Introduction new solution to change colour at a

Introduction

            The laboratory conducted used the
method of acid-base titration. Titration is the addition of precise volumes of
a solution in a burette to a measured volume of a sample solution (1). An
acid-base titration is a process of obtaining quantitative information of a
sample using an acid-base reaction by reacting with a certain volume of reactant
of a known concentration (2). The purpose of an acid-base titration is to
determine the concentration of an unknown sample by titration with a standard
solution, which could then be used to calculate the pH of the solution at the
equivalence point (3). The titrant is the solution in the burette, which is
added to the sample in the Erlenmeyer flask. The concentration of the titrant
is known, and the volume that is added to neutralize the sample is used to
calculate the concentration of the sample (1). The process of an acid-base
titration also involves the addition of an appropriate indicator, which causes
the new solution to change colour at a certain pH level to determine the
equivalence point, and therefore the point of neutralization (1). The indicator
phenolphthalein results in the titrated solution to become a light pink colour
when the solution is at its endpoint. Because the titration involved a weak
acid and a strong base, the pH of the solution in the Erlenmeyer flask will be
greater than 7, more specifically between 8 and 10 for the indicator phenolphthalein
(1). The pH value at the equivalence point is greater than 7 because the
titration of acetic acid results in the production of its conjugate base, which
then reacts with water to produce OH– ions, thus resulting in the
solution to be slightly basic (5). The equivalence point is the point when
neutralization, and the acid-base titration, are complete (1). The point at
which the equivalence point has been reached is called the theoretical endpoint
(4). The endpoint is the point when there is a sharp change in colour, which
shows that neutralization has occurred, and the acid-base titration is
complete. If too much titrant is added, the solution will be beyond the
endpoint and will result in a dark colour (1). Multiple titrations will be
carried out, and the average volume of the titrant added will be used to
calculate the concentration of the acetic acid, and the pH of the titrated
solution (2).

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            The acid-base titrations follow the
Arrhenius theory (4). This theory states that when acids and bases are
dissolved in water, they ionize or dissociate into ions. It states that an acid
is a substance that ionizes to form one or more hydrogen ions in a solution,
and a base is a substance that dissociates to form one or more hydroxide ions
in a solution (1).

The acid-base titration in this laboratory involves the reaction of a
weak acid and a strong base. The weak acid in this experiment is acetic acid,
commonly known as HC2H3O2(aq), which is the
sample (3). The volume of acetic acid is known to be 5.00 mL, but the
concentration is unknown and must be determined (3). A weak acid is an acid
that partially ionizes in water, which results in the production of some
hydrogen ions (1). On the contrary, a strong acid is an acid that ionizes
almost completely in water, which results in the production of hydrogen ions
(1). The strong base in this experiment is sodium hydroxide solution, commonly
known as NaOH(aq), which is the titrant (3). The concentration of sodium
hydroxide is known to be 0.1 mol/L, but the volume must be determined in the
titration process (3). A strong base is a base that dissociates completely in
water, which results in the production of hydroxide ions (1). On the contrary,
a weak base is a base that only partially reacts with water to partially
dissociate, which results in the production of some hydroxide ions (1).

            If the colour of the solution in the
Erlenmeyer flask becomes light pink in colour throughout, then the endpoint is
achieved and neutralization is complete, which would result in the pH of the
titrated solution to be greater than 7 (1). The pH of the titrated solution at
neutralization is assumed to be greater than 7 because the titration reaction
involves a weak acid and a strong base. This is because the titration of acetic
acid results in the production of its conjugate base, which then reacts with
water to produce OH– ions, thus resulting in the solution to be
slightly basic (5).

 

 

Experiment

Procedure:

            For further information regarding
this laboratory, please refer to Nelson Chemistry 12, as the following
investigation has been modified from the Nelson Chemistry 12 textbook (1).

To prepare a titration to determine the unknown concentration of the
acetic acid (HC2H3O2(aq)), a burette was
rinsed with distilled water and then with the standard sodium hydroxide
solution (NaOH(aq)) in order to remove any residue that may have remained in
the beaker. The rinse solution was then discarded into the waste beaker.

