Permanganometry, iodometry in analytical technique, P K MANI


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Permanganometry, iodometry in analytical technique

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Permanganometry, iodometry in analytical technique, P K MANI

  1. 1. Pabitra Kumar Mani, Assoc. Prof., ACSS, BCKV Class 13 Red–ox titrations, permanganometry, iodometry etc.
  2. 2. Redox Titrations • • • • • Basics Potassium Permanganate Potassium Dichromate Cerium IV Iodine
  3. 3. Permanganometry This valuable and powerful oxidising agent was first introduced into titrimetric analysis by F. Margueritte for the titration of iron(II). In acid solutions, the reduction can be represented by the following equation The standard reduction potential in acid solution, E0 has been calculated to be 1.51 volts; hence the permanganate ion in acid solution is a strong oxidising agent. Sulphuric acid is the most suitable acid, as it has no action upon permanganate in dilute solution. With HCl, there is a likelihood of the reaction In the HCl , permanganate can oxidize Cl- to Cl2, which can be a source of positive errors as permanganate is consumed in this reaction. (E°red Cl2/Cl-)= +1.36V
  4. 4. Permanganate titration KMnO4 Powerful oxidant that the most widely used. Eq. Wt.(=M/5): In strongly acidic solutions (1M H2SO4 or HCl, pH ≤ 1) MnO4– + 8H+ + 5e- = Mn2 + + 4H2 O Eo = 1.51 V violet color colorless manganous KMnO4 is a self-indicator. In feebly acidic, neutral, or alkaline solutions (E=M/3) MnO4– + 4H+ + 3e- = MnO2 (s) + 2H2 O Eo = 0.59 V brown manganese dioxide solid In very strongly alkaline solution (2M NaOH or Ba (OH)2) (E=M/1) MnO4– + e- = MnO42 – Eo = 0.56 V green manganate III E=M/4 (in HF or NH4HF2 Medium) MnO4– + 4e- + 6F-+ 8H+ = [MnF6] 3 – + 4H2O Trivalent Fluoro magnate anion
  5. 5. Estimation of Fe+2 In the analysis of iron ores, (solution is frequently effected in conc. HCl); the Fe+3 is reduced and the Fe+2 is then determined in the resultant solution. If Cl- is present, to prevent its oxidation in acidic medium (1-2 N) by MnO4- about 25 mL of Zimmermann and Reinhardt's solution (preventive solution) has to be used. It is prepared by dissolving 50 g of crystallised (MnSO4,4H2O) in 250 mL water, adding a cooled mixture of 100 mL conc.H2SO4 and 300 mL water, followed by 100 mL H3PO4. The manganese (II) sulphate (presence of Mn+2) lowers the oxidation potential of the MnO4- - Mn(II) couple (-1.20V) and thereby makes it a weaker oxidising agent; the tendency of the permanganate ion to oxidise chloride ion is thus reduced.( Eo of Cl-/Cl2 is much higher) [ ] ] 0.0591 MnO 4 [ H E = -1.52 log [ Mn + 2 ] 5 − + 8 positive See Vogels book
  6. 6. Determination of Nitrite: Nitrites react in warm acid solution (400C) with permanganate solution in accordance with the equation: 5NO2- + 2MnO4- + 6H+ = 5NO3- + 2 Mn2+ + 3H2O If a solution of a nitrite is titrated in the ordinary way with potassium permanganate, poor results are obtained, because the NO2- soln has first to be acidified with dil.H2SO4 . Nitrous acid is liberated, which being volatile and unstable, is partially lost. If, however, a measured volume of std. KMnO 4 soln, acidified with dil.H2SO4, is treated with the nitrite solution, added from a burette, until the permanganate is just decolorised, results accurate to 0.5-1 per cent may be obtained
