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Application of Oxidation-Reduction Titrations

1. Application of Oxidation-Reduction Titrations. Hassan.F.Askal, Ph.D. Importance of Redox Reactions in Biological systems and Pharmaceutical Applications. 2. Oxidizing agents act on human tissues by one of the following mechanisms :

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Application of Oxidation-Reduction Titrations

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  1. 1 Application of Oxidation-Reduction Titrations Hassan.F.Askal, Ph.D.

  2. Importance of Redox Reactions in Biological systems and Pharmaceutical Applications 2 • Oxidizing agents act on human tissues by one of the following mechanisms: • Germicides through liberation of oxygen in the tissues such as hydrogen peroxide • Bactericidal against anaerobes as KMnO4 (for wounds) or Cl2 (for dental therapy) through denaturing the proteins by direct oxidation. • Oxidation can have deleterious effect on the cells and tissues of human body because it produces reactive oxygen and nitrogen species, such as OH, NO, NO2 and alkoxyl radicals RO which damage cells and other body components. • Antioxidants (reducing agents) can counteract the effect of the reactive oxygen and nitrogen species and protect the other compounds in the body from oxidation.

  3. Importance of Redox Reactions in Biological systems and Pharmaceutical Applications 3 • Typical antioxidants include vitamin A, C and E; minerals such as selenium; herbs such as rosemary. • Antioxidants act by one or more of the following mechanisms: • Reaction with the generated free radicales can counteract the effect of the reactive oxygen and nitrogen species • Binding the metal ions needed for catalyzing the formation of reactive radicales • Repairing the oxidative damage to the biomolecules or induction of the enzymes that catalyze the repair mechanisms.

  4. Redox Reactions and Biological Action 4 Redox reactions are included not only in inorganic or organic drug analysis but also in the metabolic pathways of drugs. Catabolic pathways: oxidation of C-atoms of carbohydrate or lipid or protein Anabolic pathways: reduction of C-atoms of carbohydrate or lipid or protein Phase-I Metabolic Pathways (Functionalization reactions) I- Oxidative reactions Oxidation of aromatic compounds → hydroxy aromatic compounds Oxidation of olefins → Epioxides Oxidative dealkylation

  5. Redox Reactions and Biological Action 5 Examples for the oxidative dealkylation pathways: N-dealkylation, R-NH-CH3→ R-NH2 + CH2O Deamination, R-CH2-CH2-NH2→ NH3+ R-CO-CH3 N-oxide formation, R3 - N→ R3 N→O N-hydroxylation. R2NH → R2N-OH O-dealkylation. R-O-CH3→ R-OH+ CH2O S-dealkylation, R-S-CH3→ R-SH+ CH2O S-oxidation, R2S→ R2S=O + CH2O Desulphuration R2C=S→ R2C=O

  6. Redox Reactions in Biological System 6 NADPH + flavoproteinreductase RH─P-450-(Fe3+) RH─P-450-(Fe2+) RH─P-450-(Fe2+) RH (Parent drug) P-450-(Fe3+) H2O O2- ROH (oxidized product)

  7. Redox Reactions in Biological System 7 • Example of redox reaction in biological systems ethyl alcohol can be oxidized in liver into acetaldehyde • II- Reductive reactions • Reduction of aldehydes and ketones • Reduction of nitro and azo compounds • Miscellaneous reductive reactions

  8. Properties of Oxidizing Agents 8 1. Potassium permanganate (KMnO4) 2. Potassium dichromate (K2Cr2O7) 3. CericsulphateCe(SO4)2 4. Iodine (I2) 5. Potassium iodate (KIO3) 6. Sodium thiosulphate (Na2S2O3) 7. Potassium bromate(KBrO3) 8. Bromate-bromide mixture (KBrO3/KBr)

