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Transcript
ACIDS, BASES, SALTS, AND BUFFER SOLUTIONS
ADDITIONAL READING
The concepts in this experiment are also discussed in sections 3.6, 17.1-17.3, and 16.1-16.7 of Chemistry &
Chemical Reactivity by Kotz; and also discussed in chapters 16 and 17 of General Chemistry by Vining.
ABSTRACT
This experiment is divided into four parts. Throughout the experiment, a Vernier pH probe connected to the Vernier
LabQuest computer interface will be used to measure the pH of aqueous solutions.
In part A, the pH of a series of solutions of known acids and bases will be measured. After which, the hydronium
and hydroxide ion concentrations will be calculated. In part B, the pH of different aqueous salt solutions will be
explored. An estimation of pH will be made followed by a measurement. In part C, two acetic acid/sodium acetate
buffers will be prepared. After which, the corresponding buffer capacities will be determined. In part D, a series of
phosphate buffers will be prepared. The prepared buffers will correspond to specific pH values.
BACKGROUND
One way in which to calculate the pH of an acid or base is through the use of Equation (1):
pH = –log[H3O+]
(1)
Equation (1) is particularly useful for strong acids and bases because strong acids and bases completely dissociate
in water and their tendency to hydrolyze with water to reform the strong acid is negligible. A table of six strong
acids can be found in your textbook on page 609. The table is 15.3.
The strong acid HCl(aq) from table 15.3 will essentially completely ionize in an aqueous solution with virtually no
intact HCl molecules remaining. This ionization is represented by Equation (2):
HCl(aq) + H2O(l) ® H3O+(aq) + Cl-(aq)
(2)
As such, the molarity of the acid is essentially equal to the concentration of hydronium ion ([H3O+]).
For example, a 0.020 M HCl solution has a [H3O+] = 0.020 M. Using Equation (1), a pH can be calculated.
pH = –log(0.020) = 1.70
As mentioned, Equation (1) is also useful for calculating a pH of a solution of strong base. A table of six strong
bases can be found in your textbook on page 624. The table is 15.8.
The strong acid NaOH(aq) from Table 15.8 will completely dissociate in solution with virtually no intact NaOH
molecules remaining. This dissociation is represented by Equation (3).
1
NaOH(aq) ® Na-(aq) + OH-(aq)
(3)
To use Equation (1) for a strong base, though, the [H3O+] must be calculated. This calculation may be done using
the ion product constant for water (Kw), Equation (4):
Kw = 1.00 x 10–14 = [H3O+][OH–]
(4)
For example, a 0.035 M NaOH solution has [OH–] = 0.035 M.
To determine the PH of this strong base solution, Equation (4) can be used as follows:
1.00 x 10–14 = [H3O+][OH–]
1.00 x 10–14 / [OH-] = [H3O+]
1.00 x 10–14 / 0.035 = [H3O+]
2.86 x 10–13 M = [H3O+]
Now that the [H3O+] has been calculated, Equation (1) can be used as follows:
pH = -log[H3O+]
pH = -log(2.86 x 10–13 M)
pH = 12.54
It is also possible, and will be necessary at times, to calculate the [H3O+] from the pH. To do this, a mathematical
operation called an antilog (or inverse log) is used. Taking the antilog of a pH value will afford the corresponding
[H3O+]. This process is represented by Equation (5):
[H3O+] = 10–pH.
(5)
For example, the [H3O+] of an acidic solution with a pH = 3.54, can be calculated as follows:
[H3O+] = 10–pH
[H3O+] = 10–3.54
[H3O+] = 2.9 x 10–4 M
Weak acid calculations are handled in a very different way. A table of common weak acids can be found in your
textbook on page 625. The table is Table 15.9. Memorizing the table of weak acids is possible; however, a more
efficient way to remember the weak acids is as follows: If the acid in question is not one of the six strong acids, the
acid is likely a weak acid.
