CHEM 123L University of Waterloo Concentration of Stock Acetic Acid Lab Template
Fill in the template provided. 7 videos. background and procedure provided. Need reference and in-text citation.
Use the data from file: Data set TB
The data set I provided contains only digital values, so all observations must come from the experimental videos
https://drive.google.com/file/d/1Pb8w51h2diu_vLitg5qvN2fwSTtgUyhP/view?usp=sharingb
Experiment 4: pH Chemistry
Acid – Base Titrations and Buffers
Techniques introduced in this experiment:
• Use of a burette in a titration
• Use of a volumetric pipette
• LEARN content to review: Lab Techniques
o Using a pipette
o Titration using a burette
Type of post-lab report:
• Modified formal report (report sheets + formal discussion)
• LEARN content to review: Lab reports
o General instructions for lab reports
o Modified formal lab reports
Learning outcomes: By completing this experiment and report, you will practice:
• preparing a titration curve and recognizing points on a titration curve
• calculating pH values at different regions of a titration curve
• using data from a titration to determine solution concentrations
• Preparing a buffer solution and identifying factors that influence the ability of a
buffer to resist pH change
Introduction:
pH is a measure of the acidity or alkalinity of solution. By definition pH = -log [H+]. pH can be
approximated using indicator dyes or measured accurately with a pH meter. pH meters
indicate pH by measuring the electromotive force (emf) developed by a cell, using the solution
as an electrolyte. The cell generally consists of a calomel electrode and a glass membrane
electrode; immersed in the solution under test, usually at room temperature. The response of
the glass electrode is proportional to the hydrogen ion concentration in the solution. The small
electromotive force developed is fed into an electronic amplifier circuit and boosted to the
point where it can be read across a broad scale, often with great sensitivity.
The relationship between pH and the emf developed will not be covered until later years. It is
sufficient to note that emf is a linear function of pH and meter scale can be calibrated to read
either in volts or directly in terms of pH units.
pH meters require calibration with a buffer solution of known pH, ideally close to the range
under test. This will be done prior to your arrival in the lab. pH meters are valuable
instruments, the electrodes are fragile and should be handled with care. Do not allow the
magnetic stirring bar to strike the electrodes.
Acid – Base Titrations:
The acid strength of a Brønsted-Lowry acid refers to the ability of the acid to relinquish
protons. A strong acid gives up protons more readily than a weak acid. If an acid is strong and
easily loses protons, then its conjugate base does not readily accept protons. In this way, the
strength of an acid is inversely related to the strength of its conjugate base; the stronger the
acid, the weaker the conjugate base; the stronger the base the weaker the conjugate acid. (See
table 4-1). The equilibrium constant for a weak acid in water is called an acid ionization
constant, Ka. For example, phenol ionizes in water according to the equilibrium equation:
H+ + C6H5O-
C6H5OH (aq)
and its ionization constant at 25oC is:
Strongest
Acid
Strength
Weakest
Table 4-1:
Ka =
[H + ][C6 H 5O- ]
= 1.28 x 10-10
[C6 H 5OH]
perchloric acid
HClO4
ClO4-
nitric acid
hydrochloric acid
sulfuric acid
HNO3
HCl
H2SO4
NO3ClHSO4-
hydronium ion
phosphoric acid
H3O+
H3PO4
H2O
H2PO4-
acetic acid
carbonic acid
CH3COOH
H2CO3
CH3COOHCO3-
dihydrogen
phosphate ion
ammonium ion
hydrogen cyanide
hydrogen
carbonate ion
methylammonium
ion
water
methane
H2PO4-
HPO42-
NH4 +
HCN
HCO3-
NH3
CN
CO32-
perchlorate
ion
nitrate ion
chloride ion
hydrogen
sulfate ion
water
dihydrogen
phosphate ion
acetate ion
hydrogen
carbonate
hydrogen
phosphate ion
ammonia
cyanide ion
carbonate ion
CH3NH3 +
CH3NH2
methylamine
H2O
CH4
–
hydroxide ion
methide ion
OH
CH3
–
Some Conjugate Acid-Base Pairs
Weakest
Base
Strength
Strongest
Buffers:
A buffer solution is a solution that resists changes in pH. It is a system that maintains a
relatively steady H+(aq) concentration by shifting the equilibrium between a conjugate acidbase pair.