  9. 1. Potassium permanganate (KMnO4) 9 • Powerful oxidant E° = 1.52 V. In acid medium: can oxidize: oxalate, Fe2+, Ferrocyanide, As3+, H2O2, and NO2. MnO4 + 8H+ + 5e→ Mn2+ + 4H2O MnO4 + e→ MnO42 In alkaline medium: • H2SO4 is the suitable acid since HCl reacts with it:2KMnO4 + 16HCl → 2 KCl + 2MnCl2 + 5Cl2 + 8 H2O • Zimmermann’s reagent prevents reaction of KMnO4 with Cl. • Not primary standard, may contain MnO2 which catalyses auto-oxidation that is accelerated by light and acids.4MnO4 + 2H2O → 4 MnO2 + 3O2 + 4OH • It is prepared after removal of MnO2 through glass wall (filter paper contains organic matter can decompose permanganate). • Self indicator. • Standardized by: 1- Oxalate. 2- Fe2+. 3- Arsenious oxide.

  10. 2. Potassium dichromate (K2Cr2O7) 10 • Less strong than KMnO4, E°= +1.33 V. • Primary standard (highly pure and stable). • Not easily reduced by organic impurities in water. • Not reduced by HCl in dil. solution. • Not self indicator; since its orange colour is reduced to the green(Cr3+). • Used for determination of Fe2+ (Cl does not interfere); ferroin indicator.

  11. 3. CericsulphateCe(SO4)2 11 • More powerful oxidant than KMnO4 and K2Cr2O7. • Reacts in acid medium, it has 2 E° according the type of acidIn HNO3 medium E°=1.61V. and in H2SO4 E° = 1.45V • In neutral or alkaline medium its hydroxide or basic salt is ppt. • Not a primary standard (difficult to obtain pure) • Not affected by HCl or Cl on cold. • It can be used as self indicator, because its reduction product is colorlessbut it is better to use an internal redox indicator such as ferroin. • Standardized as KMnO4.

  12. 4. Iodine (I2) 12 • Slightly soluble in water and volatile. • KI added to increase solubility and decrease volatility (triiodide). • Titrations must be applied in stoppered flasks • Standard I2 prepared by: • Dissolving it in water after addition of KI (iodide is easily oxidized to iodine, specially in acid medium 4I + O2 + 4H+  2I2 + 2H2O • Mixture of iodate and iodide, solution is very stable in neutral on acidification I2 liberated. IO3 + 5I + 6H+  3I2 + H2O • Standardized by titration with Na2S2O3 using starch as indicator. I2 + 2Na2S2O3 2NaI + Na2S4O6 (sod. tetrathionate) I2 + I→ I3 (triiodide ion)

  13. 5. Potassium iodate (KIO3) 13 • Stronger oxidant than iodine. • It is obtained very pure. • It is more practical to prepare molar solution; due to variations in Eq. wt. according to the medium: • IO3 + 5I + 5H+ = 3I2 + 3H2O (0.1 N HCl) Eq. wt. = MW/5 • IO3 + 2I2 + 6H+ = 5H+ + 3H2O (4-6N HCl)Eq.wt. = MW/4

  14. 14 6. Sodium thiosulphate (Na2S2O3) • Not used as primary standard, its solution is unstable; decomposed by CO2 and certain type of bacteria. • S2O32  SO32 + S • Decomposition avoided by: using recently boiled and cold distilled water and addition of Na2CO3 or borax. • Eq. wt. = M.w./2

  15. 7. Potassiumbromate(KBrO3) 15 • Primary standard purity. • Powerful oxidizing agent in acid medium. Its solution is stable. BrO3 + 6H+ + 6e ↔ Br  + 3H2O E° = 1.44 V • Mainly used : (a) For determination of strong reductants as: As+3, Sn+2, Sb+3 & [Fe(CN)6]4 BrO3 + 3H3AsO3  Br + 3H3AsO4 • At end point, first excess BrO3 reacts with Br   Br2 which decomposes the indicator. (b) As a source of standard bromine solution: • Bromine solution is a mixture of BrO3 and Br when it is acidified gives Br2. BrO3 + 5Br  + 6H+  3Br2 + 3H2O

