Experiment 32C
FV 21Feb17
APPLICATIONS OF ACID-BASE EQUILIBRIA
MATERIALS: 50 mL buret (2), 25 mL graduated cylinder (2), 50 mL beaker (2), 150 mL beaker (2), small plastic vials
(6), stirring rods (2), plastic beakers (3), pH meter with 3 standardization buffers, stirrer and stir bar,
0.20 M CH3COOH, 0.20 M CH3COONa, 0.20 M HCl, 0.20 M NaOH; 6M HCl and 6 M NaOH in dropper
bottles, phenolphthalein, bromophenol blue and universal indicators in dropper bottles.
PURPOSE:
The purpose of this experiment is to provide a comparative study of acid-base titration curves, to examine
the properties of buffers, and to relate those concepts in terms of the chemical species involved.
LEARNING OBJECTIVES:
1.
2.
3.
4.
5.
By the end of this experiment, the student should be able to demonstrate the
following proficiencies:
Perform a pH titration.
Qualitatively describe the chemical species in important regions of a titration curve.
Define a buffer and explain how a buffer works.
Explain how an acid-base indicator is used in the laboratory.
Prepare a buffer at a specified pH.
DISCUSSION:
Acids and bases are two very important classes of chemical compounds. You should already be familiar with the
distinctions between strong acids or bases, and weak acids or bases. Those distinctions are based on their electrolyte
behaviors – strong acids/bases are essentially 100% dissociated in water, while weak acids/bases are typically solvated as
intact molecules, and only dissociate to a small extent. There are relatively few common strong acids and bases, but large
numbers of weak acids and weak bases. You should also be able to recognize the structures that commonly correspond to
weak acids and bases – carboxylic acid and amine functional groups lend acidic or basic nature to their compounds.
carboxylic acid
primary amine
Using the Bronsted-Lowry definition, an acid is a proton donor and a base is a proton acceptor, and the reaction of a
Bronsted acid and base produces a new acid and a new base as a result of the proton transfer:
CH3COOH(aq) + H2O(l) ⇌ CH3COO−(aq) + H3O+(aq)
(1)
Thus the acetate ion CH3COO− is a weak base, the conjugate base of acetic acid, CH3COOH. So, not only are amines
weak bases, but so are the conjugates of weak acids. A similar relation holds between amines and their protonated forms.
The equilibrium of equation (1) has important consequences when the weak acid/weak conjugate base system is reacted
with a strong acid or a strong base. That is the origin of the buffering behavior that we will explore in this experiment.
Weak acids and their conjugate bases do not neutralize each other, but instead may react with other acids or bases added
to the system. Thus, in a titration, the pertinent reactions for a weak acid such as acetic acid, or weak base such as acetate
ion, would be
CH3COOH(aq) + OH− (aq)  CH3COO−(aq) + H2O (l)
(2)
CH3COO−(aq)
(3)
+ H3O+(aq)  CH3COOH (aq) + H2O (l)
A titration is just a chemical reaction run in a controlled manner until stoichiometrically equal amounts of acid or base
have reacted, a situation called the ‘equivalence point’. (However, we will carry the titrations beyond the equivalence
point just to explore the curves.) Note in equations (2-3) that the result of each addition of strong base (OH−) or strong
acid (H3O+) is to convert some weak acid into its conjugate base, or weak base into its conjugate acid. At the equivalence
point, the conversion is complete. Although the general reaction between an acid and a base is often called a “neutralization
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reaction”, the resulting solution is of neutral pH only in the case of a complete strong acid - strong base reaction. For
reactions involving weak acids or bases the result is not a pH-neutral solution, because you formed the conjugate species,
which are themselves bases or acids. In all cases, the pH at the equivalence point is just the pH of the salt solution you
form by the reaction. Neutral salts (such as NaCl) form from strong acid – strong base reactions; basic salts (such as
CH3COONa) form from weak acid – strong base reactions and acidic salts (such as NH4Cl) form from weak base - strong
acid reactions.
In the titration of weak acids or bases, an interesting situation holds in the region of the “half-equivalence point”, where
you have added half the number of moles of base or acid required by the stoichiometry. The result of converting about
half of the weak acid into its conjugate base (or about half of a weak base into its conjugate acid) is the formation of a
“buffer solution”. A buffer is a solution containg a weak acid and its conjugate base in significant, comparable amounts.
