Titrations Revisited
By Drew Rutherford
Concordia College
Introduction
The first experiment today will be the titration of acetic acid in vinegar. Vinegar is a solution of
acetic acid, an organic acid of formula CH3COOH (MW = 60.0526 g/mole). In order to be sold as
vinegar, it needs to meet the FDA’s guideline of 5.00% acetic acid by mass. Knowing that the
density of vinegar is 1.04 g/mL and analyzing this solution by titration, chemists can determine the
mass percentage of acetic acid in a sample of vinegar. The titration reaction is given below:
CH3COOH + NaOH  CH3COONa + H2O
Reaction 1
A customer has purchased a vinegar solution at a local thrift store at a 20% discount and he
believes that the sample of vinegar he has purchased does not meet FDA regulations. He is suing
the thrift store for the $0.36 he feels that he has been cheated out of. The court has asked you to
analyze the sample and render your verdict. Does this sample conform to the FDA guideline?
If 2.00 mL of the thrift store vinegar required 9.73 mL of 0.150 M NaOH to reach the endpoint, then
0.00877 g of acetic acid was present.
9.73 mL NaOH x 1 liter NaOH x 0.150 mole NaOH x 1 mole CH3COOH x 60.0526 g CH3COOH
1000 mL
1 liter
1 moles NaOH
1 mole CH3COOH
= 0.0876 grams CH3COOH in sample
2.00 mL vinegar x 1.04 g vinegar = 2.08 g vinegar
1 mL
0.0876 g CH3COOH x 100 = 4.21 % CH3COOH by mass
2.08 g vinegar sample
You can find the molarity of acetic acid in vinegar by dividing the number of moles of acetic acid
in vinegar by the volume of vinegar used.
Molarity of CH3COOH = 0.00146 moles CH3COOH
0.00200 L
Molarity of CH3COOH = 0.73 M
The thrift store vinegar does not meet the FDA guidelines.
The second experiment today will be the determination of the mass percentage of sodium
carbonate in a mixture. This procedure will use a powerful variation of titration technique called
back-titration. This method is often employed to analyze unknown samples of bases. This technique
utilizes a standardized acid to overneutralize the unknown base sample. After the sample of base
has been overneutralized, a standardized base is added from the buret to the solution to titrate the
excess acid that was added. Since the total number of moles of added acid are known and the
excess acid remaining after complete neutralization of the unknown base can determined by
titration, one can determine the number of moles of bases present in the analyzed sample by the
difference between total number of moles of acid added and the number of moles of acid
remaining after neutralization. It is critical to the back titration technique to be certain to add too
much standardized acid to the unknown, but not so much that the excess acid remaining cannot
be titrated by the total capacity of one buret of standardized base.
For example, if 30.00 mL of a 0.200 M HCl solution was added to a 0.225-g sample of a mixture of
sodium carbonate and lithium chloride, the acid-base reaction between the HCl and the Na2CO3
results in the production of sodium chloride, water and carbon dioxide (see reaction below). The
lithium chloride present in the original sample does not affect the acid/base reaction. Note:
Carbonic Acid (H2CO3) is unstable in solution and will spontaneously decompose to water and
carbon dioxide.
2 HCl + Na2CO3  2NaCl + H2O + CO2-
Reaction 2
If the 0.225-g sample was completely sodium carbonate (MW=105.99), the balanced equation
indicates that we would need 0.00425 moles of HCl to react with the sodium carbonate.
STEP 1:
0.225 g Na2CO3 x 1 mole Na2CO3 x 2 moles HCl = 0.00425 moles HCl
105.99 g Na2CO3 1 mole Na2CO3
How many milliliter of our standardized 0.200 M HCl solution (that we made last week in lab) should
we add to the sodium carbonate?
STEP 2:
molarity = mole
Liters
0.200 M = 0.00425 moles HCl
liters
0.0213 liters or 21.3 mL
However, we added too much HCl to the solution. We added 30.00 mL of HCl. After the reaction is
complete, what remains in the flask?
1.
2.
3.
4.
Sodium chloride (NaCl)
Lithium chloride (LiCl)
Water (H2O)
Excess HCl that wasn’t needed for the reaction
The excess HCl remaining in this solution can then be titrated (back-titration) with a standardized
sodium hydroxide solution and the number of moles of excess acid determined. The reaction
during the back-titration is
HCl + NaOH  NaCl + H2O
Reaction 3
If the solution required 16.22 mL of 0.150 M NaOH to reach the endpoint, there were 0.00240 moles
of HCl left in solution (the excess).
Step 3:
16.22 mL x 1 liter x 0.150 moles NaOH x 1 mole HCl =
1000 mL
1 liter
1 mole NaOH
0.00243 moles HCl
How many moles of HCl did we add to the solution? We added 30.00 ml of 0.200 M HCl.
Step 4:
0.200 M HCl = moles HCl
0.03000 L
0.00600 moles HCl
A total of 0.00600 moles of HCl was added at the beginning. The excess amount of HCl (the
amount not used in the reaction of HCl with Na2CO3) was 0.00243 moles HCl.
Step 5:
total moles HCl = moles HCl used in reaction with Na2CO3 + excess HCl
0.00600 moles
= moles HCl used in reaction with Na3CO3
+
0.00243 moles HCl
0.00357 moles HCl used in reaction with Na2CO3
Reaction 2 shows the stoichiometry of the reaction of HCl with sodium carbonate.
Step 6:
0.00357 moles HCl x 1 moles Na2CO3 x 105.99 g Na2CO3 = 0.189 g Na2CO3
2 moles HCl
1 moles Na2CO3
The amount of sodium carbonate in the 0.225 g mixture was only 0.189 g. The remaining mass was
lithium chloride.
Step 7:
% Na2CO3 = 0.189 g Na2CO3 in sample x 100
0.225 g total sample mass
% Na2CO3 = 84.0 %
Experiment
General Comments
Use only deionized (DI) water for this (and any titration) lab.

