Determining the Mass Percent of
Sodium Hypochlorite (NaClO) in Bleach
Introduction
The active ingredient in most chlorine bleaches is sodium hypochlorite, NaClO. The purpose of this
two-part experiment is to determine the mass percent of sodium hypochlorite in a commercial chlorine
bleach sample. A brief description of each part of this experiment is given below.
Part 1
In the first part of this activity, you will prepare a potassium iodate (KIO3) primary standard and use
this solution to standardize a solution of sodium thiosulfate (Na2S2O3). The standardized Na2S2O3
solution will be used in Part 2 of this experiment to determine the concentration and mass percent of
NaClO in a chlorine bleach sample.
Standardization is a process in which a solution whose concentration is approximately known is
titrated with another solution whose concentration is accurately known. The solution having an
accurately known concentration is called a primary standard. In this experiment, the KIO3 primary
standard will be prepared from solid potassium iodate that has been dried overnight in a 110 C oven.
KIO3 has the characteristics required of any compound used as a primary standard. Most notably, it
can be manufactured with a high purity; it is stable in air; it has a relatively long shelf life; it has a high
molar mass (which helps reduce errors in massing); and it has a predictable and well-defined reaction
with the solution being standardized.
The chemistry that will be exploited in this part of the investigation is summarized in the net ionic
reactions shown below.
IO3- (aq) + 5 I- (aq) + 6 H+ (aq)  3 I2 (aq, brown/yellow) + 3 H2O (l)
I2 (aq, brown/yellow) + 2 S2O32- (aq)  2 I- (aq, colorless) + S4O62- (aq)
As you can see, you will be studying the chemistry that occurs when I2, formed by the reaction of
potassium iodate (KIO3) and potassium iodide (KI) under acidic conditions, reacts with sodium
thiosulfate (Na2S2O3). When KIO3 is added to KI under acidic conditions, IO3- oxidizes iodide (I-) to
iodine (I2) and the solution becomes a brown/yellow color. When Na2S2O3 is added to this solution,
the I2 is reduced back to I-. An abrupt color change (from brown/yellow to colorless) will occur when
all of the available I2 is consumed by the added Na2S2O3 (i.e., when all of the I2 has been reduced). In
a titration, this point is called the equivalence point. Since the brown/yellow solution will turn
colorless at the point where all of the available I2 has been reduced (due to reaction with S2O32-), I2 will
act as an internal indicator in the redox titration.
Part 2
In the final part of this experiment, the standardized Na2S2O3 solution will be used to determine the
mass percent of sodium hypochlorite in a commercial chlorine bleach sample. In this case, an excess
of KI will be added to a dilute bleach sample that has been acidified. Under these conditions, the
hypochlorite in the bleach will react with iodide (I-) to produce the brown/yellow I2. This solution will
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then be titrated with the standardized Na2S2O3 solution to determine the concentration of NaClO in the
diluted bleach sample and the original bleach solution. The relevant chemical reactions are given
below.
ClO- (aq) + 2 I- (aq) + 2 H+ (aq)  I2 (aq, brown/yellow) + Cl- (aq) + H2O (l)
I2 (aq, brown/yellow) + 2 S2O32- (aq)  2 I- (aq, colorless) + S4O62- (aq)
Since the density of the bleach solution is known (1.084 g/mL), the mass percent of sodium
hypochlorite in the original bleach sample can be determined.
Pre-lab
Safety
Sodium Hypochlorite is a strong oxidizer. Irritating effects increase with the strength of solution and
the time of exposure. Flush exposed skin with water for several minutes. Fumes are irritating to eyes
and mucous membranes.
DO NOT MIX CHLORINE BLEACH WITH toilet bowl cleaners, ammonia, lye, rust removers,
vinegar, or oven cleaners.
Do not mix chlorine with any other cleaning agent! The result can be a sudden release of TOXIC
chlorine gas. Inhalation can cause serious injury or death.
Potassium Iodate may cause skin irritation. As with all chemicals, wash hands thoroughly if skin is
exposed. Use the eyewash if splashed in the eyes.