Calculating the pH of a weak acid requires the use of the acid ionization constant (Ka) and the following equilibrium
expression, Equation (6):
Ka = [H3O+][A-]
[HA]
(6)
2
From Equation (6), HA is a generic term for “weak acid”, A- is a generic term for the conjugate base of the weak
acid HA, and H3O+ is the hydronium ion. The ionization of weak acid HA is represented by Equation (7):
HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
(7)
For example, using Equation 6 and the concept of Equation (7), the PH of a 0.020 M solution of benzoic acid,
C6H5COOH (a monoprotic weak acid; Ka = 6.5 x 10–5) can be determined as follows:
HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
C6H5COOH(aq) + H2O(l) ⇌ H3O+(aq) + C6H5COO-(aq)
Initial
Change
Equilibrium
(8)
C6H5COOH(aq)
H2O(l)
H3O+(aq)
0.020
Not Included
0
-x
Not Included
+x
0.020 - x
Not Included
x
Figure 1: ICE (Initial Change Equilibrum) Table for Benzoic Acid
C6H5COO-(aq)
0
+x
x
Figure 1 is an ICE table corresponding to the ionization of benzoic acid. It is useful to generate such a table when
dealing with weak acids and their calculations. The equilibrium line of the Figure 1 can be transferred to the
equilibrium expression as follows:
Ka = [H3O+][A-]
[HA]
Ka = [H3O+][C6H5COO-]
[C6H5COOH]
Ka = [x][x]
[0.020 - x]
(9)
(10)
Equation (10) can be simplified further by completing the mathematical operation of the numerator and simplifying
the denominator by assuming the x in the denominator is negligibly small. This is a logical assumption considering
that weak acids ionize to a small extent, and because the conjugate bases of weak acids hydrolyze with water to
reform the weak acid, much intact acid remains in solution.
Ka = x2
(11)
0.020
Inserting the Ka value and algebraically solving for x will afford Equations (12) and (13), respectively:
6.5 x 10-5 = x2
0.020
(12)
Ö(0.020)(6.5 x 10-5) = x
(13)
1.1 x 10-3 = x
Where from Figure 1, x = [H3O+]
3
The pH can be calculated as follows:
pH = -log[H3O+]
pH = -log(1.1 x 10-3)
pH = 2.96
The pH of a weak base can be calculated in a similar fashion. In the case of a weak base, though, the base
ionization constant (Kb) is used:
Kb =
[BH + ][OH- ]
. See pages 626-627 (Example 15.11) of Tro for a worked example.
[B]
Buffers: Solutions That Resist pH Changes
When a salt of acetic acid (like sodium acetate, NaC2H3O2) is dissolved in water, it ionizes completely, as nearly all
ionic compounds do. The acetate ion, the anion of the weak acid HC2H3O2, has a tendency to react further with
water, acting as a base (proton acceptor). This reaction is known as hydrolysis:
C2H3O2–(aq) + H2O(l) ⇌ HC2H3O2(aq) + OH–(aq)
(14)
Before NaC2H3O2 is dissolved in pure water, the water is neutral, [H3O+] = [OH-]. However, once NaC2H3O2
dissolves and dissociates in solution to form Na+ and C2H3O2– ions, a small amount of the C2H3O2– reacts further
(hydrolysis) to produce hydroxide ions. This small amount is enough to result in [OH-] > [H3O+], and the solution is
therefore basic. This is a typical example of the hydrolysis of the anion of a weak acid. The same behavior is to
be expected for anions like CO32–, PO43- which are conjugate bases of weak acids. The conjugate bases NO3- and
Cl-, however, are of strong acids. Conjugate bases of strong acids do not undergo hydrolysis.
Most solutions rapidly change pH when acid or base is added to them. However, a buffer resists pH change by
neutralizing added acid or base. Buffer solutions contain significant amounts of both a weak acid and its conjugate
base. For example, if acetic acid (HC2H3O2) were use as the acid in a buffer solution, the conjugate base required
would have to contain the acetate anion (C2H3O2–). Solutions containing just acetate anions do not exist. A cation
is necessary in order to balance the charge. In this case, to accomplish this charge balance, NaC2H3O2 is commonly
used.