A buffer solution can be made by dissolving almost equal concentrations of a weak acid and its
conjugate base (a salt of the weak acid) in water. For example, a solution of 1 M acetic acid and
1 M sodium acetate is a buffer solution. Buffer solutions can also be made from weak bases
and their conjugate acids (e.g., NH3 and NH4Cl). Buffer solutions are very important in many
fields of chemistry and biochemistry.
The pH of a buffer system can be calculated by substituting the acid and base concentrations
into the Ka expression (2) or using the Henderson-Hasselbalch equation (3).
HA
H+ + A¯
Ka =
[H+ ][A- ]
[HA]
pH = pK a + log
[A – ]
[HA]
(1)
(2)
(3)
The concentration ratio, [A-]/[HA], is called the buffer ratio. The pH of a buffer solution
depends directly on the logarithm of the buffer ratio; it increases when the ratio increases and
decreases when the ratio decreases. Observe that only the ratio of base to acid is important
and not the concentration. Thus, the numbers of moles of A¯ and HA can be substituted for
concentrations because the mole ratio in a given solution is the same as the concentration
ratio.
The ability of a buffer to resist a change in pH increases with increasing concentrations of
conjugate acid and base, since a more concentrated buffer will be able to neutralize more
added strong acid or base. The number of moles of strong acid or strong base needed to change
the pH of one litre of buffer by one unit is called the buffer capacity. The larger the buffer
capacity, the more resistant the buffer is to changes in pH.
The buffer capacity is also affected by the buffer ratio. A buffer solution is most resistant to a
change in pH if the buffer ratio is one, i.e. when the pH is equal to the pKa. The farther the ratio
deviates from one, the less buffering capacity the buffer has. Try some calculations using
equation (3) to verify this observation.
Notes:
Bronsted-Lowry acid ® substance that can donate a hydrogen ion (H+)
Bronsted-Lowry base ® substance that can accept a hydrogen ion (H+)
CH3COOH + ¯OH
acid
base
CH3COO¯
+
H2O
conjugate base conjugate acid
Lewis acid ® substance that can accept a pair of electrons
Lewis base ® substance that can donate a pair of electrons
FeBr3 +
Br2
Lewis acid Lewis base
H3O+ + OH-
2H2O
FeBr4¯ + Br +
or
H+ + OH-
H2O
Kw =[H+ ][ OH – ] = 1 x 10-14
pH = -log[H+ ]
pOH = -log[OH – ]
pKw = -log(1 x 10-14)
pH + pOH = 14
neutral solution ® [H+] = [OH ¯ ] = 1 x 10-7 ; pH = 7 @ 25 oC
acidic solution ® [H+] > 1 x 10-7 ; pH below 7
basic solution ® [H+] < 1 x 10-7 ; pH above 7
Ionization constants Ka and Kb:
General formula for an acid: HA
HA
H+ + A-
Ka =
General formula for a base: B
B + H2O
BH+ + OH-
[H + ][A - ]
[A - ]
pH = pKa + log
[HA]
[HA]
KaKb = Kw = 1 x 10-14
Kb =
strong acid [H+] , Ka , pKa ¯
weak acid [H+] ¯ , Ka ¯, pKa
[BH + ]
[OH - ][BH + ]
pOH = pKb + log
[B]
[B]
pKa + pKb = pKw = 14
strong base [-OH] , Kb , pKb ¯
weak base [-OH] ¯ , Kb ¯, pKb
The Henderson-Hasselbalch equation is just a mathematical manipulation of the Ka equation
CH3COOH
CH3COO¯ + H+
-
Ka =
[CH 3 COO ][H + ]
[CH 3 COOH]
Taking the log of both sides:
logK a = log
[CH 3 COO - ]
+ log[H + ]
[CH 3 COOH]
- log [H+ ] = - logK a + log
[CH3COO - ]
[CH3COOH]
by definition pH = -log[H +] and pKa = -log K
[CH 3 COO - ]
pH = pK a + log
[CH 3 COOH]
Experiment 4: pH Chemistry
Acid – Base Titrations and Buffers
Procedure:
Sign out from the storeroom:
Before you begin:
Make sure that your pH meter has been calibrated. Look for a graphic of an electrode beside a
bar graph on the top left of the display and a graphic on the right- hand side of the display that
looks like a “circle square root A”. If you do not see these symbols, ask your TA for help, it may
be that your meter needs to be re-calibrated.
Equipment set-up diagram:
Part A: Titration of a weak acid with a strong base
Pipette 25.00 mL of 0.1 M (record the exact molarity) acetic acid into a clean 400 mL beaker
after first placing a magnetic stirring bar on the bottom of the beaker. Add ~150 mL of
deionized water and 2 drops of the indicator phenolphthalein. Place the beaker on the
magnetic stirrer. Lower the electrode into position, approximately one inch above the bottom
of the beaker.
