REDUCTION-OXIDATION TITRATIONS

Transcription

1 REDUCTIONOXIDATION TITRATIONS A reductionoxidation 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 acidbase reactions. These include: Metal ions such as Cu 2+ 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 Mn 2+ MnO 4 MnO 2 You also need to be able to work with half equations by selecting the correct ones, combining them to form balanced overall equations. 26

2 Important halfequations for redox titrations (especially the practical work associated with this module) Table 4.1 Redox half equations As 2 O 3 + 5H 2 O 2H 3 AsO 4 + 4H + + 4e 2 C 2 O 4 2CO 2 + 2e C 6 H 8 O 6 (ascorbic acid) C 6 H 6 O 6 + 2H + + 2e Ce 4+ + e Ce 3+ 2 Cr 2 O H + + 6e 2Cr H 2 O Fe 2+ Fe 3+ + e H 2 O 2 O 2 + 2H + + 2e I 2 + 2e 2I IO 3 + 6H + + 6e I + 3H 2 O MnO 4 + 8H + + 5e Mn H 2 O NO 2 + H 2 O NO 3 + 2H + + 2e 2 2 2S 2 O 3 S 4 O 6 + 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 Pretreatment 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, Cr 3+ and Mn 2+, become reactive if oxidised to a higher oxidation state in the form of an oxyanion, Cr 2 O 7 2 and MnO 4. Excess oxidant must be removed prior to titration to prevent interference. Oxidants for the above changes include persulfate (S 2 O 8 2 ) and periodate (IO 4 ). Persulfate requires silver nitrate as a catalyst and the excess is destroyed by boiling. Periodate is removed by precipitation. 27

3 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 Potassium permanganate Purple KMnO 4 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 Mn 2+ according to the equation MnO 4 + 8H + Mn H 2 O When used in neutral or alkaline conditions it forms manganese IV oxide, an insoluble solid which is a catalyst for the autodecomposition of permanganate. This poses a problem for the shelflife of the standardised permanganate. MnO 4 + 4H + + 3e MnO 2 (s) + 2H 2 O And MnO 4 + 2H 2 O + 3e MnO 2(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 (0.51.0M) 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. 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 Mn 2+ formed during the reaction. Use the halfequations of Table 4.1 to write the full redox equation for the reaction between acidic permanganate and oxalate. 28

4 The McBride method of standardisation involved titration at 8090 o C. The high temperatures also cause the oxalate to be decomposed. The Fowler and Bright method of standardisation recommends rapid addition of 9099% of the required volume at room temperature, followed by gentle warming to about 5560 O C before completing the titration. The colour differences between the two species MnO 4, purple, and Mn 2+, colourless, mean the reaction can be selfindicating and no further indicator is required. Analysis of iron using permanganate Iron III cannot be oxidised by permanganate, however, if the Fe 3+ is reduced to Fe 2+ oxidation is then possible Use table 4.1 to write the overall equation for the reaction between permanganate and iron II. 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 Mn 2+ to the reaction Titrating cold solutions Titrating slowly Adding phosphoric acid to form a colourless soluble complex with Fe 3+ In practice ZimmermanReinhardt reagent, MnSO 4, 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 12M. Dichromate in acidic conditions reacts according to the equation below Cr 2 O H + + 6e 2Cr H 2 O 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

5 True redox indicators change colour at particular solution voltage (or potential) ie the indicator is oxidised (or reduced) at a certain potential. Examples include: Nphenylanthranilic acid Sodium diphenylamine sulfonate (in the presence of phosphoric acid) Use Table 4.1 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. Ce 4+ + e Ce 3+ Yellow colourless The solution needs to be standardised, usually against sodium oxalate or arsenic (III) oxide. Indicators of choice include ferroin and Nphenylanthranilic acid. Cerium IV works well in the oxidation of organic species such as alkanols and alkanals. Use Table 4.2 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 I 3 is formed. I 2 + I I 3 A 0.05M iodine solution has an iodide concentration of 0.15M. This can be standardised against arsenic (III) oxide or sodium thiosulfate 30

6 Iodine is volatile making its concentration very variable with time. In some of your practical work you will be instructed to work on samples individually rather than three at a time. Why? Iodine is redbrown 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 SO 2 reaction is best in acidic solution, but this risks the loss of the analyte as a partly soluble gas (H 2 S or SO 2 ). With these analytes the titration may be performed in the reverse ie the sample in the burette and the standard in the flask. Reducing Titrants Sodium Thiosulfate (Na 2 S 2 O 3 ) Analytes that exist in a fully oxidised state: Copper II Periodate (IO 4 ) 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 IO 3 + 5I + 6H + 3I 2 + 3H 2 O 31