The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation, Equation 15:
pH = pKa + log([base]/[acid]
(15)
For example, to calculate the PH of a buffer solution containing equal volumes of a 0.020 M solution of benzoic acid
(HC6H5CO2, Ka = 6.5 x 10–5) and a 0.025 M solution of sodium benzoate (NaC6H5CO2):
pH = pKa + log([base]/[acid]
pH = –logKa + log([0.0125]/[0.010])
pH = –log(6.5 x 10–5) + log([0.0125]/[0.010])
pH = –(–4.19) + 0.0969
pH = 4.29
*After mixing, the new final volume changes the concentrations
4
PRE-LAB EXERCISES – Note this set of prelab exercises is for practice. There is a similar set of prelab
exercises in Mindtap that must be completed by 8AM the day you will do experiment 7.
1. Calculate the pH (pH = –log[H+]) of the following solutions:
Show sample calculations for a and d to the right of the question
a. 0.10 M HCl solution
b. 0.010 M HCl solution
c.
0.0010 M HCl solution
d. 0.10 M NaOH solution
e. 0.010 M NaOH solution
f.
0.0010 M NaOH solution
g. 0.10 M HC2H3O2 solution (Ka = 1.8 x 10–5)
h. 0.10 M NH3 solution (Kb = 1.8 x 10–5)
2. Based on properties of weak acid/base ions, on the lines to the left of each lettered salt solution identify the
type of hydrolysis, acidic (A) or basic (B), that each salt undergoes. On the line on the right identify the
weak acid or base ion that causes the predicted type of hydrolysis.
Type of Hydrolysis
_____
_____
_____
_____
_____
_____
A.
B.
C.
D.
E.
F
Weak Base or Acid Ion
NaC2H3O2
Na2CO3
NaHSO4
NaHCO3
NH4Cl
AlCl3
sodium acetate
_________
sodium carbonate
_________
sodium hydrogen sulfate
_________
sodium hydrogen carbonate _________
ammonium chloride
_________
aluminium chloride
_________
3. Calculate the pH of a buffer made by mixing 50.0 mL of 0.10 M acetic acid (HC2H3O2) with 50.0 mL of 0.10
M sodium acetate (NaC2H3O2). The Ka for acetic acid is 1.8 x 10–5.
Use the Henderson-Hasselbalch equation:
pH = pKa + log([base]/[acid])
Name:
_________________________
Lab Day:
Room:
103
Lab Time: __________________
109
117
125
M
T
W
R
F
Instructor: __________________
5
SAFETY/HYGIENE/WASTE DISPOSAL
1. Always wear your goggles!
2. Caution: many chemicals are corrosive. You should wear gloves and be sure your skin is not exposed. Discard
the gloves after use, and wash your hands. Chemical exposure may result in an itchy sensation. If you have
any sensation of itching, burning, or tingling, thoroughly flush the area with water. Continue flushing until several
minutes after the sensation has subsided. Inform your lab instructor.
3. Never raise containers of solution, especially corrosive solutions, above shoulder level. In particular, avoid this
when filling a buret.
4. Waste solutions from this experiment should be disposed of in the waste containers designated by your lab
instructor.
5.
Sodium hydroxide is caustic and can cause skin irritations or burns. Use appropriate protective equipment,
including gloves and safety glasses/goggles.
PROCEDURE
THIS EXPERIMENT WILL BE DONE IN PAIRS
Equipment Provided in Lab:
•
•
•
•
•
•
•
Vernier LabQuest computer interface
Vernier pH Probe
Burets filled with 0.1 M HCl(aq) and 0.1 M NaOH(aq) solutions
Volumetric flasks, 10-mL, 100-mL
Beakers – various sizes
0.1 M solutions of acetic acid, sodium acetate, ammonia, phosphoric acid, sodium dihydrogen phosphate,
sodium hydrogen phosphate, sodium phosphate, and 0.5 M solutions of acetic acid and sodium acetate.