A buffer solution limits large changes in pH as strong base or acid is added, because the effect of equations (2-3) at that
point is to convert a strong base (OH−) into a weak base (CH3COO−), or a strong acid (H3O+) into a weak acid (CH3COOH)
and these, because of low degree of dissociation, don’t affect the pH as greatly as the strong electrolytes. Buffer behavior
follows the Henderson-Hasselbalch equation
𝑝𝑝𝑝𝑝 = 𝑝𝑝𝐾𝐾𝑎𝑎 + log
[𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐 𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏]
[𝑤𝑤𝑤𝑤𝑤𝑤𝑤𝑤 𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎]
(4)
where the Ka is the acid dissociation equilibrium constant for the weak acid. Note that at the half-equivalence point, the
concentrations of the weak acid and its conjugate base are the same, so the log term equals zero, and pH = pKa. (A related
equation can be derived for the weak base case, where the initial equilibrium is dictated by the base, so pOH = pKb at the
half-equivalence point. But pH = 14.00 − pOH, and pKa = 14.00 − pKb, so converting both terms just gives equation 4
again, even in the titration of a weak base. That should make sense because the concentration ratio is one in either case.)
We will be able to test both overall behavior of buffers, and the Henderson-Hasselbalch equation in this experiment.
Buffers are extremely important chemical systems, with numerous applications. The Navy uses buffered solutions to
minimize corrosion in water systems. Buffers can be made to order simply by choosing a weak acid with a pKa close to
the desired pH, and then adjusting concentrations of the weak acid and its conjugate base using the Henderson-Hasselbalch
equation. Perhaps of even more significance, your body utilizes many buffers to maintain approximately constant pH in
blood and organ systems. When biological buffers fail, the result can be serious illness or even death.
Acid-Base Indicators. Acid-base indicators are generally organic weak acids where the acid form absorbs one set of
wavelengths (hence, makes solutions appear a certain color), and the conjugate base form absorbs a different set of
wavelengths (hence, makes solutions appear a different color). They are typically very strong absorbers (high molar
absorptivity ε) so only small amounts are needed. As with all weak acids, there is an acid dissociation constant, KaInd,
associated with the equilibrium:
HIndcolor1(aq) + H2O(l) ⇌ Indcolor2−(aq) + H3O+(aq)
(5)
The equilibrium constant expression, including reference to the colors of the acid and conjugate base forms, is
K aInd = [ H 3 O + ]
−
[ Ind color
2]
[ HInd color1 ]
or, rearranged in log form
−
 [ Ind color

2]

pH = pK aInd + log

HInd
[
]
color1 

(6)
The pH of the solution to which the indicator has been added has a significant impact on the predominant form of the
indicator, and hence on the color of the solution. From the logarithmic form of equation (6), we see that if the pH of the
solution is, for example, one unit higher than the pKaInd, the concentration of the conjugate base form of the indicator,
[Ind‾], will be ten times larger than that of the acid form, [HInd]. For this case, the solution will appear to be color 2.
Similarly, a pH one unit lower than the pKaInd, will result in a ten-fold excess of the acid form, and the solution appearing
to be color 1. For pH ~ pKaInd, the solution would be a mixture of color 1 and color 2. During the course of a titration, if
the pH of the solution were to change from one unit below the pKaInd to one unit above, the solution’s color would change
from color 1 to color 2. From this, we see that a given acid-base indicator “indicates” when the pH of a solution changes
from pH values just below its pKaInd to those just above. There are hundreds of different indicators, each with a unique
value of KaInd. The selection of an indicator for use in a particular titration depends on the value of pH at which the color
change is needed. As long as the color change occurs near the equivalence point pH, on the steepest part of the titration
curve, the indicator can be used to signal the end of the titration. Generally, an indicator whose pKaInd value was close to
the pH of the equivalence point would be selected for use in a titration.
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PROCEDURE (work in groups of four; each pair does all of part A, one segment of Part B, and all of Part C):
Part A. Standardization of the pH meter.
1.
Follow the directions of your instructor regarding the titration set-up, operation of the pH meter and care and use of
the electrode.
2.