Once a titration is started, it is important that the volume of titrant in the buret is sufficient to
determine the endpoint for the trial without refilling the buret. Refilling a buret in the middle of a
titration will introduce significant errors in measurement.

Be sure to completely close the KHP container to prevent KHP from prolonged contact with the
air.

The endpoint of phenolphthalein is colorless to light pink (acidic solution to basic solution).

The endpoint of 2:3 methyl red:bromocresol indicator is light pink to light green. The complete
transition is from light pink (sometimes tannish) to light green to light blue (acidic solution to
basic solution). It is sometimes difficult to see the green endpoint since one-half a drop over will
give the blue color without observation of the light green intermediate.

Methyl red indicator changes from red to yellow (pKa = 5.4) while bromocresol green indicator
changes from yellow to blue (pKa = 4.7).

The combination of methyl red/bromocresol green indicators gives pink (red plus yellow in
acidic solution) to green (yellow plus blue basic solution) at the endpoint.

In the back titration, in order to ensure complete reaction between the HCl and the sodium
carbonate and to drive off all carbon dioxide, the reaction solution is heated to boiling on a hot
plate for 3 minutes and cooled to room temperature before the titration is conducted.

Sodium hydroxide solutions react with carbon dioxide from the air over time. This reaction
changes the concentration of standardized solutions of NaOH slightly. In order to obtain the
most accuracy, your sodium hydroxide solution should be restandardized against KHP.
Restandardize your NaOH solution:
Follow the method from last week to restandardize your NaOH solution against KHP.
Titration of Vinegar:
Obtain 10-15 mL of vinegar sample in clean, dry 50-mL beaker. Transfer 2.00 mL of vinegar sample
to a 125-mL Erlenmeyer using a volumetric pipette. Dilute with 30-40 mL of DI water. Use 3 drops of
phenolphthalein indicator and titrate with your standardized sodium hydroxide solution from last
week. Repeat at least three times. Calculate the molarity of acetic acid in the vinegar and the
mass percentage of acetic acid. Comment on whether the vinegar sample conforms to FDA
standards.
Back Titration of Na2CO3/LiCl Mixture:
Weigh out 0.200-0.250-g of the Na2CO3/LiCl unknown mixture and transfer completely to a dry, 125mL Erlenmeyer flask. Calculate the amount of your standardized HCl solution you would need to
neutralize the sample if it was completely sodium carbonate and add 2 extra milliliters of HCl to
ensure that you have an excess of HCl. Round up to a whole number. If it would take 22.3 mL, add
2 mL of HCl for a total of 24.3 mL and round up to 25.00 mL. Add that quantity of your standardized
HCl solution to the solid sample with a volumetric pipette.
Add 5 drops of 2:3 methyl red:bromocresol green indicator. Heat solution on hot plate for 3 minute.
Let the solution cool to room temperature. Titrate excess HCl in solution with your standardized
sodium hydroxide solution. Repeat at least three times. If you had a large excess of HCl, on
subsequent titrations, reduce the amount of HCl that you add. Calculate the percentage of
sodium carbonate in unknown sample.