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Preadjustment of analyte oxidation state

). 2. –. Preoxidation : Peroxydisulfate ( (NH 4 ) 2 S 2 O 8 ). Sodium bismuthate ( NaBiO. ). 3. Hydrogen peroxide (H2O2). Preadjustment of analyte oxidation state.

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Preadjustment of analyte oxidation state

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  1. ) 2 – Preoxidation : Peroxydisulfate ( (NH4)2S2O8 ) Sodium bismuthate ( NaBiO ) 3 Hydrogen peroxide (H2O2) Preadjustment of analyte oxidation state It is necessary to adjust the oxidation state of the analyte to one that can be titrated with an auxiliary oxidizing or reducing agent. Ex. Preadjustment by auxiliary reagent Fe(II), Fe(III) Fe(II) – 4 Titration Ce4+ Prereduction : Stannous chloride ( SnCl2) Chromous chloride Jones reductor (zinc coated with zinc amalgam) Walden reductor ( solid Ag and 1M HCl) http:\\asadipour.kmu.ac.ir 33 slides

  2. Reagents used in redox titration Reducing agents 1) ammonium iron(II) sulfate hexahydrate (Mohr’s salt) FeSO4(NH4)2SO4·6H2O 2) iron(II) ethylene diamine sulfate (Oesper’s salt) FeC2H4(NH3)2(SO4)2·4H2O 3) Sodium thiosulfate pentahydrate Na2S2O3·5H2O 4) Arsenic trioxide: arsenious oxide As2O3 5) Sodium oxalate and oxalic acid dihydarte Na2(COO)2 , (COOH)2·2H2O http:\\asadipour.kmu.ac.ir 33 slides

  3. Sodium thiosulfate, Na2S2O3 Thiosulfate ion is a moderately strong reducing agent that has been widely used to determine oxidizing agents by an indirect procedure that involves iodine as an intermediate. With iodine, thiosulfate ion is oxidized quantitatively to tetrathionate ion according to the half-reaction: 2S2O3 2–  S4O6 2– + 2e Eo = 0.08 Ex. 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 (-amylose) http:\\asadipour.kmu.ac.ir 33 slides

  4. Standardization of thiosulfate solution: Primary standard : potassium iodate (KIO3), K2Cr2O7, KBrO3 Titration reactions: KIO3 + 5KI + 6HCl  3I2 + 6KCl + 3 H2O I2 + 2Na2S2O3 2NaI + Na2S4O6 KIO3  3I2  6Na2S2O3·5H2O  6 Equivalent S2O32- +H+⇋ HSO3- +S(s) pH, Microorganisms, Concentration, Cu2+, Sunlight Stabilizer for sodium thiosulfate solution : Na2CO3 Na2S2O3 + H2O + CO2 Na2CO3 + H2S2O3 H2S2O3  H2SO3 + S http:\\asadipour.kmu.ac.ir 33 slides

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  6. or 16-2 Finding the end point A redox indicator is a compound that changes color when it goes from its oxidized to its reduced state. For ferroin, with E° = 1.147 V we expect the color change to occur in the approximate range 1.088 V to 1.206 V with respect SHE http:\\asadipour.kmu.ac.ir 33 slides

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  8. Starch is the indicator of choice for those procedures involving iodine because it forms an intense blue colour with iodine. Starch is not a redox indicator; it responds specifically to the presence of I2, not to a change in redox potential. Starch-Iodine Complex Structure of the repeating unit of the sugar amylose. http:\\asadipour.kmu.ac.ir 33 slides

  9. Arsenious oxide, As4O6 As4O6 + 6H2O = 4H3AsO3 H3AsO3 + I3– + H2O = H3AsO4 + 3I– + 2H+ http:\\asadipour.kmu.ac.ir 33 slides

