CHEMISTRY 163
EXPERIMENT II
pH TITRATION
pages 1 – 6 do not need to be printed
DISCUSSION
In titrimetric analysis the concentration of an unknown solution is determined by reacting it with a solution of
known concentration. Thus, the concentration of an unknown monoprotic acid HA(aq) solution can be
determined by an acid-base titration using a standard solution of a base, as for example sodium hydroxide
solution:
HA (aq) + NaOH (aq)  NaA(aq) + H2O
The stoichiometry of the reaction is given by the equation, and the object of the titration is to add enough
dissolved NaOH to just neutralize (and no more) the dissolved HA; that is, to add an amount of base equivalent
to the original acid concentration. If this can be done, the amount of acid originally present can be readily
calculated from the amount of base necessary to reach the equivalence point.
In such a titration, where the reactants and products are colorless, there must be a way of showing when enough
reagent has been added to just complete the reaction. A common method is the use of a visual indicator in the
solution being titrated. This indicates the end of the titration by a color change, or end point. A knowledge of
the chemistry allows one to pick an indicator such that the end point corresponds to the equivalence point of the
titration. An improperly chosen indicator will not correspond to the equivalence point.
The titration is usually carried out using a buret to deliver the titrant, in
this case the NaOH, to the reaction flask. The titrant is added drop-wise
until the end point ( the equivalence point), is just reached. The
calculations are best done in mole amounts, and for the titration shown
in Figure 1:
Number of moles of NaOH titrant:
32.50 ml NaOH added 
0.1006 mole
= 3.270  10-3 moles of NaOH
1000 mL
From the chemical balanced equation the stoichiometry of NaOH to HA
is 1 mole to 1 mole, so the the unknown solution of HA contained
3.270 x 10-3 moles of HA.
Fig. 1
2
INDICATORS
In an acid-base titration, the color of the indicator actually depends on the acidity (or basicity) of the solution
being titrated. The most convenient method for expressing this acidity is in terms of a "pH scale", where the
pH is the logarithm of the reciprocal of the molar hydrogen ion concentration:
pH = log 10
1
+
[H ]
= - log 10 [ H + ]
Equation (1)
where [H+] is the concentration of hydrogen ion in the aqueous solution expression in moles per liter.
Table 1
ACID BASE INDICATORS
A large number of acid base indicators are
available, and Table 1 shows the behavior of a
number of common ones. Typically, a pure
indicator shows a complete color transition over
about 2 pH units; litmus is an impure product
with a more gradual color change. These
indicators are weak acids or weak bases; the color
change occurs as they are titrated during the
course of the reaction. For this reason the
indicators must be highly colored so that only
trace amounts are necessary, like 10 -7 M, to
produce a color change with only a negligible
amount of titrant.
STRONG ACID-STRONG-BASE TITRATION
LIKE NaOH and HCl
In the titration of a hydrochloric acid solution
(strong acid) with a sodium hydroxide titrant (strong
base), the pH of the solution being titrated
continually increases.
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O
Figure 2 shows an almost sigmoid-shaped titration
curve as 25 ml of 0.1 M HCl solution is being
titrated by 0.1 M NaOH solution.
Before the equivalence point, the solution contains more acid (pH<7) than base, and past the equivalence
point, the solution contains excess added base (pH>7) than acid. At the equivalence point, all the base has
reacted with all the acid and the solution contains water and NaCl as products. Therefore the pH is that of a
"pure" NaCl solution, (pH = 7). The change in pH with added base is most rapid at the equivalence point.
For such a strong acid-base titration, methyl red is not the best indicator, because it changes color over the
pH range 4.4 to 6.3. This color change occurs before the pH 7 point. A better indicator is bromothymol blue
which changes color between 6.0-7.6 range.
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WEAK ACID-STRONG-BASE TITRATION.