0.5 M solutions of sodium acetate, sodium carbonate, sodium hydrogen sulfate, sodium hydrogen
carbonate, ammonium chloride, and aluminium chloride.
Using the pH Probes
1. The pH probes should always be immersed in a solution; either a buffer solution or DI-water. NEVER
ALLOW THE PROBES TO DRY OUT ON THE LAB BENCH.
2. The pH probes are delicate and should be treated carefully. At the base of the probe, the glass
membrane is surrounded by a plastic sheath. If the plastic surrounding the membrane breaks, the glass
membrane will break and the sensor will no longer be functional.
3. Connect the Vernier pH probe to Channel 1 of the Vernier LabQuest computer interface (which will be now
referred to as LabQuest). The pH should be displayed in a box on the LabQuest screen. The value will
fluctuate slightly. If a pH reading is not displayed, consult your lab instructor. It is not necessary to change
any other setting on LabQuest. Collected data should not be stored on/in the LabQuest. All experimental
data should be properly recorded in your laboratory notebook.
4. When using the pH probe, always rinse it with DI-water first. The probe should be carefully dried with a
Kimwipe before measuring the pH of a solution.
6
5. DI water contains dissolved carbon dioxide which will cause your pH readings to be lower than expected.
Normally pH of boiled pure water should be 7.00. It turns out that our DI water is slightly basic, therefore
we can use a buffer solution whose pH = 7.00 to calibrate the pH probe.
6. Before using the probe it must be calibrated via the following the steps:
a. Two vials of standard pH 2.00 and pH 10.00 buffer solutions can be found on your lab bench.
b. Using the stylus, tap the orange screen and select Calibrate
c. Tap Calibrate Now
d. After rinsing your probe using DI water and carefully drying, place the probe into the buffer solution
labeled pH 2.00. Make sure you hold vial during calibration. When the line voltage becomes
constant enter 2.00 for Value 1 and tap Keep.
e. After rinsing your probe in DI water and carefully drying, place the probe into the buffer solution
labeled pH 10.00. Make sure you hold vial during calibration. When the line voltage becomes
constant enter 10.00 for Value 2 and tap Keep.
f. Tap OK
g. Your screen should now display a pH reading between 9.95 and 10.05
h. Rinse your probe in DI water, carefully dry, and place the probe back into the buffer solution labeled
pH 2.00. Make sure you hold the vial. Stir gentle for two seconds. Your screen should now display
a pH reading between 1.95 and 2.05. Your pH probe is now calibrated.
7. For Parts A, B and D you, will use the advanced features of LabQuest to determine the mean pH reading.
8. Check the following settings (in upper right of the screen):
Mode: Time Based
Rate: 20 samples/s
Length: 30 s
If you need to change the settings, follow the procedure that you used in previous experiments.
Part A – Measuring the pH of Acids and Bases
EACH PAIR WILL COMPLETE THE TASKS IN BOTH SECTIONS I AND II.
1. The following stock solutions have been prepared (HCl, NaOH, HC2H3O2 and NH4OH) and can be found in a
fume hood. Note that the concentrations are approximate.
0.100 M hydrochloric acid (HCl)
0.100 M sodium hydroxide (NaOH)
0.100 M acetic acid (HC2H3O2)
0.100 M ammonium hydroxide (NH4OH)
Section I:
2. From the fume hood dispense ~15 mL of the stock 0.10 M HCl solution into a clean, dry 30-mL labeled beaker.
3. Place the pH probe into the 0.10 M HCl. Swirl the probe gently. Start collecting pH data by using the stylus to
tap the green arrow ►on the lower left of the screen. After the run is complete, determine the mean pH and
record the pH of the solution in your lab notebook. Rinse the probe by swirling the probe in a 250 mL beaker
filled with 150 mL of deionized water. Carefully use a Kimwipe to dry the probe prior to immersing in another
solution. Pour the solution form your 30 mL beaker into the waste beaker and rinse with DI-water. Dry the
beaker with a paper towel.
