Experiment 1
Acid-Base Titrations
A. Purpose
1. To learn the concept and technique of titration.
2. To standardize a sodium hydroxide (NaOH) solution against a primary
standard acid.
3. To determine the concentration of an unknown acid solution by titration with
the standardized base solution.
B. Theoretical Background
An acid reacts with a base to yield water (neutralization) and a salt. The reaction
is stoichiometric (complete) and thus feasible for study in simple quantitative
analysis. Titration is defined as the gradual addition of a measurable volume of a
solution (the titrant), to exactly react a certain amount of another substance. Thus
we can titrate an acid with a base (or vice-versa) because we have an
instantaneous stoichiometric neutralization reaction that is taking place. The point
at which all the substance (acid or base) is exactly reacted (and thus the titration
subsequently stopped), is called the equivalence point of the titration. In a simple
acid-base titration, the equivalence point is detected visually by using an acidbase indicator. An indicator is a substance (added in small amount to the titration
flask) which has the virtue of changing its color just at the point when the reactant
in the flask is completely consumed. Such a visually determined time to stop the
titration is called the end-point.
In this experiment, you will standardize a sodium hydroxide (NaOH) solution by
titration with a primary acid standard (potassium hydrogen phthalate,
abbreviated as KHP, M = 204.23 g/mol). The chemical equation of the reaction
can be written as:
HP + OH 
H2O + P2
This standardized sodium hydroxide solution is then used to determine the
concentration of an unknown (monoprotic) acid.
C. Chemicals and Setup
Sodium hydroxide solution (~ 0.1 M), KHP (solid), Unknown acid solution,
phenolphthalein indicator.
250 mL volumetric flask, 50 mL buret, 25 mL volumetric pipet, two clean 125
mL (or 250 mL) Erlenmeyer flasks, polyethylene dropping bottle, buret stand.
4
D. Procedure
 Preparation of a solution of primary standard (KHP)
1. Obtain from the stock room a sample of solid KHP. Weigh accurately (by
difference) about 3 g of KHP.
2. Transfer this weight quantitatively to a 250 mL volumetric flask. Swirl
thoroughly to dissolve the solid, then make up to the mark with distilled
water.
3. Calculate the molarity of the KHP solution that you have just prepared.
 Standardization of the NaOH solution
1. Rinse your clean buret with distilled water then with the NaOH solution.
Fill the buret with NaOH solution (the solution in the buret should be free
of air bubbles). Record the upper buret reading.
2. Pipet 25.00 mL of the KHP solution into a clean Erlenmeyer flask. Add 25
mL of distilled water (with a graduated cylinder here, there is no need for
accurate measurement, why?). Then add 2 drops of phenolphthalein
indicator.
3. Titrate the resulting solution with the NaOH solution carefully until the
indicator changes color (from colorless [acidic] to faint pink [basic]).
4. Record the lower buret reading.
5. Repeat steps (2) through (4) until the titration volumes in at least two runs
agree within 0.10 mL.
6. Carry out a blank titration. (The blank solution contains 50 mL of distilled
water and 2 drops of indicator).
 Determination of the concentration of an unknown acid
1. Pipet 25.00 mL of the unknown acid solution provided by your lab
instructor into a clean Erlenmeyer flask. Add 25 mL of distilled water and
2 drops of phenolphthalein indicator.
2. Titrate the solution with the NaOH solution carefully until the indicator
changes color.
3. Record the lower buret reading.
4. Repeat steps (2) through (4) until the titration volumes in at least two runs
agree within 0.10 mL.
5. Carry out a blank titration.
5
E. CHEM. 203 Lab Report
NAME : _______________________
EXPERIMENT No. : ______________
TITLE : __________________________________
DATE : ________________
Partner : ________________________
 Preparation of a solution of primary standard (KHP)
M 0KHP 
m KHP

M KHP  Vsoln
=
 Standardization of the NaOH solution
1
2
3
Blank
Upper reading (mL)
Lower reading (mL)
Volume (mL)
Average volume, v NaOH  v  v blank 
M 0NaOH 
M 0KHP  v KHP

v NaOH
=
 Determination of the concentration of an unknown acid
1
2
3
Blank
Upper reading (mL)
Lower reading (mL)
Volume (mL)
6
Average volume, v NaOH  v  v blank 
M 0unk 
M 0NaOH  v NaOH

v unk
=
Questions
1. A sodium carbonate (Na2CO3) primary standard solution is prepared by dissolving
10.587 g of the substance in enough distilled water, then transferring it
quantitatively to a 250 mL volumetric flask and finally making up to the mark
with distilled water. The volumetric flask is labeled T.C.  0.12 mL. Calculate the
molarity of the standard Na2CO3 solution.
2. 10.00 mL of a dilute acetic acid solution are titrated with a sodium hydroxide
(NaOH) solution of molar concentration 0.1078 M. At the equivalence point, it is
found that 13.42 mL of NaOH have been added. What is the molarity of the
unknown acetic acid solution?
7
3. Write out a balanced reaction for the neutralization of sulfuric acid (H2SO4) with
sodium hydroxide. What volume of 0.541 M NaOH is necessary to react
completely 20.0 mL of sulfuric acid whose molarity is 1.26 M?
4. Explain why we run a blank titration.
5. The equivalence point of the titration of a weak base with a strong acid occurs at
pH 5.32. Which of the following indicators is the best to detect the end-point of
the titration?
Indicator
Bromophenol blue
Methyl orange
Methyl red
Bromocresol green
Color change
yellow – blue
red – yellow
pink – yellow
Yellow – blue
pH range
3.1 – 4.6
3.3 – 4.5
4.2 – 6.3
3.8 – 5.4
Sketch the titration curve for this titration.
8