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Summary assignment 1
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Summary of 10 pages for the course unit 1 at PEARSON (unit 13 assignment)
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Finding out the acid dissociation constant (Ka) for weak acid.A weak acid is characterized by its incomplete ionization when dissolved in
water. Ethanoic acid serves as an example of a weak acid. Upon dissolution
in water, ethanoic acid reacts to form hydroxonium ions and ethanoate ions.
These ions readily engage in a reversible reaction to regenerate the original
acid and water molecules. Specifically, ethanoic acid was employed as the
weak acid in this context.
Results :
Trial run 1st run 2nd run
Initial reading/cm 0.00 0.00 0.00
Final reading/cm 14.4 13.5 13.2
Titre/cm 14.4 13.5 13.2
Mean titre (v)/cm 13.7
v/2 cm 6.85
pH of half neutralized ethanoic acid solution 1st run 4.70
pH of half neutralized ethanoic acid solution 2nd run 4.70
pH of half neutralized ethanoic acid solution 3rd run 4.72
Average pH value 4.70
Calculations :
CH3COOH(aq) CH3COO-(aq) + H+(aq)
[CH3COO](aq) [H+](aq)
Ka =
[CH3COO](aq)
At half neutralizing point:
[CH3COO](aq) = [CH3COO-](aq)
Therefore: Ka = [H+](aq)
Ka = 10pH-
Ka = 10-4.70
For ethanoic acid : Ka = 1.9 x 10-5 mol/dm-3
The accuracy of pH calibration and equipment functionality are pivotal in titration processes,
where various factors can influence outcomes and introduce errors. Not using a white tile for
colour comparison may impact results, particularly when dealing with dark colours or an
excess of added indicator. Air bubbles within solutions can distort readings, while improper
calibration of the pH meter, coupled with inadequate drying to remove excess moisture, can
lead to inaccuracies. Moreover, if the liquid within the burette funnel exceeds its intended
level, it could alter readings, potentially affecting the determination of the Ka value.
, Insufficient mixing of solutions can result in uneven distribution of reactants, and surpassing
the meniscus line in burettes may yield inaccurate volume measurements. Additionally,
adding an excessive amount of hydrochloric acid to the beaker can skew results. Inaccuracies
may also arise if the pH meter bulb fails to fully immerse in the hydrochloric acid solution.
Overfilling the pipette beyond the required 25 cm3 or exceeding the recommended three
drops of phenolphthalein can introduce further inaccuracies.
Endpoint errors, influenced by the rapid colour changes, can complicate interpretation.
Human errors, such as failing to close the burette tap properly, may alter readings by
allowing unintended drops to enter the solution. Despite published Ka values for ethanoic
acid (1.74 x 10-5 mol/dm3), discrepancies may occur due to variations in experimental
conditions. Furthermore, the Ka value for ethanoic acid is reported as 1.9 x 10-5 mol/dm3.
Comparing the pH with the experimental pH of the acidic buffer solution
Results table:
Start with Initial pH Add Final pH Change in pH
Water 7.00 0.10 mol 1.00 -6.00
HCL
Water 7.00 0.10 mol 13.00 +6.00
NaOH
Buffer 4.74 0.10 mol 4.66 -0.08
HCL
buffer 4.74 0.10 mol 4.83 +0.09
NaOH
pH = -log Ka + (conjugate base)
(Weak acid)
pH= -log (1.7x10-5) + log 1.0
2.0
pH= 4.47
Comparing with theoretical
The buffer solution displayed a close adherence to its expected behavior, revealing insights
into its functioning and limitations. As observed, when assessing the buffer action, it was
noted that the concentration of hydrogen ions decreased by an amount less than anticipated
for the quantity of base added. This phenomenon underscores the capacity of the buffer to
mitigate pH changes, as both the acid and base components of the solution interact with the
hydroxide ions generated.
When introducing a strong acid into the buffer solution, the weak base component reacts
with the hydrogen ions from the acid, forming a weak acid and effectively absorbing the
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