7 Reaction of iodine with thiosulfate Using the half equations in Table 4.1 write the balanced equation for the reaction between iodine and thiosulfate. What is the reaction ratio between the iodate ion and thiosulfate? In acidic conditions or in the presence of bacteria thiosulfate can decompose to hydrogen sulfite and elemental sulfur. Antibacterial agents such as trichloromethane or sodium benzoate should be added to maintain solution stability. The solution is also prone to layering when stored. The container should therefore be shaken thoroughly before use. When iodine is being titrated (ie present in significant quantities) an insoluble and irreversible complex forms with starch resulting in loss of iodine and therefore an artificially low endpoint value. Starch should therefore not be added to the reaction vessel until almost all the iodine has been reduced. Alternative indicators such as Vitex do not have this problem. Iodometry is the process used for the quantitative use of thiosulfate to analyse a range of oxidised species. The analyte reacts quantitatively with an excess of iodide ion to form iodine, which can then be titrated with a standardised thiosulfate solution, as shown in Fig 4.1. Analyte + Iodide Reduced analyte + Iodine Fig 4.1 The process of iodometry Iodine titrated with standardised thiosulfate The reaction does not involve a back titration as the iodide (added as KI) is not added quantitatively, simply in excess. The titration determines the moles of iodine which indicates the moles of analyte. 32

8 Write the equations for the reaction of copper II ions with iodide forming a precipitate of copper II iodide and iodine followed by the reaction of the iodine formed with thiosulfate. The excess iodide can be a problem since it can be oxidised by the oxygen in air and in solution to form iodine introducing an error. The error is known as the oxygen error and can be minimised by avoiding high concentrations of acid and titrating as soon as possible. O 2 + 4I + 4H + 2I 2 + 2H 2 O Typical iodometric analyses include: Copper II Hydrogen peroxide Hypochlorite ion For copper II analyses the formation of copper II iodide is prone to error as the iodine may adsorb onto the precipitate. Thiocyanate is added to replace iodide in the precipitate and prevent loss of iodine. Iron (II) The rapid reaction of Fe II with oxygen limits the use of Fe II as a reducing titrant as frequent standardisation is required. However it is a widely used reducing titrant as it does not produce toxic byproducts and it reacts quantitatively. Iron II is available in AR grade forms such as ammonium iron II sulfate or can be prepared by dissolving pure iron acid and reducing the solution in a Jones reductor (Zn and Hg amalgam). A major application of iron II is with dichromate for the determination of organic content of water samples. The analysis known as the chemical oxygen demand (COD) involves refluxing the sample in a excess of standard dichromate and back titrating with Fe II which has been standardised against the dichromate. Redox indicators Table 4.3 Indicator Titrants used with Colour change Starch Iodine Colourless to blue Nphenylanthranilic acid Cerium IV, dichromate Purple to colourless Ferroin Cerium IV Blue to red Diphenylamine sulfonate Dichromate Purple to colourless 33

9 Concepts 1. Why can t Fe 3+ be analysed by titration with permanganate? How could it be treated to allow titration with permanganate? 2. Explain why a true redox indicator could be used in the titration of iodine with thiosulfate, but a specific indicator, such as starch could not be used in a titration of iron (II) with permanganate 3. Give two reasons why an iodometric titration should be performed as soon as possible after addition of the iodide Applications 1. Permanganate solutions can be standardised by titrating against a primary standard oxalic acid. Calculate the molarity of a KMnO 4 solution if g of oxalic acid, (COOH) 2, is titrated to endpoint by 18.3 ml of KMnO 4 2. An iodine disinfectant solution was analysed by titration with M thiosulfate solution. The iodine solution was diluted from 10.0 to ml and 25.0 ml aliquots titrated to endpoint by 17.9 ml of the thiosulfate solution. Calculate the %w/v of the original iodine solution 3. Calculate the SO 2 content (mg/l) in 100 ml of a white wine sample which was titrated to endpoint with 14.9 ml of M I A pig iron sample weighing g is dissolved, quantitatively converted to the iron (II) state and the solution accurately made up to ml ml aliquots of this are titrated with an average of 16.4 ml of a M Ce(IV) solution. Calculate the percentage of Fe in the sample 5. Calculate the molarity of a thiosulfate solution, given the following data g of potassium iodate is dissolved and made up to 250 ml. 25 ml aliquots of the iodate are made acidic and reacted with KI. Titration of the liberated iodine required an average of 22.4 ml of thiosulfate solution. 6. The concentration of oxygen in water can be determined by iodometry. Oxygen oxidises the Mn(OH) 2 to Mn(OH) 3, and is itself reduced to water. The manganese (III) then reacts in acidic conditions with iodide to form iodine. This is titrated with thiosulfate. Calculate the concentration of oxygen in water if a 250 ml sample requires 7.8 ml of M thiosulfate in the titration. 7. Calculate the cost of titrant in a titration of 25 ml of 0.1 M iron (II) with cerium (IV), permanganate and dichromate. Costs for 500 g bottles are (NH 4 ) 2 Ce(NO 3 ) 6 $300, K 2 Cr 2 O 7 $50, KMnO 4 $ A 25 ml aliquot of a nitrite solution was reacted with 50 ml of M KMnO 4. The excess permanganate was backtitrated with 23.6 ml of M Fe 2+. Calculate the concentration of the nitrite solution (in g/l). 34