Task 2: Determining the acid dissociation constant (Ka) for a weak acid (A.P2)
Aim
The purpose of this experiment is to accurately demonstrate a reading of pH of a
half-neutralised solution of ethanoic acid and from this determine a value for Ka.
Introduction
You will first accurately determine the titre of sodium hydroxide solution required to
neutralise 25.0 cm3 of 1.0M ethanoic acid. You will then neutralise exactly half of a new
sample of 25.0 cm3 of ethanoic acid with sodium hydroxide and accurately measure the pH
of the resulting solution in order to determine Ka.
Chemicals and Equipment
➔ ethanoic acid (approx. 0.1M)
➔ sodium hydroxide solution (0.1M)
➔ safety spectacles
➔ 1 burette, 50 cm3, and stand
➔ 1 funnel, small
➔ pipette, 25.0 cm3, and pipette filler
➔ 1 beaker, 250 cm3
➔ 1 conical flask, 250 cm3
➔ distilled water
➔ pH metre (accurately calibrated)
CAUTION:
Sodium hydroxide solution is corrosive. You must wear safety spectacles throughout the experiment.
Method
1. Clamp the 50 cm3 burette to the stand. Using a funnel, rinse and fill the burette with 0.1M sodium
hydroxide solution.
2. Using a pipette filler, rinse the pipette with some of the ethanoic acid and transfer precisely 25.0
cm3 of the solution into a 250 cm3 conical flask.
3. Add 2-3 drops of phenolphthalein indicator solution to the conical flask and swirl to mix.
4. Run sodium hydroxide solution from the burette, with swirling until the solution just turns pink
and is permanent.
5. Repeat the titration until you have obtained concordant titres. Calculate the mean titre (V cm3).
6. Pipette precisely 25.0 cm3 of the ethanoic acid solution this time into a 250 cm3 beaker. Do not
add phenolphthalein indicator.
7. Run exactly V/2 cm3 of sodium hydroxide solution from the burette into the beaker of ethanoic
acid. This is the volume required to neutralise exactly half of the ethanoic acid.
8. Swirl gently to ensure that the solutions are well mixed.
9. Place an accurately calibrated pH metre into the beaker so that the tip of the probe is completely
immersed in the solution you have made.
10. When the reading is stable, read and record the final pH of the solution in the beaker.
11. Repeat the half-neutralisation steps again in order to obtain a reliable average pH reading.
, Results of the experiment
Trial run 1st run 2nd run
3
Initial reading / cm 0.00 0. 00 0.00
Final Reading / cm3 25.1 24.4 24.4
Titre (volume used) / cm3 25.1 24.4 24.4
Mean titre (V) / cm3 24.4
V/2 / cm3 12.3
pH of half-neutralised ethanoic acid solution (1st run) 4.72
pH of half-neutralised ethanoic acid solution (2nd run) 4.72
pH of half-neutralised ethanoic acid solution (3rd run, if required) 4.73
Average pH value 4.72
Calculation
The acid dissociation constant for a weak acid (Ka) is given by the equation:
+ −
𝐾𝑎 = [ [𝐻𝐴]
𝐻 ][𝐴 ]
Therefore, when a weak acid is half-neutralised, the concentration of the weak acid [HA] is equal to the
concentration of its conjugate base [A-] (ie [HA] = [A-]), so the equation for Ka can be expressed as
Ka = [H+] or pKa = pH
The equation can be rearranged such as : Ka = 10-pH
Ka = 10 -4.71
ka = 1.95 x 10-5
∴ The value of Ka is 1.95 x 10 -5
Comparing the ka for ethanoic acid with the published ka value for ethanoic acid:
Therefore, my ethanoic acid value is 1.95 x 10-5 whereas the published value of the ethanoic acid is
slightly different which is 1.74 x 10-5. This suggest that there are some reason why the value is slightly
different from the published value. This can have happened because too much solution was added,
slightly crossing the meniscus line and causing inaccurate results. This also implies that parallel errors
may also arise; these errors happen when the buret's scale is not seen from a perpendicular angle. For
example, misreading the volume of burette can affect the titration result. From above the volume of
solution appears to be lesser, while from below the volume of solution appears to be higher. Moreover,
looking in the wrong spot is also another form of measurement inaccuracy. This is because a concave
curve is formed by a solution which is also known as the meniscus line and the volume is measured at
the bottom of the curve. The volume measurement will be inaccurate if the reading is taken from the
upper regions of the curve. Therefore, to improve the error is by reading the volume of solution at eye
level and at the right angle which makes the experiment more liable and accurate. Additionally, failing
to calibrate the apparatus correctly and failing to use a white tile to examine the colour changes can
lead to contamination and may affect the final result. This is because it is possible to make mistakes
when calibrating the equipment. When starting the practical it is necessary to rinse the burette and
Aim
The purpose of this experiment is to accurately demonstrate a reading of pH of a
half-neutralised solution of ethanoic acid and from this determine a value for Ka.