Caution! Acids, such as H2SO4 are hazardous if splashed on clothing, exposed skin or in the eyes.
Prolonged exposure of the skin to even fairly dilute solutions of acid and base can cause serious burns.
If acids splash on skin or clothes, remove the affected clothing and flush the affected areas thoroughly
with cold water.
Disposal of Chemicals: All solutions may be disposed of in the sink with plenty of water.
1. Explain why chlorine bleaches should not be mixed with other household chemicals such as
ammonia. (1 pt)
2. Define the following terms: primary standard, standard solution, equivalence point. (0.3 pt)
3. Why do you think the solid KIO3 that you will use to make your primary standard will be dried
overnight in an oven? (1 pt)
4. Describe how 100 mL of 0.01 M KIO3 primary standard can be prepared from solid potassium
iodate. Write a step-by-step procedure that includes supporting calculations and any
equipment/glassware that is required. Keep in mind that the concentration of a primary standard
must be accurately known, which could influence your preparation technique and choice of
glassware. (2 pts)
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5. What is an internal indicator? What is the internal indicator in the redox titrations carried out in
this experiment? (1 pt)
6. Use the chemical reactions shown below to answer the following questions. (1 pt)
IO3- (aq) + 5 I- (aq) + 6 H+ (aq)  3 I2 (aq, brown/yellow) + 3 H2O (l)
I2 (aq, brown/yellow) + 2 S2O32- (aq)  2 I- (aq, colorless) + S4O62- (aq)
a.
What is the stoichiometric relationship between IO3- and I2?
b. What is the stoichiometric relationship between I2 and S2O32-?
c.
What is the stoichiometric relationship between IO3- and S2O32-?
d. How will knowing these relationships help you calculate the molarity of the Na2S2O3 solution
in Part 1 of this experiment?
7. A person’s blood alcohol (C2H5OH) level can be determined by titrating a sample of blood plasma
with potassium dichromate solution. The balanced equation is:
16 H+ (aq) + 2 Cr2O72- (aq) + C2H5OH (aq)  4 Cr3+ (aq) + 2 CO2 (g) + 11 H2O (l)
If 35.46 mL of 0.05961 M Cr2O72- is required to titrate 28.00 g of plasma, what is the mass percent
of alcohol in the blood? (2 pts)
Procedure
Chemicals and Solutions Used in this Experiment
Solids: Na2S2O3.5 H2O, KI, KIO3
Solutions: 0.5 M H2SO4, bleach
Part 1 - Standardizing a Na2S2O3 Solution
Discuss your answers to Pre-Lab Questions 1-7 with your lab mates. This dialog will be followed by a
class discussion. A class consensus should be reached regarding the proper method for preparing a
primary standard solution and why is it important to accurately know the concentration of the KIO3
solution.
Team Planning
In this part of the experiment, you will be standardizing a solution of Na2S2O3 with a KIO3 primary
standard. As a team, read over the Procedural Help section given below. Before proceeding, plan
your investigation. This should include a discussion of the equipment you will need, the calculations
involved in preparing your primary standard, the calculations for preparing the Na2S2O3 solution to be
standardized, and also the proper method for determining the concentration of the Na2S2O3 solution
using your titration data. Consult with your TA to see if your plan seems reasonable. Record your
plan in your lab notebook along with any data that you are given or that you collect (mass and identity
of reagents used, reagent concentrations, volume of reagents added, etc.).
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Read the general procedure for the redox titration below, and carry out your group’s plan to
standardize the Na2S2O3 solution. Record all pertinent data, observations, and calculations in your lab
notebook. All data should be reported with the appropriate number of significant figures. In addition,
the proper units should be included on all of your measurements and calculated quantities.
1. Prepare 100 mL of 0.01 M KIO3 primary standard according to the procedure that your group has
agreed upon.
2. Prepare 250 mL of a 0.1 M Na2S2O3 solution.
3. Fill a clean buret with the Na2S2O3 solution that your group prepared and record the initial volume
to the nearest 0.01 mL.
