Determination of the Sodium Hypochlorite Concentration of Commercial Bleach, Lab Reports of Chemistry

Download Determination of the Sodium Hypochlorite Concentration of Commercial Bleach and more Lab Reports Chemistry in PDF only on Docsity! Skyline College Chemistry 210 Laboratory Manual (Spring 2016 Revision - JM) 1 Experiment 21. Determination of the Sodium Hypochlorite Concentration of Commercial Bleach PURPOSE In this experiment, you will determine the concentration of household bleach by means of an oxidation- reduction titration involving iodine. This is an example of a classic “wet bench” qualitative analysis. BACKGROUND Nature and Production of Bleach. Liquid household bleaches usually contain approximately 3-8% sodium hypochlorite (NaOCl) by weight; 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. Many at laundry novices discover the hard way that bleach splashed on clothes leads to white spots. Liquid laundry bleach is prepared commercially by electrolysis of a cold, stirred solution of sodium chloride. Chlorine gas is produced at the anode: 2 Cl– (aq) ! Cl2(g) + 2e– Oxidation and hydroxide ion is formed at the cathode: 2 H2O(l) + 2e– ! 2 OH– (aq) + H2(g) Reduction In solution, the chlorine gas and hydroxide ion react to give hypochlorite ion: C12(g) + 2 OH– (aq) ! OCl– (aq) + H2O(l) + Cl– (aq) Preparation of a Sodium Thiosulfate Solution. In this experiment, you will prepare 500 mL of 0.1 M sodium thiosulfate solution. You will use this solution for the final titration of I3 – ion. Its concentration should be close to 0.1 M but will not be precisely that concentration. Among the reasons that a sodium thiosulfate solution cannot be prepared accurately by weight is that the solid sodium thiosulfate pentahydrate (Na2S2O3·5H2O) effloresces (forms gaseous products) easily and, as a consequence, its composition is not exact. In addition, fresh solutions of sodium thiosulfate tend to undergo a slight change in concentration because of the presence of sulfur- consuming bacteria in distilled water or plastic storage bottles; this decomposition can be suppressed if the water is boiled before the solution is prepared and if a tiny amount of sodium bicarbonate is added to the solution. Accordingly, you must standardize the sodium thiosulfate solution. We will use primary-standard-grade potassium iodate (KIO3) to standardize the solution of sodium thiosulfate. However, before considering the standardization procedure, you should be aware of some additional facts concerning thiosulfate. First, when a solution of thiosulfate is acidified, the resulting unstable thiosulfuric acid (H2S2O3) slowly decomposes to yield elemental sulfur, sulfur dioxide, and water: H2S2O3(aq) ! S(s) + SO2(g) + H2O(l) Skyline College Chemistry 210 Laboratory Manual (Spring 2016 Revision - JM) 2 In spite of this decomposition, thiosulfate can be used successfully to titrate triiodide in an acid solution if the solution is so thoroughly mixed during a titration that no excess of thiosulfate is present at any time. Second, thiosulfate solutions can be oxidized by dissolved atmospheric oxygen; for this reason, solutions being titrated (and the water used to prepare the sodium thiosulfate solution) are often boiled to remove the dissolved oxygen. Part A. Standardization of a Sodium Thiosulfate Solution. In the procedure for the standardization of sodium thiosulfate, the species actually being titrated is the triiodide ion (I3 -), which is formed stoichiometrically from the reaction between a known (weighed) amount of potassium iodate and an excess of iodide ion: KIO3(s) + 8 KI(s) + 6 HCl(aq) ! 3 I3 – (aq) + 3 H2O(l) + 6 KCl(aq) + 3 K+ (aq) (Reaction 1) The triiodide ion produced in the above reaction is then titrated with the solution of the sodium thiosulfate (the same reaction that will be used to titrate the I3 – formed by the sodium hypochlorite reaction). The I3 - (aq) ion is yellow in dilute solution, and dark red-brown when concentrated. I3 – (aq) + 2 Na2S2O3(aq) ! 