4. Fill your burette with 0.10 M HCl. Dispense 1.00 mL of the 0.10 M HCl solution into a 10-mL volumetric flask.
Using you wash bottle, carefully fill the flask exactly to the line with DI-water. Mix thoroughly and pour the
solution into your clean, dry 30-mL beaker labeled. Using the stylus, tap the file cabinet icon to set up for run
2. Using the pH probe, measure the mean pH of the solution. Record the data in your laboratory notebook.
Rinse and wipe the probe as before.
5. Dispense 1.00 mL of the 0.10 M HCl solution into a 100-mL volumetric flask. Using you wash bottle, fill the
flask exactly to the line with DI-water. Mix thoroughly and pour ~15 mL into a clean, dry 30-mL beaker labeled
Using the stylus, tap the file cabinet icon to set up run 3. Using the pH probe, measure the mean pH of the
solution. Record the data in your laboratory notebook. Rinse and wipe the probe as before.
7
6. Obtain ~15 mL of a 0.10 M acetic acid solution from the fume hood and place it into a clean, dry 30-mL beaker.
Using the stylus, tap the file cabinet icon to set up run 4. Using the pH probe, measure the mean pH of the
solution. Record the data in your laboratory notebook. Pour the solution into your waste beaker, rinse the
beaker with DI-water and dry with a paper towel.
Section II:
7. Follow steps 2 – 5 using 0.10 M NaOH instead of the 0.10 M HCl solution. Record the mean pH for each
solution into your lab notebook. Be sure to rinse the PH probe and carefully dry the probe using a Kimwipe.
each time prior to immersing it into another solution.
8. Obtain ~15 mL of a 0.10 M ammonia solution (NH4OH) into a clean, dry 30-mL beaker. Using the stylus, tap
the file cabinet icon to set up the new run. Using the pH probe, measure the mean pH of the solution. Record
the data in your laboratory notebook. Pour the solution into your waste beaker, rinse the beaker with DI-water
and dry with a paper towel.
Asides:
9. Before executing Section II, calculate the concentrations of the H3O+ and OH– ions for stock solution of the
acid (0.100 M) and base (0.100 M).
10. The rinse water from the pH probe can be dumped down the drain. Refill your 250-mL beaker with 150 mL of
fresh DI- water. Re-calibrate your probe, if necessary. This procedure will be performed as your move through
each part of the experiment.
11. When the probe is not in use, it should be kept moist by being submerged in DI water.
Part B – Measuring the pH of Salt Solutions
1. The 0.50 M solutions of each of the following salts have been prepared, and can be found in vials:
A.
B.
C.
D.
E.
F.
NaC2H3O2
Na2CO3
NaHSO4
NaHCO3
NH4Cl
AlCl3
sodium acetate
sodium carbonate
sodium hydrogen sulfate
sodium hydrogen carbonate
ammonium chloride
aluminium chloride
2. Using your pH probe, measure and record the mean pH of each salt solution. Between each measurement,
remember to rinse the PH probe with distilled water and carefully dry the probe with a Kimwipe. These solution
will represent runs 5 - 10. Do not discard these solutions.
3. Arrange the salts in order of increasing pH (lowest to highest) using their assigned letters. How do the measured
pH values compare with your predictions in the Pre-Lab exercise? Completing the hydrolysis equations will
help you understand why salts are either acidic or basic
8
Part C – Investigating Buffer Capacity
EACH PAIR WILL COMPLETE THE ENTIRETY OF PART C
Buffer I
Buffer I should be prepared by mixing 10.0 mL of 0.10 M acetic acid with 10.0 mL of 0.10 M sodium acetate in a
100-mL labeled beaker. The solutions can be found in dispensing burets on the long window sill. Measure and
record the initial pH.
