Name Period Date. Lab 9: Analysis of Commercial Bleach
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1 Name Period Date Lab 9: Analysis of Commercial Bleach Introduction Many common products are effective because they contain oxidizing agents. Some products, which 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 (or NaOCl). Commercial bleaches are made by bubbling chlorine gas into a sodium hydroxide solution. Some of the chlorine is oxidized from the molecular form (Cl 2 ) to the hypochlorite ion, ClO -. Some of the molecular form is also reduced to the chloride ion, Cl -. This type of reaction, where the same type of element is both oxidized and reduced, is called a disproportionation reaction. The solution remains strongly basic. The net ionic chemical equation for the process is: Cl 2 (g) + 2OH - (aq) ClO - (aq) + Cl - (aq) + H 2 O(l) The amount of hypochlorite ion present in a solution of bleach is determined by an 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. In acid solution, hypochlorite ions oxidize the iodide ions to form iodine, I 2. 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) + I 2 (aq) + H 2 O(l) from: HCl bleach KI solution (2) Iodine is only slightly soluble in water. It dissolves very well in an aqueous solution of iodide ion, in which it forms a complex ion called the triiodide ion. Triiodide is a combination of a neutral I 2 molecule with an I - ion. The triiodide ion is yellow in dilute solution, and dark-brown when concentrated. I 2 (aq) + I - (aq) I 3- (aq) iodine iodide triiodide (3) The triiodide is titrated with a standard solution of thiosulfate ions, which reduces the iodine back to iodide ions: 2- I 3- (aq) + 2 (aq) 3I - 2- (aq) + S 4 O 6 (aq) from Na 2 solution called diothionate ion 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 I 3 - ion as the method of determining the end point, but this is not a very sensitive procedure. The addition of starch to a solution that contains iodine or triiodide ion forms 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 a solution, which contains a great deal of iodine, the complex that forms may not be reversible. Therefore, the starch is not added until shortly before the end point is reached, when the solution has faded to a light yellow. The quantity of thiosulfate used in step (3) is directly related to the amount of hypochlorite initially present. 1
2 Materials Bleach, containing NaClO Hydrochloric acid, HCl, 3 M (in the hood) Sodium thiosulfate solution, Na 2, M Volumetric flask, 100-mL with stopper Erlenmeyer flask, 125-mL or 250-mL Potassium iodide, KI (s) Starch solution, 2% Graduated cylinder, 25-mL Transfer pipet, 5-mL Buret Ring stand Buret clamp Pipet bulb Safety Alert Concentrated bleach is damaging to the eyes, skin, and clothing. Hydrochloric acid is also hazardous. Both give off strong vapors. If you spill either solution on yourself, wash off with lots of water. Neutralize hydrochloric acid spills with baking soda Adding hydrochloric acid to bleach may cause chlorine gas to be given off. Carry out step 3 in a fume hood. Always use a pipet bulb when pippetting. Never pipet by mouth. Wear Chemical Splash Goggles and a Chemical-Resistant Apron. Procedure 1. Dilute the concentrated bleach. Use a pipet bulb and a 5-mL transfer pipet to measure 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. 2. Measure the potassium iodide. Weigh out approximately 2 g of solid KI. This amount is a large excess over that which is needed. 3. Oxidize the iodide ion with hypochlorite ion. Using a graduated cylinder, transfer 25 ml of the dilute bleach into an Erlenmeyer flask. Add the KI and about 25 ml of distilled water. Swirl to dissolve the KI. Working in a fume hood and slowly with swirling, add approximately 2 ml of 3 M HCl. The solution should be a dark yellow to redbrown from the presence of I 3 - complex. 4. Titrate the iodine. Rinse the buret with distilled water and then three times with small portions of your 0.10 M sodium thiosulfate solution. Then, fill up the buret with the 0.10 M sodium thiosulfate solution. Titrate the bleach-iodine mixture in your Erlenmeyer flask until the iodine color becomes yellow. Add one dropperful of starch solution. The blue/black color of the starch-iodine complex should appear. Continue the titration until one drop of Na 2 solution causes the blue color to disappear. Record the final buret reading. 5. Repeat Repeat the titration beginning with step 2 two more times. Disposal The solutions may safely be flushed down the drain with a large excess of water. After you have poured out a solution, leave the faucet running. 2
