REDUCTION-OXIDATION TITRATIONS - cffet.net

1 Chapter 4. REDUCTION-OXIDATION TITRATIONS . A REDUCTION-OXIDATION reaction (Redox reaction) involves the transfer of an electron(s) from species in the reaction to another. Redox reactions allow TITRATIONS of analytes that cannot be determined by acid-base reactions. These include: Metal ions such as Cu2+. Anions such as nitrite Neutral or weakly acidic organic species such as Vitamin C. You should revise the following terms and concepts which were introduced in your Introductory chemistry classes, as they form the basis of the theory of Redox reactions Terminology: oxidation reduction Oxidant Reductant Oxidising agent Reducing agent oxidation number Half equation Overall equation You should be able to determine oxidation numbers and be able to use these to determine if a species has been oxidised or reduced during a reaction. Determine the oxidation number of the metal in the following species: Mn Mn2+. MnO4- MnO2. You also need to be able to work with half equations by selecting the correct ones, combining them to form balanced overall equations.

2 26. Chapter 4. Important half-equations for redox TITRATIONS (especially the practical work associated with this module). Table Redox half equations As2O3 + 5H2O 2H3 AsO4 + 4H+ + 4e C2O42- 2CO2 + 2e C6H8O6 (ascorbic acid) C6H6O6 + 2H+ + 2e Ce4+ + e Ce3+. Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O. Fe2+ Fe3+ + e H2O2 O2 + 2H+ + 2e I2 + 2e 2I- IO3- + 6H+ + 6e I- + 3H2O. MnO4- + 8H+ + 5e Mn2+ + 4H2O. NO2- + H2O NO3- + 2H+ + 2e 2S2O32- S4O62- + 2e In Redox TITRATIONS a strong oxidising agent such as permanganate is able to oxidise more species than a weaker oxidising agent such as iodine. Sample Pre-treatment Many species that cannot be analysed directly by redox can have their oxidation state changed leading to a species which can be analysed by a redox titration. Examples include: Metal ions that are essentially inert in their cation form, Cr3+ and Mn2+, become reactive if oxidised to a higher oxidation state in the form of an oxyanion, Cr2O72- and MnO4-. Excess oxidant must be removed prior to titration to prevent interference.

3 Oxidants for the above changes include persulfate (S2O82-) and periodate (IO4-). Persulfate requires silver nitrate as a catalyst and the excess is destroyed by boiling. Periodate is removed by precipitation. 27. Chapter 4. reduction of the analyte can be done in a number of ways: Jones reductor is a column containing an amalgam of zinc and mercury through which the analyte, such as iron (III) is passed. There are a number of problems involved in this procedure including the reduction of other species in the sample and the production of hydrogen gas with acidic solutions. Walden reductor in which the zinc amalgam is replaced by silver granules and dilute HCl. The method reduces fewer analytes but is costly. Homogeneous reduction methods use the reducing agent to dissolve the sample. The excess reducing agent must be destroyed prior to analysis. Tin II chloride is a common reducing agent, excess is destroyed by addition of mercury II chloride. OXIDISING TITRANTS.

4 Potassium permanganate Purple KMnO4- is a strong oxidising agent and is a common redox titrant. It is not a primary standard even though it is solid, available in AR grade and has a relatively high formula weight. When used in acidic conditions it is reduced to colourless Mn2+ according to the equation MnO4- + 8H+ Mn2+ + 4H2O. When used in neutral or alkaline conditions it forms manganese IV oxide, an insoluble solid which is a catalyst for the auto-decomposition of permanganate. This poses a problem for the shelf-life of the standardised permanganate. MnO4- + 4H+ + 3e MnO2 (s) + 2H2O. And MnO4- + 2H2O + 3e MnO2(s) + 4OH- Preparation of a solution of permanganate involves heating on a steam bath for about an hour and then filtering the solution before storing in a dark coloured glass storage bottle as light promotes the decomposition of the solution. Sulfuric acid ( ) is the acid of choice for a redox reaction with permanganate. Using hydrochloric acid poses a problem as the permanganate may react with the chloride ion forming chlorine gas.

5 Standardisation of permanganate may be against the primary standards, sodium oxalate or arsenic III oxide (very toxic), although a number of secondary standards could also be used. The reaction between oxalate and permanganate is initially slow, but is catalysed by the presence of Mn2+ formed during the reaction. Use the half-equations of Table to write the full redox equation for the reaction between acidic permanganate and oxalate. 28. Chapter 4. The McBride method of standardisation involved titration at 80-90oC. The high temperatures also cause the oxalate to be decomposed. The Fowler and Bright method of standardisation recommends rapid addition of 90-99% of the required volume at room temperature, followed by gentle warming to about 55-60OC. before completing the titration. The colour differences between the two species MnO4-, purple, and Mn2+, colourless, mean the reaction can be self-indicating and no further indicator is required. Analysis of iron using permanganate Iron III cannot be oxidised by permanganate, however, if the Fe3+ is reduced to Fe2+ oxidation is then possible Use table to write the overall equation for the reaction between permanganate and iron II.

