When Life Gives you Lemons – Make Lemonade!
DETERMINING BUFFER CAPACITY
1. Introduction:
In this experiment you will work in groups of 4 or 5, split into team “A” and team
“B”. Each team will have specific tasks to complete; however, every group member will
be responsible for all analysis questions on their own and keeping track of the data.
Part of the data analysis will include you plotting the volume of base added
(NaOH) versus the recorded pH like the sample plots we drew in class with our “What is
a buffer?” notes.
2. Background:
A buffer solution is one that is resistant to change in pH when small amounts of
strong acid or base are added. For example, when 0.01 mole of strong acid or base are
added to distilled water, the pH drops to 2 with the acid and rises to 12 with the base. If
the same amount of acid or base is added to an acetic acid – sodium acetate buffer, the
pH may only change a fraction of a unit.
Buffers are important in many areas of chemistry. When the pH must be
controlled during the course of a reaction, the solutions are often buffered. This is often
the case in biochemistry when enzymes, proteins, or DNA is being studied, like a gel
electrophoresis. Our blood is buffered to a pH of 7.4. Variations of a few tenths of a pH
unit can cause illness or death. Acidosis is the condition when pH drops too low.
Alkalosis results when the pH is higher than normal.
Two species are required in a buffer solution. One is capable of reacting with
OH- (a proton donor – a weak acid) and the other will react with H+ (H3O+) (a proton acceptor – a weak base).
The two species must not react with each other.
Many buffers are prepared by combining a weak acid and its conjugate (acetic acid and
sodium acetate) or a weak base and its conjugate (ammonia and ammonium chloride).
In general, the pH range in which a buffer solution is effective is +/- one pH unit on
either side of the pKa. The Henderson–Hasselbalch equation provides the information
needed to prepare a buffer.
Do not worry about the The Henderson–Hasselbalch equation; you will see that
more in later biology and chemistry classes in this class I will give you the protocol
(SOP) for any buffers you need to make.
There is a limit to the amount of acid or base that can be added to a buffer
solution before one of the components is used up. This limit is called the buffer capacity
and is defined as the moles of acid or base necessary to change the pH of one liter of
solution by one unit.
3. Objective:
- Explain the plot of pH vs. volume to show buffering capacity of a solution.
- Graphically represent the difference between a buffered and un-buffered
solution.
4. Materials:
pH meter
10ml graduated cylinder
250ml beaker
Stir bar and plate
Graph paper
Disposable pipette
Deionized water
Powders lemonade drink mix
0.1 M NaOH
0.01M Citric acid
pH buffer 7 and 10
blue tray to coral your supplies for Day 2
5. Day 1 Procedure:
Team “A”: Calibrating the pH meter.
1. You will calibrate your pH meter following the SOP: Calibrating a Digital pH Meter you
have in your lab book, again please refer to manufactures calibration specifications. Use
the pH standard buffers provided (pH 10 and pH 7),
2. Read the SOP and the manufactures instructions before moving on.
3. Calibrate your pH meter.
4. When you have finished obtain 50ml of 0.1M NaOH from the stock bench.
5. Test the pH of the NaOH with the pH meter following the SOP: Using a Digital pH
Meter and record in data section. Remember to rinse with deionized water after test.
6. Obtain 40ml of 0.01M citric acid and add to a 250ml beaker then add 60ml of
deionized water, label “citric acid”.
Team “B” Calibrate a disposable pipette
1. Obtain a plastic, disposable pipette and a 10ml glass graduated cylinder.
2. Squeeze plunger on the pipette and fill with deionized water (25°C).
3. Carefully, count the number of drops it takes to fill the glass graduated cylinder until
the meniscus of the water is exactly at 5ml mark.
4. Record in data section. & Do Analysis Question #1 - calculate the volume per drop
which, in turn, will be needed in the calculation of the buffer capacity.
5. Now that you have calibrated your pipette, obtain 1.0g of the powdered lemonade
drink and put in a 250ml beaker and label lemonade drink mix.
6. Add 100ml of deionized water to beaker and place on your blue tray.
ALL:
Coral your supplies on a blue tray for tomorrow.
Add lab papers to your lab notebook leaving space for the graph before the analysis.
Set-up graph for analysis question #2.
Read through and discuss the procedures for Day 2 with your team mates.
6. Day 2 Procedure:
Buffer Capacity of Lemonade drink mix. Work as a group to complete the following.