If an acetic acid solution (weak acid) is titrated with sodium hydroxide solution (strong base) the
stoichiometry is given by the following equation:
CH3COOH(aq) + NaOH(aq)  CH3COONa(aq) + H2O
The pH titration curve for this weak-acid-strongbase titration (Figure 3, showing 25 mL of 0.1 M
CH3COOH being titrated by 0.1 M NaOH.) differs
from that of the previous case. This solution is
basic (pH>7) at the equivalence point. This is so
because the dissolved salt CH3COONa present at
the equivalence point, is the source of a weak
base, CH3CO2-(aq), as shown by the equation
below
CH3COONa (aq)  Na+(aq) + CH3CO2-(aq)
The CH3CO2-(aq) produced, being a weak base,
reacts with water as shown by the equation below, producing
OH-(aq), thus making the solution basic.
CH3CO2-(aq)+ H2O  CH3CO2H(aq) + OH-(aq)
In this titration, methyl red would give a premature end point with gradual color change. For this solution of
acetic acid and sodium hydroxide, the appropriate indicator is phenolphthalein, which changes color over
the pH range 8.0 to 9.6.
PURPOSE OF EXPERIMENT
The purpose of this experiment, with the use of the Vernier Lab Quest, is to:
1. Practice technique called standardization; EXACT concentration of an aqueous solution is determined
2. To determine the molarity of an aqueous solution of Acetic acid CH3CO2H (aq) by titrating
with a standardized sodium hydroxide solution and plotting a titration curve of pH vs volume of NaOH
3. To determine the molarity of an aqueous solution of HCl by titrating with a standardized
sodium hydroxide solution and plotting a titration curve of pH vs volume of NaOH
4. To calculate pH of solution after addition of amounts of titrant using Handerson-Hasselbalch equation
and make an estimation of the ionization constant of acetic acid
5. Compare the two titration curves of strong acid and weak acid
6. To plot a titration curve of a polyprotic acid, if time permits
4
KHP EQUATION Structure & Titration potassium hydrogen phthalate. (KHP)Titration in Water
Weak acid-Strong Base
O
O
C
C
Na+
O-
OH
+ H 2O
+ NaOH
O-
O-
C
C
K+
K+
O
O
Potassium hydrogen phthalate
(KHP), M.W. = 204.22 g
Figure 6.
Determine the following based on your experimental data:
1.
2.
The Molarity of the unknown CH3COOH in the vinegar solution from the concentration of the
standard NaOH solution and then the percent by weight of CH3COOH in the vinegar.
An experimental value of the acid dissociation constant Ka for CH3COOH from the pH at different
point in the titration, using your graph.
HINT: Acetic acid (the acid in Vinegar) has this equilibrium in water:
CH3COOH  CH3COO- + H+
+
Acid dissociation constant, K a = [ H ][ CH 3 COO ]
[ CH 3 COOH]
1
pH = pKa + log [CH 3COO ]
Or, Handerson-Hasselbalch equation:
[CH 3COOH ]
At 25% equivalence: (when ¼th of the NaOH required to reach equivalence point is delivered.)
CH3COOH
¾ remains

1
pH = pKa + log [CH 3COO ] = pKa + log
[CH 3COOH ]
CH3COO+
¼ neutralized
[1 ]
4
[3 ]
4
= pKa + log
H+
[1]
= pKa – 0.477
[3]
Calculate pKa and then find Ka, based on the fact that: pKa = - log Ka
At 100% equivalence: The hydrolysis equilibrium of acetate occurs at this point.

CH3COO- + H2O  CH3COOH + OH-1 Base dissociation constant, K b = [ OH ][ CH 3 COOH]
1
[ CH 3 COO ]
Separately determine [OH-1], [CH3COO-] and [CH3COOH] at this point to calculate Kb Then find Ka
by this relationship: Ka = Kw / Kb
5
PROCEDURE
A. PREPARATION AND STANDARDIZATION OF A 0.3?? M NaOH SOLUTION
Prepare about 300-mL of approximately 0.3?? M NaOH by diluting X mL of 6.0 M NaOH (dilute) with
distilled water in a flask to be kept stoppered when not in use. M1V1 = M2V2. Calculate X and check with
instructor. Shake the flask after its preparation in order to make sure the solution is homogeneous. Be
careful to keep this NaOH solution uncontaminated and stoppered throughout the rest of the experiment;
You will find the exact concentration of this NaOH solution to a high degree of accuracy by titrating
against the precisely known weight of a standard (potassium hydrogen phthalate, KHP, a very pure
compound). This process is thus called standardization.
