PREPARATION AND STANDARDIZATION OF A SODIUM HYDROXIDE SOLUTION
(POTASSIUM HYDROGEN PHTHALATE METHOD)
Solid sodium hydroxide and sodium hydroxide solutions are not a very stable
O
substances. They adsorbs both water and carbon dioxide from the air, changing
C OK
the both the purity of the solid and the concentration of a solution. As a result,
C OH
accurate standard sodium hydroxide solutions can not be prepared directly from
its solid since its purity is never guaranteed but rather, an approximate solution
O
is prepared and then standardized against a pure substance or a solution of
KHC8H4O4
known concentration. In this experiment, you will prepare a solution of sodium
hydroxide and standardize it by titrating it against a measured mass of potassium hydrogen
phthalate, KHC8H4O4. Note its structure above. Potassium hydrogen phthalate is often used to
standardize a strong base such as sodium hydroxide since it is stable, can be obtained in high
purity, and does not readily absorb water.
In a similar experiment, a student measured 0.171 gram of potassium hydrogen phthalate into a
small beaker. He then titrated the potassium hydrogen phthalate against a sodium hydroxide
solution he had prepared earlier. The initial reading of the volume of NaOH solution in the
syringe was 1.00 ml and 0.20 ml at the end of the titration. Using this information, he was able
to determine the concentration of his sodium hydroxide solution.
The first step in the calculation is to determine the number of moles of potassium hydrogen
phthalate reacted.
moles of KHC8 H 4 O 4 =
mass of KHC8 H 4 O 4
0.171 g
=
= 0.000838 mol
formula mass of KHC8 H 4 O 4 , 204g / mol
204g / mol
Now that we know the number of moles of KHC8H4O4, we can find the number of moles of
NaOH reacted with the KHC8H4O4 and the concentration of the NaOH solution. First, we must
find the volume of NaOH solution reacted. 1.00 ml - 0.20 ml = 0.80 ml NaOH solution. 0.80 ml
divided by 1000 ml per liter = 0.00080 liter of NaOH solution.
KHC8 H 4 O 4 (aq) + NaOH(aq)  KNaC8 H 4 O 4 (aq) + H 2 O()
0.000838 mol KHC8 H 4 O 4 x
Molarity =
1 mol NaOH
= 0.000838 mol NaOH
1 mol KHC8 H 4 O 4
# moles NaOH
0.000838 mol
=
= 1.05 M NaOH
Volume NaOH solution
0.00080 liters
This same calculation is repeated for each of your trials and the average concentration
1
calculated.
In this experiment you will measure your skill by determining the molecular mass of the
KHC8H4O4, using a titration curve developed from your trials using the following technique.
A pH titration is an important tool in analytical
14
chemistry since it allows the experimenter to better
understand the titration process and provides
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important information regarding the nature of the
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acid and base reacted. Furthermore, a pH titration
provides a more accurate end-point or “equivalence
8
point” than can be obtained using only an indicator.
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Once the “equivalent point” is known, pH titrations
can be used to determine the pKa of the acid, a
4
value related to the strength of the acid. pH
2
titrations also allow the experimenter to better
0.0
0.2
0.4
0.6
0.8
1.0
determine the best choice of an acid-base indicator
used in conventional titrations. The adjacent graph
demonstrates a typical pH titration and some of its A typical pH titration of a weak monoprotic acid
important features.
vs a strong base
The most important data point on the graph is the equivalence point. At the equivalence point,
the number of moles of aqueous hydrogen ion, H+(aq), have been neutralized by exactly the
same number of moles of aqueous hydroxide ion, OH-. The equivalence point may be used to
determine the stoichiometry of the reaction and even the molar mass of the acid. The graph has
its maximum slope at the equivalence point. That is, the addition of a small volume of titrant
produces the greatest change in the pH of the solution. This concept is important since it is the
basis of the graphical analysis of the data.
Another important point on a pH titration graph is at one-half the equivalence point. At one-half
the equivalence point, exactly one-half of the acid in the solution has been neutralized giving a
solution with an equal concentration of both the acid and its conjugate base. Since the
concentration of each of the materials is the same, then the pH of this solution equals the pKa of
the acid as predicted by the Henderson-Hasselbach equation:
pH  pK a  log
[Base  ]
[Acid]
That is, since the concentration of both the acid and its conjugate base are equal, their ratio is 1,
and the log of 1 equals 0. Therefore, the log [conjugate base] over the log [conjugate acid] term
drops out of the expression and the pH of the solution coincides with the pKa. This phenomenon
occurs only at one-half the equivalence point.
2
The propose of this experiment is to perform a pH titration with of weak acidic salt , KHC8H4O4,
, and determine its pKa.
A good approximation of the equivalence point in a pH titration may determined graphically
14
using the following method. While not always
perfect, this method allows the experimenter to
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determine both the equivalence point and pKa of the
10
acid quickly.
