1. a Write the  balanced equations for the following oxidation–reduction reactions that you will encounter in this experiment:  a. The reduction of ClO− by I− in acidic solution  ? b. The reduction of I2 by S2O32−in acidic solution

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Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
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Chapter1: Chemical Foundations
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1. a Write the  balanced equations for the following oxidation–reduction reactions that you will encounter in this experiment:  a. The reduction of ClO− by I− in acidic solution  ?

b. The reduction of I2 by S2O32−in acidic solution    

      2.a.  What conversion factor will enable you to calculate the number of moles of ClO– ions from the number of moles of S2O32− ions used in the titration?  

b.  A 10.0-mL sample of aqueous NaOCl is treated with excess KI in an acidic solution. The quantity of iodine that is liberated is such that 28.02 mL of 0.0250 M Na2S2O3 solution must be added to cause the disappearance of the dark blue color due to the starch indicator. What is the molarity of the solution of NaOCl?  

388
Cleaning and Filling your Buret
1. Instructions for using a buret can be found in the Introduction section of this manual. Clean
your buret according to the directions, and fill it with some of the Na₂S₂O3 solution.
3.
4.
Doing the Trial Titration
1. Pipet 10.0 mL of the diluted bleach solution into a clean 125-mL Erlenmeyer flask.
2.
Add one of the samples of solid KI to the flask, followed by 20 mL of distilled water from a
graduated cylinder and 20 drops of 2 M HCl. Swirl to obtain a homogeneous solution.
5.
6.
7.
8.
CAUTION: Undiluted bleach and hydrochloric acid can cause
chemical burns and ruin your clothes. In addition, bleach can be
especially irritating to the eyes. If you spill either of these
solutions on you, wash the contaminated area thoroughly and
immediately report the incident to your laboratory instructor.
You may require further treatment. Be sure to wear approved
chemical splash goggles.
3.
4.
5.
Experiment 21B
6.
7.
8.
9.
Record the initial buret reading to the closest 0.01 mL.
Place the flask under the buret with the capillary tip inside the mouth of the flask. Insert a piece of
white paper under the flask.
Doing the Exact Titrations
1. Repeat Steps 1 through 4 of the procedures for the trial titration.
2.
Subtract 2 mL. from the volume found in the trial titration. Rapidly add this new volume to the
flask from the buret.
Add 40 drops of the 0.2% starch indicator to the flask.
Rinse the walls of the flask with distilled water from a plastic wash bottle.
Continue the titration on a drop-by-drop basis. Swirl the flask rapidly after each addition.
Remember that the endpoint will occur when one drop results in a colorless solution.
While swirling, add increments of about 1 mL of the Na₂S₂O3 solution to the flask. Continue until
the brown color fades to yellow.
Add 40 drops of a 0.2% starch solution to the flask.
Continue to add 1-mL portions of the Na₂S₂O3 solution until one addition causes the solution to
become colorless.
Record the final buret reading to the nearest 0.01 mL, and calculate the approximate volume of
the Na₂S₂O3 solution required for the titration.
Record the final buret reading to the nearest 0.01 mL.
Repeat the procedure with a second sample of the diluted bleach solution.
If the volumes at the endpoints differ by more than 0.15 mL. (about 3 drops) or some other amount
specified by your instructor, repeat the titrations with additional samples of the diluted bleach
until the required precision is obtained.
Calculate and record the molarity of the diluted bleach solution. Obtain the mean molarity.
