Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
PURPOSE:
To standardize a solution of 0.10 M sodium hydroxide (NaOH) solution with potassium
hydrogen phthalate (KHC8H4O4), by a pH titration.
PRINCIPLES:
Most shelf reagents, such as 0.10 M sodium hydroxide, could not be kept at a previously
determined accurate concentration, because the concentration of these reagents change in time,
due to exposure to the environment.
In order to obtain a reagent of accurately known concentration, expressed to four significant
figures, the concentration of the reagent must be determined by reacting it with a known amount
of another reagent. The entire procedure by which the molarity of a solution of one substance
(NaOH) is obtained from an accurately known amount of another substance, commonly referred
to as a primary standard, is called standardization. The preferred method commonly used for
the standardization of NaOH is an Acid – Base titration with potassium hydrogen phthalate
(KHC8H4O4, thereafter abbreviated as KHP) used as a primary standard.
Potassium hydrogen phthalate (KHP) is a soluble salt, and as such is completely dissociated in
aqueous solution.
K+(aq)
HP- (aq)
Hydrogen phthalate ion
The hydrogen phthalate ion, HP is a weak acid and it undergoes partial ionization:
KHP(aq)
HP- (aq) + H2O(l)
+
H3O+(aq) + P2-(aq)
The addition of NaOH to this equilibrium system will cause the OH- ions to combine with the
hydronium (H3O+) ions to form water.
OH-(aq) + H3O+(aq)
H2O(l)
The decrease in the concentration of hydronium (H3O+) ions will cause the equilibrium system to
shift to the right and, as such, more of the weak acid (hydrogen phthalate ion, HP-) will ionize.
Successive additions of NaOH will continually remove H3O+ ions, shift the ionization
equilibrium of the weak acid (hydrogen phthalate ion, HP-) to the right and force the weak acid
into complete ionization.
The situation can be summarized in the equilibrium table below:
Equation:
HP- (aq) +
H3O+(aq)
H2O(l)
Stress:
+
P2-(aq)
decreases
Shift:
New Equilibrium: decreased
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decreased
decreased
1
increased
Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
The net result is the complete neutralization of potassium hydrogen phthalate (KHP) by the
NaOH solution, as shown in the equations below:
KHP(aq)
+
NaOH(aq)
KNaP(aq)
+
H2O(l)
OH- (aq)
P2- (aq)
+
H2O(l)
Net Ionic Equation:
HP- (aq)
+
Note that according to the stoichiometry of the reactions involved here:
Number of moles of KHP = Number of moles of HP- = Number of moles of NaOH
Or simply:
Number of moles of KHP = Number of moles of NaOH
The number of moles of KHP is known, if an accurately determined mass of KHP is used for the
titration and, as such, the number of moles of NaOH used to completely neutralize the weak acid
is also known.
The challenge is to experimentally determine by titration the exact volume of NaOH required to
completely neutralize the known mass of the weak acid (potassium hydrogen phthalate, KHP).
This is achieved by slowly adding the NaOH solution of unknown concentration to the solution
containing an accurately determined mass of KHP, while the pH is monitored with either an
indicator (a substance whose color depends on the pH) or a pH meter. As the acid and the base
combine, they neutralize each other. At the equivalence point – the point in the titration when
the number of moles of base is equal to the number of moles of acid – the titration is complete.
Two common methods are available to experimentally determine the exact volume of NaOH
required to reach the equivalence point of this acid – base titration.
Method I:
Determining the Equivalence Point by the use of an “Indicator”
The exact volume of NaOH required to reach the equivalence point is determined by the use of
an indicator. With an indicator, we rely on the point where the indicator changes color – called
the end point – to determine the equivalence point (when the amount of acid equals the amount
of base). With the correctly chosen indicator, the end point of a titration (indicated by the color
change) occurs at the equivalence point.
After the indicator is added to the solution containing the primary standard, the NaOH solution
of unknown concentration is delivered carefully from a buret, until the indicator changes color.
The indicator chosen will have one color before the reaction is complete (before the equivalence
point is reached) and another color when the completion occurs (the equivalence point is
reached). The color change of the solution signals the exact completion of the reaction.
