH(aq) needed to prepare exactly 100.0 mL of a 0.10
M solution.
Data: In addition to recording the pH values and volume measurements as acid and base are added, also
record the actual mass of solid measured out (watch sig figs!!) on the balance.
Questions.
Internet Activity Questions. Make sure the computer’s volume is turned on, then go to the following
website to view an animation on the acetic acid/acetate ion (CH3COOH/ CH3COO-) buffer system:
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf
Click on the “Add Strong Acid, H+” button to see what occurs at the molecular level when H+ is added to a
CH3COOH/ CH3COO- buffer system. Circle the correct answer for each question.
1. The strong acid (H+) is neutralized by:
a) CH3COOH
b) CH3COO-
2. As H+ is added, [CH3COOH]:
a) increases
b) decreases
3. As H+ is added, the [CH3COO-]:
a) increases
b) decreases
Click on “Add Strong Base, OH-” button to see what occurs at the molecular level when OH- is added to a
CH3COOH/ CH3COO- buffer system. Circle the correct answer for each question.
4. The strong base (OH-) is neutralized by:
a) CH3COOH
b) CH3COO-
5. As OH- is added, [CH3COOH]:
a) increases
b) decreases
6. As OH- is added, [CH3COO-]:
a) increases
b) decreases
General Questions
1. Which of the following combinations could form buffer solutions? Circle all that apply.
a) HCN and NaCN
b) NaClO4 and HClO4
c) LiOH and HNO3
d) HOCl and LiOCl
e) NH3 and NH4I
f) KOH and KNO2
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2. Complete and balance the equations for these reactants: CH3COOH(aq) and NaOH(aq).
molecular: ___________________________________________________________
ionic: _______________________________________________________________
net ionic:
___________________________________________________________.
3. Complete and balance the equations for these reactants: NaCH3COO(aq) and HCl(aq).
molecular: ___________________________________________________________
ionic: _______________________________________________________________
net ionic:
_____________________________________________________________
Answer the open-ended questions based on trends you noticed in this experiment. Be specific in your answers
and refer to the chemicals involved. Use the chemical’s name (or formula), don’t just say “acid” or “base.”
4. What happened to the buffer’s pH when a small amount (less than 5 mL) of HCl(aq) and NaOH(aq) was added?
For each added substance, indicate if the pH increases or decreases by a large or a small amount.
5. What happened to the water’s pH when a small amount (less than 5 mL) of HCl(aq) and NaOH(aq) was added?
For each added substance, indicate if the pH increases or decreases by a large or a small amount.
6. Explain why the water and buffer respond differently to the addition of small amounts of HCl(aq) and
NaOH(aq). Be specific about the chemicals and reactions involved for each system.
7. What happened to the pH when a large amount (more than 5 mL) of HCl(aq) and NaOH(aq) was added to the
buffer? Why was there a sudden change? How is this different than the results you observed when only a small
amount of acid or base was added? Be specific about what chemicals and reactions were involved.
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Chemical Representation Drawings
Draw the ions and/or compounds present for each of the solutions below. The first two are completed as
examples. Keep weak acids undissociated in the pictures. First determine whether each solution is an
acid, base, or salt!
K+
HF
K+
HF
FF-
2 moles of HF(aq)
2 moles of KF(aq)
2 moles of HCl(aq)
2 moles of NaOH(aq)
2 moles of CH3COOH(aq)
2 moles of NaCH3COO(aq)
.\
The following example shows which substances are present in solution when NaOH is added to the the HFKF buffer system described on page 1.
Example. Draw the ions and/or compounds present when two moles of an HF(aq)/KF(aq) buffer solution
combine with one mole of NaOH(aq):
→
+
HF(aq)/KF(aq) system
2 moles weak acid (keep
weak acid together) and
2 moles soluble salt (ions!)
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NaOH(aq)
1 mole strong base
(ions!)
products from reaction
1 HF reacted with OH- to form 1 H2O and
F-. This leaves 1 HF unreacted, 2 more F- ions
and the spectator ions: Na+ and K+ .
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1. The neutralization of NaOH(aq) by CH3COOH(aq). (Refer to the total ionic reaction on page 6.)
Draw the ions and/or compounds present when two moles of a CH3COOH(aq)/ NaCH3COO(aq) buffer
solution combine with one mole of NaOH(aq):
→
+
CH3COOH(aq)/ NaCH3COO(aq)
NaOH(aq)
products from reaction
2. The neutralization of HCl(aq) by NaCH3COO(aq). (Refer to the total ionic reaction on page 6.)
Draw the ions and/or compounds present when two moles of a CH3COOH(aq) / NaCH3COO(aq) buffer
solution react with one mole of HCl(aq):
→
+
CH3COOH(aq)/ NaCH3COO(aq)
HCl(aq)
products from reaction
Calculations
For the following calculations calculate the pH values like these were homework problems; pretend
you do not know the pH. Complete these calculations in your lab notebook! Include the question
number and question/calculation description with each one.
1. Calculate the pH of the original CH3COOH/CH3COO- buffer solution you prepared. (Hint: what
concentration of each component of the buffer was used?)
2. Calculate the pH after 2.00 mL of 0.50 M NaOH(aq) is added to your buffer solution. Make sure you
use the correct volume of buffer, the 30 mL you actually used in the experiment, not the total volume
you prepared.
3. Calculate the pH after 2.00 mL of 0.50 M NaOH(aq) is added to 30.0 mL of water.
4. Calculate the pH after 8.00 mL of 0.50 M HCl(aq) is added to your buffer solution. Make sure you use
the correct volume of buffer, the 30 mL you actually used in the experiment, not the total volume you
prepared.
5. Calculate the pH after 8.00 mL of 0.50 M HCl(aq) is added to 30.0 mL of water.
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