Lab 6: Acid-Base Equilibria: Determination of Ka and Investigation of Buffers
Objectives:
1. To determine the equilibrium constant, Ka, for a weak acid.
2. To study the buffer capacity of a buffer.
Introduction:
Part A: Determination of the Dissociation Constant of a Weak Acid
For a general weak acid of the form HA that is dissolved in water, the following equilibrium is set
up.
HA (aq)
H+(aq) + A-
(6-1)
where HA (aq) is the weak acid, H+(aq) is the hydrogen ion and A- is the conjugate base. The
equilibrium constant Ka, for this reaction is called the acid dissociation constant of the acid HA.
(6-2)
At the half equivalence point, half the acid has been converted to the salt. Therefore, [A -] = [HA]
and equation (6-2) becomes
[H+] = Ka
(6-3)
pH = pKa.
(6-4)
or
Part B - Preparation and Properties of a Buffer
The Henderson-Hasselbalch equation is used for the calculation of the pH or composition of a
buffer solution. With mixtures consisting of weak acids and their salts the convenient
approximations [HA] = total acid concentration, and [A- ] = salt concentration, can often be made.
Hence an acetic acid/acetate buffer solution containing 0.1 M acetic acid and 0.05 M sodium
acetate would have a pH of 4.4.
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pH = pKa + log [conjugate base] = 4.7 + log (0.05) = 4.7 + (-0.3) = 4.4
[conjugate acid]
(0.1)
(6-5)
Buffer solutions behave as follows: when a strong acid is added, the H + from that acid combine
with a portion of the anion to form undissociated acid, thereby removing most of the added H +
from the solution
H+ (aq) + A- (aq)
HA (aq)
When strong base is added part of the undissociated acid reacts to form anions
OH- (aq) + HA (aq)
A- (aq) + HOH (l)
In Part A, we saw that at the half equivalence point, where [A-] = [HA] and equation (6-2)
becomes
[H+] = Ka
(6-3)
Using the Henderson-Hasselbalch equation (6-5)
(6-6)
When [A-] = [HA], equation (6-6) reduces to
pH = pKa + log 1
(6-7)
pH = pKa.
(6-8)
and
This means that an acid is half dissociated when the pH of the solution is numerically equal to the
pKa of the acid.
Click on this link
to view a buffer solution
(Flash plugin required)
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Apparatus:
1.
2.
3.
4.
5.
6.
7.
8.
9.
Burettes – 2x per pair of students
100 mL beakers – 3x per student (check lockers)
100.0 mL volumetric flask – 1 per pair of students
bulbs
Stir bar
Stir plate
pH meter – 1 per pair of students
100 mL graduated cylinder
50 mL graduated cylinder
Solutions:
1. pH 4 and pH 7 buffers
2. Part A: 0.20 M NaOH - 50 mL per pair of students – in dispenser
3. Part B:
o 0.10 M NaOH - 80 mL per pair of students
o 0.10 M HCl - 80 mL per pair of students
o 0.100 M acetic acid - 10 mL per pair of students – in dispenser
o 0.100 M sodium acetate - 10 mL per pair of students – in dispenser
o Phenolpthalein indicator
o Metacresol purple or bromocresol green indicator
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Procedure:
Part A - Determination of the Dissociation Constant of a Weak Acid
1. Weigh out 1.0 g of the unknown acid into a clean 250 mL Erlenmeyer flask. Measure 100
mL of distilled water in a graduated cylinder and use it to dissolve the acid with mixing in
the flask.
2. Measure 50 mL of the acid solution with a graduated cylinder into a new, clean 250 mL
Erlenmeyer flask.
3. Add 2 to 3 drops of phenolphthalein to the 50 mL acid in the new flask. Titrate to the
phenolphthalein end point with 0.2 M NaOH. The light pink colour should persist for
approximately 30 seconds when the end point is reached.
4. Calibrate the pH meter as demonstrated.
1. Immerse pH electrode into the pH 7 buffer solution.
2. Adjust the Standardize knob to display pH=7.00 .
3. Rinse the electrode with distilled water and blot dry.
4. Immerse the pH electrode into the pH 4 buffer solution.
5. Adjust the Temperature knob to display pH=4.00.
5. Mix the original 50 mL of acid solution with the neutralized solution and measure the pH
of the solution.