This method will be used in the next experiment (“Standardization of a Sodium Hydroxide
Solution with a Primary Standard”).
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
Method II:
Determining the Equivalence Point by a “pH Titration”
This method concentrates on the pH changes that occur during the titration and will be the
method we use in this experiment.
A plot of the pH of the solution during a titration (shown below) is known as a titration curve or
pH curve.
Before any base (NaOH) is added to the solution, the pH is low (as expected for a solution of a
weak acid). As small amounts of NaOH are added, the added NaOH converts the stoichiometric
amount of the weak acid (Hydrogen phthalate ion, HP-) into its conjugate base
(phthalate ion, P2--). The solution now contains significant amounts of both a weak acid
(hydrogen phthalate ion, HP-) and its conjugate base (phthalate ion, P2--).
The solution is now a buffer.
As more NaOH is added, it converts more of the weak acid (Hydrogen phthalate ion, HP -) into
its conjugate base (phthalate ion, P2--).
The point of inflection in the middle of the curve is the equivalence point. Notice that the pH
changes very quickly near the equivalence point (small amounts of added base cause large
changes in pH).
At the equivalence point all of the weak acid ((hydrogen phthalate ion, HP-) has been converted
into its conjugate base (phthalate ion, P2--) and the solution is no longer a buffer, since it no
longer contains significant amounts of both a weak acid and its conjugate base. Instead, the
solution contains only a weak base (phthalate ion, P2--).
As a result, the pH at the equivalence point is NOT neutral, but basic.
The titration of a weak acid by a strong base always has a basic equivalence point because at the
equivalence point all of the weak acid has been converted into its conjugate base, resulting in a
weakly basic solution.
Beyond the equivalence point, the solution is basic because the weak acid (Hydrogen phthalate
ion, HP-) has been completely neutralized and excess base is added to the solution.
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
While the procedure for determining the equivalence point by from a titration curve (pH
measurements during the titration) is much slower than the one employing an indicator, this
method has more versatile uses than the standardization using an indicator.
A “pH Titration”, also referred to as a “Titration Curve” can be used to:
1. Standardize a solution of 0.10 M sodium hydroxide (NaOH) solution, as performed in
this experiment,
2. Determine the Acid Ionization Constant of a Weak Acid, and
3. Allow a chemist to choose an appropriate indicator for subsequent titrations of similar
samples with the same reagent, as illustrated below.
The figure above shows that phenolphthalein (color change between pH = 8.7 – 10) would be the
correct indicator for the titration of a weak acid with a strong base since, as previously explained,
the equivalence point for such a titration occurs in the range of a basic pH.
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
PROCEDURE:
1. Prepare the KHP solution
 Using the analytical balance, accurately weigh (from a vial) between 0.4g and 0.6g a
sample of solid KHP into a 250 mL beaker, by using the weighing bottle technique.
Photo by Andrew Huertas
Photo by Andrew Huertas




Place the vial on the analytical balance. Read and record its mass.
Place a clean (does not have to be dry) 250 mL beaker on the centigram
balance, placed adjacent to the analytical balance.
Do no record the mass of the empty beaker. Tare the centigram balance
so that the reading on the centigram balance is 0.00 g.
 Dispense a few crystals of the KHP from the vial directly into the beaker by
taping gently the vial against the beaker.
 Follow the mass readings on the centigram balance, as they increase upon the
addition of the solid KHP. When you reach the desired mass range
(0.4g – 0.6g), stop adding KHP. When this is the case, it means that
you have transferred between 0.4g and 0.6 g of KHP from the vial into the
beaker.
 Cap the vial and reweigh it on the analytical mass. Read and record this mass.
 The accurate mass of KHP transferred into the beaker is the difference
between the mass of the vial before transferring some of the solid sample into
the beaker (Mass of vial + KHP) and the mass of the vial after transferring the
solid sample into the beaker (Mass of vial – KHP)
o ONLY THE CAPPED VIAL WILL BE PLACED ON THE
ANALYTICAL BALANCE, BUT NOT THE FLASK.
o THE TWO MASSES OF THE VIAL (BEFORE AND AFTER
TRANSFER) MUST BE RECORDED IN FOUR SIGNIFANT
FIGURES.