6 For many samples containing Fe such as steel, hydrochloric acid is usually the acid that has been used for the dissolving process. This can cause problems as chlorine gas and hypochlorous acid may be formed. Changes to the method allow iron to be determined. Changes include: Diluting the sample Adding Mn2+to the reaction Titrating cold solutions Titrating slowly Adding phosphoric acid to form a colourless soluble complex with Fe3+. In practice Zimmerman-Reinhardt reagent, MnSO4, phosphoric acid and sulfuric acid is added after preliminary reduction of the iron. Potassium dichromate Dichromate is a weaker oxidising agent than permanganate and can therefore oxidise fewer species. This has the benefit of decreasing unwanted reactions in the titration. Dichromate will only react with chloride in acidic conditions and at chloride levels in excess of 1-2M. Dichromate in acidic conditions reacts according to the equation below Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O.

7 Yellow/orange green Potassium dichromate is a primary standard and can be prepared directly from the solid by accurate weighing. It can be standardised against a pure iron II salt or metallic iron converted to iron II. The strong colour of both the oxidised and reduced form of dichromate mean that an external indicator is required for a titration. 29. Chapter 4. True redox indicators change colour at particular solution voltage (or potential) ie the indicator is oxidised (or reduced) at a certain potential. Examples include: N-phenylanthranilic acid Sodium diphenylamine sulfonate (in the presence of phosphoric acid). Use Table to write the overall equation for the reaction between dichromate and ironII. Cerium (IV). Cerium IV, often as the sulfate, has an oxidising strength between permanganate and dichromate. It is extremely stable in solution (in the presence of 1 M sulfuric acid) and is not affected by storage or light. It is however, 100 times more expensive than permanganate.

8 Ce4+ + e Ce3+. Yellow colourless The solution needs to be standardised, usually against sodium oxalate or arsenic (III) oxide. Indicators of choice include ferroin and N-phenylanthranilic acid. Cerium IV works well in the oxidation of organic species such as alkanols and alkanals. Use Table to write the overall equation for the reaction between cerium IV and iron II. Iodine Iodine is a weak oxidising agent, only reacting with species that are easily oxidised. Solution conditions are normally neutral or acidic since iodine reacts with hydroxide. Where iodine is used to directly titrate the analyte the reaction is termed iodimetric. Iodine has a solubility in water of only g/100 mL of water. In the presence of an excess iodide ions soluble triiodide ion I3- is formed. I2 + I- I3 - A iodine solution has an iodide concentration of This can be standardised against arsenic (III) oxide or sodium thiosulfate 30. Chapter 4. Iodine is volatile making its concentration very variable with time.

9 In some of your practical work you will be instructed to work on samples individually rather than three at a time. Why? Iodine is red-brown and forms colourless iodide ion in its redox reaction. With experience this endpoint can be identified without the presence of an indicator. The endpoint is enhanced with the addition of a starch based indicator. Starch forms an intense blue colour in the presence of excess iodine. Starch is not a true redox indicator but a specific indicator for iodine. Iodine can be used to titrate easily oxidised species without reacting with other species. It is common for analysis of : Sulfur dioxide Sulfide Ascorbic acid Back titration may be used if the analyte reacts with oxygen in solution. This occurs in the case of Vitamin C which decomposes when in solution. With sulfide and SO2 reaction is best in acidic solution, but this risks the loss of the analyte as a partly soluble gas (H2S or SO2). With these analytes the titration may be performed in the reverse ie the sample in the burette and the standard in the flask.

10 Reducing Titrants Sodium Thiosulfate (Na2S2O3). Analytes that exist in a fully oxidised state: Copper II. Periodate (IO4-). Hypochlorite (ClO-). cannot be titrated with oxidising titrants such as permanganate or dichromate. They can be titrated with reducing agents such as sodium thiosulfate. Many other reducing agents are affected by oxidation in the air and the standardised figure is only valid for a short period of time. Thiosulfate reacts quantitatively only with iodine, however thiosulfate is standardised using potassium iodate as the primary standard. The iodate does not react with the thiosulfate, but in the presence of an acidic solution and iodide ions it forms iodine which can be titrated. Formation of iodine IO3- + 5I- + 6H+ 3I2 + 3H2O. 31. Chapter 4. Reaction of iodine with thiosulfate Using the half equations in Table write the balanced equation for the reaction between iodine and thiosulfate. What is the reaction ratio between the iodate ion and thiosulfate?