1. Accurately weigh out two samples of KHP, around 0.8?? g each, on a weighing paper. Record the
weight to three significant figures on data sheet.
2. For each flask, carefully transfer a sample into 125 ml Erlenmeyer flask. Measure 25.0-mL distilled
water. Wash the weighing paper with a small amount of this water into the flask. Be sure to get all the
KHP off the weighing paper and into the 125-ml Erlenmeyer flasks. After washing the KHP down into
the flasks add the remaining water and swirl the flasks to totally dissolve the KHP. All the KHP
particles must be totally dissolved in solution.
3.
Add three drops of phenolphthalein to each flask, stopper, and set aside
EQUIPMENT Set up the equipment using the Vernier Lab Quest equipment and the following:
Lab Quest, Drop Counter, pH meter, ring stand, 3-prong clamp, stir plate, stirrer bar, flash drive,
10-mL graduated cylinder, waste beaker, 2 stoppers, 10.00-ml volumetric pipette & bulb,
LAB QUEST PROCEDURES
1. Set up Lab Quest, connect AC adapter, turn on
2. Lower Drop Counter on ring stand, attach 3-prong clamp to ring stand, connect Drop Counter to
the DIG/SONIC 1 channel on side of Lab Quest
2. Put together the value assembly for the 60-mL reservoir for the titrant, be sure both values are in the
“closed” position, attach reservoir to the 3-prong clamp, position the reservoir tip over the slot in the
Drop Counter. The drops must pass through the opening without touching anything
WAIT, when instructed, you will insert the pH meter through the round hole on the surface
of the Drop Counter and connect to channel 1 on the Lab Quest
3. CALIBRATE DROP VOLUME (mL) be sure drops fall through slot without any obstruction
a. Place a beaker below the slot of the Drop Counter
b. Fill the reagent reservoir with titrant
c. Open top value, open bottom value, adjust the drop flow by the top value, about 1-2 drops/sec
d. Let small amount of titrant to pass through the slot, close the bottom value only. Remove beaker
6
e. Place 10-mL graduated cylinder below the slot. Readjust equipment if needed
f. Lab Quest screen: select SENSORS then CALIBRATE (Drop Counter), CALIBRATE NOW
g. Open bottom value to start drops. Be sure drops pass through slot. In upper right part of screen you
should see the drop count
h. Continue till graduated cylinder is filled between 9.0 & 10.0-mL
i. Close bottom value and STOP Lab Quest program
j. Record the volume from the graduated cylinder to the nearest 0.1-mL
k. Touch box on screen, use the key pad, enter this volume value, and enter OK
l. At top of screen choose EQUATION, see Drops/mL, record this value. In the event you lose
your data, you can manual re-enter this value instead of doing another calibration
m. Continue with the experiment
4. STANDARDIZE NaOH SOLUTION Titrate KHP with NaOH solution
a. Place flask on stir plate and add stir bar. Adjust Drop Count above the flask. Fill reservoir with
titrant. Adjust speed of stir bar
b. Start Lab Quest. Open bottom value to slowly add NaOH solution to the flask
c. Add drops till one drop turns the solution permanently pink. Close bottom value. Record the volume
on the data sheet
d. Refill reservoir and titrate the second KHP sample
e. Calculate molarity of sodium hydroxide solution for each, average, and record the average
Results need to be within 0.5%
[NaOH] = (mols KHP)/{NaOH delivered (L)}
SAVE THE STANDARDIZED NaOH SOLUTION FOR THE REST OF THE EXPERIMENTS
B. ACETIC ACID: WEAK ACID-STRONG BASE TITRATION
PROCEDURE – Concentration of Acetic Acid in Vinegar Solution using stansardized NaOH soln
Obtain about 30-ml of an unknown vinegar solution in a DRY beaker or flask. This will be used for
pipetting the 10.00-ml of vinegar that you need for the titrations.
For all titrations, use a clean 250-ml BEAKER.