Transcribed Image Text:388 Cleaning and Filling your Buret 1. Instructions for using a buret can be found in the Introduction section of this manual. Clean your buret according to the directions, and fill it with some of the Na₂S₂O3 solution. 3. 4. Doing the Trial Titration 1. Pipet 10.0 mL of the diluted bleach solution into a clean 125-mL Erlenmeyer flask. 2. Add one of the samples of solid KI to the flask, followed by 20 mL of distilled water from a graduated cylinder and 20 drops of 2 M HCl. Swirl to obtain a homogeneous solution. 5. 6. 7. 8. CAUTION: Undiluted bleach and hydrochloric acid can cause chemical burns and ruin your clothes. In addition, bleach can be especially irritating to the eyes. If you spill either of these solutions on you, wash the contaminated area thoroughly and immediately report the incident to your laboratory instructor. You may require further treatment. Be sure to wear approved chemical splash goggles. 3. 4. 5. Experiment 21B 6. 7. 8. 9. Record the initial buret reading to the closest 0.01 mL. Place the flask under the buret with the capillary tip inside the mouth of the flask. Insert a piece of white paper under the flask. Doing the Exact Titrations 1. Repeat Steps 1 through 4 of the procedures for the trial titration. 2. Subtract 2 mL. from the volume found in the trial titration. Rapidly add this new volume to the flask from the buret. Add 40 drops of the 0.2% starch indicator to the flask. Rinse the walls of the flask with distilled water from a plastic wash bottle. Continue the titration on a drop-by-drop basis. Swirl the flask rapidly after each addition. Remember that the endpoint will occur when one drop results in a colorless solution. While swirling, add increments of about 1 mL of the Na₂S₂O3 solution to the flask. Continue until the brown color fades to yellow. Add 40 drops of a 0.2% starch solution to the flask. Continue to add 1-mL portions of the Na₂S₂O3 solution until one addition causes the solution to become colorless. Record the final buret reading to the nearest 0.01 mL, and calculate the approximate volume of the Na₂S₂O3 solution required for the titration. Record the final buret reading to the nearest 0.01 mL. Repeat the procedure with a second sample of the diluted bleach solution. If the volumes at the endpoints differ by more than 0.15 mL. (about 3 drops) or some other amount specified by your instructor, repeat the titrations with additional samples of the diluted bleach until the required precision is obtained. Calculate and record the molarity of the diluted bleach solution. Obtain the mean molarity.
21B. The Strength of a Laundry Bleach
Introduction
Hypochlorous acid (HCIO) is one of the important chlorine oxoacids (Ebbing/Gammon, Section
21.9). Solutions of sodium hypochlorite (NaOCI), a salt of that acid, are sold as laundry bleach. The
hypochlorite anion (CIO) is a strong oxidizing agent, but not as strong as ClO₂, CIOs, and ClO4.
Purpose
In this experiment, you will use an oxidation-reduction titration to determine the quantity of NaOCI
in a commercial bleach.
Concept of the Experiment
As you will see, you can determine the quantity of the CIO ion in a solution through two oxidation-
reduction reactions. First, a known quantity of this anion is reduced to CI ions in an acidic solution,
using excess potassium iodide. The I ions are oxidized to I₂ in this reaction. The solution that results is
brown because that is the color of I₂ in water. Second, the I2 is reduced to L during a titration with
asolution of sodium thiosulfate. Thiosulfate anions (S₂O) are oxidized to tetrathionate ions (S₂0²)
in this reaction.
You may have found this description somewhat confusing, but it should become clear after you
have balanced the equations for the two oxidation-reduction reactions in the Prelaboratory
Assignment. Moreover, the relationship between the original quantity of ClO ions and the quantity
of S₂0 ions used in the titration will be evident after these equations are balanced.
Although you could use the disappearance of the color due to aqueous I₂ to detect the endpoint of the
titration, this technique would not be very sensitive. Instead, you will use starch as an indicator.
Starch reacts with 12 to form a dark blue color. This reaction is reversible. Consequently, the blue
color fades during the course of the titration as I is consumed. The endpoint occurs when one drop of
the Na₂S₂O3 solution causes the color to change from blue to colorless. A trial titration will enable you
to locate the endpoints of subsequent titrations more easily.
Procedure
Getting Started
1.
2.
Obtain a 10-mL transfer pipet and a 50-mL buret.
Obtain about 70 mL of a 0.0250 M solution of Na₂S₂O3 and 3 samples of solid KI. Each of
these samples should have a volume of about 1 cm³. They can be stored on pieces of waxed
paper.
3.
A buret containing the bleach solution will be available for general use. This solution must be
diluted, however. Record the initial buret reading to the nearest 0.01 mL. Carefully add about 3