Obtain ~60 mL of an unknown NaOH solution (for part A) from the front bench. After properly
preparing the burette, fill it with unknown NaOH and prepare it for titration. Place it in
position to titrate the NaOH solution into the CH3COOH solution, as shown in the diagram.
Read and record the burette reading and pH, both to two decimal places. Add NaOH to the acid
slowly. Initially take burette and pH readings after every 2.00 - 3.00 mL addition.
As the equivalence point approaches the pH will rise more rapidly; add smaller (~ 0.10 mL)
volumes of NaOH before each pH reading. Start to reduce the volume of each titrant
addition at a pH of approximately 5.5. Note the volume at which the indicator changes
colour. Carry the titration beyond the endpoint, continue with small volume additions until
the pH plateaus*, and note how pH changes in this region (add 40.00 mL NaOH in total).
(* when the pH is changing only in the second decimal place, you have reached a plateau)
Part B: Titration of a weak base with a strong acid
Pipette 25.00 mL of 0.1 M (record the exact molarity) ammonia into a clean 400 mL beaker
after first placing a magnetic stirring bar on the bottom of the beaker. Add ~150 mL of
deionized water and 2 drops of the indicator methyl red. Place the beaker on the magnetic
stirrer. Lower the electrode into position, approximately one inch above the bottom of the
beaker.
Obtain ~60 mL of an unknown HCl solution (for part B) from the front bench. After rinsing a
burette, fill it with unknown HCl and prepare it for titration. Place it in position to titrate the
HCl solution into the ammonia solution.
Read and record the burette reading and pH, both to two decimal places. Add HCl solution to
the base slowly. Initially take burette and pH readings after every 2.00 – 3.00 mL addition.
As the equivalence point approaches the pH will drop more rapidly; add smaller (~ 0.10 mL)
volumes of HCl before each pH reading. Start to reduce the volume of each titrant addition
at a pH of approximately 8.5. Note the volume at which the indicator changes colour. Carry
the titration beyond the endpoint, continue with small volume additions until the pH plateaus*,
and note how pH changes in this region (add 40.00 mL HCl in total).
(* when the pH is changing only in the second decimal place, you have reached a plateau)
Part C: Preparation and Dilution of a Buffer
Use a graduated cylinder to measure volumes. (This will mean that volume measurements in
Part C are not as accurate as Parts A and B, as we are only looking for trends in the data.)
• Obtain 100 mL of 1 M acetic acid and 100 mL of 1 M sodium acetate solutions from the
stock bottles. Record the actual concentrations of these solutions as indicated on the stock
bottles.
•
Measure the pH of each of these solutions using the pH meter and record the values on
your data sheet.
• Prepare a buffer solution by mixing 100 mL 1 M acetic acid with 100 mL 1 M sodium acetate
solution. Stir.
• Prepare a 1/5 (one part in 5 parts) diluted buffer solution by adding 40 mL of buffer to 160
mL deionized water. Stir.
• Prepare a 1/50 (one part in 50 parts) diluted buffer solution by adding 20 mL of the 1/5
buffer to 180 mL deionized water. Stir.
• Measure the pH of the 3 buffer solutions. Save the buffer solutions.
Addition of strong acid to buffered and unbuffered solutions:
• Clean your burette, rinse and then fill with 3 M HCl
• With a graduated cylinder, measure 75 mL of the original buffer solution into a 200 mL
Berzelius beaker and place the stir bar in this solution. Set up your burette, stirrer and pH
probe as in part A/B.
• Using the burette, add 1 mL of 3 M HCl to the buffer, read and record the pH. Continue
adding acid in 1 mL increments until 15 mL of HCl has been added. Record the pH after
each 1 mL addition of HCl.
• Repeat the addition of HCl three more times following the instructions in the previous
step. For these three trials, use:
§
75 mL of the 1/5 diluted buffer, then,
§
75 mL of 1/50 diluted buffer and finally
§
75 mL of deionized water
• Record the pH after every 1 mL addition until a total of 15 mL of HCl has been added for
each trial.
Addition of strong base to buffered and unbuffered solutions:
• Clean the burette again, rinse and then fill with 3 M NaOH
• With a graduated cylinder, measure 75 mL of the original buffer solution into a 200 mL
Berzelius beaker and place the stir bar in this solution. Set up your burette, stirrer and pH
probe as before.
