1 / 32

Reduction- Oxidation Reactions 5th lecture

Reduction- Oxidation Reactions 5th lecture. Ceric as titrant: Ce 4+. Ceric as titrant: Ce 4+. Properties. Ce 4+ salts are strong oxidants in H 2 SO 4 Ce 4+ + e  Ce 3+ Yellow Colorless .

lane-lowery
Download Presentation

Reduction- Oxidation Reactions 5th lecture

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Reduction- Oxidation Reactions • 5thlecture

  2. Ceric as titrant: Ce4+

  3. Ceric as titrant: Ce4+ Properties • Ce4+ salts are strong oxidants in H2SO4 • Ce4+ + e  Ce3+ • Yellow Colorless • Although it could be used as self indicator it is preferable to use ferroin as indicator especially in case of det. of ferrous salts. • Ce4+ cannot be used in neutral or alkaline solution due to • hydrolysis to hydrated ceric oxide • They have wide range of oxidising power but they don’t oxidise HCl even in presence of Fe2+ salts • Ce4+ forms more stable complexes than Ce3+ • Ceric salts are much more stable than MnO4-

  4. Ceric as titrant: Ce4+ Preparation and standardization of Ce4+ solu Prepared from primary standard Ce(NO3)6 (NH4)2 in conc H2SO4 or in 72% HClO4. If using other salts it should be standardized • Against arsenious trioxide: • 2Ce4+ + H3 AsO3 + H2O  2Ce3+ + H3AsO4+ 2H+ • (2) Against oxalate • In both cases , the reaction is slow it requires heat • to 50°C, using ICl as catalyst and ferroin indicator 2Ce4++ H2C2O4 ↔ 2Ce3+ + 2CO2 + 2H+

  5. Ceric as titrant: Ce4+ Applications • Direct titrations: • determination of reducing agents • Fe2+, AsO33- , C2O42-, H2O2, I-, • Fe(CN)64- using ferroin indicator • Color change from red to pale blue • [ Fe (CN)6]4-+ Ce4+  Ce3++ [ Fe (CN)6]3- • Advantages: • Better than MnO4- as it is less subject to interference of organic matter • It is preferable to be used instead of MnO4- in the determination of Fe2+ since we can use HCl. H2O2 + 2Ce4+  2Ce3+ + 2H+ + O2

  6. Ceric as titrant: Ce4+ Applications of Ce4+ (b) Back titrations: Determination of polyhydroxy alcohols, aldehydes, hydroxy acids. example: glycerol, citric acid C3H8O3+8Ce4++3H2O  3HCOOH+8Ce3++8H+ The excess Ce4+ is titrated against sodium oxalate or AsO33- using ICl as catalyst and ferroin as indicator at 50oC.

  7. Potassium dichromate as titrant

  8. Potassium dichromate as titrant Properties • It is a primary standard due to the stability of its solution and is obtainable in high purity • Its oxidation potential is lower than KMnO4 and Ce4+ so it is limited in use • It does not oxidise Cl- into Cl2, oxalic acid , ferrocyanide Its main application is the direct and indirect determination of Fe2+ ion

  9. Potassium dichromate as titrant • It can not serve as a self indicator reagent Cr2O72-(Orange) + 14H+ + 6e  2Cr3+ (green) + 7H2O • Many redox indicators are unsuitable: • because of their high oxidation potential, and • because of the deep green colour of Cr3+ which • causes the colour change of the indicator to be • less clear • The indicators usually used are: • diphenyl amine sulphonic acid. • 4,7-dimethyl 1, 10 phenanthroline ferrous.

  10. Potassium dichromate as titrant Applications 1-Determination of Fe2+ Iron (internal indicator) E° Fe3+/Fe2+ 0.77 E° diphenylamine 0.76 E° ferroin 1.06 Is there need for H3PO4?? H3PO4 Titrate with Cr2O72- E0 = 1.33 v diphenyl amine H2SO4 Fe2+ Fe3+ Role of H3PO4 or F-: • decreases the Fe3+/Fe2+ system potential so that Fe2+ ion will • be oxidized before the indicator • and to remove the dark colour of Fe3+ ion giving a more clear • colour change.

  11. Potassium dichromate as titrant Applications • 1-Determination of Fe2+ Iron external indicator • Ferricyanide:Fe2+ is titrated with dichromate in acidic medium.Occasionally remove a drop from the solution and add it to ferricyanide solu. a blue color of ferrous ferricyanide is formed. At the E.P. No more Fe2+ is present so no blue color is formed. • Diphenylcarbazide: After oxidation of Fe2+ to Fe3+, the first exx of dichromate oxidizes the indicator and gives a red color.

  12. Applications Potassium dichromate as titrant 2- Determination of some oxidising agents Add a measured exx of Fe2+ ion and back titrate the exx. using Cr2O72- and diphenylamine as indicator.