Step 1. Prepare your graph by plotting the pH of
the solution verses the volume of base added to the
solution. Use as large and best quality graph paper
available to you. It is not necessary to connect the
points at this time.
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6
4
2
0.0
Step 2. Draw the best straight line through both
linear regions of the graph where buffering occurs.
The best line represents the general direction of the
data, includes as many point as possible, and has as
many points above the line as below it. Extend this
line beyond the data points so that it may be used in
the next step, determining the equivalence point.
Note the drawing above.
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0.4
0.6
0.8
1.0
0.2
0.4
0.6
0.8
1.0
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12
10
8
Step 3. Draw a line through the portion of the
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graph with the greatest slope. Measure the length of
this line between the two lines drawn in Step #1,
4
and find its center. The center of this line is
approximately the equivalence point. Drop a line
2
0.0
perpendicular to the x-axis at this point to determine
the volume of titrant added to reach the equivalence
point. Once this value is known, the concentration
of a monoprotic acid may be determined from the relationship:
Vacid x M acid  Vbase x M base
Step 4. Once the equivalence point has been determined, the pKa of the acid can be obtained.
At a volume of one-half the equivalence point volume, the pH of the solution equals the pKa of
the acid. In the adjacent example, the equivalence point volume is 0.80 ml giving a value of 0.40
ml at one-half the equivalence point. The pH of the solution when 0.40 ml of base has been
added to the solution is 5.4. Consequently, the pKa of this acid would be 5.4 and the Ka would
10-5.4 or 4 x 10-6.
3
Materials:
2-1.0 ml syringes, 30-ml or 50-ml beaker, ring stand and clamp,1-microtip 1-ml pipet, magnetic
stirrer and stir bar (optional), pH probe, Potassium hydrogen phthalate, solid sodium hydroxide,
spatula, a 25-ml graduated cylinder, 100-ml beaker,3-microtip polyethylene transfer pipets, 0.5%
phenolphthalein solution, 30-50 ml beaker, wash bottle, polyethylene storage bottle.
Procedure:
1. If necessary, attach tip extenders to both hypodermic syringes. It is
important to first rinse the syringe with the solution being measured
before filling them to be certain that the solution inside the syringe
has the same concentration as the solution being studied. Fill the
base syringe with sodium hydroxide solution prepared in next
procedure step. Note that syringes are color-coded. Partly fill a
microtip polyethylene transfer pipet with phenolphthalein solution,
the acid-base indicator in this experiment. Attach a tip extender to
the hypodermic syringe. It is important to rinse and prepare the
syringes as directed by your instructor. Fill the base syringe with the sodium hydroxide
solution. Partly fill a microtip polyethylene transfer pipet with phenolphthalein solution, the
acid-base indicator in this experiment. Record the initial volume of the NaOH solution in
the syringe to the nearest 0.01 ml.
2. Tare a clean, dry 100 ml beaker. Add between 0.9 gram and 1.1 grams of solid sodium
hydroxide to the beaker and record the mass of the solid to the precision of your balance.
Add 10-15 milliliters of distilled or deionized water to the solid and stir or swirl the mixture
until the sodium hydroxide pellets dissolves completely. Transfer the solution to a 25-ml
graduated cylinder. Add distilled or deionized water to the graduated cylinder until the total
volume reads 25.0 ml. Pour the solution back and forth a few times between the beaker and
graduated cylinder a few times to insure that the solution is mixed throughly. It may be
necessary to cool the solution in the graduated cylinder by running cold tap water around the
outside until its temperature is approximately room temperature. This step is important since
its volume will change as it cools causing its concentration to increase. Transfer the solution
to a small flask or storage bottle to act as a source of the sodium hydroxide solution for this
and subsequent experiments.
2. Tare the small beaker to be used as a reaction vessel. Add between 0.15 gram and 0.17 gram
of potassium hydrogen phthalate( an Acidic Salt) and mass to the precision of your balance.
Add a few milliliters of distilled or deionized water from your wash bottle and swirl the
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mixture until most of the solid dissolves. Add one drop of the phenolphthalein indicator to
the potassium hydrogen phthalate solution in the reaction vessel and swirl.
3. Throughly rinse the pH electrode with distilled water and loosely clamp it on the ring stand.
Lower the pH probe into the solution until the glass bulb of the electrode is totally immersed in
the solution. It may be necessary to add more distilled water at this time. If you have a magnetic
stirrer (or plastic coated paper clip) , place a small stir bar into the solution and adjust the
position of the pH probe so the glass member of the electrode is completely immersed in the
solution and doesn’t touch the stir bar. Mix the solution in the beaker by swirling or by stirring
with the magnetic stirrer. Finally, clamp the pH probe securely to the ring stand.