  16. O H O H B r B r d a r k + 3 B r 2 + 3 H B r B r P h e n o l 2 , 4 , 6 - T r i b r o m o p h e n o l Uses of bromine solution 16 Determination of primary aromatic amines, phenols, quinolines, acetylsalicylic acid, sulphanilic acid, etc. BrO3 + 5 Br  + 6 H+ 3 Br2 + 3 H2O Known ex. standard • The excess Br2 is determined by the reaction with iodide: • Br2 + 2I  I2 + 2 Br  &I2 + 2 Na2S2O3  Na2S4O6 + 2 I • Chloroform is added to dissolve TBP & as indicator.

  17. Uses of bromine solution 17 2- Indirect determination of some cations such as Mg+2, Al+3 Mg+2 + C9H6NOH (oxine) → Mg(C9H6OH)2 + 2H+ The ppt is separated, dissolved in acid Mg(C9H6NO)2 + 2H+→ 2C9H6NOH + Mg+2 2C9H6NOH + 4Br2→ 2C9H5ONBr2 + 4 HBr Known excess standard + 2KI → 2KBr + I2 ≠StandardNa2S2O3 Using starch as ind.

  18. Uses of bromine solution 18 3- Determination of unsaturation: ─CH=CH─ + Br2→ Br ─CH─CH─Br + 2 HBr Known excess standard ≠StandardNa2S2O3 using starch as ind. → 2KBr + I2 + 2KI Determination of unsaturation in oils and fats using bromine-dioxane. DioxaneDioxanetetrabromide

  19. Examples of Applications of Redox Titrations 19 • Free elements • a. Metallic and reduced iron b. Zinc powder c. Sulphur d. Free halogens 2. Determination of peroxides: a. H2O2 b. ZnO2 c. Higher oxides of heavy metals. • 3. Determination of cations A. Determination of iron: 1. Ferrous 2. Ferric 3. Substances that reduce Fe+3 to Fe+2 4. Substances that oxidize Fe+2 to Fe+3 5. Ferro and ferricyanides B. Determination of copper salts C. Determination of HgCl2 D. Determination of cations that form insoluble oxalates.

  20. Examples of Applications of Redox Titrations 20 • 4. Determination of Anions A. Determination of soluble oxalates • B. Determination of sulphide, sulphite, thiosulphate, • sulphate & persulphates. C. Determination of halides, chlorate and hypochlorite. D. Determination of nitrite and nitrate 5. Determination of aldehydes. 6. Determination of moisture content (Karl-Fischer reagent) 7. Bromometric determination of phenolic compounds. • Determination of organic pharmaceutical compounds • 9- Analysis of mixtures.

  21. Free elements 21 a. Metallic iron (Fe) It can be determined by dissolving it in ferric chloride solution and the produced ferrous chloride formed is titrated with standard permanganate in the presence of Zimmermann reagent. Iron oxides do not interfere. ≠ St. KMnO4(self indicator) in the presence of Zimmermann reagent

  22. Free elements (cont.) 22 b. Zinc powder (Zn) is determined by reaction with ferric sulphate and the produced ferrous iron, is acidified with dilute H2SO4 and titrated with standard KMnO4. This method determines the free zinc and not zinc oxide, because the oxide will not reduce ferric salt + H+ ≠ St. KMnO4(self indicator) in the presence of Zimmermann reagent

  23. Free elements (cont.) 23 c. Sulphur (S) Sulphur is converted by refluxing with Na2SO3 to Na2S2O3 which is then titrated with standard iodine solution. I2 + 2Na2S2O3 2NaI + Na2S4O6 The excess Na2SO3, which also reacts with I2, is converted by addition of formaldehyde into formaldehyde bisulphite. ≠ Stand. I2(Starch indicator)

