AP Chemistry—Chapter 15: Additional Aspects of Equilibrium
Buffers Lab
BACKGROUND INFORMATION:
A buffer is a solution that resists change in pH upon addition of acid, addition of base, and upon
dilution. Buffer solutions contain significant quantities of both partners of a Bronsted-Lowry conjugate
acid-base pair. To understand how buffers accomplish this, it will be necessary to review some BronstedLowry acid-base chemistry.
According to Bronsted and Lowry an acid is a proton (an H + ) donor. A base is defined
as a substance that can accept a proton. When an acid gives up a proton, a species that
can accept a proton, a base, is formed from the acid. That base is called the conjugate
base of the acid.
HA H+ + A–
HA = acid; A– = conjugate base
For example, acetic acid (CH3COOH) loses a proton to form acetate (CH3COO–) which will be
the conjugate base. The acetate is a potential proton acceptor and, as such, it must be considered to be a
base. The acetate is considered to be the conjugate base of acetic acid, and acetic acid is considered to be
the conjugate acid of the base, the acetate ion. In a similar manner sodium bicarbonate (HCO 3–) may lose
a proton to become sodium carbonate (CO32–). Sodium bicarbonate is the acid, and sodium carbonate is
the conjugate base. Finally, note the sodium dihydrogen phosphate and disodium monohydrogen
phosphate system. The sodium dihydrogen phosphate (H 2PO4–) is the proton donor, the acid, while its
product, disodium monohydrogen phosphate (HPO 42– ), is the proton acceptor and hence, the conjugate
base.
Although we frequently represent the proton in water (aqueous) solution as H +, actually the
proton combines with water to form ions like H3O+, H5O2+ , and H7O3+. The main form is the hydronium
ion, H3O+.
Acids and bases can be divided into two broad categories, strong and weak. Strong acids lose
their acidic protons virtually 100 % and strong bases accept protons virtually 100 %. Weak acids and
bases lose and gain protons respectively less than 10 %. All the acids and bases described above would be
considered to be weak.
To form a buffer one may mix a weak acid with its conjugate base or a weak base with its
conjugate acid. The mixture that results will resist any attempt to change the pH and will act as a buffer.
In this theory we will refer to a buffer as a mixture of a weak acid and its conjugate base. Suppose we mix
the acetic acid (CH3COOH) with its conjugate base, acetate, (CH3COO–). There will be present in the
mixture both an acid and a base that would tend to react with any added acid or base.
Suppose that a source of H+ is added. The conjugate base, acetate, will react as follows:
CH3COO– + H3O+ H2O + CH3COOH
Suppose that a source of OH - is added. The weak acid, acetic acid, will react as follows:
CH3COOH + OH–
H2O + CH3COO–
In either case, the invading species is not allowed to change the pH of the solution.
One must use weak acids in buffers so the conjugate bases will have a tendency to react with protons. If a
strong acid were mixed with its conjugate base, the conjugate base would have no tendency to react with
protons.( A strong acid reacts 100 % to lose its protons, so there must be no tendency of the reaction to go
in reverse to pick up protons.) This would leave the solution susceptible to attack by protons.
In this experiment you will make several mixtures of weak acids with their conjugate bases to create
buffers. You will test these by adding both acid and base to the mixtures. You will also test distilled
water and single components of the buffers in the absence of their conjugates to see how they hold up to
the challenge of added acid and base.
When working with a buffer one must be concerned with two major questions 1) over what pH range will
the buffer work and 2) what is the capacity of this buffer to resist pH change?
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AP Chemistry—Chapter 15: Additional Aspects of Equilibrium
Buffers Lab
Buffers do not buffer only at a pH of 7.0. Some do, but others buffer in the acidic range of pH's, and
others buffer in the basic range of pH's. The initial pH of a buffer depends upon two factors , 1) the
strength of the weak acid or weak base and 2) the ratio of weak acid to its conjugate base. Considering the
first fact, the stronger the weak acid the more acidic will be the pH of the buffer; the weaker the weak
acid the more basic the initial pH. After the range is determined by the strength of the weak acid
component, the actual initial pH is determined by the ratio of the weak acid to the conjugate base. The pH
of the buffer will be a bit more basic if more of the conjugate base is present than the weak acid, and it
will be a bit more acidic if more of the weak acid is present than the conjugate base.
The buffer capacity depends upon two factors. The concentration of the buffer is one major factor. In
general the more concentrated the buffer, the more ingredients are available to attack added H+ and OH–
ions. The second factor relates to what is being added to the buffer and how much of each component,
acid and conjugate base, is available to react. Suppose that a buffer had 100 times as much acetic acid as
it had acetate. This buffer could resist a challenge by base because there would be plenty of acetic acid to
react with the base. It would not, however, be able to resist an attack by acid, because there would be
relatively little acetate to react with the acid.
