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Analysis of bleach by redox titration

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					Redox titration lab      v2’’     Dr. Breinan          Chemistry                                 p. 1
                                        Analysis of bleach by redox titration
Introduction: Bleaches are one of many products that contain chemicals that work as oxidizing agents (others
  include hair coloring, scouring powders, and toilet bowl cleaners). The active ingredient in laundry bleach is
  sodium hypochlorite, (specifically the hypochlorite ion, ClO¯). This ion is readily reduced to the chloride ion
  under many conditions. The amount of sodium hypochlorite in bleach can be determined by a series of redox
  reactions involving forms of iodine and culminating with titration by a solution containing the thiosulfate ion,
  S2O32-.. These unbalanced reactions are described here:

  (1) A large excess amount of the iodide ion is added to the bleach and the solution is acidified. The following
    reaction occurs.
                           H+ (aq) + ClO¯(aq) + I¯ (aq) ---> Cl¯ (aq) + I2 (aq) + H2O (l)

  (2) The excess iodide can form a polyatomic ion (called “triiodide”) with the iodine formed in reaction (1).
    This helps the iodine dissolve and gives the solution a color that varies from yellow to dark red-brown,
    depending on how concentrated it is. The formation of triiodide is shown here:
                                              I2 (aq) + I¯ (aq) --> I3¯ (aq)

  (3) Finally, the triiodide ion undergoes titration by a solution of thiosulfate ions of known molarity. This
    redox reaction reforms iodide and creates dithionate ions....
                                    I3¯ (aq) + S2O32¯(aq) ---> I¯ (aq) + S4O62¯

   The color of the triiodide ion can serve as its own indicator. When it is all reduced by the thiosulfate, it will
appear colorless again. However the color change can be made much more striking by use of a starch indicator.
Starch forms a dark blue-black complex with either iodine or triiodide. The disappearance of the triiodide by
titration will be signaled by the loss of color from this complex. Care must be taken in this procedure
however... starch added to a high concentration of iodine may lead to an irreversible complex... one in which a
color change would not occur. For this reason, the starch indicator in this lab will not be added until shortly
before the end point is reached. Under these conditions, the starch complex will disappear as the triiodide is
removed by titration with thiosulfate.

CAUTIONS:
- The reaction of bleach with hydrochloric acid can release poisonous chlorine gas. Perform this part of the
  reaction in a fume hood with the exhaust fan running.

Objectives:
- To determine the percentage of sodium hypochlorite in bleach by a redox titration analysis.
- To practice proper technique for titrating with burets

Pre-lab:   complete on a separate sheet.       Review lab procedures for burets and valves.
1. For reaction (1) in the introduction:
 a) identify the species oxidized and reduced; show the initial and final values of changing oxidation #’s
 b) write balanced half-reactions for the oxidation and reduction that are occurring
 c) use the half-reactions to balance the net ionic equations given (I know you can do these by inspection, but I
  want you to practice the method even if it is longer)
 d) repeat steps a) - c) for reactions (2) and (3).
 e) which reaction might be called the reverse of a disproportionation? Explain.
2. According to the introduction and instructions for the lab, what is the chemical of known molarity in this lab?
  Of unknown molarity?
3. According to the equations given in the introduction, determine the mole ratio that relates the known and
  unknown chemicals (you must use all three equations and two other ratios).
4. What is the biggest danger in this lab? What safety precautions will you take?
5. Prepare a data table for this lab.
Redox titration lab     v2’’     Dr. Breinan          Chemistry                                 p. 2
6. (a) Outline briefly the steps you must follow to use a buret in a titration (see titration sheet).
(b) Outline briefly the steps you must follow to use a pipet with an ASE bulb (see lab procedures).