4. Pipet 25 mL of your group's primary standard solution into a 125 mL Erlenmeyer flask. Add ~2 g
of solid KI to this solution. Swirl the solution to dissolve the KI. Add 10 mL of 0.5 M H2SO4 to
the flask.
5. Titrate the KIO3/KI/H2SO4 solution with Na2S2O3 until the solution turns colorless. Be sure to stir
the reaction or swirl continuously. Record the final buret volume in your notebook.
6. Use your titration data to determine the molar concentration of the Na2S2O3 solution.
7. Repeat the titration two more times and determine an average Na2S2O3 concentration.
8. Do not dispose of the standardized Na2S2O3 solution as it will be used in Part 2 of this lab to
determine the mass percent of NaClO in a commercial bleach product.
Part 2 - Determining the Mass Percent of NaClO in a Commercial Chlorine Bleach
In this part of the experiment, the Na2S2O3 solution that was standardized in Part 1 will be used to
determine the mass percent of sodium hypochlorite in a commercial chlorine bleach sample. In this
case, an excess of KI will be added to a dilute bleach sample that has been acidified. Under these
conditions, the hypochlorite in the bleach will react with iodide (I-) to produce the brown/yellow I2.
This solution will then be titrated with the standardized Na2S2O3 solution to determine the
concentration of NaClO in the diluted bleach sample and the original bleach solution. The relevant
reactions are shown below.
ClO- (aq) + 2 I- (aq) + 2 H+ (aq)  I2 (aq, brown/yellow) + Cl- (aq) + H2O (l)
I2 (aq, brown/yellow) + 2 S2O32- (aq)  2 I- (aq, colorless) + S4O62- (aq)
The mass percent of sodium hypochlorite in the original bleach sample will be determined using the
calculated NaClO concentration in the bleach solution and the density of the bleach solution, 1.084
g/mL.
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Team and Class Discussion
Discuss your answers to Pre-Lab Questions 8 and 9 with your group. This group dialog should be
followed by a class discussion if time permits and groups are progressing through this experiment at a
similar pace. Before you begin this part of the lab, group (and/or class) consensus should be reached
regarding the concepts covered in these Pre-Lab questions.
Discuss the stoichiometric relationship between the number of moles of Na2S2O3 you will titrate into
the bleach sample and the number of moles of NaClO in the bleach.
Team Planning
Use the information in the Procedural Help section to help you formulate a plan to determine the mass
percent of NaClO in a commercial bleach sample. Your plan should include the necessary equipment,
the amounts and types of reagents you will require in your investigation, and the approach you will use
to determine the mass percent of NaClO in your bleach sample. Your proposal should also include
how your team plans to share the work. When your team has agreed on a strategy, write it down in
your lab notebook and review it with your TA. You may begin this part of the experiment as soon as
your plan has been approved.
As you implement your investigation, record all pertinent data, observations, and calculations in your
lab notebook. All data should be reported with the appropriate number of significant figures, and the
proper units should be included on all of your measurements and calculated quantities. Your notebook
should also indicate the tasks for which you were personally responsible.
1. Pipet 10 mL of a bleach sample into a 100 mL volumetric flask and dilute to the mark with
deionized water.
2. Pipet 25 mL of the diluted bleach into a 250 mL Erlenmeyer flask.
3. Add ~2 g of KI, 10 mL of 0.5 M H2SO4, and 20 mL of deionized water to the flask.
4. Titrate the dilute bleach sample with your standardized Na2S2O3 solution.
5. Use your titration data and reaction stoichiometry to determine the molar concentration of NaClO
in the 25 mL dilute bleach sample.
6. Determine the molar concentration of NaClO in the original bleach sample.
7. Determine the mass percent of NaClO in the commercial bleach solution. Assume that the density
of the bleach solution is 1.084 g/mL.
8. Repeat the titration and the associated calculations two more times.
9. Calculate an average value for the mass % of NaClO in the commercial bleach solution that was
investigated.
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