  10. Reagents used in redox titration Oxidizing agents 1) Potassium permanganate KMnO4 : Permanganometry 2) Ceric sulfate / Ceric ammonium sulfate Ce(SO4)2·2(NH4)2SO4·4H2O : Cerimetry 3) Potassium dichromate K2Cr2O7 : Dichrometry 4) Iodine I2 : Iodimetry, Iodometry 5) Potassium iodate KIO3 : Iodatimetry 6) Potassium bromate KBrO3 : Bromatimetry http:\\asadipour.kmu.ac.ir 33 slides

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  12. Permanganate titration Oxidation with permanganate : Reduction of permanaganate KMnO4 Powerful oxidant that the most widely used. 1) In strongly acidic solutions (1M H2SO4 or HCl, pH  1) MnO4–+ 8H+ + 5e = Mn2 + + 4H2 O Eo = 1.51 V KMnO4 is a self-indicator. 2) In feebly acidic, neutral, or alkaline solutions MnO4– + 4H+ + 3e = MnO2 (s) + 2H2 O Eo = 1.695 V 3) In very strongly alkaline solution (2M NaOH) MnO4– + e = MnO42 – Eo = 0.558 V   http:\\asadipour.kmu.ac.ir 33 slides

  13. Permanganate titration Duration of colour in end point (30 seconds) MnO4– + 3Mn2+ + 2H2O  5MnO2 + 4H+ K=1*1047 Stability of aqoues solution of MnO4- MnO4– + 2H2O  4MnO2 (s) + 3O2 (g) +4OH-   http:\\asadipour.kmu.ac.ir 33 slides

  14. Standardization of KMnO4 solution Potassium permanganate is not primary standard, because traces of MnO2 are invariably present. Standardization by titration of sodium oxalate (primary standard) : 2KMnO4 + 5 Na2(COO)2 + 8H2SO4 = 2MnSO4 + K2SO4 + 5Na2SO4 + 10 CO2 + 8H2O 2KMnO4 5 Na2(COO)2  10 Equivalent http:\\asadipour.kmu.ac.ir 33 slides

  15. 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. • Manganesse dioxide 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. • Filter the liquid through a sintered-glass filter crucible to remove solid MnO2. • Transfer the filtrate to a clean stoppered bottle freed from grease with cleaning mixture. • Protect the solution from evaporation, dust, and reducing vapors, and keep it in the dark or in diffuse light. • If in time managanese dioxide settles out, refilter the solution and restandardize it. http:\\asadipour.kmu.ac.ir 33 slides

  16. http:\\asadipour.kmu.ac.ir 33 slides

  17. Applications of permanganometry • H2O2 • 2KMnO4 + 5 H2O2 + 3H2SO4 = 2MnSO4 + K2SO4 + 5O2 + 8H2O • (2) NaNO2 • 2NaNO2 + H2SO4 = Na2SO4 + HNO2 • 2KMnO4 + 5 HNO2 + 3H2SO4 = 2MnSO4 + K2SO4 + 5HNO3 + 3H2O • (3) FeSO4 • 2KMnO4 + 510 FeSO4 + 8H2SO4 = 2MnSO4 + K2SO4 + 5Fe2(SO4)3 + 8H2O • (4) CaO • CaO + 2HCl = CaCl2 + H2O • CaCl2 + H2C2O4 = CaC2O4 + 2HCl (excess oxalic acid) • 2KMnO4 + 5 H2C2O4 + 3H2SO4 = 2MnSO4 + K2SO4 + 10CO2 + 8H2O (back tit) • (5) Calcium gluconate • [CH2OH(CHOH)4COO]2Ca + 2HCl = CaCl2 + 2CH2OH9CHOH)4COOH • (NH4)2C2O4 + CaCl2 = CaC2O4 + 2 NH4Cl • CaCl2 + H2SO4 = H2C2O4 + CaSO4 • 2KMnO4 + 5 H2C2O4 + 3H2SO4 = 2MnSO4 + K2SO4 + 10CO2 + 8H2O http:\\asadipour.kmu.ac.ir 33 slides

