Hello, below is a chemistry assignment. Instead of givingyou my data, then you in turn show me the equation and compute thesolutions. I would like if you would instead show me the equationsneeded to perform calculations and the thought process forquestions. I will input my own data into provided equations. Ithink, that maybe, by me inputting the data thatthis will be a better study approach for myself.As always, thanks for the helpAnalysis of Bleach, Sodium Hypochlorite Commercial bleaches are
Hello, below is a chemistry assignment. Instead of giving you my data, then you in turn show me the equation and compute the solutions. I would like if you would instead show me the equations needed to perform calculations and the thought process for questions. I will input my own data into provided equations. I think, that maybe, by me inputting the data that this will be a better study approach for myself. As always, thanks for the help Analysis of Bleach, Sodium Hypochlorite Commercial bleaches are made by the electrolysis of aqueous sodium chloride solutions. In the process sodium hydroxide, dichlorine gas, and dihydrogen gas are formed. If the sodium hydroxide and the dichlorine are allowed to mix, sodium hypochlorite is generated. Some of the dichlorine is oxidized to the hypochlorite ion present in the solution, while some is reduced to the chloride ion in a reaction called a disproportionation reaction. Liquid household bleaches usually contain approximately 3-8% sodium hypochlorite (NaOCl) by mass; it is the oxidizing power of the hypochlorite ion that is responsible for the beneficial action of bleach. Hypochlorite performs its bleaching function by oxidizing stains (or dyes) to produce colorless, soluble, or gaseous species. As mentioned, liquid laundry bleach is prepared commercially by electrolysis of a cold, stirred solution of sodium chloride. Dichlorine gas is produced at the anode: 2Cl-(aq) Ã Cl2(g) + 2eâ€“ (Oxidation)(1) and hydroxide ion is formed at the cathode: 2 H2O(l) + 2eâ€“ Ã 2 OHâ€“(aq) + H2(g) (Reduction)(2) The overall process is: Cl2(g) + 2NaOH(aq) Ã NaOCl(aq) + NaCl(aq) + H2O(l)(3) The amount of hypochlorite ion present in a solution of bleach is determined by a redox titration. In this experiment, the titration involves iodide and thiosulfate ions. Iodide ion, I-, can be oxidized by almost any oxidizing agent. The three reactions in the analysis are: NaOCl(aq) + 2KI(aq) + 2HCl(aq) Ã I2(aq) + NaCl(aq) + 2KCl(aq) + H2O(l)(4) I2(aq) + KI(aq) ÃŸ Ã KI3 (aq) (5) 2Na2S2O3(aq) + I2(aq) Ã Na2S4O6(aq) + 2NaI(aq)(6) The bleach is a good oxidizing agent and will oxidize iodide ion in an acidic solution as shown in Equation 4. The resulting diiodine can be titrated with thiosulfate. An excess of KI is added to the titration mixture for several reasons. First, the actual amount of hypochlorite in solution is unknown so enough KI must be added to react with all the bleach, which is being analyzed. Second, since diiodine is somewhat volatile, it will be stabilized by the presence of excess iodide due to the formation of the triiodide ion (I3-). The presence of the triiodide ion gives a dark orange to reddish brown color to the solution. This is shown in Equation 5. Once the bleach has been converted to the triiodide, the free diiodine in solution from Equation 5 can be titrated with standardized thiosulfate according to Equation 6. As the diiodine is consumed, the equilibrium in Equation 5 shifts to produce more I2 until all of the diiodine is used. The reaction that occurs during the titration is given by Equation 6. The titration of highly acidic solutions of diiodine with thiosulfate yields quantitative results provided air oxidation of iodide is minimized and the thiosulfate is added slowly to prevent its decomposition. Once the iodide is oxidized to diiodine, it must be titrated immediately with the thiosulfate as I2 is volatile. The end point in the titration is readily determined by means of a starch solution. Starch is used as an indicator because it reacts with I2 to form a dark color. The dark color will fade during the titration as I2 is consumed. The end point occurs when one drop of the Na2S2O3 solution causes the disappearance of the last trace of I2 and the solution changes from dark to colorless. Since starch is partially decomposed in the presence of a large excess of diiodine, the indicator is never added to a diiodine solution until the bulk of that substance has been reduced. The change in color of the diiodine solution from red to a faint yellow signals the proper time for the addition of the starch indicator. In order to determine the bleach concentration, the concentration of the thiosulfate must be known and unfortunately, it is not a good primary standard for accurate work because the solid cannot be dried without decomposing it. Therefore, the thiosulfate must be standardized before use. Several excellent primary standards are available for the standardization of thiosulfate solutions. In general, these are oxidizing