7. Household vinegar is a solution of acetic acid. A solution of diluted vinegar is made by taking 10.0 mL of
  household vinegar and diluting it to 250.0 mL in a volumetric flask. The diluted vinegar is titrated with
  0.0500 M sodium hydroxide. It takes 16.7 mL of the sodium hydroxide solution to neutralize 25.0 mL of the
  diluted vinegar.
 a) Write the Bronsted-Lowry reaction that takes place.
 b) Find the molarity of acetic acid in household vinegar. (Hint: find the molarity of the diluted vinegar first,
   then take into account the dilution)


Procedure:
A. Making the unknown solution to be titrated against:
A1 Your bleach should be diluted by about a factor of 20 before you use it. You may be given a diluted
  sample with the contents specified. Alternatively, you may mix your own. Using a buret or pipet, add
  about 5 mL (record the amount exactly) of pure bleach to a 100 mL volumetric flask and dilute with
  distilled water. In either case you will need approximately 65 mL out of the total to clean and fill your buret.
 Check the bleach bottle for the known percent sodium hypochlorite to use in calculations.

B. Preparing your known titrating solution.
B1 Your teacher will provide you with a standard solution of 0.10 M sodium thiosulfate (Na2S2O3). Take
   about 60-65 mL in a small beaker to rinse and prepare your buret.

C. The titration procedure. (you should complete 3 trials from your filled burets)
 C1. Make your initial buret readings
 C2. Obtain about 0.8 to 1 g of potassium iodide in the beaker or flask you will use for the titration (this
   amount is a large excess... do not record the mass).
 C3. Release about 15 mL of the diluted bleach solution into your titration beaker or flask. Mix to dissolve the
   KI. You will see a yellowish color as the KI dissolves.
 C4. CAUTION: HCl is corrosive. Also, poisonous chlorine gas may be released in this step! Working in the
   fume hood with the hood on, add about 1 mL (about 1 dropperful) of 3M HCl to the flask. The deep red
   color of the I3¯ ion will appear.
 C5. For this system, you must start the titration before adding the indicator. Begin the titration by releasing
   some of the sodium thiosulfate solution into the titration flask. This can be done fairly rapidly at first,
   while swirling the flask. Temporarily stop when the dark red has turned to a light orange-yellow (the exact
   color is not important)
 C6. Add one dropperful of starch indicator solution. The dark color of the iodine-starch complex should
   appear
 C7. Carefully resume titrating to the end point with the thiosulfate solution. If you go past the endpoint, you
   may back-titrate. Record the final volumes.
 C8. Repeat steps C1 to C7 twice more with the solutions left in your buret. Between each trial do the
   following:
   - Do NOT alter your burets... you should be able to use them three times each
   - Dump the titrated (neutralized) solution from the last trial down the drain and rinse your titration flask ...
      it need not be dry

D. Disposal and cleanup-
 D1. Chemicals: place each excess chemical in the appropriate waste beaker.
 D2. Burets: If a buret brush is available, brush the burets first. Rinse several times with water.
 D3. After you have cleaned up, determine the volumes of each solution used in each trial by subtraction. Give
    a neat copy of this data and the volume of undiluted bleach you used to your teacher.
Redox titration lab    v2’’    Dr. Breinan        Chemistry                              p. 3


Processing:
1. Present your data and use it to....
 a) calculate the molarity of hypochlorite in household bleach
 b) find the percent by mass of sodium hypochlorite if the density of bleach is 1.08 g/mL (this
    corresponds to the number often printed on the label of bleach)
 c) find your percent error in the mass % (ask your instructor if it wasn’t available during the lab)
2. Assume that you did the calculations exactly as above, but there were some unknown mistakes in
  your procedure. What would the following mistakes mean (would the result you calculated above be
  too high, too low, or just right)? Explain.
  a) The pipet or buret used to measure the bleach (step A1) was rinsed with distilled water just before
    being used to measure the commercial bleach
  b) you accidentally dumped in 2 g of KI (step C2)
  c) Some of the iodine that formed during the titration vaporized from solution

3. The production of bleach involves a “disproportionation” reaction. This means that the same
  chemical (chlorine in this case) is both oxidized and reduced (in this case in a basic solution). The
  products are the hypochlorite ion and chloride ions. Write half reactions for the oxidation and
  reduction that are occurring and use these to write a balanced net ionic equation for the formation of
  bleach.

4. What advantage was there to diluting the bleach in this experiment?

				
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