  24. Free elements (cont.) 24 d. Free halogens Iodine can be determined by direct titration with sodium thiosulphate solution. Bromine or chlorine displaces iodine from potassium iodide. I2 + 2Na2S2O3 2NaI + Na2S4O6 (sod. tetrathionate) Br2 + 2KI  I2 + 2KBr Cl2 + 2KI  I2 + 2KCl ≠ Stand. Na2S2O3 (Using starch as indicator)

  25. 2. Determination of peroxides 25 1. Hydrogen peroxide A. as reducing agent By direct  KMnO4. (self indicator) 2MnO4 + 5H2O2 + 6H+→ 2Mn+2 + 5O2 + 3H2O By direct  Ce+4 (using ferroin indicator) 2Ce+4 + H2O2→ 2Ce+3 + O2 + 2H+ B. as oxidizing agent (iodometrically)H2O2 + 2H+ + 2I→ I2 + 2H2O 2. Zinc peroxide ZnO2 + 2H+→ H2O2 + Zn+2 KMnO4or + KI  I2  S2O3-2. ≠ Stand. Na2S2O3 (Using starch as indicator)

  26. 2. Determination of peroxides (cont.) 26 Organic peroxides 1. Carbamide peroxide • Topical antiseptic and disinfectant solution • H2N-CO-NH2..H2O2→ H2N-CO-NH2 + H2O2 • is assayed for H2O2 content iodometrically • 2. Hydrous benzoyl peroxide • Keratolytic and keratogenic agent for acne. • It can be determined iodometrically.(C6H5COO)2 + 2I + 2H+  2C6H5COOH + I2

  27. 3. Determination of oxides 27 1. Higher oxides of manganese and heavy metals as MnO2, PbO2, Pb3O4 a. Iodometrically MnO2 + 4HCl  MnCl2 + Cl2 + 2H2O Cl2+ 2KI  I2  S2O32 b. Indirect titration with reducing agents MnO2 +2Fe 2+ + 4H+ Mn2+ + 2Fe 3+ + 2H2O MnO2 + C2O42- + 4H+ Mn2+ + 2CO2 + 2H2O 2MnO2 + 2AsO33- + 4H+2Mn2+ + 2AsO43- + 2H2O Stand. KMnO4(self indicator) ≠ Known excess standard

  28. 3. Determination of oxides (cont.) 28 Oxides of As or Sb (As2O3, As2O5, Sb2O3 and Sb2O5) • As2O5 and Sb2O5 can oxidize I  I2 or • I2 oxidizes As2O3 andSb2O3  As2O5 and Sb2O5 • depending on the pH of the medium: • In presence of much acid as HCl or HI, • The Eº of As5+/As3+ or Sb5+/Sb3+  (it oxidizes I  I2) As2O3 + 2I2 + 2H2O ← As2O5 + 4I + 4H+ Sb2O3 + 2I2 + 2H2O ← Sb2O5 + 4I + 4H+ • In slightly acid or neutral medium, • The Eºof As5+/As3+ or Sb5+/Sb3+  (I2 oxidizes As2O3 andSb2O3  As2O5 and Sb2O5) (iodometric) ≠ Stand. Na2S2O3 (starch as indicator) (Iodimetric)

  29. 3. Determination of oxides (cont.) 29 Addition of mild alkali (NaHCO3) keep the oxidation going on. As2O3 ≠ 2I2 + 2H2O →As2O5 + 4I + 4H+ Sb2O3 ≠ 2I2 + 2H2O →Sb2O5 + 4I + 4H+ using starch as indicator Na2CO3 or NaOH (pH>8) can not be used as they react with iodine forming hypoiodite and iodate which react with both oxides to form iodinum ions I2 + OH  IO + I 3IO  IO3 + 2Iself oxidation reduction --------→ --------→ NaHCO3 NaHCO3