            A waste beaker was placed underneath
the burette. The valve of the burette was closed during this step to ensure
that the standardized sodium hydroxide solution within the burette did not leak
out of the burette. A funnel was used to fill and adjust the burette with the
standardized sodium hydroxide solution slowly so that the solution does not
spill over the sides of the burette. The valve and burette tip were analyzed
and cleared of any air bubbles that existed. The initial volume level of the
standardized sodium hydroxide solution in the burette was recorded to the
neatest 0.1 mL so that the calculations were as accurate as possible.

            To prepare the sample solution, 5.00
mL of the acetic acid whose concentration is unknown was measured using a
graduated cylinder. The acid was then transferred to a clean, dry Erlenmeyer
flask. 20.00 mL of distilled water was measured using the graduated cylinder,
and then put into the Erlenmeyer flask with the acetic acid. For each
measurement, a pipette was used near the end of the measurement to add small
amounts of the liquids to ensure that the measurements were as accurate as
possible. 2-3 drops of the phenolphthalein indicator were then added to the
contents in the Erlenmeyer flask. This step was imperative in the laboratory as
the phenolphthalein indicator is what made the titrated solution change colour,
and therefore show when the acid-base titration was complete.

            The Erlenmeyer flask was then placed
underneath the burette, so that when the burette valve was opened, the
standardized sodium hydroxide solution would enter the Erlenmeyer flask and
react with the acetic acid. A white sheet of paper was placed beneath the flask
to see any colour changes in the Erlenmeyer flask clearly.

            The titrant, which is the
standardized sodium hydroxide solution, was slowly added to the acid in the
Erlenmeyer flask while the Erlenmeyer flask’s contents were being swirled. When
the pink colour of the solution in the Erlenmeyer flask started to take longer
time periods for the pink colour to fade away, very small amounts of the titrant
were slowly added to the Erlenmeyer flask to ensure accuracy of the endpoint.
This step was taken to ensure that too much base was not added to the solution,
which would have caused the solution in the Erlenmeyer flask to turn a dark
pink colour, meaning the endpoint has been over shot.

The titration process was stopped when the endpoint was reached in the
laboratory. The endpoint was reached when the light pink colour of the solution
observed in the Erlenmeyer flask spread throughout the solution rather than
fading after the flask was swirled. The final burette reading was recorded to
the nearest 0.1 mL.

            The same acid-base titration
procedure as above was performed two more times, for a total of three trials of
the acid-base titration process. The observations for each trial were recorded
in the observation table, so that measurements could be used from the three
trials to measure the average volume of the sodium hydroxide solution.

            The titrated solutions were disposed
of in a central waste beaker, as instructed by the teacher. The burette was
rinsed with distilled water to ensure that any remaining residue was removed.
The burette valve was then left in an open position to dry.

Observations

This observations table shows the volumes in mL of the titrant and
sample, which are acetic acid (HC2H3O2(aq))
and sodium hydroxide solution (NaOH(aq)) respectively, added to the Erlenmeyer
flask in order to for the endpoint to be reached and neutralization to be
complete. At this point, the colour of the titrated solution in the Erlenmeyer
flask also became light pink in colour throughout.

           The following information in the
Observation Table was obtained throughout the duration of the laboratory. Refer
to Appendix A for Table 1.

Table 1: Final volumes of acetic acid and sodium
hydroxide solution used in the acid-base titration when neutralization is
complete

 

Substance

Trial
#1

Trial
#2

Trial
#3

HC2H3O2(aq)

5
mL

5
mL

5
mL

NaOH(aq)

44.4
mL

43.5
mL

43.8
mL

 

Discussion

            In the first trial of the acid-base
titration, 44.4 mL of the sodium hydroxide solution was added to the 5 mL of
acetic acid in order for there to be a pink colour change throughout the
titrated solution, indicating the endpoint is reached and neutralization has
occurred. To reach the endpoint and neutralization of the titrated solution,
43.5 mL of the sodium hydroxide solution was added to the 5 mL of acetic acid
in the second trial, and 43.8 mL of the sodium hydroxide solution was added to
the 5 mL of acetic acid in the third trial.

