Experiment 5 - Oxidation-Reduction Titration Iodimetry

Experiment 5: Oxidation Reduction Titration Iodimetry dela Cruz, Marie Giecel V. Pabilane, Alma L. Group 6, Chem 27.1, SEJ1, Ms. Noime Walican January 29, 2011 I. Abstract Iodimetry is an oxidation-reduction titration which uses iodine, an oxidizing agent that reacts rapidly with a variety of strong reductants, as titrant, with starch as the indicator. On the other hand, iodometry is a redox titration wherein I2 is being produced in the system. Iodometric methods are used in the determination of many organic and inorganic substances, with Na2S2O3 as titrant (commonly used). The concentrations may be expressed as %(w/w) of the ion or metal desired. This amount of I2 present can be calculated by the stoichiometry of the titrant used with I2 according to the balanced redox reaction. Then, the amount of the desired ion or metal will be computed using the ratio of this ion with I2 according to the balanced reaction. In this experiment, iodometric method is done to determine the bleaching power (which is the concentration of Cl-) of a bleach. The experimental result for the concentration of Cl- in the sample bleach was 1.788x10-4 g/ml. II. Keywords: iodometry, iodimetry, redox titrations, oxidant, reductant III. Objectives In this experiment, the students must be able to: (1) Prepare a standard solution of Na2S2O3 and; (2) Determine the strength of bleaching agent by oxidation – reduction titration. IV. Introduction and Principles Oxidation – reduction titrations involves the titration between an oxidizing agent and a reducing agent. Iodimetry is a redox titration that uses iodine as the titrant. Although iodine’s oxidizing power is less than that of ceric, permanganate and dichromate ions, it is widely used as redox titrant because of its ability to react rapidly with many strong reductants, and the availability of a good indicator for iodine which is starch. Starch forms a deep-blue complex with iodine, but not with iodide. The half reaction for the reduction of iodine is given by: I2(s) + 2e-  2IIodine also has some disadvantages. It is not very soluble in water but dissolves readily in an iodide solution forming triiodide ion. Also, iodine solutions are not very stable. It can be oxidized by oxygen in acidic medium. Iodometry, which is the one performed in this experiment, uses iodide ion, not as titrant, but to produce iodine in the system. Then, the iodine is consumed in the titration. The reagent most often used as titrant for iodine is sodium thiosulfate (Na2S2O3). Its reaction with iodine is given by: I2 + S2O32-  2I- + S4O62Thiosulfate ion is a moderately strong reductant. It is readily soluble in water and is stable in Chem 27.1. Oxidation – Reduction Titration Iodimetry

both neutral and basic conditions. It is one of the few reducing agents that are not oxidized by air. It has been widely used to determine oxidizing agents by an indirect procedure, involving iodine as intermediate. The scheme used to determine the oxidizing agents involves addition of excess KI to slightly acidic analyte. The reduction of the analyte produces a stoichiometrically equivalent amount of iodine, and then the produced iodine is titrated with standard solution of sodium thiosulfate as indicated in the reaction above. Redox titration curve is a plot of electrode potential vs. volume of titrant used. The electrode potential is a log function of the concentration of titration of reactants and products and it refers to the potential of the electrode immersed in the solution relative to standard hydrogen electrode. V. Methodology and Materials A. Preparation and Standardization of 0.1M Na2S2O3 solution The weight of Na2S2O3(H2O)5 crystals needed to prepare 500ml of 0.1M Na2S2O3 solution was calculated. This amount of crystal was dissolved in 100ml previously boiled distilled water in a beaker. It was diluted to 500ml and 0.20g sodium carbonate was added as preservative. The solution was stored in a clean reagent bottle. In each of the three 500-ml Erlenmeyer flasks, 0.10-0.15 g of primary standard grade potassium dichromate was placed. Each sample was dissolved in 50ml previously boiled distilled water and 4ml of 1:2 sulfuric acid was added. Five grams of KI dissolved in 5ml distilled water was added and the flask was swirled. The flask was covered with watch glass and allowed to stand for 3 minutes to allow the reaction to complete. Page 1 of 3

The solution was diluted with 50ml doubledistilled water and titrated with thiosulfate solution until the brown color of iodine had almost disappeared. Five milliliters of starch solution (prepared in procedure B) was added and titration was continued until the last drop of titrant removed the blue color of starch iodine complex giving a clear emerald-green solution. Three trials were done and the molarity of sodium thiosulfate was calculated in each. The average deviation should be 1-3ppt. B. Preparation of Starch Indicator In 5ml double-distilled water, 0.5 g starch was dissolved. This was added to 100ml boiling water and was boiled for another 2 minutes.