1. RECORD your unknown number on your DATA Sheet.
2. Refill the reservoir with the NaOH solution
3. Obtain a 10.00-ml volumetric pipette
4. Using the volumetric pipette, prepare a vinegar solution for titration as follows:
5. Pipet exactly 10.00-ml of the unknown vinegar solution from the 30-ml you have set aside, and put it
into the 250-ml beaker. To this 10.00-ml vinegar solution add some distilled water (15-20 ml) to
increase its volume in the beaker.
6. Add three drops of phenolphthalein to the beaker containing the 10.00-ml vinegar solution.
7. Put a magnetic stir bar in the beaker containing the 10.00-ml vinegar solution.
8. Insert the pH meter electrode (properly rinsed!) into this diluted vinegar solution, allowing clearance
above the bottom of beaker so that stirrer bar will not come into contact and break the electrode
7
9. Stir this vinegar solution at a moderate speed, without splashing
10. Start Lab Quest program, open bottom value, start titrating with the standardized NaOH solution
11. Titrate each vinegar sample till the graph flattens out at the top
12. OBSERVE the RATE OF CHANGE OF pH, and how strong a temporary coloring you are getting
from the phenolphthalein indicator.
13. Save graph on flash drive
14. Calculate the molarity of the acetic acid in the unknown vinegar solution. Record on the Data sheet.
C. HCl: STRONG ACID-STRONG BASE TITRATION
Prepare a 10.00-ml HCl solution as you did for the vinegar solution. Use Bromothymol blue as an
indicator this time. Record the titration results as you did with vinegar.
Continue this titration as you did in Part B - VINEGAR
Save the graph when finished.
D. POLYPROTIC ACID-STRONG BASE TITRATION
Prepare a 10.00-ml ACID solution as you did before.
Record the titration curve as you did before.
Save the graph when finished.
REQUIREMENTS
Make sure all your graphs are properly labeled, all values required recorded
Be sure all calculations are properly shown
Rinse pH probe with distilled water and PUT THE pH ELECTRODE INTO its original pH 7 buffer
solution when you have finished the experimentation
Clean and rinse all equipment before storing
PUT THE pH ELECTRODE INTO the original pH 7 buffer solution when you have finished
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Name______________________________________ 163 Titrations
DATA SHEET
Drop/mL _____________
Molarity NaOH ____________
Molecular Wt. KHP ____________
calculation:
A. Data for the Standardization of NaOH
Run 1
Run 2
Mass of KHP
_______________ g
_______________ g
volume NaOH at end point
_______________ mL
_______________mL
Show calculations for the molarity of sodium hydroxide for Run 1 and Run 2
Run 1
Run 2
Molarity of NaOH
_____________ M
Average molarity of NaOH
________________M
______________ M
B. pH Titration of Weak Acid Unknown Numbe _______
Run 1
Run 2 (optional)
Vol of CH3COOH
_______________ mL
_______________ mL
volume NaOH at end point
_______________ mL
_______________mL
9
Show calculations for the molarity of weak acid for Run 1 and Run 2
Run 1
Run 2
Molarity of CH3COOH
_____________ M
Average molarity of CH3COOH
________________M
______________ M
Acid dissociation constant of acetic acid from experimental work. Show calculations below.
(Note Equation (1), page 2, for calculation of H+ concentration and the HINT which is on page 5-6)
At 25% of equivalence
At 50% of equivalence
At 75% of equivalence
At 100% of equivalence
10
pH (Show how to find
this in your graph)
[H+]
Calculated Ka
At 25% of equivalence
At 50% of equivalence
At 75% of equivalence
At 100% of equivalence
QUESTION
Is Methyl Red a suitable indicator for the NaOH-vinegar titration? Why or Why not?
C. pH Titration of Strong Acid
Unknown Number _________
Run 1
Run 2 (optional)
Vol of HCl
_______________ mL
_______________ mL
volume NaOH at end point
_______________ mL
_______________mL
Show calculations for the molarity of strong acid for Run 1 and Run 2
Run 1
Run 2
Molarity of HCl
Average molarity of HCl
D. pH Titration of Polyprotic Acid
record and save titration curve for acid
_____________ M
______________ M
________________M