            Refer to Appendix B for all
calculations of the laboratory. The three volumes of sodium hydroxide solution were
used to find the average volume, which was 43.9 mL. The concentration of OH–(aq)
is 0.1 mol/L, because it has the same concentration as NaOH(aq). The volume of OH–(aq)
was also calculated to be 0.0439 L. Since the H+(aq) ion and OH–(aq)
ion have a 1:1 mole ratio, the concentration of the original acetic acid was
calculated to be 0.878 mol/L. The Ka constant for acetic acid used
in calculations was determined to be 1.8 x 10-5, and the Kb
constant was calculated to be 5.5556 x 10-10. The concentration of
the conjugate base was calculated to be 0.08978 mol/L. An ICE table was created
to organize and calculate the concentrations of C2H3O2–(aq),
HC2H3O2(aq), and OH–(aq). The
concentration of OH–(aq) was solved to be 7.0624 x 10-6,
which was used to find the pH level of the titrated solution. The calculations
of the titrated solution at the endpoint showed that the pOH level was 5.15 and
the pH level was 8.85.

The indicator used in the acid-base titration was phenolphthalein, which
results in the titrated solution to become light pink in colour throughout the
solution at the endpoint. Phenolphthalein is particularly good for titrations
involving a weak acid and a strong base because it will change colour at a pH
in the range of 8 to 10, which is greater than 7 (1). Since the laboratory
calculations showed that the pH of the titrated solution was 8.85, this
calculation applies to the calculation of the pH level that should be
calculated, in the range of 8 to 10. The pH of the titrated solution is 8.85,
greater than 7, because the titration of acetic acid results in the production
of its conjugate base, which then reacts with water to produce OH–
ions, thus resulting in the solution to be slightly basic (5).

According to a lab report based on the titration of acids and bases, the
writer explains that certain indicators correspond to specific ranges of pH
levels (11). The choice of indicator is important as it is important to reduce
indicator error in the laboratory (11). For instance, if the pH of the titrated
solution at its endpoint is in the range of 8-10, like it was in this
laboratory (8.85), then phenolphthalein will be a more accurate indicator than alizarin
yellow, which undergoes a colour change when the pH is in the range of 10-12
(11). Also, because the titration involves a weak acid and a strong base, it is
likely that the pH of the titrated solution will only be slightly basic.

            The findings in this laboratory led
to both a concentration for acetic acid and pH of the titrated solution that
make sense. The concentration of the acetic acid determined in the laboratory
was 0.878 mol/L. This makes sense as the actual concentration of 5% acetic acid
is 0.87 mol/L (6). As well, the pH of the titrated solution should be slightly
basic because it involves a weak acid and a strong base. Therefore, the pH of
8.85 is accurate and coincides with the phenolphthalein indicator, which turned
light pink in colour throughout when the titrated solution reached a pH in the
range of 8-10 (1). Refer to Appendix B for more details on calculations.

In a similar laboratory conducted that demonstrated an acid-base titration
between acetic acid and sodium hydroxide solution, the pH level measured was
9.3 (12). Two indicators in separate trials were used, one being phenolphthalein,
and the other being bromothymol blue (12). The lab results showed that
phenolphthalein was a more accurate indicator, as it changes colour when the pH
value is in the range of 8-10, while bromothymol blue changes colour in the pH
range of 6 to 8 (12). Therefore, it was correct to use phenolphthalein as an
indicator in the class laboratory since the pH level at the endpoint was 8.85,
which is within the range of the colour change expected by this indicator.

            The findings in this laboratory
coincide with the hypothesis stated above, whereby the titrated solution
experienced a colour change and became light pink in colour throughout the
solution; when the endpoint is achieved and neutralization is complete, the pH
of the titrated solution was greater than 7 as it was calculated to be 8.85.
The production of the conjugate base of acetic acid, which reacts with water to
produce OH– ions, resulted in the solution to be slightly basic because
the pH of 8.85 is slightly greater than 7 (5).

            All steps and procedures completed
in this laboratory were done to the best of human ability. All measurements and
transfer of substances were performed as accurately as possible. Also, all
variables were used with a high degree of control to ensure accuracy. Distilled
water was used throughout the laboratory in diluting the acetic acid, and
washing the apparatuses, instead of tap water (2). This is because distilled
water is a pure substance with a pH that is neutral. As well, tap water
contains many ions and minerals that will disrupt the titration process while
distilled water is pure and does not contain minerals and other ions that will
interfere with the acid-base titration.