Standardize each of the pH meters with the three buffer solutions provided (pH = 4, 7 and 10). It is NOT necessary
to pour the buffers out of their vials; simply make the measurements in the original containers. After calibration, keep
the tip of the pH electrode immersed in deionized water when not in use.
3.
Thoroughly rinse and fill a 50 mL buret with deionized water. Be sure to clear any air bubbles beneath the stopcock.
Use this as the source when measured volumes of water are required in Parts B and C of the procedure.
Part B. Titration of a Weak Acid or Weak Base (two pairs work together; one pair does Segment I, the other does
Segment II).
Segment I – Titration of Acetic Acid Solution
1.
Obtain ~ 40 mL of 0.20 M CH3COOH solution in a 50 mL beaker. Rinse a 25 mL graduated cylinder with water
and then rinse with the 0.20 M CH3COOH solution, discarding the rinses each time. Measure out 20.0 mL of the
CH3COOH solution and add it to a clean 150 mL beaker. Use the water buret to measure out 20.0 mL of deionized
water and add it to the beaker. Add 2 drops of phenolphthalein indicator solution to the beaker. Add a stir bar
and place the beaker on the magnetic stirrer. Place the pH electrode in the solution and stir gently. The electrode
should remain in the holder while it is in the solution.
2.
Rinse a second buret with water and then rinse with 0.20 M NaOH solution, discarding the rinses each time. Fill
up the buret with NaOH solution, remove any air bubbles from beneath the stopcock, and refill the buret. Mount
the buret in a buret clamp and arrange the buret so that additions of titrant can be made while the pH electrode
remains in the solution in the beaker.
3.
Record the actual concentrations of the CH3COOH and NaOH solutions. Record the pH of the solution before
any titrant is added.
4.
Begin the titration, taking pH readings and noting the color or color changes at approximately 1.0 mL intervals
(or 0.5 mL near the equivalence point, as the pH is changing rapidly). The pH reading may drift somewhat, but
you should only wait about 10 seconds after each addition before taking a pH reading. Continue the titration,
recording pH and color until you have added a total of 35 mL of titrant.
5.
After you have completed the titration, remove the pH electrode from the titration beaker and place the electrode
in a small beaker of deionized water. (It will be needed in Part C.) Retrieve the stir bar, flush the titrated solution
down the drain and wash your titration beaker and the base-containing buret.
6.
Exchange your Part B data with the group working on Segment II, so that all groups have all the data on both
titrations.
Segment II – Titration of Sodium Acetate Solution
1.
Obtain ~ 40 mL of 0.20 M CH3COONa solution in a 50 mL beaker. Rinse a 25 mL graduated cylinder with
water and then rinse with the 0.20 M CH3COONa solution, discarding the rinses each time. Measure out 20.0
mL of the CH3COONa solution and add it to a clean 150 mL beaker. Use the water buret to measure out 20.0 mL
of deionized water and add it to the beaker. Add 2 drops of bromophenol blue indicator solution to the beaker.
Add a stir bar and place the beaker on the magnetic stirrer. Place the pH electrode in the solution and stir gently.
The electrode should remain in the holder while it is in the solution.
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2.
Rinse a second buret with water and then rinse with 0.20 M HCl solution, discarding the rinses each time. Fill
up the buret with HCl solution, remove any air bubbles from beneath the stopcock, and refill the buret. Mount
the buret in a buret clamp and arrange the buret so that additions of titrant can be made while the pH electrode
remains in the solution in the beaker.
3.
Record the actual concentrations of the CH3COONa and HCl solutions. Record the pH of the solution before any
titrant is added.
4.
Begin the titration, taking pH readings and noting the color or color changes at approximately 1.0 mL intervals
(or 0.5 mL near the equivalence point, as the pH is changing rapidly). The pH reading may drift somewhat, but
you should only wait about 10 seconds after each addition before taking a pH reading. Continue the titration,
recording pH and color until you have added a total of about 35 mL of titrant.
5.
After you have completed the titration, remove the pH electrode from the titration beaker and place the electrode
in a small beaker of deionized water. (It will be needed in Part C.) Retrieve the stir bar, flush the titrated solution
down the drain and wash your titration beaker and the acid-containing buret.
6.
Exchange your Part B data with the group working on Segment I, so that all groups have all the data on both
titrations.
Part C. Buffers
1.