Part B - Preparation and Properties of a Buffer
Recipe for making a buffer solution:
Buffer Solution
0.100 M acetic acid
0.100 M sodium acetate
A
10.00 mL
50.00 mL
B
20.00 mL
40.00 mL
C
30.00 mL
30.00 mL
D
40.00 mL
20.00 mL
E
50.00 mL
10.00 mL
1. Your instructor will assign a buffer solution ID to your group. Using the above recipe,
prepare 60.0 mL of the assigned buffer. The solutions are in bottle-top dispensers that
are set to 10.0 mL. Record the pH of the prepared buffer. Cover the beaker with a watch
glass.
2. Record the exact concentrations of the NaOH and HCl solutions which are approximately
0.10 M. Acclimatize and fill a buret with NaOH and another burette with HCl. Label the
burettes. Do not take more than 60 mL of either reagent to start.
3. Measure 20.0 mL of your buffer into two clean and dry 100 mL beakers.
4. Put a magnetic stirrer in one and insert the pH electrode. Check that the probe is
immersed and not touching the spinning stir bar. Add three drops of phenolphthalein
indicator. Record the initial pH of the solution.
5. Titrate with small amounts of NaOH and record the pH and volumes after waiting about
20 seconds after each addition. Continue with small amounts until the pH has changed
1.0 pH unit above the pKa.
6. Continue adding NaOH with smaller amounts until the basic equivalence point is reached
(it will be close to the colour change of the indicator). Record the pH and volumes after
each addition after mixing for 20 seconds. Aim to record a pH and volume reading for
every 0.2 pH change.
7. Add NaOH a few mL beyond the equivalence point and record the volume and pH. You
can stop when pH > 12.
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8. Using a new buffer sample, repeat the experiment using 0.10 M HCl using either
metacresol purple or bromocresol green as an indicator with the second 20 mL sample of
buffer. The colour change of the indicator may not coincide with the actual equivalence
point but it can be used as a signal as the equivalence point is reached. You can stop
when pH < 2.
9. Your instructor will provide you with the class data for all the assigned buffer solutions for
carrying out data analysis.
Datasheet:
Part A - Determination of the Dissociation Constant of a Weak Acid
Measured pH of the solution ___________________________________________
Part B - Preparation and Properties of a Buffer
Assigned Buffer Solution:
__________________________________________
Initial pH of the buffer solution ___________________________________________
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Titration with NaOH
Strong Acid/Strong Base Titration
Volume of Acid (mL)
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pH
Strong Acid/Strong Base Titration
Volume of Acid (mL)
pH
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Titration with HCl
Strong Acid/Weak Base Titration
Volume of Acid (mL)
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pH
Strong Acid/Weak Base Titration
Volume of Acid (mL)
pH
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Postlab Questions:
Part A: Determination of the Dissociation Constant of a Weak Acid
1. Calculate the Ka of the unknown acid in Part A.
Part B - Preparation and Properties of a Buffer
1. Your instructor will provide you with the class data on the website to download.
2. Graph the pH titration data for your buffer solution in Part B. Combine the data from the NaOH
and HCl titrations on the same graph. The y-axis is pH. The x-axis is volume of reagent added.
Use 1 mL as the large scale on the x-axis. Place the y-axis with an increasing pH scale near the
middle of the graph paper. Plot the titration volume of HCl on the left side of the y-axis with
increasing volume going right to left. Plot the titration volume of NaOH on the right side of the yaxis with increasing volume going from left to right. Draw a smooth curve through the data points
and label the following:




initial pH of the buffer solution
the basic equivalence point
the acidic equivalence point
the pKa
3. Download the class data on the website. If you can, on the same graph, graph the titration data
for the different buffer solutions (A to E). Use different colours to distinguish the different buffer
solutions. Create a legend to identify coloured curves to the different buffer solutions.
4. The buffer capacity of a buffer solution is the amount of acid or base that can be added to a
volume of a buffer solution before its pH changes significantly. Estimate the buffer capacity of the
different buffer solutions for:
(a) NaOH, and (b) HCl. Explain.
5. From the class data, which buffer solution has the greatest buffer capacity for (a) NaOH, and
(b) HCl. Explain.
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