            In addition, multiple titrations
were completed using the exact same procedure each time. The difference between
each of the three volumes found in the laboratory are very minimal, each of which
did not exceed the difference of 0.5 mL with the average volume of the sodium
hydroxide solution. Therefore, there was very little variation in the
measurement of the volume of sodium hydroxide solution added in the acid-base
titration. None of the amounts measured in any given trial were outliers from
the measurements of the average and other trials conducted. By averaging the
measurements in the three trials, a more accurate measurement is calculated,
which results in higher accuracy and less errors from occurring. The
measurements of the sodium hydroxide solution in three trials each follow the
10% variation test. In the first trial, the difference between the amount of
sodium hydroxide solution added, and the average amount of sodium hydroxide
solution added, was 0.5 mL, which was a 1.14% difference. In the second trial,
the difference calculated was 0.4 mL, which was a 0.91% difference. In the
third trial, the difference calculated was 0.1 mL, which was a 0.23%
difference. Refer to the Appendix B for these calculations.

            A major source of error was the
transfer of liquids between each measuring apparatus, the Erlenmeyer flask, and
the burette. When the substances were transferred, some droplets were left in
the pouring apparatus, while a small amount of the substances spilled over in
the pouring process. This resulted in the volume of each substance to be
slightly different, which would directly affect the concentration, volume, and
amount in moles of the substances in the remaining process of the laboratory.

            Misjudging the colour change of the
titrated solution was also a source of error (8). The endpoint required a very
specific volume of the sodium hydroxide solution to be added, whereby a few
droplets over the endpoint would result in the inaccurate calculation of the
amount of titrant added to reach the endpoint and neutralization.

Another source of error was the reaction of the sodium hydroxide solution
with the air surrounding it (9). The sodium hydroxide solution was exposed to
air both at the top of the burette, and when the burette valve was open. The
sodium hydroxide directly captured carbon dioxide from the air, which is a
process known as air capture (10). This reaction resulted in the formation of
sodium carbonate and water (9). This reaction means that the sodium hydroxide
solution will lose its strength over time, which will impact the results
calculated, whereby the volumes of sodium hydroxide solution calculated in the
three trials, and the average volume, would be inaccurate (9). Therefore, the
pH of the titrated solution resulted in a pH lower than it should be. This is
because the strength of the strong base in the acid-base titration became
weaker.

            If this experiment were to be run
again, there are several steps that would be taken in order to ensure a greater
level of accuracy in the measurements and calculations involved in the
acid-base titration. The introduction of a white paper held behind the
measuring apparatus when reading the volume inside the apparatus would be
useful. This would allow for more accurate readings of the amount of each
substance being measured. As well, more control of the valve on the burette
when adding the titrant to the sample is necessary to make sure that one does
not overshoot the endpoint and therefore inaccurately measure the endpoint and
point of neutralization. A final step that would be taken would be the thorough
cleaning of the burette, the Erlenmeyer flask, and other apparatus used, to
ensure that there are no side reactions while the substances are being
measured, and while titration is in progress (7). This also includes rinsing
the burette more thoroughly with distilled water (7).

Conclusion

            The laboratory was conducted with
success. The acid-base titration process is a precise process that involves the
addition of precise volumes of sodium hydroxide solution to the acetic acid
sample, which allowed for the calculation of the concentration of the acetic
acid, and the pH of the titrated solution at the endpoint and neutralization.
The concentration of acetic acid was determined to be 0.878 mol/L. This can be
compared to the actual concentration of 5% acetic acid, which is 0.874 mol/L.
The difference in concentration is therefore 0.004 mol/L, 0.46% in difference,
which is very minimal. This means that the laboratory was executed accurately
and with success. As well, the indicator used in this acid-base titration was phenolphthalein,
which was very effective in making the titrated solution experience a colour
change from colourless to light pink when the endpoint was reached and
neutralization occurred. The pH of the titrated solution in the Erlenmeyer was
calculated to be 8.85. This agrees with the pH level that results in a colour
change of phenolphthalein, which is in the range of 8 and 10. The pH level of
the titrated solution determined was greater than 7 (8.85), which applies to
the understanding that with a weak acid and a strong base, the titration of
acetic acid results in the production of its conjugate base, which then reacts
with water to produce OH– ions, thus resulting in the solution to be
slightly basic (5). Conclusively, this laboratory was effective in reinforcing
important concepts including acid-base titration, weak and strong acids and
bases, indicators, and laboratory techniques. It will therefore be beneficial
for continued laboratory exercises and study in this course.