  13. Potassium dichromate as titrant Applications 3-reducing agents 4-Organic compd 5-Pb2+ Na2SO3 glycerol PbO Addmeasured exx of Cr2O72- in presence of: Sulphuric acid Sulphuric acid glacial HAC ICl as catalyst The excess dichromate is titrated iodometrically Cr2O72- + 6I- + 14 H+  2Cr3+ + 3I2 + 7 H2O 3 SO32- + Cr2O72- + 8H+  3 SO42- + 2Cr3+ + 4H2O 3C3H8O3 + 7 Cr2O72- + 56 H+  14Cr3+ + 9 CO2 + 40H2O 2Pb2+ + Cr2O72- + H2O  2PbCrO4↓ (ppt)+ 2H+

  14. Iodine as oxidant

  15. Properties: Iodine as oxidant The iodine/iodide half reaction is I2 + 2e  2I- (Eo = +0.535V) I- can be oxidized by systems I2 can oxidize systems of of higher oxidation potential lower oxidation potential MnO4-/Mn2+ Sn4+/Sn2+ Cr2O72-/Cr3+ S4O62-/S2O32- ClO3-/Cl- S/S2- ↑ E° ↓ E° • Iodometric method • Indirect titration • Add KI to oxidizing agents, • equivalent I2 is libarated and • titr with Na2S2O3 • To determine oxidizing agents • Iodimetric method • Direct titration with I2 • To determine reducing • agents

  16. Properties: Iodine as oxidant • Systems having oxidation potentials near to that of iodine/iodide e.g AsO43-/AsO33-, Fe3+/Fe2+ Their reactionswith Iodine is directed forward or backword by control of experimental conditions. i.e. Change in oxidation potential 1-the pH of the medium 2-addition of complexing agents 3-addition of precipitating agents

  17. Factors affecting the potential of I2/I- system: 1-Effect of pH: The potential of: AsO43-/AsO33-= +0.57 I2/2I- = +0.54 To determine arsenite sample using Iodine the pH of the solution should be adjusted to 8.3 by adding NaHCO3 I2 + AsO33- + H2O  2I- + AsO43- + 2H+ E AsO43-/ AsO33-=Eo – 0.059 / 2 log [AsO33- ] / [AsO43-][H+]2 • ↓ [H+] by addition of NaHCO3↓the oxidation • potential of AsO43- / AsO33- system. • NaHCO3 reacts with H+ giving CO2 and H2O shifting the reaction • to the right and prevent reversibility. • At higher pH if using NaOH, I2 reacts with OH- producing OI- so • consuming more I2. Also OI- has oxidizing properties which • differ than I2.

  18. Iodine as oxidant 0.059 2 2-Effect of Complexing agents: I2 + 2 e  2I- E= = Eo - Log [I-]2 / [ I2] • When HgCl2 is added to the I2/I- system it forms [HgI4]2-Thus: • removing the I- ions from the share of the reaction, • minimizing its concentration, • increasing the ratio of I2 / [I-]2 • increasing the oxidation potential of I2 /2I- system • So I2 could determine AsO33-.

  19. Another example Iodine as oxidant E° Fe3+/Fe2+= 0.77V E° I2/ 2I- = 0.54V Fe3+ + e  Fe2+ E Fe2+ / Fe3+ = 0.559 - How to determine Ferrous salts using Iodine? • When pyrophosphate, EDTA or F- is added to the Fe3+/Fe2+ system it form [FeF6]3- or [Fe(PO4)6]3- • Thus: • removing the Fe3+ ions from the share of the reaction, • minimizing its concentration, • decreasing the ratio of Fe3+/ Fe2+ • lowering the oxidation potential of Fe3+/ Fe2+ below that of I2/2I- system.

  20. Iodine as oxidant 3- Effect of precipitating agents • Fe(CN)63- + e  Fe (CN)64- • minimizing conc of ferrocyanide • increasing ferri/ferro potential • So Ferri/Ferro system can oxidize I- to I2 E° Ferri/Ferro= 0.36V E° I2/ 2I- = 0.54V To determine [Fe(CN)6]3- ion iodometrically;Zn2+ should be present: it precipitate Zn2[ Fe(CN)6] ion E = Eo -

  21. Iodine as oxidant How to determine Cu2+ salts using KI ? • E° Cu2+/Cu+ = 0.46 • E° I2/2I- = 0.54 • It is expected that I2 oxidizes Cu+ (cuprous), however, Cu2+ (cupric) oxidizes I- • Procedure: Cu2+ is treated with KI and the liberated I2 is titrated with S2O32- • 2Cu2+ + 4I-  I2 + Cu2I2 ↓ • The precipitation of Cu2I2 increases the oxidation potential of • Cu2+ /Cu+ • E = E0 - 0.059 log [Cu+] • 1 [Cu2+]So Cu2+ oxidizes I- to I2 • I2 tends to be absorbed on Cu2I2 so the reaction with S2O32- is incompltete so add SCN- near the end point to form Cu2(SCN)2 which has no tendency to adsorb I2.