mL of the solution to a 100-mL volumetric flask. (A 100-mL graduated cylinder may be used if
a volumetric flask is not available.) Record the final buret reading to the nearest 0.01 mL, and
calculate the volume of undiluted bleach used. Add distilled water to the flask until the bottom
of the meniscus coincides with the etched line on the flask. Add the last 0.5 mL with a medicine
dropper. Insert a stopper in the flask, and mix the solution thoroughly.
4.
Be careful in your handling of the solutions used in this experiment. Obtain directions from your
laboratory instructor regarding proper disposal of all solutions used during this experiment.
Transcribed Image Text:21B. The Strength of a Laundry Bleach Introduction Hypochlorous acid (HCIO) is one of the important chlorine oxoacids (Ebbing/Gammon, Section 21.9). Solutions of sodium hypochlorite (NaOCI), a salt of that acid, are sold as laundry bleach. The hypochlorite anion (CIO) is a strong oxidizing agent, but not as strong as ClO₂, CIOs, and ClO4. Purpose In this experiment, you will use an oxidation-reduction titration to determine the quantity of NaOCI in a commercial bleach. Concept of the Experiment As you will see, you can determine the quantity of the CIO ion in a solution through two oxidation- reduction reactions. First, a known quantity of this anion is reduced to CI ions in an acidic solution, using excess potassium iodide. The I ions are oxidized to I₂ in this reaction. The solution that results is brown because that is the color of I₂ in water. Second, the I2 is reduced to L during a titration with asolution of sodium thiosulfate. Thiosulfate anions (S₂O) are oxidized to tetrathionate ions (S₂0²) in this reaction. You may have found this description somewhat confusing, but it should become clear after you have balanced the equations for the two oxidation-reduction reactions in the Prelaboratory Assignment. Moreover, the relationship between the original quantity of ClO ions and the quantity of S₂0 ions used in the titration will be evident after these equations are balanced. Although you could use the disappearance of the color due to aqueous I₂ to detect the endpoint of the titration, this technique would not be very sensitive. Instead, you will use starch as an indicator. Starch reacts with 12 to form a dark blue color. This reaction is reversible. Consequently, the blue color fades during the course of the titration as I is consumed. The endpoint occurs when one drop of the Na₂S₂O3 solution causes the color to change from blue to colorless. A trial titration will enable you to locate the endpoints of subsequent titrations more easily. Procedure Getting Started 1. 2. Obtain a 10-mL transfer pipet and a 50-mL buret. Obtain about 70 mL of a 0.0250 M solution of Na₂S₂O3 and 3 samples of solid KI. Each of these samples should have a volume of about 1 cm³. They can be stored on pieces of waxed paper. 3. A buret containing the bleach solution will be available for general use. This solution must be diluted, however. Record the initial buret reading to the nearest 0.01 mL. Carefully add about 3 mL of the solution to a 100-mL volumetric flask. (A 100-mL graduated cylinder may be used if a volumetric flask is not available.) Record the final buret reading to the nearest 0.01 mL, and calculate the volume of undiluted bleach used. Add distilled water to the flask until the bottom of the meniscus coincides with the etched line on the flask. Add the last 0.5 mL with a medicine dropper. Insert a stopper in the flask, and mix the solution thoroughly. 4. Be careful in your handling of the solutions used in this experiment. Obtain directions from your laboratory instructor regarding proper disposal of all solutions used during this experiment.
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2.a.  What conversion factor will enable you to calculate the number of moles of ClO– ions from the number of moles of S2O32− ions used in the titration?  

b.  A 10.0-mL sample of aqueous NaOCl is treated with excess KI in an acidic solution. The quantity of iodine that is liberated is such that 28.02 mL of 0.0250 M Na2S2O3 solution must be added to cause the disappearance of the dark blue color due to the starch indicator. What is the molarity of the solution of NaOCl?  

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