  30. 3. Determination of oxides (cont.) 30 As2O3 andSb2O3 can be determined by direct reaction with potassium bromate in acid medium  As2O5 and Sb2O5using methyl orange as irreversible redox indicator 3As2O3 ≠ 2BrO3→As2O5 + 2Br  3Sb2O3 ≠ 2BrO3→Sb2O5 + 2Br  BrO3 + 5Br  + 6H+→ 3Br2+ 3H2O Br2+ methyl orange → irreversible oxidation (bleaching color) H+ H+

  31. 3. Determination of oxides (cont.) 31 PbO can be determined by dissolving the sample in glacial acetic acid  lead acetate and precipitating it as lead oxalate PbO + CH3COOH →Pb (CH3COO)2 + 2H+ Pb (CH3COO)2 + H2C2O4→PbC2O4 + 2 CH3COOH Filter and wash the ppt ≠ Stand. KMnO4(self indicator) Known excess standard

  32. 4. Determination of Cations A. Iron 32 I. Ferrous salts:ferrous sulphate, ammonium sulphate (Mohr’s salt), carbonate, ferrous sulphide, Fe+2 KMnO4after addition of Zimmermann’s reagent, self indicator. Fe+2 K2Cr2O7after addition of H3PO4 and internal redox (diphenylamine) or external indicator. Fe+2  Ce+4, irreversible methyl red indicator. Fe+2  I2in presence of F or PO43 using starch

  33. 4. Determination of Cations A. Iron (cont.) 33 II. Ferric salts: by three methods • Ferric (Fe3+)reduced to (Fe2+) by pre-reductants 1- SnCl2 +Fe+3→ SnCl4 + Fe+2 KMnO4 after addition of Zimmermann’s reagent, self indicator SnCl2 + 2HgCl2→ Hg2Cl2  + SnCl4 2- Zn and H2SO4: 2Fe+3+ Znº→ Zn+2 + 2Fe+2  KMnO4 H2SO4 accelerates the reduction by Znºand the unreactedZnº is removed by filtration • 3- Amalgamated Znº(Znº + HgCl2). Excellent reducing agent; (Jones reductor).

  34. 4. Determination of Cations A. Iron (cont.) 34 • Iodometrically: Fe+3/Fe+2 system E°= 0.77 • I2/2I system E°= 0.54 • You must  difference between the 2 systems • either by  I conc. by addition of xss I or  I2 conc. by extraction with immiscible solvent as CHCl3 or CCl4. • Fe+3 + I→ Fe+2 + I2  S2O3-2using starch • Direct titration with titanous chloride • usingmethylene blue (irreversible) orthiocyanate as indicators (the end point is colorless due to disappearance of either the blue color of MB or the red color of Fe(SCN)3 which is formed at the beginning) • FeCl3 ≠ TiCl3 FeCl2 + TiCl4 standard

  35. 4. Determination of Cations A. Iron (cont.) 35 III. Reductants that reduce Fe+3→ Fe+2 : 1- SnCl2 +Fe3+→ SnCl4 + Fe2+ KMnO4 2- Znº + 2Fe3+ → Zn+2 + 2Fe2+  KMnO4 3- Feº + 2Fe3+ → 3Fe2+  KMnO4 IV. Oxidants that oxidize Fe+2→ Fe+3 : 1- K2S2O8 +2Fe2+ + 2H+→ 2Fe3+ + 2KHSO4 2- KClO3+ 6Fe2+ + 6H+ → 6Fe3+ + KCl + 3H2O 3- MnO2+ 2Fe2+ + 4H+ → 2Fe3++ Mn2+ + 2H2O Persulphate Stand. KMnO4 or K2Cr2O7(self indicator) (Ferroinind.) ≠ Chlorate Known excess standard