Clean and rinse two 25 mL graduated cylinders. Rinse one with 0.20 M CH3COOH, and the other with 0.20 M
CH3COONa, discarding the rinses. Label these and use them for your reagent in the steps that follow. Use the
buret containing deionized water to measure out any water volumes required.
2.
Measure out 20.0 mL of CH3COOH and pour it into a clean plastic beaker. From your buret, add 20.0 mL of
deionized water to the beaker. Add 3 drops of universal indicator, stir the solution and record the color.
Measure the pH using a pH meter. Record your data in the Data table.
a. Divide the contents of the beaker equally into two plastic vials.
b. To one vial, add 2 drops 6 M HCl and stir to mix. Record the color, and measure the pH.
c. To the second vial, add 2 drops 6 M NaOH and stir to mix. Record the color, and measure the pH.
3.
Measure out 20.0 mL of CH3COONa and pour it into a clean plastic beaker. From your buret, add 20.0 mL of
deionized water to the beaker. Add 3 drops of universal indicator, stir the solution and record the color.
Measure the pH using a pH meter. Record your data in the Data table.
a. Divide the contents of the beaker equally into two plastic vials.
b. To one vial, add 2 drops 6 M HCl and stir to mix. Record the color, and measure the pH.
c. To the second vial, add 2 drops 6 M NaOH and stir to mix. Record the color, and measure the pH
4.
Measure out 20.0 mL of CH3COOH and pour it into a clean plastic beaker. Measure out 20.0 mL of
CH3COONa and pour it into the same plastic beaker. Add 3 drops of universal indicator, stir the solution and
record the color. Measure the pH using a pH meter. Record your data in the Data table.
a. Divide the contents of the beaker equally into two plastic vials.
b. To one vial, add 2 drops 6 M HCl and stir to mix. Record the color, and measure the pH.
c. To the second vial, add 2 drops 6 M NaOH and stir to mix. Record the color, and measure the pH
Clean-Up
1. Empty all vessels except the three colored buffer solutions used to standardize the pH meter. All solutions can
go down the drain.
2. Rinse all glassware with deionized water
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Name ____________________________________________
Section ___________________
Partner ___________________________________________
Date ______________________
DATA SECTION
Experiment 32C
Part B – Weak Acid/Base Titrations
Segment I
Segment II
conc. CH3COOH ___________ M
conc. CH3COONa ___________ M
conc. NaOH
conc. HCl
vol.
NaOH
(mL)
0.00
___________ M
Titration of CH3COOH Solution
vol.
pH
color
NaOH
pH
(mL)
color
vol.
HCl
(mL)
0.00
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___________ M
Titration of CH3COONa Solution
vol.
pH
HCl
pH
color
(color)
(mL)
color
Part C – Buffers
initial
#
Solution
1
20 mL of 0.20 M CH3COOH +
20 mL of H2O (buret)
2
20 mL of 0.20 M CH3COONa +
20 mL of H2O (buret)
3
20 mL of 0.20 M CH3COOH +
20 mL of 0.20 M CH3COONa
calculated
pH*
measured
pH
observed
color
after 6 M HCl
measured
pH
color
after 6 M NaOH
measured
pH
*After collecting all of your data, show your work for the calculation of pH for the three solutions here.
Ka(CH3COOH) = 1.76 x10−5:
a. solution 1
b. solution 2
c. solution 3
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color
DATA ANALYSIS:
Part B
1.
Use Excel to create a plot of pH vs. volume NaOH solution for the CH3COOH titration of Segment I.
a. On the titration curve itself, identify the equivalence point and the half-equivalence point. Record the
pH and volume of titrant of those points below.
equivalence point
vol titrant (mL)
pH
2.
half-equivalence point
vol titrant (mL)
pH
Use Excel to create a plot of pH vs. volume HCl solution for the CH3COONa titration of Segment II.
a. On the titration curve itself, identify the equivalence point and the half-equivalence point. Record the
pH and volume of titrant of those points below.
equivalence point
vol titrant (mL)
pH
3.
half-equivalence point
vol titrant (mL)
pH
For comparison, create a third plot of pH vs. volume NaOH solution for the strong-acid-strong base titration, using
HCl-NaOH data provided to you. (Your Instructor may ask you to plot the two acid titrations on the same chart to
make the differences more obvious.)
