Analysis of Commercial Bleach Essay

Custom Student Mr. Teacher ENG 1001-04 10 May 2016

Analysis of Commercial Bleach

Introduction: Many commercial products are effective because they contain oxidizing agents. Some products that contain oxidizing agents are bleaches, hair coloring agents, scouring powders, and toilet bowl cleaners. The most common oxidizing agent in bleaches is sodium hypochlorite, NaClO (sometimes written NaOCl). Commercial bleaches are created by bubbling chlorine gas into a sodium hydroxide solution (remember this from your “funky redox rxns”?). Some of the chlorine is oxidized to the hypochlorite ion, ClO- and some is reduced to the chloride ion, Cl- (a disproportionation reaction). The solution remains strongly basic. The chemical equation for the process is: Cl2(g) + 2OH- (aq) —> ClO- (aq) + Cl- (aq) + H2O (l)

The amount of hypochlorite ion present in a solution of bleach can be determined by oxidation-reduction titration. One of the best methods is the iodine-thiosulfate titration procedure. Iodide ion, I-, is easily oxidized by almost any oxidizing agent (It has many electrons to lose!). In an acid solution, hypochlorite ions oxidize iodide ions to form iodine, I2. The iodine that forms is then titrated with a standard solution of sodium thiosulfate. The analysis takes place in a series of steps:

1.Acidified iodide ion is added to hypochlorite ion solution, and the iodide is oxidized to iodine. 2H+ (aq) + ClO-(aq) + 2I-(aq) —> Cl-(aq) + I2(aq) + H2O (l) 2.Iodine is only slightly soluble in water. It dissolves very well in an aqueous solution of iodide ions, in which it forms a complex ion called the triiodide ion. Triiodide is a combination of a neutral I2 molecule with an I- ion. The triiodide ion is yellow in dilute solution, and dark red-brown when concentrated. I2(aq) + I-(aq) —> I3-(aq)

3.The triiodide is titrated with a standard solution of thiosulfate ions, which reduces the iodine back to iodide ions: I3-(aq) + 2S2O32-(aq)—> 3I-(aq) + S4O62-(aq)
During this last reaction the red-brown color of the triiodide ion fades to yellow and then to the clear color of the iodide ion. It is possible to use the disappearance of the color of the I3- ion as a method of determining the end point, but this is not a very sensitive procedure. Addition of starch to a solution that contains iodine or triiodide ion forms a reversible blue complex. The disappearance of this blue colored complex is a much more sensitive method of determining the end point. However, if the starch is added to the solution that contains a great deal of iodine, the complex that results may not be reversible. Therefore, the starch is not added until shortly before the end point is reached. The quantity of thiosulfate used in step (3) is directly related to the amount of hypochlorite initially present. Materials:

Transfer pipet, 5-mL, and bulb
Buret, Buret stand, and clamp
Small beaker

Volumetric flask, 100 mL w/ stopper
Erlenmeyer flask, 125 mL, or 250 mL
25 mL graduated cylinder

SAFETY ALERT – Concentrated bleach and hydrochloric acid are both damaging to skin, eyes, and clothing and they give off strong vapors. If you spill either solution on yourself, wash off with lots of water. Neutralize hydrochloric acid spills with baking soda.

1. Dilute the concentrated bleach
Use a pipet bulb and a 5-mL transfer pipet to measure out 5.00 mL of a commercial bleach solution into a 100 mL volumetric flask. Dilute to the mark with distilled water, stopper and mix well by inverting repeatedly.

2. Measure the potassium iodide.
Weigh out approximately 2 g solid KI. This is a large excess over that which
is needed.

3. Oxidize the iodide ion with hypochlorite ion.
Carefully measure, using graduated cylinder, 25.00 mL of the dilute bleach into an Erlenmeyer flask. Add the solid KI and about 25 mL of distilled water. Swirl to dissolve the KI. Slowly, with swirling, add approximately 2 mL of 3 M HCL. The solution should turn dark yellow to red-brown from the presence of the I3- complex.

SAFETY ALERT – Adding HCl to bleach may cause chlorine gas to be given off. Avoid smelling. You may be asked to do this step in the hood.

Titrate the iodine
Obtain about 55 mL of sodium thiosulfate solution (beaker). This should be enough for the entire experiment including cleaning. Clean and prepare the buret appropriately (as you’ve done in earlier labs). Record the initial buret reading. Titrate with a standard 0.10 M sodium thiosulfate solution until the iodine color becomes light yellow. Add one dropper of starch solution (Don’t add this at the beginning of the titration). The blue color of the starch-iodine complex should appear. Continue the titration until one drop of Na2S2O3 solution causes the blue color to disappear. Record the final buret reading.

If time allots, repeat the titration beginning with step 2.
The solutions may be safely flushed down the drain with a large excess of water.

Data Table: Design your own! Only pertinent measurements for calculations are necessary. Calculations:
1. Use the equations given to determine the mole ratio between moles of sodium thiosulfate and moles of sodium hypochlorite. SHOW WORK.
2. If you did more than one trial, calculate the average volume of Na2S2O3 needed for the titration of 25.00 mL of diluted bleach.
3. Use the average volume and molarity of Na2S2O3 to determine the molarity of the diluted bleach.
4. Calculate the molarity of the commercial (concentrated) bleach.
5. Assume that the density of the commercial bleach is 1.08 g/mL. Calculate the percent by mass of NaClO in the commercial bleach.

6. You will be given the actual percent by mass NaClO in commercial bleach according to the label. Calculate the percent error value (assuming the label is correct).
Discussion (post lab questions):
1. Write balanced redox half reactions for the equations (1) and (3) from the introduction. For each half reaction, identify what is being reduced and what is being oxidized.
2. The I3- ion is sort of a strange one. What is the oxidation number for each I in this ion? Draw out the Lewis structure for this ion. The I’s are in one chain, I—I—I (not a 3 membered ring!).
3. The reaction with thiosulfate ions produces the dithionate ion, S4O62-. Assuming O produces it’s usual -2 oxidation state, calculate the oxidation state on the sulfur. The fact that the oxidation state comes out a fraction should tell you that the S’s in the ions do not all have the same oxidation state. In fact, the overall oxidation number is an average of the oxidation numbers on the sulfurs that make up the ion. With this in mind, what might be the logical break down of oxidation numbers on the four sulfurs?
4. In this analysis, an “aliquot,” or diluted fraction of the initial solution is used for the titration. What advantage is there in diluting the original solution for this analysis?
5. How would each of the following laboratory mistakes affect the calculated value of the percent NaClO in the commercial bleach (too high, too low, no change). EXPLAIN YOUR REASONING. Be sure to use your calculations as evidence and re-state the error.

a. In step 1, the pipet was rinsed with distilled water immediately before being used to measure the commercial bleach solution.
b. In step 2, 3 g of KI was used instead of 2.
c. In step 3, some of the iodine that formed vaporized from solution.

6. Explain errors that you think might have affected your results (include in error analysis).

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