  36. 4. Determination of Cations A. Iron (cont.) 36 V. Ferrocyanide By direct titration  KMnO4or  Ce4+ 1- 5[Fe(CN)6]4-  MnO4- +8H+→ 5[Fe(CN)6]3-+ Mn2++3H2O (self indicator) 2- [Fe(CN)6]4-  Ce4+→ [Fe(CN)6]3-+ Ce3+ (ferroin indicator) Ferricyanide 1- By iodometrical titration 2[Fe(CN)6]3- + 2I-→ 2[Fe(CN)6]4- + I2  S2O3-2(starch) we add H2SO4 and ZnSO4(to precipitate Zn2[Fe(CN)6]) to ↑↑ the E° of [ferri]/[ferro] to oxidize iodide to iodine. 2- [Fe(CN)6]3- + pre-reductants→ [Fe(CN)6]4-  MnO4- Remove xss e.g. Sulphite or sulphide or sod. peroxide

  37. 4. Determination of Cations B. Determination of copper salts 37 • Iodometrically: Cu+2/Cu+ system E°= 0.15 • I2/2I system E°= 0.54 • Expected • Actually • 2Cu+2 + 4I→ 2CuI2 • The instability of CuI2 results in formation of Cu2I2 ppt which  Cu+→ ↑ E° of Cu+2/Cu+ system to become able to oxidize I → I2 I2 oxidize Cu+ → Cu+2 The reverse occurs due to instability of CuI2 → Cu2I2  + I2  S2O3-2 using starch unstable We add KSCN… Why? • In presence of tartarate or citrate I2 oxidizes Cu+ to cu+2 as these anions form stable complexes with Cu+2

  38. 4. Determination of Cations C. Determination of HgCl2 38 HgCl2 is reduced firstly to Hg° by HCHO in Ca(OH)2 medium. Hg2+ + HCHO + 3OH-  Hg° + HCOO- + H2O Hg° + I2  HgI2  + 2I-  [HgI4]2- known Excess stand colorless complex red ppt ≠ Stand. Na2S2O3 (starch as indicator)

  39. 4. Determination of Cations D. Cations form insoluble oxalates. 39 (Ca2+, Ba2+Sr2+, Mg2+, Cd2+, Bi3+, Zn2+, Ni2+, Co2+& Pb2+) Ca2++ H2C2O4→Ca C2O4 + xss. H2C2O4 (i) The washed precipitate is dissolved in dil. H2SO4 and ≠ KMnO4 at 60°C. or (ii) The excess oxalic acid in the filtrate and washing is back titrated with standard KMnO4 at 60°C. known Excess stand Filter and wash the ppt then follow one of the 2 ways Oxidizing substance such as K2S2O8, KClO3, NO3-, MnO2can be determined by treatment with a known excess oxalic acid and the residual oxalic acid is then back titrated with standard KMnO4.

  40. 5. Determination of Anions A. Soluble oxalates. 40 Soluble oxalates KMnO4 or Ce+4 in presence of sulphuric acid and heating to 60°C. The reaction is slow at start becomes rapid after formation of reduction product (Mn+2 or Ce+3).

  41. Sources of errors in iodine titrations 41 • (A) Errors due to iodine (I2) • I2, being volatile, is subjected to loss. So, iodimetric and iodometric titrations should be carried out in a glass-stoppered flask. • Iodine may undergo changes to hypoiodite and iodide • I2 + H2O  HIO + HI • (B) Errors due to iodide ion • Iodides are easily oxidized by atmospheric oxygen to iodine, specially in acid medium. • 4I- + O2 + 4H+  2I2 + 2H2O • (C) Errors due to starch • Starch may be decomposed by micro-organisms, these decomposition products give non-reversible reddish colour with iodine that masks the true end point.

  42. Sources of errors in iodine titrations 42 Also glucose which may result from hydrolysis of starch can give error as it is reducing agent, Boric acid may be added as a preservative or better starch solution is freshly prepared. (D) Errors due to thiosulphate Thiosulphate is unstable being slowly decomposed with the precipitation of sulphur: S2O32- + H+  HSO3- + S Decomposition may be caused by CO2 (in distilled water) or by bacterial action. Therefore, the solution should be prepared with recently boild and cooled dist. water and Na2CO3 is added to prevent bacterial activity (pH = 9-10).