  7. 7. Preparation of 0.1 N potassium permanganate solution KMnO4 is not pure. Distilled water contains traces of organic reducing substances which react slowly with permanganate to form hydrous managnese dioxide. MnO2 promotes the autodecomposition of permanganate. 1) Dissolve about 3.2 g of KMnO4 (mw=158.04) in 1000ml of water, heat the solution to boiling, and keep slightly below the boiling point for 1 hr. Alternatively , allow the solution to stand at room temperature for 2 or 3 days. 2) Filter the liquid through a sintered-glass filter crucible to remove solid MnO2. 3) Transfer the filtrate to a clean stoppered bottle freed from grease with cleaning mixture. 4) Protect the solution from evaporation, dust, and reducing vapors, and keep it in the dark or in diffuse light. Preserve it in amber – coloured glass bottle. 5) Standardise from time to time. If in time managanese dioxide
  8. 8. Ordinary distilled water is likely to contain reducing substances (traces of organic matter, etc.) which will react with the KMnO4 to form MnO2. The presence of the manganese dioxide is very objectionable because it catalyses the autodecomposition of the permanganate solution on standing 4 MnO4- +2H2O = 4 MnO2 +3O2 +4 OHPermanganate is inherently unstable in the presence of Mn+2 ions: 2MnO4- +3Mn2+ + 2H2O = 5 MnO2 + 4H+ Potassium permanganate solutions may be standardised using Primary standards : arsenic(III) oxide or sodium oxalate Secondary standards : metallic iron etc.
  9. 9. Standardization of KMnO4 solution Standardization by titration of sodium oxalate Na2C2O4.2H20 (primary standard) (Fowler and Bright) : C2O42- = 2CO2 + 2 e- E° red = +0.77V 2KMnO4 +5 Na2(COO)2 +8H2SO4 = 2MnSO4 +K2SO4 +5Na2SO4 +10 CO2 + 8H2O The reaction between oxalic acid and potassium permanganate can be represented as: 2KMnO4 + 5 H2C2O4 +3H2SO4 = 2MnSO4 +K2SO4 +10 CO2+ 8H2O In ionic form the reaction can be represented as: 2MnO4- + 5 C2O4 2- + 16H+ = 2Mn2+ + 10 CO2 + 8H2O This titration is carried out in warm conditions (60 oC). The reaction at room temperature is slow because of the equilibrium nature of this reaction. CO2 is highly soluble in water and thus heating removes all dissolved CO2 out of the solution driving the reaction in forward direction. Also at low temperature, the reduction of permanganate may not be complete producing Mn(III) (in the form [Mn(C2O4)3]3-). The formation of this species introduce errors in titrations as no. of electrons utilized here are different as compared to production of Mn2+.
  10. 10. Standardization of KMnO4 solution by Arsenic(III) oxide This procedure of H.A.Bright, which utilises As(III) oxide as a primary stand. and KI or potassium iodate (KIO3) as a catalyst for the reaction, is convenient in practice and is a trustworthy method for the standardisation of permanganate solns. As2O3 weighed, dissolved in 3N NaOH, H2SO4(4N) added, a drop of very dilute KIO3 added as catalyst and titrated by MnO4-.