Refill your buret containing the 0.10 M HCl solution
Buffer II
Buffer II should be prepared by mixing 10.0 mL 0.20 M acetic acid with 10.0 mL of 0.20 M sodium acetate in a 100
mL labeled beaker. The solutions can be found in dispensing burets on the long window sill. Measure and record
the initial pH.
Refill your buret containing the 0.10 M NaOH solution
1. Record the initial volume reading of either 0.10 M HCl or 0.10 M NaOH solutions in your buret. Dispense these
solution in 1.00-mL increments into your buffer solution. Use the pH probe to stir the solution gently. When
the change in pH approaches 0.80 units; slowly dispense either the 0.10 M HCl or 0.10 M NaOH solutions dropwise from the buret. Continue to add the HCl or NaOH until the pH of the solution changes by 1 full pH unit.
Stop dispensing the HCl or NaOH. Record the final volume reading of HCl or NaOH from the buret and also
calculate and record the change in volume of dispensed solutions. Dispose of the solution in your waste beaker.
Again, between each analysis, remember to rinse the pH probe with DI water and carefully dry the probe using
a Kimwipe.
2. Dispose of waste beaker solution in the appropriate waste container. Refill your 250-mL beaker with fresh DIwater. When the probe is not in use, it should be kept moist by being submerged/stored keep in DI-water.
Part D – Investigating Phosphate Buffers
PRIOR TO PREPARING THE BUFFER CORRESPONDING TO YOUR ASSIGNMENT PH RANGE, HAVE YOUR
LAB INSTRUCTOR APPROVE THE TWO PHOSPHATE SOLUTIONS YOU HAVE CHOSEN TO PREPARE YOUR
BUFFER.
1. Your instructor will assign each pair of students one of the three phosphate buffers to prepare. Make sure you
record your assigned pH range (2 - 3, 7 - 8, or 11 - 12) in your laboratory notebook.
2. Phosphoric acid is a triprotic acid (3 acidic hydrogens) having the following three respective ionization
constants: Ka1 = 7.5 x 10–3, Ka2 = 6.2 x 10–8, Ka3 = 4.2 x 10–13. Phosphate buffers with pH values near 2, 7, and
12 can be prepared using the appropriate phosphate conjugate acid-base pairs. Decide which of the following
2 solutions below you would use in order to prepare a phosphate buffer assigned by your instructor. Predict,
using the Henderson-Hasselbalch equation, what this pH should be.
A.
B.
C.
D.
0.10 M phosphoric acid (H3PO4)
0.10 M sodium dihydrogen phosphate (NaH2PO4; the monobasic salt which contains H2PO4– ions)
0.10 M sodium hydrogen phosphate (Na2HPO4; the dibasic salt which contains HPO42– ions
0.10 M sodium phosphate (Na3PO4; the tribasic salt which contains PO43– ions)
3. All these solution can be found in dispensing beakers on the long window sill. Using a graduated cylinder,
obtain 10.0 mL of the two solutions that you have chosen and your lab instructor has approved. Use the white
lined graduated cylinder to measure out the acid, while the blue lined should be used to measure the
9
conjugate base. Pour the solution from your graduate into a clean/dry beaker labeled. Carefully, without
banging the probe or splashing solution, mix the solution using your pH probe. Using the stylus, tap the file
cabinet icon to set up the new run. Measure and record the mean pH of this buffer. Compare the experimental
pH to the predicted value. Did you select the correct 2 solutions? If not, re-think this and try again.
Clean Up:
1. Rinse the pH probe with distilled water. Dispose of your waste beaker solution in the appropriate waste
container. Place the pH probe into a beaker containing clean DI water. Rinse and dry all beakers and flasks
and return them to their original location. Do NOT drain the HCl and NaOH solutions from the burets.