a. On the titration curve itself, identify the equivalence point. Record the pH and volume of titrant of that
point below. (Note that the half-equivalence point has no special significance in strong acid-strong base
titrations.)
equivalence point
vol titrant (mL)
pH
4. Qualitatively compare the titration curves of the weak acid (question 1 above) and the strong acid (question 3 above)
in terms of the following. Briefly explain the reason for the differences.
a. initial pH (before any NaOH was added)
b. pH at the equivalence point
c. steepness of the curve near the equivalence point
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5. Based on the pH values at the half-equivalence points, determine the value of Ka for acetic acid using the data from
both CH3COOH and CH3COONa titrations. Average those values. (See the discussion below equation (4).)
a. Ka of acetic acid based on data from the titration of acetic acid solution, CH3COOH
b. Ka of acetic acid based on data from the titration of sodium acetate solution, CH3COONa.
average Ka ____________
6. Would bromophenol blue be a suitable indicator for the titration of acetic acid solution with NaOH? Why or why not?
Part C
1. Considering only the relative amounts of the species acetic acid CH3COOH and acetate ion CH3COO−, which solution
(#1, 2 or 3) of Part C is most similar in composition to
a. the initial solution in the acetic acid titration of Part B?
_______
b. the equivalence point in the acetic acid titration of Part B? Explain.
_______
c. the half-equivalence point in the acetic acid titration of Part B? Explain.
_______
d. the equivalence point in the sodium acetate titration of Part B? Explain
_______
2. Using your experimental evidence, which solution(s) (#1, 2 or 3) of Part C act(s) like a buffer solution? Explain your
answer.
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3. Explain how a buffer works to resist large pH changes. Provide enough information to indicate your understanding of
buffers.
4. a. How would you prepare 500.0 mL of a HOCl/NaOCl buffer with pH = 7.89? Assume that you want the total buffer
concentration (i.e., [weak acid] + [conjugate base]) to be 0.1000 M. How many moles of NaOCl and HOCl would
be needed? (Review Example 17.4 on pp. 640-641 in the Kotz textbook.) What mass of NaOCl would provide
that many moles of NaOCl? What volume of 1.000 M HOCl solution would provide that many moles of HOCl?
hypochlorous acid, HOCl
molar masses, g/mol: HOCl, 52.46;
Ka = 2.9x10−8
NaOCl, 74.44
moles NaOCl needed =
moles of HOCl needed =
mass of NaOCl needed =
volume of 1.00 M HOCl needed =
b. If 0.0010 moles of NaOH are added to the buffer in question 5a, what would happen to the pH (increase a little,
decrease a little, or stay the same)? Explain your answer (no calculation is required here but drawing a
picture or writing chemical reactions might be helpful).
Calculate the pH of the solution in question 5b. (Review Example 17.5, pp. 642-643, of the Kotz textbook)
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Name ________________________________________
Section _______
Date _________
PRE-LAB QUESTIONS
Experiment 32C
1. Conjugates of weak acids and bases:
a) Write the chemical formula for the conjugate base of propenoic acid, C3H4O2. ________________
b) Write the chemical formula for the conjugate acid of propylamine, C3H9N.
________________
2. The pH at the equivalence point in a titration is the pH of the salt solution formed in the titration. What pH is expected
at the equivalence points in the following titrations? (circle)
a. HCl(aq) + NH3(aq)
a. pH < 7
b. pH = 7
c. pH > 7
b. HClO4(aq) + CsOH(aq)
a. pH < 7
b. pH = 7
c. pH > 7
c. C6H5COOH(aq) + KOH(aq)
a. pH < 7
b. pH = 7
c. pH > 7
3a. In Part B, step 1, 20.0 mL of 0.200 M CH3COOH is mixed with 20.0 mL of deionized water. What is the concentration
of the CH3COOH acid in the solution that results?
b. What is the pH expected for the solution formed in question 3a? Ka for CH3COOH is 1.76 x 10−5.
c. What volume of 0.195 M NaOH solution will be required to titrate the solution of question 3a to the equivalence point?
4. At the half-equivalence point in the titration of a weak acid, pH = pKa. For the titration of benzoic acid (C6H5COOH)
with NaOH, the pH = 4.190 at the half-equivalence point. What is the Ka of the benzoic acid?
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