  18. Oxidation with Ce4+ Ce4+ + e = Ce3+ 1.7 V in 1 N HClO4 yellow colorless1.61 V in 1N HNO3 1.47 V in 1N HCl 1.44 V in 1M H2SO4 Indicator : ferroin, diphenylamine Preparation and standardization: Ammonium hexanitratocerate, (NH4)2Ce(NO3)6, (primary standard grade) Sodium oxalate. http:\\asadipour.kmu.ac.ir 33 slides

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  20. O OH CH3 CH3 O OH Applications of cerimetry (1) Menadione (2-methylnaphthoquinon: vitamin K3) HCl, Zn Reduction 2 Ce(SO4)2 • Iron • 2FeSO4 + 2 (NH4)4Ce(SO4)4 = Fe2(SO4)3 + Ce2(SO4)3 + 4 (NH4)2SO4 http:\\asadipour.kmu.ac.ir 33 slides

  21. Oxidation with potassium dichromate Cr2O72– + 14H+ + 6e = 2Cr3+ + 7H2O Eo = 1.36 V K2Cr2O7 is a primary standard. Indicator : diphenylamine sulphonic acid http:\\asadipour.kmu.ac.ir 33 slides

  22. Ex. Redox titration( hydroquinone vs dichromate standard solution ) Cr2O72– + 14H+ + 6e  2 Cr3+ + 7 H2O Eo= 1.33 HO OH  O O + 2H+ + 2e Eo= 0.700 3 3 HO OH + Cr2O72– + 8H+ 3 O O + 2 Cr3+ + 7 H2O Eo= Eocathode – Eoanode = 1.33 – 0.700 = 0.63 V K = 10 nEo/0.05916= 10 6(0.63) / 0.05916 = 10 64 redox indicator : diphenylamine colorless to violet Very large : quantitative : complete reaction http:\\asadipour.kmu.ac.ir 33 slides

  23. Iodimetry and iodometry • 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. I2 + V.C→ 2I- + …… I- + Cu2+→ I2+ Cu+ I2 + S2O32- → 2I- + S4O62- http:\\asadipour.kmu.ac.ir 33 slides

  24. standard I3- 1) Iodine only dissolves slightly in water. Its solubility is enhanced by interacting with I- 2) 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 I3- by quantitative reverse disproportionation: 3) Cu2++4I- 2CUI + I2 http:\\asadipour.kmu.ac.ir 33 slides

  25. Stability of I2 Solutions • In 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. I2 + OH-⇌ IO- + I- + H+ 3IO-⇌ IO3- + 2I- http:\\asadipour.kmu.ac.ir 33 slides

  26. Iodimetry http:\\asadipour.kmu.ac.ir 33 slides

  27. iodometry http:\\asadipour.kmu.ac.ir 33 slides

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  29. Bromatimetry BrO3– + 5Br– + 6H+ 3Br2 + H2O 2I– + Br2 I2 + 2Br– I2 + 2 S2O32– 2I–+ S4O62– http:\\asadipour.kmu.ac.ir 33 slides

  30. Addition reactions http:\\asadipour.kmu.ac.ir 33 slides

  31. Determining water with the Karl Fisher Reagent The Karl Fisher reaction : I2 + SO2 + 2H2O  2HI + H2SO4 For the determination of small amount of water, Karl Fischer(1935) proposed a reagent prepared as an anhydrous methanolic solution containing iodine, sulfur dioxide and anhydrous pyridine in the mole ratio 1:3:10. The reaction with water involves the following reactions : C5H5N•I2 + C5H5N•SO2 + C5H5N + H2O  2 C5H5N•HI + C5H5N•SO3 C5H5N+•SO3–+ CH3OH C5H5N(H)SO4CH3 Pyridinium sulfite can also consume water. C5H5N+•SO3–+ H2O C5H5NH+SO4H– It is always advisable to use fresh reagent because of the presence of various side reactions involving iodine. The reagent is stored in a desiccant-protected container. The end point can be detected either by visual( at the end point, the color changes from dark brown to yellow) or electrometric, or photometric (absorbance at 700nm) titration methods. The detection of water by the coulometric technique with Karl Fischer reagent is popular. http:\\asadipour.kmu.ac.ir 33 slides

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