  11. 11. Titration of K2Cr2O7 with Mohr’s salt. K2Cr2O7 a strong oxidizing agent (E°red = +1.33V) but, not as strong oxidizing agent as permanganate (E°red = +1.51V). It is widely used in redox titrations because of several advantages over permanganate. Unlike KMnO4, K2Cr2O7 is available in high purity and is highly stable upto its melting point. Its aqueous solutions are not attacked by organic matter and thus composition of aqueous solution does not change on keeping. The aqueous solutions are quite stable towards light. It is an excellent primary standard and its standard solutions can be prepared by direct weighing of an amount of it and dissolving in a known volume of distilled water. K2Cr2O7 acts as oxidizing agent in acidic medium only: The neutral aqueous solution of K2Cr2O7 is 1:1 equilibrium mixture of dichromate and chromate, a consequence of hydrolysis of dichromate ions. Cr2O72– + H2O = 2 CrO42– + 2H+ Orange yellow Chromate ions are weaker oxidizing agent than dichromate. Thus
  12. 12. [ ] 0.0591 Cr2 O 7 [ H E =E log [Cr +3 ] 6 0 −2 (10 )(10 0.0591 E = −1.33 log -2 2 6 [10 ] −3 0.0591 E = −1.33 log10 - 27 6  0.0591  E = −1.33 + 27 x   6  E = − .06V 1 ] + 14 −2 ) 14
  13. 13. E =E [ Fe ] - 0.0591log [ Fe ] +3 0 +2 0.003 E = -0.771 - 0.0591log 0.15 E = -0.671V [ ] 0.0591 MnO 4 [ H + ] 0 E=E log [ Mn + 2 ] 5 − 8 ( 0.02)( 1.00 ) 0.0591 E = -1.51 log ( 0.005) 5 E = -1.52 V 8
  14. 14. E Mn = E Fe 0.05 = -0.771 - 0.0591log = 0.79V 0.103
  15. 15. Methods Involving Iodine • Iodimetry: a reducing analyte is titrated directly with iodine (to produce I−). • Iodometry, an oxidizing analyte is added to excess I− to produce iodine, which is then titrated with standard thiosulfate solution. • Iodine only dissolves slightly in water. Its solubility is enhanced by interacting with I- • A typical 0.05 M solution of I2 for titrations is prepared by dissolving 0.12 mol of KI plus 0.05 mol of I2 in 1 L of water. When we speak of using iodine as a titrant, we almost always mean that we are using a solution of I plus excess I−.
  16. 16. The direct iodometric titration method (Iodimetry) refers to titrations with a standard solution of iodine. The indirect iodometric titration method (Iodometry) deals with the titration of iodine liberated in chemical reactions. The normal oxidation potential of the reversible system: 2I- ⇋ I 2 + 2e in most iodometric titrations, when an excess of iodide ion is present the tri-iodide ion is fromed I2 (aq) + I- ⇋ I3since iodine is readily soluble in a solution of iodide. The half-cell reaction is better written: I3- +2e ⇋ 3I-
  17. 17. and the standard oxidation potential is -0.5355 volt. Iodine or the tri-iodide ion is therefore a much weaker oxidising agent than potassium permanganate, potassium dichromate, and cerium(IV) sulphate. In most direct titrations with iodine (iodimetry) a solution of iodine in potassium iodide is employed, and the reactive species is therefore the triiodide ion I3-. Strictly speaking, all equations involving reactions of iodine should be written with I3-; rather than with I2 e.g.
  18. 18. I2 I2 I2 I2 + + + + 2S2O3- → 2I- + S4O6-2 H2S → S + 2I- + 2H+ SO3-2 +H2O → 2I- + SO4-2 SnCl2 +2HCl → 2I- + SnCl4 + 2H+ The normal oxidation potential of the iodine-iodide system is independent of the pH of the solution so long as the latter is less than about 8; at higher values iodine reacts with hydroxide ions to form iodide and the extremely unstable hypoiodite, the latter being transformed rapidly into iodate and iodide by self-oxidation and reduction:
  19. 19. By suitable control of the pH of the solution, it is sometimes possible to titrate the reduced form of a substance with iodine, and the oxidised form, after the addition of iodide, with sodium thiosulphate. Thus with the arsenite-arsenate system: H3 ASO3 + I2 + H2O ⇋ H3AsO4 + 2 H+ + 2Ithe reaction is completely reversible. At pH values between 4 and 9, arsenite can be titrated with iodine solution. In strongly acid solutions, however, arsenate is reduced to arsenite and iodine is liberated. Upon titration with sodium thiosulphate solution, the iodine is removed and the reaction proceeds from right to left