3 I– (aq) + Na2S4O6(aq) + 2 Na+ (aq) (Reaction 2) The concentration of the thiosulfate can then be calculated through the stoichiometry of the 2-step process. Potassium iodate is advantageous for this standardization over other standard oxidants (such as potassium dichromate) because it forms a colorless reduction product (I- ion) and because it reacts almost instantly in a solution which is only slightly acidic. But it can decompose to KI and O2, we will want to work with a fresh bottle. This latter fact is significant because iodide ion is rapidly oxidized by atmospheric oxygen in a strongly acidic solution. Part B. Determination of Hypochlorite (OCl-) Concentration. Because hypochlorite is an oxidizing agent, its quantity or concentration in solution can be determined by means of a paired set of oxidation-reduction reactions. First, hypochlorite is reduced quantitatively and rapidly to chloride ion in an acidic aqueous medium by excess iodide ion, which is simultaneously converted to triiodide ion (in the presence of the excess iodide): NaOCl(aq) + 3 KI(s) + 2HCl(aq) ! I3 – (aq) + 3 KCl(aq) + H2O(l) + Na+ (aq) (Reaction 1) Second, the triiodide ion (which is stoichiometrically equivalent to the original amount of hypochlorite) is determined titrimetrically with a standard solution of sodium thiosulfate, the thiosulfate ion being converted to tetrathionate ion: I3 – (aq) + 2 Na2S2O3(aq) ! 3 I– (aq) + Na2S4O6(aq) + 2 Na+ (aq) (Reaction 2) The endpoint of the titration is detected with the aid of starch, which forms an intensely colored purple complex with triiodide ion that disappears at the endpoint. We use this two-step quantitative reaction to obtain a good color change at the endpoint. Otherwise, we could not identify when the reaction was complete. Skyline College Chemistry 210 Laboratory Manual (Spring 2016 Revision - JM) 5 LAB REPORT: Please provide a FULL lab report in the style of other experiments in this course (by this point you shold know how). DATA / CALCULATIONS / RESULTS Provide the data, carry out the calculations, and report the results as described in the procedure. DISCUSSION QUESTIONS 1. Why is NaOCl first reacted with iodide ions to form triiodide ions for titration with S2O3 2- ions, instead of titrating the hypochlorite directly? 2. Write the balanced chemical equations for the 2 reactions in the standardization process. 3. Write the balanced chemical equations for the 2 reactions in the titration of the bleach solution. 4. Calculate relative ranges for your [Na2S2O3] and [NaOCl]. Discuss your precision for each part. 5. Calculate your percent difference for the mass % NaOCl bleach you determined and that reported by the manufacturer. Discuss sources of any discrepancy. 6. Error Analysis Questions for Bleach Titration: a. A student rinsed a pipet with distilled water immediately before being used to measure the commercial bleach solution sample. Would this error affect the percent of NaOCl (too high, too low, or no change)? b. 3 grams of KI were used instead of two grams. Would this error affect the percent of NaOCl (too high, too low, or no change)? c. Discuss any additional sources of error. CONCLUSIONS A. Average molarity of the sodium thiosulfate solution. B. Average molarity of hypochlorite in bleach. C. Mass percentage of NaOCl in the bleach solution. Skyline College Chemistry 210 Laboratory Manual (Spring 2016 Revision - JM) 6 Bleach Titration Experiment PRELAB QUESTIONS Name: 1. Calculate the mass of sodium thiosulfate pentahydrate (Na2S2O3 • 5 H2O) needed to prepare 500 mL of an approximately 0.10 M solution. Answer this question in the PRELAB in your lab notebook. 2. When should the starch indicator be added in the titration and why is it added then? 3. A 3.00-mL bleach (NaOCl is the active ingredient) sample is mixed with potassium iodide and acid to completely form triiodide (I3 -) ions. You then titrate the I3 − ions with a thiosulfate (S2O3 2−) solution you have previously determined to be 0.1306 M. The endpoint is reached when 31.52 mL of the thiosulfate solution is added. What is the molar concentration of the bleach solution? What is its mass percent? Note: This question is very similar to part B of the lab. See previous page for the densities of commercial bleach.