Computations: 1: [(0.1095 g K2Cr2O7 / 294.19 g/mol)(mol Cr2O72- / mol K2Cr2O7)(3 mol I2 / mol Cr2O72-)(2 mol S2O32- / 1 mol I2)] / (0.01383 L) = 0.1615 M 2: [(0.1024 g K2Cr2O7 / 294.19 g/mol)(mol Cr2O72- / mol K2Cr2O7)(3 mol I2 / mol Cr2O72-)(2 mol S2O32- / 1 mol I2)] / (0.01348 L) = 0.1549 M 3: [(0.1025 g K2Cr2O7 / 294.19 g/mol)(mol Cr2O72- / mol K2Cr2O7)(3 mol I2 / mol Cr2O72-)(2 mol S2O32- / 1 mol I2)] / (0.01327 L) = 0.1575 M B. Determination of Bleaching Power Trial

V, Na2S2O3 (mL) 15.75 16.3 16.15 15.81 15.8

% Cl- (g/mL)

C. Analysis of the Unknown Fifty milliliters of liquid bleach was placed in a 250-ml volumetric flask and was diluted to mark. Ten milliliters of this was put in each of the five 250-ml Erlenmeyer flasks. Then, 50ml of distilled water was added to it. Three grams KI, 8ml of 1:6 sulfuric acid and 3 drops of 3% ammonium molybdate (catalyst) were added. The flasks were covered with watch glass and let stand for 3 minutes to allow reaction to complete. The liberated iodine was titrated with standardized thiosulfate solution until brown color of iodine has almost disappeared. Five milliliters of starch indicator was added and titration was continued until blue color disappeared. The % Cl in each trial was calculate, assuming the density of liquid bleaches is 1.0 g/mol.

Ave = 1.788 x 10-4 g/mL

VI. Results

2: n, S2O32- = (0.1580 M)(0.01630 L) = 2.5754 x 10-3 mol

Reactions Involved: Standardization 14H+(aq) + Cr2O72-(aq) + 6I-(aq) → 2Cr3+(aq) + 3I2(aq) + 7H2O(l) 2S2O32-(aq) + I2(aq) → 2I-(aq) + S4O62-(aq) Bleaching power 2S2O32-(aq) + I2(aq) → 2I-(aq) + S4O62-(aq) 2H+(aq) + ClO-(aq) + 2I-(aq) → I2(aq) + Cl-(aq) + H2O(l)

A. Trial 1 2 3

Standardization of Na2S2O3 W, K2Cr2O7 (g) 0.1095 0.1024 0.1025

V, Na2S2O3 (mL) 13.83 13.48 13.27

M, Na2S2O3 (M) 0.1615 0.1549 0.1575

Ave = 0.1580 M Chem 27.1. Oxidation – Reduction Titration Iodimetry

1 2 3 4 5

1.764x10-4 1.826x10-4 1.809x10-4 1.771x10-4 1.770x10-4

Computations: n, S2O321: n, S2O32- = (0.1580 M)(0.01575 L) = 2.4885 x 10-3 mol

3: n, S2O32- = (0.1580 M)(0.01615 L) = 2.5517 x 10-3 mol 4: n, S2O32- = (0.1580 M)(0.01581 L) = 2.49798 x 10-3 mol 5: n, S2O32- = (0.1580 M)(0.01580 L) = 2.4964 x 10-3 mol M, ClO1: M ClO- = [(2.4885 x 10-3 mol S2O32-)(mol I2 / mol S2O32-)(mol ClO- / mol I2)] / [(0.010 L)(250 mL / 10mL)] = 4.977 x 10-3 M 2: M ClO- = [(2.5754 x 10-3 mol S2O32-)(mol I2 / mol S2O32-)(mol ClO- / mol I2)] / [(0.010 L)(250 mL / 10mL)] = 5.1508 x 10-3 M 3: M ClO- = [(2.5517 x 10-3 mol S2O32-)(mol I2 / mol S2O32-)(mol ClO- / mol I2)] / [(0.010 L)(250 mL / 10mL)] = 5.1034 x 10-3 M Page 2 of 3

