Iodometric and Iodimetric. Titrations. Module 3. Iodometric Titration of Copper. Module 4. The Sample: Copper wire. Module 5. The Procedure for the Iodometric. Iodometry. The term “iodometry” describes the type of titration that uses a standardised sodium thiosulfate solution as the titrant, one of the few stable reducing. Iodine, the reaction product, is ordinary titrated with a standard sodium Solutions of sodium thiosulfate are conveniently standardised by titration of the iodine.
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Iodometry, known as iodometric titration, is a method of volumetric chemical analysis, a redox . Create a book · Download as PDF · Printable version. involve the potentiometric titration of aqueous iodine with sodium thiosulfate using . This procedure describes the analysis of hydrogen peroxide by iodometric. Iodometry and iodimetry are two common titration methods used in analytical chemistry. These two types of titrations are based on oxidation.
What is Iodometry As mentioned above, Iodometry is an indirect method. The technique of Iodometry is commonly used in experiments where the amount of oxidizing agents in a water body needs to be quantified. What happens here is, an excess amount of Iodide solution typically Potassium Iodide is mixed with a sample of the water that needs to be tested.
Due to the oxidizing agents present in the water body, the Iodide ions get oxidized to Iodine, while the oxidizing agents get reduced. This is the initial redox reaction. Then the produced Iodine is titrated with a reducing agent such as sodium thiosulfate solution. Here, the Iodine reduces to Iodide ions while the thiosulfate ions get oxidized further.
This is the second redox reaction and it is the reaction used for the titration. This is performed in the presence of a starch indicator to make it easier to recognize the end point. Iodine forms a deep-blue colour complex with starch and as the Iodine breaks down to Iodide ions, the colour disappears.
In strongly alkaline or acidic solutions the oxidation of the thiosulfate does not proceed by a single reaction. In the former, the thiosulfate ion is oxidized to sulfate as well as to the tetrathionate. In the latter, the thiosulfuric acid formed undergoes an internal oxidation-reduction reaction to sulfurous acid and sulfur. Both of these reactions lead to errors since the stoichiometry of the reactions differs from that shown above for the thiosulfate as a reducing agent.
The control of pH is clearly important.
In many cases the liberated iodine is titrated in the mildly acidic solution employed for the reaction of a strong oxidizing agent and iodide ion. In these cases the titration of the liberated iodine must be completed quickly in order to eliminate undue exposure to the atmosphere since an acid medium constitutes an optimum condition for atmospheric oxidation of the excess iodide ion.
The iodine that is liberated can be titrated in the usual manner with standard thiosulfate solution. The reaction involving cupric ion and iodide takes place quantitatively since the cuprous ion formed as result of the reduction is removed from the solution as a precipitate of cuprous iodide.
Iron interferes since iron III ions will oxidize iodide. If arsenic and antimony are present, they will provide no interference at this pH if they are in their higher oxidation states. Under these conditions the tin does not interfere with the analysis. Sources of Error The following are the most important sources of error in the iodometric method: 1. Loss of iodine by evaporation from the solution. This can be minimized by having a large excess of iodide in order to keep the iodine tied up as tri-iodide ion.
It should also be apparent that the titrations involving iodine must be made in cold solutions in order to minimize loss through evaporation. Atmospheric oxidation of iodide ion in acidic solution. In acid solution, prompt titration of the liberated iodine is necessary in order to prevent oxidation. Starch solutions that are no longer fresh or improperly prepared.
The indicator will then not behave properly at the endpoint and a quantitative determination is not possible. Preparation of a 0. Place it in your 1 liter beaker and boil the water for at least 5 minutes. Dissolve the thiosulfate in the hot water and then cool this solution with the aid of an ice bath to room temperature.
Then add the carbonate and stir until it is completely dissolved. Transfer the solution to your plastic 1 liter bottle. When not in use store this bottle in the darkness of your equipment cabinet as the decomposition of thiosulfate is catalyzed by light. Blank Determination Potassium iodide may contain appreciable amounts of iodate ion which in acid solution will react with iodide and yield iodine.
The liberated iodine would react with thiosulfate and thereby cause the apparent molarity of the thiosulfate to be too low. The following procedure allows for the determination of a blank correction which will properly correct for any iodate that might be present.
Prepare a solution of exactly 2. If a blue-black color appears right after mixing, use the thiosulfate solution in the buret to determine the volume of solution required to cause the color- to disappear. This volume must be subtracted from the standardization and analyses volumes. If the potassium iodide is completely iodate-free no color will of course develop and no blank correction is necessary. Dissolve the iodate in 75 mL of distilled water.
Cover the flasks with parafilm and store them. Rinse and fill your buret with the solution. Add 2. If a blank correction is required add exactly 2. If no blank determination is required, the exact amount of KI is not crucial but should be close to 2 g. Then add 10 mL of 1 M HCl to one of the solutions.
It will turn a dark-brown color. Immediately titrate it with the thiosulfate solution. When the color of the solution becomes very pale yellow add 5 mL of starch indicator. Continue the titration until the blue color of the starch complex just disappears.
Follow the same procedure with each of the other two solutions, first adding the HCl then titrating. Correct your titration data for buret error and if necessary apply the blank correction. Calculate the molarity of the Na2S2O3 solution. Results should agree to within 0. If you do not achieve that kind of precision, titrate additional samples.
Dissolution of the Brass Sample The following procedures in this section make use of the hot plates in the fume hoods.
In the presence of iodine, the thiosulphate ions oxidise quantitatively to the tetrathionate ions. To determine the concentration of the oxidising agents, an unknown excess of potassium iodide solution is added to the weakly acid solution.
The iodine, which is stoichiometrically released after reduction of the analyte, is then titrated with a standard sodium thiosulphate solution Na2S2O3. This suspension is a watery solution of starch with a few drops of bactericide added to prevent decomposition, as this would stop the starch behaving as an indicator.
Once the bond between the iodine I2 and the helical chain of beta-amylose is formed it turns an intense blue.
Colour of the starch solution in the presence of I2. In the presence of I- ions the starch solution is colourless.
Source: Istituto comprensivo di Tubirgo Important considerations Iodometric titration needs to be done in a weak acid environment which is why we need to remember that: 1.
Sodium thiosulphate needs a neutral or weak acid environment to oxidise with tetrathionate in an alkaline solution we would get sulphate oxidation ; 3. In a strong acid environment thiosulphate decomposes to S2; 4. Oxidation is a chemical process which is catalysed by various factors presence of oxygen, levels of unsaturation in the oil, presence of metals, temperature and leads to the formation of hydroperoxides.