Add about 100 mL of D.I. water to the beaker containing the KHP.
You may use the 100 mL mark on the beaker to estimate the volume of D.I. water
added. No need to use a graduated cylinder.
Place a stirring magnet in the beaker.
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Chemistry 102
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EXPERIMENT 9
pH TITRATION

Warm the solution gently with swirling (on a hot plate with low setting at about 400))
until the KHP is completely dissolved.
 Place the beaker on a stirring plate and set the stirring plate so as to provide a gentle
and uniform mixing of the KHP solution. Stirring is important because in the next
step the pH will tend to drift, unless a completely homogeneous solution of KHP is
achieved.
2. Fill the buret with 0.1 M NaOH
 Use a small beaker and a funnel to transfer the NaOH solution into the buret.
 Rinse the buret with three portions of NaOH solution, coating the barrel each time
before emptying out the solution
 After filling the buret (use a small beaker) with the NaOH solution, flush out the
tip thoroughly, making sure no bubbles are caught in the tip.
 Check the stopcock for leaks.
 Set the level of the NaOH solution at the 0.00 mL mark
3. Calibrate and clamp the pH meter
 Perform a two-point calibration of the pH meter with pH = 7.01 and pH = 4.01
buffers.
 Immediately immerse the electrode
into the KHP solution by solidly
clamping it at the appropriate height.
The electrode should be immersed in
the solution, but should be high
enough to allow stirring magnet to
stir the solution.
Keep the pH meter turned “OFF”
until you complete the set up.
Photo by Andrew Huertas
4. Clamp the buret
 Use a buret clamp to clamp the buret. The buret should be clamped so that its tip is
within the beaker but above the surface of the solution.
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
5. Proceed with the first titration
 Before you proceed with the titration, make sure that:
 the solid sample of KHP is completely dissolved,
 good stirring is provided by the stirring magnet and the stirring plate,
 the pH meter is properly clamped and inserted into the solution, and
 the tip of the buret is within the beaker but above the surface of the
solution.
 Record the initial buret reading (should be 0.00 mL)
 Turn the pH meter “ON” and read and record the initial pH of the KHP solution
before any of the NaOH solution has been added.
 Begin the titration by adding successive portions of about 1 mL of the NaOH
solution. Read and record the buret reading and the pH after each addition.
Recall that all buret readings should be read and recorded to the nearest 0.01 mL
(two decimals).
o At the beginning of the titration, while you are adding 1 mL of NaOH, the
pH will change very little (less than 0.3 pH units)
o When the pH begins to increase by more than about 0.3 pH units, decrease
the portions of NaOH that you add to about 0.2 mL.
o When the pH begins to increase again by less than about 0.3 pH units after
the addition of NaOH, increase the portions of NaOH that you add to
about 1 mL of NaOH,
o Continue the titration until the pH is about 11.5 – 12.0
o Discard the solution through a funnel in order to catch the stirring magnet.
o Wash out the beaker and the stirring magnet with plenty of tap water
followed by a few rinses with D.I. water.
Steps throughout the
pH titration
Before the
Equivalence Point
Around the
Equivalence point
Passed the
Equivalence point
Change in pH
(Δ pH)
Less than 0.3 units
Additions of
NaOH
About 1mL
More than 0.3 units
About 0.2 mL
Less than 0.3 units
About 1 mL
6. Repeat steps 1 through 5 with a second sample of solid KHP
7. Plot separate graphs of the pH versus the volume of NaOH used.
 Clearly label each graph.
 Determine the equivalence point for each titration from the graph.
 Determine the volume of NaOH at the equivalence point.