Acidic Buffer.
1. Place the pH electrode in the 250ml beaker with the lemonade drink mix. Add a stir
bar and place on stir plate set to low. Be cautious to keep stir bar from contacting the
electrode and be sure to keep the heat knob in the off position.
2. Begin adding the 0.1M NaOH drop-wise to the Lemonade drink mix using the
calibrated disposable pipette. After every drop record the pH before proceeding with an
additional drop.
3. Continue to add 0.1M NaOH drop-wise until pH has increased by at least 2. When
this titration is completed dispose of mix down drain after retrieving stir bar. Make sure
to rinse pH electrode with deionized water.
Un-buffered Cirtric Acid
1. Place the pH electrode in the 250ml beaker labeled “Citric Acid.” Add a stir bar and
place on stir plate set to low. Be cautious to keep stir bar from contacting the electrode
and be sure to keep the heat knob in the off position.
2. Begin adding the 0.1M NaOH drop-wise to the un-buffered citric acid using the
calibrated disposable pipette. After every drop record the pH before proceeding with an
additional drop.
3. Continue to add 0.1M NaOH drop-wise until pH has increased by at least 2. When
this titration is completed dispose of mix down drain after retrieving stir bar. Make sure
to rinse pH electrode with deionized water.
Data:
Table 1: Volume per drop Disposable pipette
Total Drops added to Volume reached in
Vol. of water/ # of
Graduated cylinder
graduated cylinder
drops
Figure 1: Experiment Set-up
ml/ drop
Table 2: Lemonade Mix Titration
Table 3: Citric Acid Titration
Concentration of
NaOH
pH of NaOH
pH of
Lemonade
Drops of NaOH
M
Total vol. NaOH
added (ml)
pH
reading
Concentration
of NaOH
pH of NaOH
pH of Citric
Acid
Drops of NaOH
1
1
2
2
3
3
4
4
5
5
6
6
7
7
8
8
9
9
10
10
11
11
12
12
13
13
14
14
15
15
16
16
17
17
18
18
19
19
20
20
Vol. NaOH
added at equiv.
point
Moles NaOH
added
Vol. NaOH
added at equiv.
point
Moles NaOH
added
ml
mol
n = # of Moles V = volume M = molarity
be sure to convert mL into L for the volume
M
Total vol. NaOH
added (ml)
pH
reading
ml
mol
formula in symbols
n
= V * M
formula in units
# of moles = L * moles/L
Analysis and Graphing the Titrations
1. The number of drops to volume using the ml/drops value team “B” found.
2. On a separate piece of graph paper, graph pH (y-axis) vs. vol. (ml, x-axis) for the
lemonade mix and the Citric acid. It may help to use different colors to
for each plot (Lemonade and citric acid).
(Label axes with titles and units, title the graph, make a key, and use a ruler!)
3. Using your graph and data table, determine the equivalence point on the Lemonade
vs. NaOH graph line. The equivalency point is where the greatest pH change or “pH
spike” occurred, usually immediately beyond the buffer capacity region.
Using a yellow highlighter, highlight this region on your graph.
4. Find the point on the graph where pH spike occurs;
5. Find the volume of NaOH you added that is right after the pH spike.
6. Find the volume of NaOH just before the pH spike.
______mL
______mL
7. Add the volumes in #5 and #6 and divide the sum by two; this will be volume of
NaOH added at the equivalence point (when the spike occurred), record in data table.
8. Now find the moles of NaOH added by using the molarity of the NaOH and the
volume of NaOH added at equivalence point (convert to L). Moles = Vol. x M
9. Now repeat steps 2-6 for the citric acid.
Conclusions:
(You may paste in these questions in your lab notebook, but there is NOT enough room
to write these answers here. You will need to write each answer in complete sentences in
your lab notebook.)
1. Between the two solutions, which required more NaOH to reach equivalence point?
2. Based on your data, which solution is buffered? and which is not? Explain.
3. Looking at both curves on the graph, compare and contrast the two against each other.
Describe the shape of each graph? Does either graph have any “flat” regions? Explain.
4. Find another group and compare your buffered and un-buffered graphs to theirs.
a. Do you have similar shaped graph, if no how do they differ?
b. Are your buffering capacities with in 10% of each other?
c. If no, propose one (1) reason why.
5. List any questions you have about pH or buffers
or tell me what you enjoyed about his lab,
or tell me what you were frustrated with about this lab.