4. Begin the titration by recording both the initial syringe reading and pH of the acid/water
mixture before any base has been added on your data table. Add two drops of base, swirl or stir
the mixture, and wait until the pH reading stabilizes. Potassium hydrogen phthalate dissolves
slowly and therefore the solution must be mixed for some time after each addition to be certain
that the reaction is complete. Record both the pH and syringe reading. Continue this process of
adding base in two drop increments, stirring, and recording both the syringe reading and the pH
of the solution. Be certain that you observe the color of the solution and note at what pH the
color of the phenolphthalein just changes. Eventually, the pH of the solution will begin to rise
much faster than originally observed. At this point, add the base in one drop portions. Towards
the end of the titration where change in the pH has slowed, the two drop rate can be resumed.
Add base until exactly 1.00 ml of sodium hydroxide solution has been added to the beaker.
5. Throughly rinse and dry the reaction vessel, refill the NaOH syringe, and repeat the
experiment as time permits.
6. Throughly rinse all glassware and the syringe before returning them to their storage area.
Discard the phenolphthalein pipet. Save the standardized sodium hydroxide solution for
future experiments if directed to do so by your instructor.
Calculations:
Q1. Calculate the number of moles of potassium hydrogen phthalate, KHC8H4O4, reacted with
the sodium hydroxide solution from the mass reacted and its formula mass, 204 g/mol.
Q2. From the number of moles of potassium hydrogen phthalate reacted and the volume of
NaOH solution consumed, find the molarity of the sodium hydroxide for each of your trials.
Q4. Calculate the average molarity of your sodium hydroxide solution.
Q5. (a) Using the mass of sodium hydroxide added to the graduated cylinder, what is the
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expected molarity of the sodium hydroxide solution prepared in this experiment?
(b) How do you account for the discrepancy between your experiment value from the
titration and the value calculated from the mass of sodium hydroxide used to prepare the
sodium hydroxide solution?
Q6. Determine the volume of sodium hydroxide added to the beaker containing the acid salt
solution by subtracting the syringe reading at that point from the initial syringe reading.
For example, assuming a base syringe full of solution read 1.00 ml and the data point value
was 0.44 ml, then the volume of base added to beaker would be the difference or 0.56 ml.
Q7. Plot the pH of the solution verses the volume of sodium hydroxide added to the beaker.
Use the largest and highest quality graph paper available to you. It is not necessary to
connect the points.
Q8. Determine the equivalence point as describe in the introduction to this experiment.
Q9. Determine the pH of the acid solution at one-half the equivalence point. The pKa of a weak
acid is the pH at exactly one-half the equivalence point.
Q10. What is the Ka of your weak acid? This is calculated from the pKa determined in question
#Q9?
Q11. Find the average pKa of these values.
Q12. Why are you able to use phenolphthalein as an indicator to find the equivalence point or
endpoint of a titration instead of a pH meter? Explain you answer using the data obtained
in this experiment.
Q13. What do you believe is the main source of error in this experiment? How could that error
be corrected?
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DATA:
Mass of solid sodium hydroxide:____________grams
Trial #1
Mass of potassium hydrogen
phthalate, KHC8H4O4
Trial #2
g.
g.
Initial volume of NaOH
solution
ml.
ml.
Final volume of NaOH
solution
ml.
ml.
Volume of NaOH reacted
ml.
ml.
Moles of KHC8H4O4 reacted
mol
mol
Moles of NaOH reacted
mol
mol
M.
M.
Concentration of the NaOH
solution, in moles/liter
Average concentration of the NaOH solution
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M.
Trial II
Trial I
Syringe Reading
pH
Syringe reading
Syringe Reading
pH
pH
Syringe reading
1.
ml
26.
ml
1.
ml
26.
ml
2.
ml
27.
ml
2.
ml
27.
ml
3.
ml
28.
ml
3.
ml
28.
ml
4.
ml
29.
ml
4.
ml
29.
ml
5.
ml
30.
ml
5.
ml
30.
ml
6.
ml
31.
ml
6.
ml
31.
ml
7.
ml
32.
ml
7.
ml
32.
ml
8.
ml
33.
ml
8.
ml
33.
ml
9.
ml
34.
ml
9.
ml
34.
ml
10.
ml
35.
ml
10.
ml
35.
ml
11.
ml
36.
ml
11.
ml
36.
ml
12.
ml
37.
ml
12.
ml
37.
ml
13.
ml
38.
ml
13.
ml
38.
ml
14.
ml
39.
ml
14.
ml
39.
ml
15.
ml
40.
ml
15.
ml
40.
ml
16.
ml
41.
ml
16.
ml
41.
ml
17.
ml
42.
ml
17.
ml
42.
ml
18.
ml
43.
ml
18.
ml
43.
ml
19.
ml
44.
ml
19.
ml
44.
ml
20.
ml
45.
ml
20.
ml
45.
ml
21.
ml
46.
ml
21.
ml
46.
ml
22.
ml
47.
ml
22.
ml
47.
ml
23.
ml
48.
ml
23.
ml
48.
ml
24.
ml
49.
ml
24.
ml
49.
ml
25.
ml
50.
ml
25.
ml
50.
ml
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pH