  43. 5. Determination of Anions b. S2-, SO32-, S2O32-,SO42- & S2O82- . 43 • Direct titration • SulphideS2 ≠ I2 → 2I + Sº  • Use dilute soln. to decrease the inclusion of I2 by Sº • Thiosulphate2S2O32 ≠ I2 → S4O62 + 2I • Back titration. • 3.SulphiteSO32 + I2 + H2O → SO42 + 2I + 2H+ • S2 + I2 → 2I + Sº  • 4.SulphateSO42 + BaCrO4→ BaSO4 + CrO42(filtrate) • 2CrO42 (filtrate) + 2H+  Cr2O72 + H2O St. S2O32(starch indicator) ≠ Known xss. standard Iodometrically +2KI  I2  S2O32 (starch)

  44. 5. Determination of Anions C. Determination of halides, ClO3 and ClO 44 Chlorates (ClO3) & Hypochlorites (ClO) Iodometrically Chlorate Hypochlorite ≠ Stand. Na2S2O3 (starch as indicator)

  45. 5. Determination of Anions C. Determination of halides (cont.) 45 • e.g.1: Chlorine in bleaching powder + 4H+  Cl2 + Ca2+ + 2H2O • Acetic acid used for acidification not HCl, since calcium chlorate present as a result of hypochlorite decomposition will react slowly with I and liberate iodine. • e.g.2: N-chloro-organic compounds as water disinfectant, these in water release hypochlorous acid (HOCl) which is the active germicidal species. ClO + 2I + 2H+  Cl + I2 + H2O (iodometrically) (calcium hypochlorite Ca(OCl)2 + basic chloride). CaCl2.Ca(OH)2.H2O

  46. 5. Determination of Anions C. Determination of halides (cont.) 46 • Organically combined iodine compound • Thyroxine, diiodohydroxyquinoline & chiniofonThese are converted into iodide by: • A.Iodoxyl • B. Thyroxine • I + 3Br2 + 3H2O → IO3 + 6Br + 6H+ ZnI2 filter, cool ≠ IO3- I Znº + gl. HOAc Reflux ignition at 700°C K2CO3 xss Br2 removed by phenol. + xss I  I2  S2O32 C.Oxygen flask combustion method Iodide determined iodometrically.

  47. Oxygen flask combustion method 47 • The compound is wrapped in a filter paper attached to pt. wire and sealed in the stopper of a special oxygen filled flask containing dilute Na2S2O5 (sodium metabisulphite) solution. • Combustion is complete within 30 sec at 1200°C. • The resulting iodine is absorbed (reduced) by Na2S2O5 forming iodide

  48. 5. Determination of Anions d. Determination of nitrites and nitrates 48 Nitrites nitrite sample is added to an excess of acidified KMnO4 soln. (the tip of the pipette containing the nitrite solution should be below the surface of the liquid during addition in order to stop volatility of the unstable nitrous acid, liberated from acidified nitrite solution). Stand. KMnO4 or K2Cr2O7(self indicator) (Ferroinind.) Stand. Fe2+(self indicator) ≠ ≠ Nitrates Known excess standard Known excess standard

  49. Determination of Aldehydes 49 • e.g. formalin (formaldehyde), glucose, and lactose. • 2R-CHO + I2 + 2NaOH → 2R-COOH + 2 NaI • Sucroseafter acid hydrolysis (inversion) of sugar • C12O22O11 + H2O → H+→ C5H11O5-CHO + C5H11O5CO • (glucose) (fructose) I2 + 2NaOH → NaOI + NaI + H2OC5H11O5-CHO + IO→ C5H11O5-COOH + I • Change in oxidation number = 2 (Glucose  Sucrose) ≠ St. S2O32(starch indicator) Kn. Xss. St. └→ acidification → I2 Gluconic acid

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