3 Data Buret Reading 1 st Run 2 nd Run 3 rd Run Initial Final Volume Used Calculations Please show all work in the space below. 1. Use the equations given in the Introduction to determine the number of moles of sodium thiosulfate that are equivalent to one mole of sodium hypochlorite. Convert from one mole of sodium hypochlorite to moles of sodium thiosulfate using a series of mole ratios from the equations. 2. Calculate the average volume of Na 2 needed for the titration of ml of diluted bleach. 3. Use the average volume and molarity of Na 2 to determine the number of moles of Na Using the mole ratios, calculate the number of moles of NaClO used in the titrations. 5. Using the number of moles of NaClO and the volume of bleach used, find the molarity of NaClO in the bleach. 3
4 6. Using the molarity of the diluted bleach, calculate the molarity of the commercial bleach. Use the dilution equation. 7. Assume that the density of the commercial bleach is 1.08 g/ml. Calculate the percent by mass of the NaClO in the commercial bleach. 8. The commercial bleach bottle says that it has a mass percent of 6.0%. Calculate the percent error in your experimental value. 9. Calculate the average deviation of the three values you obtained for the volume of Na 2 solution. 4
5 Discussion Please write in complete sentences. 1. Define oxidation and reduction. 2. Write balanced oxidation and reduction half-reactions for equations (1), (2) and (3). (1) (2) (3) 3. For each equation above, identify what substance is actually being oxidized or reduced. (1) (2) (3) 4. In this analysis, an aliquot, or a diluted fraction of the initial solution is used for the titration. What advantage is there in diluting the original solution for the analysis? 5
6 5. The reaction with thiosulfate ions produces the dithionate ion, S 4 O a. Calculate the oxidation number of sulfur in this ion. b. Do you think that the sulfur atoms in the ion will all have the same oxidation number? c. What might the oxidation numbers be? 6. How would each of the following laboratory mistakes affect the calculated value of the molarity of NaClO in the commercial bleach (too high, too low, no change)? Explain each in detail. 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, three grams of KI was used instead of two grams. c. In step 3, some of the iodine that formed, vaporized from the solution. 7. What is the major source of error in this experiment? Explain. 6
7 Name Period Date Lab 9: Analysis of Commercial Bleach Preliminary Laboratory Assignment Pull up our AP Chemistry website. Find the link for Redox Titration Simulation and pull up the file. Follow the instructions below to practice the titration process and calculations that will follow. Show all of your work in an organized way to get full credit. 1. Practice performing a redox titration between KMnO 4 and Fe +2. a. On the left-side of the screen, labeled by #1, choose the appropriate reaction. b. On the right-side of the screen, push the slider bar up until the titration is complete. You should be watching for a color change. If you over-shoot the color change, click the repeat button at the bottom to run the same titration over again. If you think you are close, use the drop-wise button to add only a small amount at a time. c. Once you have finished the titration, fill in the data table below and complete the calculations. Molarity of KMnO 4 Volume of KMnO 4 Volume of Fe +2 i. Using the molarity and volume of KMnO 4, calculate the number of moles of MnO 4 - in the reaction. ii. Using the balanced chemical equation (shown in the middle of the page), calculate the number of moles of Fe +2 in the reaction. iii. Using the moles and volume of Fe +2, calculate the molarity of the Fe +2 solution. iv. Type your calculated molarity in the corresponding box in the page and press OK. Be sure it says correct. If not, look for errors and ask for help if you cannot find any. 2. Follow the same process for the third reaction type- between I 2 and -2. Calculate the molarity of -2. Show your data and all your calculations in the space below: Molarity of I 2 Volume of I 2 Volume of -2 7
8 Name Period Date 1. What is meant by a titration? Lab 8: Analysis of Commercial Bleach Preliminary Laboratory Assignment 2. A solution of household vinegar (a mixture of acetic acid and water) is to be analyzed. A pipet is used to measure out 10.0 ml of the vinegar, which is placed in a 250-mL volumetric flask. Distilled water is added until the total volume of solution is 250 ml. A 25.0-mL portion of the diluted solution is measured out with a pipet and titrated with a standard solution of sodium hydroxide. The neutralization reaction is as follows: HC 2 (aq) + OH - (aq) C 2 - (aq) + H 2 O(l) It is found that 16.7 ml of M NaOH is needed to titrate 25.0 ml of the diluted vinegar. Calculate the molarity of the diluted vinegar. Find the moles of NaOH, then the moles of HC 2, then the molarity of HC Calculate the molarity of the household vinegar. Use the dilution equation. 4. The household vinegar has a density of 1.05 g/ml. Calculate the percent by mass of acetic acid in the household vinegar. Break the molarity (from #3) into moles of HC 2 and liters of solution. Use the moles of HC 2 to find the mass of HC 2, and then find the mass of the total solution using the volume of solution and density. 8
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