Introduction
You will first accurately determine the titre of sodium hydroxide solution required to
neutralise 25.0 cm3 of 1.0M ethanoic acid. You will then neutralise exactly half of a new
sample of 25.0 cm3 of ethanoic acid with sodium hydroxide and accurately measure the pH
of the resulting solution in order to determine Ka.
Chemicals and Equipment
➔ ethanoic acid (approx. 0.1M)
➔ sodium hydroxide solution (0.1M)
➔ safety spectacles
➔ 1 burette, 50 cm3, and stand
➔ 1 funnel, small
➔ pipette, 25.0 cm3, and pipette filler
➔ 1 beaker, 250 cm3
➔ 1 conical flask, 250 cm3
➔ distilled water
➔ pH metre (accurately calibrated)
CAUTION:
Sodium hydroxide solution is corrosive. You must wear safety spectacles throughout the experiment.
Method
1. Clamp the 50 cm3 burette to the stand. Using a funnel, rinse and fill the burette with 0.1M sodium
hydroxide solution.
2. Using a pipette filler, rinse the pipette with some of the ethanoic acid and transfer precisely 25.0
cm3 of the solution into a 250 cm3 conical flask.
3. Add 2-3 drops of phenolphthalein indicator solution to the conical flask and swirl to mix.
4. Run sodium hydroxide solution from the burette, with swirling until the solution just turns pink
and is permanent.
5. Repeat the titration until you have obtained concordant titres. Calculate the mean titre (V cm3).
6. Pipette precisely 25.0 cm3 of the ethanoic acid solution this time into a 250 cm3 beaker. Do not
add phenolphthalein indicator.
7. Run exactly V/2 cm3 of sodium hydroxide solution from the burette into the beaker of ethanoic
acid. This is the volume required to neutralise exactly half of the ethanoic acid.
8. Swirl gently to ensure that the solutions are well mixed.
9. Place an accurately calibrated pH metre into the beaker so that the tip of the probe is completely
immersed in the solution you have made.
10. When the reading is stable, read and record the final pH of the solution in the beaker.
11. Repeat the half-neutralisation steps again in order to obtain a reliable average pH reading.
, Results of the experiment
Trial run 1st run 2nd run
3
Initial reading / cm 0.00 0. 00 0.00
Final Reading / cm3 25.1 24.4 24.4
Titre (volume used) / cm3 25.1 24.4 24.4
Mean titre (V) / cm3 24.4
V/2 / cm3 12.3
pH of half-neutralised ethanoic acid solution (1st run) 4.72
pH of half-neutralised ethanoic acid solution (2nd run) 4.72
pH of half-neutralised ethanoic acid solution (3rd run, if required) 4.73
Average pH value 4.72
Calculation
The acid dissociation constant for a weak acid (Ka) is given by the equation:
+ −
𝐾𝑎 = [ [𝐻𝐴]
𝐻 ][𝐴 ]
Therefore, when a weak acid is half-neutralised, the concentration of the weak acid [HA] is equal to the
concentration of its conjugate base [A-] (ie [HA] = [A-]), so the equation for Ka can be expressed as
Ka = [H+] or pKa = pH
The equation can be rearranged such as : Ka = 10-pH
Ka = 10 -4.71
ka = 1.95 x 10-5
∴ The value of Ka is 1.95 x 10 -5
Comparing the ka for ethanoic acid with the published ka value for ethanoic acid:
Therefore, my ethanoic acid value is 1.95 x 10-5 whereas the published value of the ethanoic acid is
slightly different which is 1.74 x 10-5. This suggest that there are some reason why the value is slightly
different from the published value. This can have happened because too much solution was added,
slightly crossing the meniscus line and causing inaccurate results. This also implies that parallel errors
may also arise; these errors happen when the buret's scale is not seen from a perpendicular angle. For
example, misreading the volume of burette can affect the titration result. From above the volume of
solution appears to be lesser, while from below the volume of solution appears to be higher. Moreover,
looking in the wrong spot is also another form of measurement inaccuracy. This is because a concave
curve is formed by a solution which is also known as the meniscus line and the volume is measured at
the bottom of the curve. The volume measurement will be inaccurate if the reading is taken from the
upper regions of the curve. Therefore, to improve the error is by reading the volume of solution at eye
level and at the right angle which makes the experiment more liable and accurate. Additionally, failing
to calibrate the apparatus correctly and failing to use a white tile to examine the colour changes can
lead to contamination and may affect the final result. This is because it is possible to make mistakes
when calibrating the equipment. When starting the practical it is necessary to rinse the burette and