• Using the burette, add 1 mL of 3 M NaOH to the buffer, read and record the pH. Continue
adding acid in 1 mL increments until 15 mL of NaOH has been added. Record the pH after
each 1 mL addition of NaOH.
• Repeat the addition of NaOH three more times following the instructions in the
previous step. For these three trials, use:
§
75 mL of the 1/5 diluted buffer, then,
§
75 mL of 1/50 diluted buffer and finally
§
75 mL of deionized water
• Record the pH after every 1 mL addition until a total of 15 mL of NaOH has been added
for each trial.
Data Set TB 09
Part A
Part B
[Acetic Acid] = 0.1003 M
[Ammonia] = 0.1034 M
TB 09
Vol of NaOH added
(mL)
0.00
2.04
4.11
6.04
8.14
10.01
12.34
14.15
16.27
18.45
19.78
20.03
20.14
20.19
20.28
20.42
20.51
20.64
20.72
20.83
20.94
21.05
21.13
21.21
21.29
21.41
21.55
21.62
21.73
21.80
21.91
21.99
TB 09
Vol of HCl added
(mL)
0.00
2.32
3.76
5.38
6.35
10.69
12.66
14.54
17.23
18.89
20.53
20.56
20.61
20.88
20.92
20.97
21.13
21.18
21.25
21.34
21.46
21.52
21.68
21.74
21.79
21.85
21.91
22.04
22.13
22.24
22.33
22.47
pH
3.49
3.92
4.17
4.39
4.57
4.72
4.88
5.02
5.14
5.32
5.50
5.52
5.54
5.54
5.56
5.58
5.60
5.63
5.67
5.68
5.69
5.73
5.76
5.78
5.80
5.82
5.83
5.86
5.87
5.89
5.90
5.90
pH
10.65
10.15
9.98
9.75
9.59
9.42
9.27
9.11
8.91
8.69
8.34
8.30
8.28
8.19
8.14
8.13
8.08
8.04
8.00
7.95
7.90
7.85
7.79
7.72
7.64
7.54
7.31
7.07
6.82
6.59
6.39
6.16
22.04
22.10
22.14
22.23
22.33
22.42
22.51
22.62
22.70
22.84
22.93
23.01
23.09
23.21
23.39
23.45
23.54
23.60
23.69
23.78
23.87
23.99
24.09
24.16
24.23
24.35
24.48
24.53
24.64
24.71
24.82
24.91
25.01
25.10
25.23
25.29
25.42
25.49
25.56
25.62
5.91
5.92
5.92
5.94
5.97
5.98
5.99
6.02
6.03
6.05
6.06
6.09
6.11
6.14
6.16
6.23
6.27
6.29
6.34
6.38
6.44
6.48
6.55
6.67
6.74
6.80
6.93
7.02
7.32
8.67
9.24
9.85
10.11
10.29
10.45
10.60
10.72
10.81
10.85
10.91
22.56
22.63
22.71
22.76
22.89
23.00
23.08
23.21
23.34
23.49
23.58
23.69
23.81
23.98
24.10
24.20
24.99
25.92
26.92
27.99
29.91
31.00
33.00
34.98
37.08
39.03
40.01
6.02
5.71
5.36
4.96
4.04
3.83
3.70
3.59
3.48
3.40
3.33
3.27
3.22
3.18
3.13
3.11
2.93
2.78
2.66
2.57
2.49
2.39
2.30
2.24
2.15
2.10
2.05
25.74
25.80
25.90
26.00
26.15
26.29
26.78
28.98
32.23
34.97
36.87
10.92
10.98
11.04
11.08
11.09
11.13
11.16
11.49
11.88
12.02
12.07
40.04
12.09
Part C
Data set
[NaOH] (M)
[HCl] (M)
[Acetic Acid] (M)
initial pH of acetic acid
[Sodium acetate] (M)
initial pH of sodium acetate
TB 09
3.0
3.0
0.9986
2.24
1.027
7.93
Addition of HCl
TB 09
Data set
Volume of HCl added
(mL)
0
1
2
3
4
5
6
7
8
9
1/1 buffer
1/5 buffer
1/50 buffer
pH
pH
pH
DI water pH
4.70
4.59
4.50
4.44
4.35
4.25
4.14
4.07
3.95
3.84
4.72
4.38
3.81
2.19
1.52
1.28
1.12
1.00
0.91
0.84
4.69
2.14
1.79
1.45
1.34
1.18
1.07
0.95
0.87
0.81
5.96
1.52
1.42
1.28
1.13
1.05
0.96
0.88
0.82
0.76
10
11
12
13
14
15
3.68
3.44
3.03
2.10
1.52
1.25
0.78
0.74
0.69
0.65
0.62
0.60
0.76
0.72
0.68
0.64
0.62
0.59
0.70
0.66
0.64
0.59
0.58
0.55
Addition of NaOH
TB 09
Data set
Volume of NaOH added
(mL)
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
1/1 buffer
1/5 buffer
1/50 buffer
pH
pH
pH
DI water pH
4.73
4.81
4.89
4.95
5.01
5.14
5.23
5.31
5.39
5.49
5.62
5.81
6.05
6.73
12.24
12.76
4.75
5.21
5.56
12.14
12.69
12.93
13.04
13.13
13.21
13.29
13.38
13.45
13.48
13.50
13.52
13.53
4.74
11.64
12.82
12.99
13.09
13.18
13.25
13.31
13.36
13.42
13.46
13.51
13.54
13.57
13.59
13.61
5.86
12.73
12.97
12.94
13.14
13.23
13.34
13.42
13.49
13.54
13.59
13.67
13.73
13.81
13.85
13.89
What could go wrong in pH measurement?
