How does altering the temperature of a system affect the acid dissociation constant of acetic acid in solution?
Introduction
The yield of a chemical product can be maximised using Le Chatelier’s
principle, which states that “a change… in an equilibrated system will
stimulate a response that partially offsets the change” (Pulido,
Chappell & McDuff, 2019). This experiment will explore this principle
in relation to the dissociation of the weak acid acetic acid (𝐶𝐻3 𝐶𝑂𝑂𝐻)
in water, by investigating how changes to the temperature of the
system affects the extent of the forward reaction, measured by
recording the change in pH of the system as equilibrium is established,
and utilising this to determine the value of the acid dissociation
constant (Ka) at each temperature.
Results
Alana Kirn
East Doncaster Secondary College
Unit 3 Chemistry, 2019
Table 1: Effect of Temperature on the Ka value for Acetic Acid
Temperature
Acid Dissociation Constant (Ka)
50°C
1.4089x10-4 M
40°C
8.9376x10-5 M
32°C
5.2565x10-5 M
19°C
1.8222x10-5 M
10°C
9.6429x10-6 M
Graph 1: Effect of Temperature on the Ka value for Acetic Acid
Aim
The aim of this investigation is to determine the effect that
temperature has upon the value of the acid dissociation constant (Ka)
value of acetic acid solution.
Hypothesis
As temperature increases from standard laboratory conditions of 25°C,
it is expected that the change in pH as equilibrium is established will
be greater, and thus, the acid dissociation constant will increase. This
is because the endothermic forward reaction should be favoured to
partially oppose the increase in temperature.
Methodology
In this investigation, 3.0M acetic acid (𝐶𝐻3 𝐶𝑂𝑂𝐻) was reacted with
water, which acted as a base, according to the following equation:
𝐶𝐻3 𝐶𝑂𝑂𝐻 (𝑎𝑞) + 𝐻2 𝑂(𝑙) ⇄ 𝐶𝐻3 𝐶𝑂𝑂− (𝑎𝑞) + 𝐻3 𝑂+ (𝑎𝑞)
The temperature was manipulated using a water bath to produce five
distinct temperatures ranging between 10°C and 50°C. The change in
pH over the course of the reaction at each of these was then recorded
using digital pH probes, and this information utilised to calculate
concentration of hydronium ions in solution. Hence, the acid
dissociation constant could be calculated using the following formula:
[𝐶𝐻3 𝐶𝑂𝑂− ][𝐻3 𝑂+ ]
𝐾𝑎 =
[𝐶𝐻3 𝐶𝑂𝑂𝐻]
Acetic acid can cause skin and eye irritation. To minimise this risk,
safety glasses and gloves should be worn. Care should also be taken
when heating the reaction mixture.
Discussion
Acids react by donating protons to a base, producing a conjugate base
and acid respectively (Commons & Commons, 2016). This process,
referred to as dissociation, is a chemical equilibrium that can occur to
different extents depending on the relative strength of the acid and
base, and can be represented numerically by the acid dissociation
constant, denoted by the symbol Ka (Flowers, Theopold & Langley,
2019). In this investigation, the acid dissociation constant of acetic acid
(𝐶𝐻3 𝐶𝑂𝑂𝐻) was calculated at various temperatures. It was found that
this value decreased exponentially with decreasing temperature, with a
value of 1.4089×10-4 M calculated at the highest reached temperature
of 50°C, as opposed to 9.6429×10-6 M at 10°C. Furthermore, at
temperatures greater than 32°C, the pH decreased from its initial
value, suggesting that the reaction proceeded in the forward direction,
while at lower temperatures the reverse reaction was favoured.
This can be explained by considering Le Chatelier’s principle, which
states that when a change is made to an equilibrated system, the
system will favour one direction to partially oppose this change. Hence,
as temperature is increased or decreased, equilibrium will shift to
consume or produce heat respectively to balance the temperature
change. The disassociation of acetic acid in an endothermic process;
there is an overall decrease in temperature of the system over the
course of the reaction (Bylikin, Horner, Murphy & Tarcy, 2014).
Thus, an increase in temperature from standard laboratory conditions
of 25°C is expected to cause equilibrium to shift to the right, as
observed, in order to partially consume the excess heat. Likewise, the
observed favouring of the reverse reaction with a decrease in
temperature can also be expected, as the reverse, exothermic reaction
would replace some of the lost heat.
Potential sources of error within this investigation may result from
inaccurate measurements of pH, as the digital probe utilised recorded
the pH of distilled water as approximately 4.52, rather than the
expected value of 7.0. This suggests that the recorded pH values were
much lower than their true values due to a calibration error. Inaccurate
recording of temperature, imprecise measurements of reactants and
variations from the assumed concentration of 3.0M acetic acid may
have further impacted the validity of results. These errors could be
reduced by utilising more accurate measuring apparatus, such as
volumetric glassware. Furthermore, the reactants may not have been
equally dispersed throughout the beaker when the initial pH value was
recorded. This would result in an inconsistent concentration of acetic
acid throughout the reaction mixture, resulting in an inaccurate pH
reading. This could be improved by stirring the mixture.
To provide a greater insight into the effect of temperature on the acid
dissociation constant of acetic acid, further experimentation could be
undertaken, utilising a wider range of temperatures. Furthermore, the
experiment could be repeated using various other weak acids, in
addition to strong acids, strong bases and weak bases to provide a
point of comparison.
Conclusion
In this investigation, the effect of temperature on the dissociation
reaction of acetic acid with water was explored. As hypothesised, it
was found that the acid dissociation constant increased with
temperature, with values of 9.6429×10-6 M and 1.4089×10-4 M
observed at 10°C and 50°C respectively.
Bibliography
Pulido, K., Chappell, C., & McDuff, A. (2019). Effect of Temperature on Equilibrium. Retrieved from
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/
Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Le_Chatelier's_Principle/
Effect_Of_Temperature_On_Equilibrium_Composition
Bylikin, S., Horner, G., Murphy, B., & Tarcy, D. (2014). Chemistry Course Companion (pp. 195-196).
Oxford: Oxford University Press.
Flowers, P., Theopold, K., & Langley, R. (2019). 6.4: Equilibria Involving Weak Acids and Bases.
Retrieved from https://chem.libretexts.org/Courses/University_of_California_Davis/UCD_Chem_
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