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2.1: Day 1 Procedures - Standardization of Sodium Thiosulfate

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  • Introduction

    General References – These are provided to improve your understanding of the techniques and afford practical hints that may help you avoid mistakes that may prove costly in terms of laboratory time.

    Joy Michaud, A Citizens Guide to Understanding and Monitoring Lakes and Streams, (1994) Washington State Department of Ecology, Chapter 4, pp. 45-58. Available online as a PDF document:

    • Measuring Mass and Volume     MHS, Chapter 5, pp. 52-64
    • Pipets, Transfer of Liquids
    • Standard Reducing Agents     SWH, Chapter 20, pp. 511-513
    • Standardizing Thiosulfate Solutions     SWH, Chapter 20, pp. 513-514
    • Applications of Thiosulfate Solutions     SWH, Chapter 20, p. 514
    • Dissolved Oxygen     SWH, Chapter 23, pp. 681-682
    • Phosphate Ion Determination     SWH, FIA, pp. 1059-1061
    • UV & VIS Spectroscopy     MHS, Chapter 25, pp. 465-475

    Required Videos: Digital Laboratory Techniques Manual

    • #1 Volumetric Techniques
    • #2 Titration
    • #7 Filtration
    • #11 Balance
    • #13 Automatic Pipet

    Dissolved Oxygen (DO) Determination7

    This experiment will sketch out the procedure for determining the dissolved oxygen (DO) levels in water samples obtained from the Charles River. This experiment uses the azide modification of the iodometric Winkler titration method.8,9 The procedure was first written up by a graduate student in 188810 and has since become the standard for determination of dissolved oxygen in sewage, streams, rivers, and various water systems.

    The experimental method is based on the oxidation of Manganese (II) from a manganous sulphate solution to a higher trivalent/tetravalent oxidation state. The resulting oxidation in the presence of base uses oxygen as the oxidizing agent and results initially in the formation of a white precipitate and later in the formation of a brown precipitate. The reaction scheme for the initial oxidation, which involves the addition of manganous sulfate solution and alkaline base, is as follows:

    \[ \rm Mn(SO_4)_{(aq)} + 2 KOH_{(aq)} \rightarrow \underset{\text{white ppt}}{Mn(OH)_{2 \space (s)}} + K_2SO_4\]

    \[ \rm 2 Mn(OH)_{2 \space\space (s)} + O_{2 \space\space (aq)} \rightarrow \underset{\text{brown ppt}}{2 MnO(OH)_{2 \space\space (aq)}}\]

    There are different perspectives in the literature as to how exactly the oxidized manganese brown ppt should be represented. Some have established it as trivalent manganese in the form of Mn(OH)3 others have indicated that hydrated MnO2 could also be the brown color.11 In the next part of the experiment with the addition of acid and KI in the alkaline potassium iodide azide solution the oxidized brown precipitate is acidified and goes back into solution simultaneously as the manganic ion is reduced back to manganous and elemental Iodine is generated via the oxidation of I- in the acidic medium. The amount of iodine, which is generated, is proportional to the amount of oxygen, which is present in the original sample. The reaction for the acidification and reduction is as follows:

    \[ \rm MnO(OH)_{2 \space\space (s)} + 2 H_2SO_{4 \space\space (aq)} \rightarrow Mn(SO_4)_{2 \space\space (aq)} + 2 H_2O_{(l)}\]

    \[ \rm Mn(SO_4)_{2 \space\space (aq)} + 2 KI_{(aq)} \rightarrow MnSO_{4 \space\space (aq)} + K_2SO_{4 \space\space (aq)} + I_{2 \space\space (aq)}

    In the final stage of the azide modification of the (Winkler) titration sodium thiosulfate is added. The sodium thiosulfate reacts with elemental iodine to produce sodium iodide. At the moment that all of the elemental Iodine has been converted the solution turns from yellow to clear. A starch indicator is used to capture the dramatic color change at the endpoint. The reaction is as follows:

    \[ \rm 2Na_2S_2O_{3 \space\space (aq)} + I_{2 \space\space (aq)} \rightarrow Na_2S_4O_{6 \space\space (aq)} + 2 NaI_{(aq)}\]

    The net overall ionic equation for the scheme presented is as follows:

    \[ \rm O_{2 \space\space (aq)} + 4 S_2O_3^{2-} + 4 H^+ \rightarrow 2 S_4O_{6 \space\space (aq)}^{2-} + 2 H_2O_{(l)} \]

    From this net reaction we can readily see that 4 moles of thiosulfate are required for each mole of oxygen.

    Day 1 - Standardization of Sodium Thiosulfate Solution

    TAs Preparation of 0.025XX M sodium thiosulfate solution12

    The TAs will prepare a solution of approximately 0.025XX M Na2S2O3 as follows: Mass out 6.205 g of Na2S2O3. 5H2O and dissolve it in 800 mL of hot distilled water.13 Add 1.5 mL of 6N NaOH to slow down bacterial decomposition and dilute the solution to 1 Liter in a volumetric flask. The solution bottle should be closed and stoppered immediately. Store the solution in the refrigerator until ready to use. Mix the dilute sodium thiosulfate solution very thoroughly by vigorous shaking with repeated inversions for several minutes each time the solution is used.14

    Standardization of approximately 0.025XX M sodium thiosulfate solution with potassium bi-iodate solution15

    The TAs will distribute a dry weighing bottle to each student containing approximately 0.1 gram of reagent grade potassium bi-iodate KH(IO3)2 , which has been previously dried in a 103-105 0 C drying oven for 1.5 hours or overnight.16

    At the start of the lab, each student should remove the weighing bottle from the oven and let it cool in a small desiccator charged with calcium chloride. Leave the stopper off the weighing bottle until the first time the dessicator is opened up after the KH(IO3)2 has cooled. Observe the precautions involved in the use of desiccators mentioned in Chapter 6 of the Techniques Manual.