  22. To reverse the reaction i.e. To allow iodine to oxidize cuprous. Add tartarate or citrate which forms with cupric a stable complex so decreasing the oxidation potential of Cu2+/Cu+ E = E0 - 0.059 log Cu+ 1 Cu+2

  23. Iodine as oxidant Titration methods: Since iodine may be either reduced or produced by oxidation Direct Iodimetric method Indirect Iodometric method Iodine for determination of reducing agents I- is added to oxidizing agents,the librated I2 is titr. with Na2S2O3 Titrating agent Added near the end of titr (when the brown color of I2 becomes pale) Added at the beginning of titr. Indicator (Starch) permanent blue color disappearance of blue color E.P.

  24. Iodine as oxidant Iodine Na2S2O3 Na2S2O3 oxidant + KI→I2 Reductant + starch E.P. Add starch Colorless E.P.

  25. Iodine as oxidant Detection of the end point in iodine titrations: 1- The use of starch: • Starch is used in the form of colloidal Solu giving a deep blue adsorbtion complex with traces I2 • In exx I2 an irreversible blue adsorption complex is formed which is not changed • Starch consists of  amylase and amylopectin I2 gives blue adsorption complex with  amylase. • In strong acid medium: starch hydrolyses giving products which give with iodine non reversible reddish color masking the end point change.

  26. Iodine as oxidant Detection of the end point in iodine titrations: 1- The use of starch: • Starch can not be used in alcoholic solu.because alcohol hinders the adsorption of I2 on starch • The sensitivity of the blue color decreases with temperature due to gelatinization of starch and volatility of Iodine • Starch indicator solution must be freshly prepared when it stands decomposition takes place and its sensitivity is decreased. A preservative can be added

  27. Iodine as oxidant Detection of the end point in iodine titrations: 2- Use of organic solvent (CHCl3 or CCl4) • In presence of alcohol or conc acids, organic solvents are recommended as indicators. • These solvents dissolve iodine to give intensely • coloured purple solution, so that a trace of I2 gives an • intense colour, and the end point will be the appearance • Or disappearance of the colour in the organic solvent layer. • I2 is soluble in CHCl3 or CCl4 90 times more than in H2O • It is important that the mixture be shaken well near the end point in order to equilibrate the iodine between the aqueous and organic phases to enable aqueous S2O32- to react with I2 in CHCl3

  28. Sources of error in iodimetry Iodine as oxidant • A- Error due to I2: • I2 is volatile especially at high temp and at a low • Conc of I- ion so: ●Use stoppered glass containers • ●Avoid elevated temp & cool during titratn • ●Moisten the stopper with I- • I-+I2→ I3- (triiodide) less volatile and more stable • (2) I2 conc is changed if the solution gets in contact with • rubber, organic matter, dust, SO2, H2S • (3) I2 may undergo disproportionation into HOI and I- • I2 + H2O  HOI + I- + H+ • To overcome this difficulty the solution may be acidified • to shift the reaction to the left.

  29. Sources of error in iodimetry Iodine as oxidant • B- Error due to I- ion: • I- ion is liable to atmospheric oxidation. • This is catalysed by light, heat, Cu2+, NO gas • The medium must be completely free from O2 so introduce CO2 (add little NaHCO3). • In titration which needs standing for time, standing should be away from light. • If we need acid medium, never use HNO3,it contains nitrous oxide. 4H+ + 4I- + O2  2I2 + 2H2O

  30. Sources of error in iodimetry Iodine as oxidant C- Error due to S2O32- ion: • Thiosulphate is affected by pH, the most favourable pH is 7 till pH 9 • Under these conditions:S2O32- is oxidized to S4O62-,where every 2 S2O32- is oxidized by 1 I2to S4O62- (tetrathionate) 2S2O32-+I2 S4O62-+2I- • Under acidic conditions: thiosulphate is changed to bisulphite (HSO3-) with the precipitation of S. Every 2 HSO3- is oxidized by 2 I2to 2HSO4- • Therefore, The consumed I2 in acid medium is double that consumed in neutral medium.

  31. Sources of error in iodimetry Iodine as oxidant C- Error due to S2O32- ion: • In pH>9:I- changes to IO- (hypoiodite) oxidizing S2O32- to SO42- which is an incomplete reaction • Thiosulphate is decomposed during storage by thiobacteria, so: ●boiling water is used as a solvent, ●preservatives e.g. sodium benzoate, CHCl3, or HgI2 may be added. ●The pH is adjusted by adding borax, Na2CO3 or NaHCO3 to about pH 9 which inhibits bacterial action.

  32. Sources of error in iodimetry Iodine as oxidant D- Error due to starch: Starch may be decomposed by microorganisms into products e.g. glucose causes error due to its reducing action other products gives nonreversible reddish color with I2 which masks the true end point. To avoid this, preservatives e.g. H3BO3 and formamide are added.

More Related