You’ve heard the claims on commercials “Our product makes clothes whiter and brighter than the other brands…” So, whom are you going to believe?  Well, the good news is that you don’t have to take anyone’s word for the oxidizing power of your bleach.  Over the next two lab periods, you will be investigating that for yourself.  According to industry standards, the oxidizing power of bleach products is reported in terms of “available chlorine.”  This is a convenient term, but it does not reflect the true chemical form of the oxidizing agent.  The reason for this is that the active ingredient in most cheap household bleaches is sodium hypochlorite (NaClO) not chlorine (Cl2).  

          Another common type of bleach, the solid,  non-chlorine, color-safe variety contains reagents such as sodium perborate (NaBO2·H2O2·H2O) which produce hydrogen peroxide when dissolved in water.  For this reason, these products are often referred to as oxygen bleaches.  Clorox 2 is a common example of oxygen bleach.

            Both types of bleach work by oxidation and chlorination, which means they accept electrons from other compounds. (If you need to review this topic, read section 4.4 in Chang.)  Colored compounds, including organic compounds that we use for clothing dyes and those that cause stains on clothing, usually contain double and triple bonds.  These bonds absorb light of a particular frequency, which causes them to appear as a particular color to our eyes. Bleaches reduce these multiple bonds to single bonds.  When this occurs, light will be reflected by the material rather than absorbed.  The reflection of all wavelengths is seen as white by the human eye, so our clothes appear whiter and brighter.

    Interestingly, one of the best brightening and “whitening” agents is “laundry bluing.”  This material is a concentrated intensely blue liquid that looks like it would stain anything a pernicious shade of blue!  So, why doesn’t this blue liquid stain your clothes?  Why does it make your white clothes look even whiter?  Does it change into a white substance when it contacts fabric?  No! “Laundry bluing” actually changes your whole shirt slightly blue !   It turns out that the human eye sees a bluish-white tint as whiter than pure white!  “Laundry Bluing” is not a popular component of the average laundry process anymore.  But, it is a fascinating window into how what we perceive visually is not always straight forward.  For a more complete explanation into this subject, visit the following web site.

For more information about how bleach works, try these interesting web sites.

            Bleach has other uses, many of which have only been recently investigated.  Studies have shown that solutions of sodium hypochlorite can inactivate HIV as well as destroy the protein coats of some common allergens, rendering them incapable of causing an allergic reaction.  For more information, visit the following web sites.

Bleach inactivation of HIV:

Allergens and bleach:

Chlorine chemistry plays an integral part in many products you use every day.  Among these are; table salt, vinyl (PVC), drugs, water purification, refrigerants, fire extinguishers, disinfectants, self-deodorizing sock….  The list is extensive and the history is fascinating.  You can visit the following site for a quick overview and timeline.

Chlorine chemistry:



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.

DON'T MIX Chlorine Bleach with:

 Toilet Bowl Cleaners 



Rust Removers


Oven Cleaners

Don't 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 eye wash 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: Dispose of any extra KIO3 solution in the waste oxidants container.  All test solutions may be disposed of in the sink with plenty of water.


In this week’s lab, you will be standardizing a solution of sodium thiosulfate (Na2S2O3).   Standardization is a process in which a solution, whose concentration is approximately known, is titrated with a solution whose concentration is accurately known.  You will prepare an accurately known solution of a primary standard, potassium iodate (KIO3).   This compound has the characteristics required of any compound used as a primary standard; high purity, stability when exposed to the air, long shelf life, high molar mass (which helps reduce errors in massing), and a predictable and well-defined reaction with the solution being standardized.  The KIO3 you will use has been dried overnight in a 110°C oven. 

You were previously introduced to this type of titration in Chem 184, when you investigated the concentration of acetic acid in vinegar.  Essentially, you were standardizing a weak acid, accurately determining its concentration, by using a NaOH solution whose molarity was accurately known.  (It had been standardized prior to your lab with potassium hydrogen phthalate (KHP), another compound often used as a primary standard.)

            Your written prelab assignment is to answer the questions below and be prepared to discuss them at the beginning of your lab period. 

a) How would you prepare a 100 mL primary standard solution of 0.01 M KIO3? Write out all the steps involved and be sure to show your calculations. (Remember, this solution’s concentration must be accurately known, so the container used in preparing it should be specified in your answer.)

 b) Why do you think the KIO3 will be dried overnight in the oven before you use it to make your solution?

Laboratory Part I:

As a team, take a few minutes to compare your prelab answers and come to a consensus about which ones are the most correct.  Choose one member from your team to write these answers on the board.  Discuss these answers as a class before proceeding with your individual team investigations.  What is the class consensus regarding the proper method for preparing a primary standard solution? Why is it important to accurately know the concentration of the KIO3?