    For the standard potassium bi-iodate solution, 0.0021 M: mass out 0.0818 g of dry KH(IO3)2 from the weighing bottle which had been previously heated for at least 1.5 hours and has now cooled in your dessicator into 50 mL of warm distilled water17 and dilute to 100 mL in a volumetric flask. The solution will be warm and can be titrated warm. Estimate all weights to ±0.1 mg (0.0001 g) and record all data immediately in you lab notebook.

    For the preparation of an aqueous starch solution: Dissolve 0.5 g of soluble starch and 0.05 g salicylic acid preservative18 by adding a few mL of distilled water to make a paste and dissolve in 25 mL of hot distilled water. The starch solution should be prepared fresh on the day you are going to use it. Keep it warm on hot plate & add hot to your solution.

    For the standardization titration: Take 100 mL of freshly mixed thiosulfate solution 0.025XX M and pour it into a beaker and keep it stirred and well mixed. The stock solution must be mixed several times before you draw off the 100 mL. Sodium thiosulfate solutions have a tendency to come out of solution when sitting for a period of time and are described as being perishable. Obtain a 50 mL burette and use a few mL of the thiosulfate solution to clean the burette. Then fill the burette with the freshly mixed thiosulfate solution letting it run down the sides slowly to avoid any bubbles forming on the inside of the burette. If you do get a bubble, tap the burette lightly on the lab bench or flick the burette with your finger to drive any bubbles to the surface. Make sure the tip of the burette is filled with thiosulfate solution and not air. For the standardization titration prepare three separate 250 mL Erlenmeyer flasks. Prior to the start of each titration19 add 2.0 g of potassium iodide (KI) into 100 mL of distilled water then add a few drops of concentrated sulfuric acid (DO NOT ADD SULFURIC ACID DIRECTLY TO KI AS I2 (g) WOULD ESCAPE). Pipette out and add 25.0 mL of the warm potassium bi-iodate solution then add 75 mL of distilled water for a total volume of approximately 200 mL. Immediately start the titration and titrate the liberated iodine in each flask with the thiosulfate titrant, stirring constantly. When the solution becomes a pale yellow color add 1.0 mL of the freshly prepared hot aqueous starch solution (15 to 20 drops), which changes the color of the solution from pale yellow to blue.20 Continue the titration until the color changes from blue to colorless. Record the volume of the thiosulfate used from the buret. Disregard any change back to the blue color after the endpoint has been reached.21 Repeat the titration with additional samples. Volumes should agree to within 5%.

    Each time you fill the burette with fresh solution, rinse the burette 3 times with 2mL of the new solution. Discard each wash into the appropriate waste container. Tilt the burette to allow the entire inner surface of the burette to come into contact with the liquid. After rinsing out the burette, fill it with the Na2S2O3 solution. Expel air bubbles trapped below the stopcock by fully opening the stopcock a second or two. If this is unsuccessful, see your TA for additional advice.

    The titration may be carried out rapidly at first, but the endpoint should be approached carefully. The endpoint should be sharp and easily located to within a fraction of a drop. The endpoint is taken as the first distinct colorless solution that persists for 10 seconds or more after thorough mixing. The color is not permanent and may fade back to the blue in a matter of minutes, which should be disregarded.

    Make all burette readings by estimating to the nearest 0.01 mL, allowing time for drainage. The tendency of liquids to stick to the walls of the burette can be diminished by draining the burette gradually. A slowly drained burette will provide greater reproducibility of results. Run a sufficient number of titrations to assure a precise and presumably accurate standardization. Record the final buret readings from each trial and subtract from the initial readings on the buret to quantify the amount in mL of thiosulfate used. The standardization titration should be repeatable to within 5%. The balanced equations for the standardization reactions are as follows:

    \[ \rm KH(IO_3)_{2 \space\space (aq)} + 10 KI_{(aq)} + 6 H_2SO_{4 \space\space (aq)} \rightarrow 6 I_{2 \space\space (aq)} + 6 H_2O_{(l)} + 5 K_2SO_{4 \space\space (aq)} + KHSO_{4 \space\space (aq)}\]

    \[ \rm 6 I_{2 \space\space (aq)} + 12 S_2O_3^{2-} \rightarrow 12 I^- + 6 S_4O_6^{2-}\]

    From the equations, you now have the stoichiometry of the reactions and should now be able to calculate the Molarity of the Thiosulfate solution.

    If three titrations do not result in the desired precision, it will be necessary to conduct additional titrations. With your notebook pages turned in at the end of the day, include a table giving the calculated molarity of the Na2S2O3 from each titration, calculate the average, standard deviation and the 95% confidence limits of the mean. No error propagation necessary for the thiosulfate standardization calculations.

    Record your calculated molarity for each titration on your TAs class data sheet before leaving the lab for the day. The TAs will average the results from the entire team and present each team member with the team average.

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