  20. 20. Preparation and Standardization of Solutions Two important sources of error in titrations involving iodine are: (a) loss of iodine owing to its appreciable volatility; and (b) acid solutions of iodide are oxidised by oxygen from the air: + 4I + O2 + 4H ⇋ 2I2 + 2 H2O • Acidic solutions of I3- are unstable because the excess I− is slowly oxidized by air: • In neutral solutions, oxidation is insignificant in the absence of heat, light, and metal ions. At pH ≳ 11, triiodide disproportionates to hypoiodous acid (HOI), iodate, and iodide. • An excellent way to prepare standard I3- : is to add a weighed quantity of potassium iodate to a small excess of KI. Then add excess strong acid (giving pH ≈ 1) to produce I2 by quantitative reverse disproportionation: IO3- + 5I- + 6H+ ⇋ 3I2 + 3H2O
  21. 21. Fact File 1: Introduction to iodometric and iodimetric titrations Third: Iodometric titration 2 Cu 2+ + 4I- → 2CuI + I2 Analyte of unknown concentration I2 + 2S2O32- → 2I- + S4O62- Titrant -standrard solutions: sodium thiosulfate -known concentration
  22. 22. Starch-Iodine complex Starch solution(05~ 1%) is not redox indicator. The active fraction of starch is amylose, a polymer of the sugar αD-glucose ( 1,4 bond). The polymer exists as a coiled helix into which small molecules can fit. In the presence of starch and I–, iodine molecules form long chains of I5– ions that occupy the center of the amylose helix. ••••[I I I I I]– ••••[I I I I I]– •••• Visible absorption by the I5– chain bound within the helix gives rise to the characteristic starch-iodine color.
  23. 23. Structure of the repeating unit of the sugar amylose. Schematic structure of the starchiodine complex. The amylose chain forms a helix around I6 unit. View down the starch helix, showing iodine, inside the helix.
  24. 24. Starch-Iodine Complex • • • Starch is the indicator of choice for those procedures involving iodine because it forms an intense blue complex with iodine. Starch is not a redox indicator; it responds specifically to the presence of I2, not to a change in redox potential. The active fraction of starch is amylose, a polymer of the sugar α-d-glucose. In the presence of starch, iodine forms I6 chains inside the amylose helix and the color turns dark blue
  25. 25. Common Titrant for Oxidation Reactions Iodine (Solution of I2 + I-) I3- is actual species used in titrations with iodine K = 7 x 102 Either starch of Sodium Thiosulfate (Na2S2O3) are used as indicator I3- I3- + S2O32- Before endpoint I3- + Starch Before endpoint At endpoint
  26. 26. SOLUTION
  27. 27. Determination of Cu+2 : 2Cu+2 + 4I- → CuI + I2 Acetic acid buffer pH ~4.5 or better NH4HF2 buffer. In presence of free mineral acid, at pH<3, dissolved O 2 liberate I2 from I- also. The elments which interferes with the iodometric determination are iron, arsenic and antimony, Trivalent iron is reduced by iodide: 2Fe3+ + 2I- ⇋ 2Fe2+ + I2 but by addition of excess of fluoride, the iron(III) is converted into the complex [FeF6]3-, which yields so small a concn of Fe+3 ions that it has no oxidising action upon the
  28. 28. DETERMINATION OF THE AVAILABLE CHLORINE IN Bleaching powder the hypochlorite solution or suspension is treated with an excess of a solution of potassium iodide, and strongly acidified with acetic acid: Ca(OCl)+ KI +HAc → CaCl2 + I2 + H2O + KAc The liberated iodine is titrated with standard sodium thiosulphate solution. Determination of hypochlorite in bleaches [CaCl(OCl)H2O]: OCl– + 2I– + 2H+ → Cl– + I2 + H2O (unmeasured excess KI) I2 + 2 S2O3 2– → 2I– + S4O6 2– Indicator: soluble starch (β-amylose)
  29. 29.