            Recall that in the vinegar titrations you performed last semester, phenolphthalein was used as an indicator to visualize the endpoint.  Your solution turned pink when the moles of OH- ions from the NaOH were equal to the moles of protons from the acetic acid, because that reaction had the 1:1 stoichiometry shown below.

NaOH + HCl ® NaCl + H2O

During the next two lab periods, you will take advantage of the fact that aqueous I2 produced in your reactions is colored.  That species will be used as an internal indicator.  Where will the I2 come from?  The reaction being used in these investigations is designed to produce it! Iodide can be oxidized to iodine in an acidic solution.  In this lab, the oxidizing agent will be the potassium iodate (KIO3) you are using to standardize a sodium thiosulfate (Na2S2O3) solution.  Next week, the oxidizing agent will be the sodium hypochlorite (NaClO) in household bleach.   The oxidation of iodide to iodine is shown below.

                                IO3- (aq) + 5 I-(aq)  + 6  H+(aq) ® 3 I2 (aq, yellow)  +  3 H2O  

The I2 can now be reduced back to I- by titrating it with the prepared Na2S2O3 from the buret.  When all of the I2 has been reduced (at the stoichiometric point), the solution will be colorless.  An unbalanced equation for this reaction appears below.

   ? I2(aq, brown/yellow) + ? S2O32-(aq)  ®  ? I-(aq, colorless) + ? S4O62-(aq) 

Each team will obtain 25 mL of the prepared KIO3 in a 100 mL beaker.   Use the chart below to determine how much of the prepared Na2S2O3 your group should add to the KIO3 solution.  Each group should mix these solutions well, then add 10 mL of the prepared 0.5 M H2SO4 and about 2 grams KI to the solution in the beaker.  Report your observations on the blackboard.  As a class, discuss what these results mean.


Team 1

Team 2

Team 3

Team 4

Team 5

KIO3 Solution

25 mL

25 mL

25 mL

25 mL

25 mL

NaS2O3 Solution

5 mL

10 mL

15 mL

20 mL

25 mL

How many moles of each reagent did you use? What is the stoichiometry of this reaction?   How will knowing this relationship help you calculate the molarity of the NaS2O3.

As a team, read over the procedure for Laboratory Part II, the standardization titration.  Before proceeding, plan your investigation.  This should include a discussion of the equipment you will need, calculations involved in preparing your primary standard, calculations for preparing the Na2S2O3 solution to be standardized, and also the proper method for determining the accurate molarity of your Na2S2O3, using your titration data. Consult with your GTA to see if your plan seems reasonable.  Record your plan in your lab notebooks along with any data (mass of reagents used, volume of reagents added, known molarities of reagents, etc.) that you are given or that you collect.  Guidelines for keeping your laboratory notebook can be obtained by clicking here or by following the link on the Chem 188/189 home page.

Laboratory Part II:

Carry out your group’s plan to standardize the Na2S2O3, which you formulated from the written procedural outline.  Be careful to record all the pertinent data and observations in your lab notebook along with all calculations for molarity.  Your notebook should also indicate the tasks for which you were personally responsible.  (For additional pointers on keeping a laboratory notebook, click here.)  Finally, write your report.  The questions in italics are points to consider as you write up your results and conclusions.  Remember to include the relevant chemical equations in both your notebook and your report.

Procedure for Laboratory Part II: Standardization of Na2S2O3

1.  Prepare 100 mL of 0.01 M KIO3 primary standard according to the procedure that your group has agreed upon.

2.  Using an Erlenmeyer flask, prepare 250 mL of a Na2S2O3 solution that is approximately 0.1  M. Why is it not necessary to use a volumetric flask for this solution?

3.   Fill a clean buret with your Na2S2O3 solution and record the volume to ± 0.02 mL. Why should the volume be recorded to  ± 0.02 mL  ?

4.  Pipet 25 mL of your primary standard solution into a 125 mL Erlenmeyer and to this add about 2 g of solid KI. Dissolve this by swirling, and then add about 10 mL of 0.5 molar H2SO4.  Why is a volumetric pipet used to measure out the KIO3, but approximate amounts are sufficient for the KI and acid?

5.  Titrate the KIO3 with the Na2S2O3 solution until it becomes colorless. Be sure to stir the reaction or swirl continuously.

6.  Repeat the titration two more times and use your three calculated molarities to get an average molarity.  This molarity should have four significant figures.  Why?

7.  Store your solution in a clean, completely dry, stoppered volumetric flask.  You will use this standardized Na2S2O3 to determine the “available chlorine”  (the industrial expression of oxidizing power) in commercial bleach during your next lab period.  Label your container.