 Attach ALL your graphs to your Report Form
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
8. For ALL samples of KHP used for the titrations record and calculate:
 The Molar Mass of potassium hydrogen phthalate (KHC8H4O4), KHP (g/mol)
 The number of Moles of KHP added
 The Number of Moles of NaOH that have reacted with KHP
 The Volume of NaOH added to react completely with the KHP
 The Molarity of NaOH for each titration (mol/L)
 The Mean Molarity of NaOH (mol/L)
______________________________________________________________________________
IMPORTANT NOTES
This experiment is run individually and the reported molarity is graded as an unknown on a
sliding scale for 30 points. In order to earn a good grade, it is in your interest to put in the
maximum effort to report an accurate molarity for the NaOH solution.
In order to allow students to obtain an accurately determined molarity (and a good grade) by
performing several trials, the experiment is scheduled to be done in two laboratory sessions.
It is expected that on the first day you will be able to do no more than two trials.
The guidelines given below will assist you in deciding how many additional trials (if any) you
wish to perform during the second laboratory session.
Guidelines for an accurate standardization
1. A proper standardization requires a minimum of three trials that agree within plus or
minus 0.5% of each other.
 If you perform an additional third titration on the second lab session and your
three titrations yield three molarities that agree within plus or minus 0.5% of each
other, you may conclude that the Mean Molarity that you will be reporting is
reasonable accurate. If this is not the case, in order to obtain an accurate Mean
Molarity, a fourth titration needs to be performed.
 If you perform two additional titrations on the second lab session and three out of
your total of four titrations yield three molarities that agree within plus or minus
0.5% of each other, you may ignore one of your molarities and calculate the Mean
molarity based on the three titrations that agree within plus or minus 0.5% of each
other.
2. A quick check of the reproducibility of your titrations is to calculate the ratio of
g KHP/mL NaOH to four significant figures. This ratio should vary only in the last
significant figure.
3. Repeat the standardization process until three trials agree within these limits.
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
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Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
NAME: ________________________________
Date: _____________
SAMPLE 1
Mass of vial + KHP: _______________ g
Mass of vial – KHP
_______________ g
Mass of KHP:
_______________ g
Buret
reading
(mL)
Vol.
added
(mL)
0.00
pH
Δ pH
Buret
reading
(mL)
Initial pH: _______________ *
(enter this value in the table)
Vol.
added
(mL)
*
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pH
Δ pH
Buret
reading
(mL)
Vol.
added
(mL)
pH
Δ pH
Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
SAMPLE 2
Mass of vial + KHP: _______________ g
Mass of vial – KHP
_______________ g
Mass of KHP:
_______________ g
Buret
reading
(mL)
Vol.
added
(mL)
0.00
pH
Δ pH
Buret
reading
(mL)
Initial pH: _______________ *
(enter this value in the table)
Vol.
added
(mL)
*
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pH
Δ pH
Buret
reading
(mL)
Vol.
added
(mL)
pH
Δ pH
Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
SAMPLE 3
Mass of vial + KHP: _______________ g
Mass of vial – KHP
_______________ g
Mass of KHP:
_______________ g
Buret
reading
(mL)
Vol.
added
(mL)
0.00
pH
Δ pH
Buret
reading
(mL)
Initial pH: _______________ *
(enter this value in the table)
Vol.
added
(mL)
*
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pH
Δ pH
Buret
reading
(mL)
Vol.
added
(mL)
pH
Δ pH
Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
SAMPLE 4
Mass of vial + KHP: _______________ g
Mass of vial – KHP
_______________ g
Mass of KHP:
_______________ g
Buret
reading
(mL)
Vol.
added
(mL)
0.00
pH
Δ pH
Buret
reading
(mL)
Initial pH: _______________ *
(enter this value in the table)
Vol.
added
(mL)
*
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pH
Δ pH
Buret
reading
(mL)
Vol.
added
(mL)
pH
Δ pH
Chemistry 102
______________________________________________________________________________
EXPERIMENT 9
pH TITRATION
Sample 1
Sample 2
Mass of sample (g)
Molar Mass of KHP
( 204.2 g/mol)
Number of moles of
KHP added
Number of moles of NaOH that
have reacted with the KHP
Volume of NaOH added to react
completely with the KHP (mL)
Volume of NaOH added to
react completely with the KHP (L)
Molarity of NaOH (mol/L)
Mean Molarity of
NaOH
(mol/L)
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Sample 3
Sample 4