(These were the errors that should have been in Mobius, but cannot be uploaded while the videos are
active)
Titration:
-
-
Forgetting to add indicator: you will not be able to determine the indicator end point and the
concentration of titrant, or compare it to the concentration you find graphically
Adding titrant too fast / in too large volume / missing the equivalence point: if done in the lab
this mistake cannot be recovered from, you would simply need to go back to the beginning and
repeat the titration
Not recording volume or pH to 2 decimal places: this will limit the number of significant digits
you can include in your post-lab report answers
Not continuing to add titrant past the indicator end point: to produce a full titration curve you
must titrate beyond the equivalance point, in our case to a total of 40 mL of titrant added.
Not stirring: thorough stirring is needed during the titration to ensure that the solution is well
mixed, and that the pH is uniform throughout the solution. Stirring also helps to speed up the
reaction meaning that the pH will stabilize and the pH meter will produce a reading faster.
Buffers:
-
-
Adding strong acid to the full volume of buffer instead of 75 mL: you will not exceed the
buffering capacity for some of the buffers (thus you will not see what we are hoping to prove)
and will not have any buffer remaining for the next step (adding strong base to a 75 mL volume
of buffer)
Accurately measuring the volume of strong acid or strong base added to the buffer solutions:
this is not really an error, more of a time waster, since in this part we are only looking at trends.
The volumes could be recorded very accurately, but it is not necessary here.
pH measurement:
-
-
-
pH probe not immersed far enough into the test solution: for most pH probes to read correctly,
the tip will need to extend about 1 cm into the test solution, this ensures that the porous glass
tip and the ceramic junction are both in solution.
Forgetting to press the Read button: the pH meter only takes a new reading when this button is
pushed, it must be done every time you wish to record a ph value
Not allowing the pH meter to reach a stable value: the pH meters used in the first year
chemistry lab will have a flashing decimal point when the meter is actively making a pH
measurement. When the value is stable and ready to record, the pH meter will beep, and the
decimal point will stop flashing.
Pressing any other button on the pH meter: pressing the Cal button will initiate a calibration
sequence, and if done during a titration may mean that you will need to start over; pressing the
power button will turn the meter off, but it should maintain it’s original calibration when turned
back on, so no harm done.
Sample calculations for pH:
(Refer to the graph on the third page)
① Before addition of NaOH, the pH is determined by the dissociation constant of the weak
acid.
CH3COOH
0.1 – x
CH3COO- + H+
x
x
[CH3 COO- ][H+ ]
x2
Ka =
=
= etc…
[𝐶𝐻! 𝐶𝑂𝑂𝐻]
0.1-x
② After addition of NaOH: For every mole of base (OH-) added the moles of CH3COOH
decreases and the moles of CH3COO- increases proportionally due to the reaction of OHwith H+.
OH
CH3COOH
0.1 – x – OH-
CH3COO- + H+
x + OH-
H2O
x
[CH3 COO- ][H+ ] [𝑥 + 𝑂𝐻" ][𝐻# ]
Ka =
=
[𝐶𝐻! 𝐶𝑂𝑂𝐻]
0.1 - x - 𝑂𝐻"
Assume that the effect of self-dissociation is much smaller than the shift in the equilibrium
caused by the addition of NaOH (x