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4: M ClO = [(2.49798 x 10 mol S2O )(mol I2 / mol S2O32-)(mol ClO- / mol I2)] / [(0.010 L)(250 mL / 10mL)] = 4.99596 x 10-3 M 5: M ClO- = [(2.4964 x 10-3 mol S2O32-)(mol I2 / mol S2O32-)(mol ClO- / mol I2)] / [(0.010 L)(250 mL / 10mL)] = 4.9928 x 10-3 M

of ClO- to HOCl-. In basic bleaching solutions, hypochlorite content is usually around 5 to 6 % NaOCl. In the determination of the bleaching power of the unknown, the %Cl was obtained by determining the number of moles of S2O32-, and then determining the molarity of ClO- from the moles of S2O32-. After computing for the molarity of ClO-, the %Cl can be determined.

% Cl- : 1: (4.997 x 10-3 mol ClO- / L)(mol Cl- / mol ClO-)(35.45 g Cl- / mol Cl-)(1 L / 1000 mL) = 1.764 x 10-4 g/mL 2: (5.1508 x 10-3 mol ClO- / L)(mol Cl- / mol ClO-) (35.45 g Cl- / mol Cl-)(1 L / 1000 mL) = 1.826 x 10-4 g/mL 3: (5.1034 x 10-3 mol ClO- / L)(mol Cl- / mol ClO-) (35.45 g Cl- / mol Cl-)(1 L / 1000 mL) = 1.809 x 10-4 g/mL 4: (4.99596 x 10-3 mol ClO- / L)(mol Cl- / mol ClO-) (35.45 g Cl- / mol Cl-)(1 L / 1000 mL) = 1.771 x 10-4 g/mL 5: (4.9928 x 10-3 mol ClO- / L)(mol Cl- / mol ClO-) (35.45 g Cl- / mol Cl-)(1 L / 1000 mL) = 1.770 x 10-4 g/mL VII. Discussion In the first part of the experiment, the oxidizing agent used was potassium dichromate (K2Cr2O7) to accelerate the reduction of the aqueous iodide (I-) to iodine (I2). Addition of potassium iodide (KI) to the solution of iodine causes the formation of the I3- complex (Ka = 570) in acidic conditions. Without proper handling, iodine can be lost by vaporization. Therefore, preventive measures such as the covering of watch glass must be exercised to prevent loss of iodine. In the standardization, the molarity of the thiosulfate solution was obtained by dividing the weight of the dried dichromate by the molecular mass of potassium dichromate. The quotient was then stoichiometrically converted to moles of S2O32-, and then divided by the volume of sodium thiosulfate obtained through titration. Bleaching is process of whitening, or removing the natural color of textile fibers and many other materials through the treatment with chemicals or by exposure to the sun, heat or water. Sodium hypochlorite is an oxidizing bleaching agent used in household laundering and sanitation. To remain effective as a bleaching agent, hypochlorite must remain alkaline (ph > 9.0) to suppress the hydrolysis Chem 27.1. Oxidation – Reduction Titration Iodimetry

VIII. Conclusion and Reccomendations In this experiment, the %Cl content of the unknown bleach was determined. It was found that the average content of the unknown bleach was 1.788x10-4 g/mL. Possible sources of error include oxidation and loss of iodine by vapor, insufficient acidity of standard iodine solution which can cause incomplete reduction of dichromate by iodide. It is recommended that the experimental procedure should be followed strictly. Use protective gloves to protect hands since some reagents are iritants, and wear lab gown and mask. Proper disposal of wastes and used reagents must also be observed. IX. References

1. Hargis L.(1988).Analytical Chemistry. Prentice-Hall Inc.New Jersey.

2. Skoog, D.,West, D, Holler, F., Crouch, S.

3.

(2004).Fundamentals of Analytical Chemistry 8th ed.Thomson Learning Brooks/Cole.Singapore. Pierce, W.C., D.T. Sawyer, and E. L. Haenisch. “Iodometry and Iodimetry”. Quantitative Analysis, 4th ed. 1958. John Wiley and Sons: New York, pp291-307.

I hereby certify that I have given substantial contribution to this report.

DELA CRUZ, MARIE GIECEL V. PABILANE, ALMA L.

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