agents that liberate an equivalent amount of diiodine when treated with an excess of iodide ion. The resulting solution containing the diiodine is titrated with the thiosulfate. Two commonly used primary standards are potassium dichromate and potassium iodate. Since the dichromate ion is carcinogenic, potassium iodate will be used. The reactions in the standardization are the same reactions occurring in the analysis of bleach. A known amount of potassium iodate is measured and reacted with excess iodide in acid to produce diiodine (Equation 7). The diiodine is titrated with the thiosulfate (Equation 6) and its concentration is calculated from stoichiometry. KIO3(aq) + 5KI(aq) + 6HCl(aq) Ã 3I2(aq) + 3H2O(l)(7) Procedure: Iodometric Method Standardization of Sodium Thiosulfate Obtain 100 mL of the sodium thiosulfate for standardization. Setup a burette using a ring stand and burette clamp. Drain the burette and add some distilled water to check for leaks. Add about 5 mL of the thiosulfate solution to the burette and wash the walls and drain the solution through the tip. Discard. Repeat two more times. Fill the burette with the thiosulfate solution including the tip (make sure there are no air bubbles in the tip) and record the initial volume to 0.01 mL. Record the mass of a clean 250 mL Erlenmeyer flask. The flask may be wet on the inside but the outside of flask must be dry. Add the required amount of the standard KIO3 solution (see instructor) to the flask and record the mass of the flask and solution. Record the mass percent concentration of the KIO3 solution. Add 25 mL of distilled water and 1.0 g KI. Once the KI dissolves, add 10 mL of 1.0 M HCl and titrate immediately with the thiosulfate solution until the color of the solution becomes pale yellow. At this point add 5 mL of the starch indicator and titrate to the disappearance of the blue color. Record the final volume of thiosulfate solution to 0.01 mL. Repeat the titration two more times. Note: The burette should never be refilled during a titration as this increases the error in volume. If the burette does not contain enough solution to complete a titration, refill before beginning. Analysis of Bleach Obtain 40 mL of the commercial bleach, a 250 mL volumetric flask, and a 20 mL pipette. Pipette 20 mL of the bleach and transfer it to the volumetric flask. Fill the flask with distilled water to the mark. Cover the flask and mix the solution well by inverting and shaking 15 times. Obtain a second burette and prepare it with the diluted bleach solution. After rinsing it 3 times, fill the burette and tip with the diluted bleach solution. Record the initial volume to 0.01 mL. Add about 25 mL of the bleach solution in the burette to a 250 mL Erlenmeyer flask. Record the final volume of bleach to 0.01 mL. Record the initial volume of thiosulfate to 0.01 mL. Add 25 mL of distilled water and 1.0 g of solid KI. Once the KI dissolves, add 10 mL of 1.0 M HCl and titrate immediately with the standardized Na2S2O3 solution until the color of the solution becomes pale yellow. Add 5 mL of starch and titrate to the disappearance of the blue color. Record the volume of the thiosulfate to 0.01 mL. Repeat the analysis two more times. Discard all solutions as directed by the instructor. Rinse the burettes with distilled water, multiple times, and refill with distilled water. Store the burette completely filled with distilled water and stoppered. Page BreakAnalysis of Bleach, Sodium Hypochlorite Name Partnerâ€™s Name Data Standardization of Na2S2O3 Solution Mass Percent of Potassium iodate Sample 1Sample 2Sample 3 Mass of Erlenmeyer Flask Mass of Erlenmeyer Flask and KIO3 solution Initial Volume of Na2S2O3 Final Volume of Na2S2O3 Calculations (Attach a separate page showing the calculations) Mass of KIO3 solution Mass of pure KIO3 used Moles of pure KIO3 used Moles of I2 produced Moles of Na2S2O3 present Volume of Na2S2O3 used Molarity of Na2S2O3 Average Molarity of Standardized Na2S2O3 Solution Data Analysis of Commercial Bleach Density of Commercial Bleach Volume of Commercial Bleach used Diluted Volume Sample 1Sample 2Sample 3 Initial Volume of Diluted NaOCl Final Volume of Diluted NaOCl Initial Volume of Na2S2O3 Final Volume of Na2S2O3 Calculations (Attach calculation sheet) Volume of Na2S2O3 used Moles of Na2S2O3 present Moles of I2 present Moles of NaOCl present Volume of NaOCl used Molarity of NaOCl Average Molarity of Diluted NaOCl Solution Molarity of Commercial Bleach Page BreakQuestions 1.What were your results? What did you learn in this experiment? 2.Using the density of commercial bleach and your molarity for the commercial bleach, calculate the mass percent NaOCl for the commercial bleach? Obtain the reported mass percent for the bleach used from the instructor and determine the percent error using the reported mass percent as the accepted value. 3.Why is it necessary to titrate the solution immediately following the addition of KI? Explain. 4.Why is it important to use excess KI? Explain.