In this experiment two buffers will be made and tested, as well as, distilled water and some individual
ingredients from which the buffers will be made. The pH will be measured using Universal Indicator.
This indicator is a mixture of pH indicators. Over a pH range from 3 to 11 the color of this mixture will
vary from red to orange to yellow to green to blue to purple. Cards showing the colors at various pH’s
will be available to compare against. If the color of the solution is between two shades as seen on the
cards, one may declare that the pH is half way between the two.
PRELABORATORY QUESTIONS:
1. Define pH, and describe the pH scale and what the pH scale measures.
2. What is meant when one says that a buffer has a high buffer capacity?
3. What is meant when one says that acetic acid is a stronger acid than benzoic acid?
4. The blood buffer system is made up of H2CO3 and HCO3–. Describe with the use of equations how this
system responds to added H3O+ and to added OH–.
5. How does a conjugate base differ from the acid, HB?
6. How can HCO3– function as both the acidic component of the HCO 3– /CO32– buffer system and as the
basic component of the H2CO3/HCO3– buffer system?
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AP Chemistry—Chapter 15: Additional Aspects of Equilibrium
Buffers Lab
7. What two factors are involved in determining the exact pH of a buffer?
PROCEDURE:
1. Label seven medium test tubes #1 through #7.
2. Mix 5 mL of 0.1 M sodium carbonate (Na2CO3) and 5 mL of 0.1 M sodium bicarbonate (NaHCO 3) in
a small beaker. This is a CO32–/HCO3– buffer.
3. Divide the solution equally between test tubes # 1 and # 2.
4. Mix 5.0 mL of 0.1 M sodium monohydrogen phosphate (Na2HPO4) and 5.0 mL of 0.1 M sodium
dihydrogen phosphate (NaH2PO4). This is a HPO42–/H2PO4– buffer.
5. Divide this solution equally between test tubes # 3 and #4.
6. Put 5 mL of 0.1 M NaHCO3 in test tube # 5, 5 mL of 0.1 M NaH2PO4 in test tube # 6, and 5 mL of
cooled, boiled distilled water in test tube #7.
7. Add 3 drops of Universal indicator to EACH OF THE SEVEN TEST TUBES.
8. Estimate the pH of each solution by comparing your resulting colors to the colored cards at your lab
table. Record your results in the data table.
9. Add 1 drop of 1 M hydrochloric acid (HCl) to tubes #1, #3, #5, #6, and #7. Shake the solutions, note
the color changes, and estimate the pH. Record your results.
10. Add 2 drops of 1 M sodium hydroxide (NaOH) to the tubes that received the hydrochloric acid (#1,
#3, #5, #6, and #7), shake the solution, and again estimate the pH.
Compare tubes #1, #2, and #7 with one another, and #3, #4, and #7 with one another. What do you
observe?
Also note the behavior of tubes #5 and #6 as HCl is first added and then as excess NaOH is added. What
did you observe?
DATA SUMMARY:
Tube #
Contents of Tube
1
2
3
4
5
6
7
CO32–/HCO3–
CO32–/HCO3–
HPO42–/H2PO4–
HPO42–/H2PO4–
NaHCO3
NaH2PO4
Distilled H2O
Initial pH
pH after adding
HCl
pH after adding
NaOH
Now that you have collected the bulk of the information you will check the buffer capacity of buffer tubes
#1 and #3. Add additional drops of 1 M HCl to test tube #1 until the pH is at 2. Count the number of
drops of HCl required to reach this point. Now, add additional drops of 1 M HCl to test tube #3 until the
pH is at 2. Count the number of drops of HCl required to reach this point. Record the number of drops of
each.
Tube #1 _____________
Tube #3 _____________
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AP Chemistry—Chapter 15: Additional Aspects of Equilibrium
Buffers Lab
RESULTS:
1. What was the purpose for adding 2 drops of NaOH to the test solution in the second part of the
experiment? Why did we not just add 1 drop of NaOH?
2. What is the purpose of a buffer? Does a buffer always hold the pH of a solution at pH 7?
3. Based on your experimental evidence which buffer system, CO 32–/HCO3– or HPO42–/H2PO4–, had the
greatest buffering capacity? What is your evidence for this conclusion?
4. Which solution showed the largest change in pH for the addition of 1 drop of HCl? Was this system
buffered?
5. Write equations for the reaction of the CO32–/HCO3– buffer reacting with an acid and a base.
6. Write equations for the reaction of the HPO42–/H2PO4– buffer reacting with an acid and a base.
7. Is 0.1 M NaHCO3 a good buffer? Explain.
8. Is 0.1 M NaH2PO4 a good buffer? Explain.
9. Calculate what the initial pH of Test Tubes 1-7 should have been. How close were your estimates
using Universal Indicator Solution?
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