2. Delete any stored data from the computer interface by using the stylus to tap the “File” menu and selecting
“New”. Discard any stored data. The LabQuest should be left turned on for the next lab section. If the lab is
scheduled to end at 4:50 PM or 8:20 PM then the LabQuest should be turned off by pressing the silver button
on the top left of the device.
3. Have your lab instructor sign your notebook and submit the carbon copies of your notebook pages.
10
Suggested Setup for Lab Notebook (All lab data goes in the notebook; do not record lab
data here)
PART A – Measuring the pH of Acids and Bases
Team A
SOLUTION
pH
0.100 M HCl
0.0100 M HCl
0.00100 M HCl
0.100 M HC2H3O2
CALCULATED [H3O+] (M)
CALCULATED [OH–] (M)
CALCULATED [H3O+] (M)
CALCULATED [OH–] (M)
Team B
Solution
0.100 M NaOH
0.0100 M NaOH
0.00100 M NaOH
0.100 M NH4OH
PH
Show how you calculated the [H+] and [OH–] for the 0.100 M solution of the acid (Team A) or base (Team B). Note
that when you take the antilog (inverse log) of a number, the answer should contain one less significant figure.
PART B – Measuring the pH of Salt Solutions
Record the pH for the following aqueous salt solutions:
________
________
________
________
________
________
A.
B.
C.
D.
E.
F.
NaC2H3O2
Na2CO3
NaHSO4
NaHCO3
NH4Cl
AlCl3
sodium acetate
sodium carbonate
sodium hydrogen sulfate
sodium hydrogen carbonate
ammonium chloride
aluminium chloride
Arrange salt solutions in order of increasing pH (most acidic to most basic) using their assigned letters.
Complete and balance the following hydrolysis equations involving acidic or basic ions from the salt solutions:
CO32–(aq) + H2O(l) ⇌
HSO4–(aq) + H2O(l) ⇌
HCO3–(aq) + H2O(l) ⇌
NH4+(aq) + H2O(l) ⇌
11
PART C – Investigating Buffer Capacity
Buffer A
Buffer B
Final pH reading of buffer
__________
__________
Initial pH reading of buffer
__________
__________
Change in pH reading of buffer
__________
__________
Final buret reading
__________
__________
Initial buret reading
__________
__________
Volume of HCl dispensed to change pH by 1 unit
__________
__________
PART D – Investigating Phosphate Buffers
Assigned Ka value
______
pH Range
Solutions used to prepare phosphate buffer
________
Predicted pH of phosphate buffer
________
Measured pH of phosphate buffer
________
Acid
Conjugate Base
____________ and ____________
Show how you calculated the predicted pH of your phosphate buffer:
12
Experiment 7 Post-Lab Report Questions (to be submitted to lab instructor when the
corresponding MindTap report is due)
Your name:
Instructor:
1. Show calculation of [H3O+] and [OH–] from the pH of 0.010 M HCl you measured in part A.
Be sure to include units, and to report the correct number of significant figures.
2. Complete and balance the hydrolysis reactions for the reactions listed
a. CO32– + H2O =
b. HSO4– + H2O =
c. HCO3– + H2O =
d. NH4+ + H2O =
e. AlCl3 + H2O =
(hint: AlCl3(aq) is more accurately described as Al(H2O)63+)
3. Show calculation of the calculated pH of your phosphate buffer (part D). Be sure to include
units, and to report the correct number of significant figures.
13
4. Are aqueous solutions of sodium chloride acidic, basic, or neutral? Explain.
5. In part D, you prepared a phosphate buffer using two solutions. Why are two solutions
required?
6. Define the term “buffer capacity”. Which buffer solution from part C had higher capacity, A or
B?
The concentrations of acid and base in buffer B are twice the concentrations of buffer A (buffer
B is twice as strong as buffer A). Show that the ratio of the volume of titrant required to raise
the pH of each buffer 1 unit (VB/VA) is directly related to the ratio of the buffer strengths
(strengthB/strengthA)
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