Reduction- Oxidation Reactions 5th lecture

Ceric as titrant: Ce 4+

 Although it could be used as self indicator it is preferable to use ferroin as indicator especially in case of det. of ferrous salts. Ceric as titrant: Ce 4+ Properties  Ce 4+ salts are strong oxidants in H 2 SO 4 Ce 4+ + e  Ce 3+ Yellow Colorless  They have wide range of oxidising power but they don’t oxidise HCl even in presence of Fe 2+ salts  Ce 4+ cannot be used in neutral or alkaline solution due to hydrolysis to hydrated ceric oxide  Ceric salts are much more stable than MnO 4 -  Ce 4+ forms more stable complexes than Ce 3+

Ceric as titrant: Ce 4+ Preparation and standardization of Ce 4+ solu Prepared from primary standard Ce(NO 3 ) 6 (NH 4 ) 2 in conc H 2 SO 4 or in 72% HClO 4. If using other salts it should be standardized (1) Against arsenious trioxide: 2Ce 4+ + H 3 AsO 3 + H 2 O  2Ce 3+ + H 3 AsO 4 + 2H + (2) Against oxalate In both cases, the reaction is slow it requires heat to 50°C, using ICl as catalyst and ferroin indicator 2Ce 4+ + H 2 C 2 O 4 ↔ 2Ce CO 2 + 2H +

Ceric as titrant: Ce 4+ (a)Direct titrations: determination of reducing agents Fe 2+, AsO 3 3-, C 2 O 4 2-, H 2 O 2, I -, Fe(CN) 6 4- using ferroin indicator Color change from red to pale blue [ Fe (CN) 6 ] 4- + Ce 4+  Ce 3+ + [ Fe (CN) 6 ] 3- Advantages:  Better than MnO 4 - as it is less subject to interference of organic matter  It is preferable to be used instead of MnO 4 - in the determination of Fe 2+ since we can use HCl. Applications H 2 O 2 + 2Ce 4+  2Ce H + + O 2

Ceric as titrant: Ce 4+ Applications of Ce 4+ (b) Back titrations: Determination of polyhydroxy alcohols, aldehydes, hydroxy acids. example: glycerol, citric acid C 3 H 8 O 3 +8Ce 4+ +3H 2 O  3HCOOH+8Ce 3+ +8H + The excess Ce 4+ is titrated against sodium oxalate or AsO 3 3- using ICl as catalyst and ferroin as indicator at 50 o C.

Potassium dichromate as titrant

 It is a primary standard due to the stability of its solution and is obtainable in high purity  Its oxidation potential is lower than KMnO 4 and Ce 4+ so it is limited in use  It does not oxidise Cl - into Cl 2, oxalic acid, ferrocyanide Its main application is the direct and indirect determination of Fe 2+ ion Properties

 Many redox indicators are unsuitable: because of their high oxidation potential, and because of the deep green colour of Cr 3+ 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. Potassium dichromate as titrant  It can not serve as a self indicator reagent Cr 2 O 7 2- (Orange) + 14H + + 6e  2Cr 3+ (green) + 7H 2 O

Potassium dichromate as titrant Applications 1-Determination of Fe 2+ Iron (internal indicator) diphenyl amine H 2 SO 4 Fe 2+ Titrate with Cr 2 O 7 2- E 0 = 1.33 v Fe 3+  decreases the Fe 3+ /Fe 2+ system potential so that Fe 2+ ion will be oxidized before the indicator  and to remove the dark colour of Fe 3+ ion giving a more clear colour change. Role of H 3 PO 4 or F - : E° Fe3+/Fe E° diphenylamine 0.76 H 3 PO 4 E° ferroin 1.06 Is there need for H 3 PO 4 ??

Potassium dichromate as titrant Applications 1-Determination of Fe 2+ Iron external indicator  Ferricyanide:Fe 2+ 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 Fe 2+ is present so no blue color is formed.  Diphenylcarbazide: After oxidation of Fe 2+ to Fe 3+, the first exx of dichromate oxidizes the indicator and gives a red color.

2- Determination of some oxidising agents Add a measured exx of Fe 2+ ion and back titrate the exx. using Cr 2 O 7 2- and diphenylamine as indicator. Potassium dichromate as titrant Applications

Potassium dichromate as titrant 3-reducing agents 4-Organic compd 5-Pb 2+ Na 2 SO 3 glycerol PbO Add measured exx of Cr 2 O 7 2- in presence of: Sulphuric acid Sulphuric acid glacial HAC ICl as catalyst The excess dichromate is titrated iodometrically 3 SO Cr 2 O H +  3 SO Cr H 2 O Cr 2 O I H +  2Cr I H 2 O Applications 3C 3 H 8 O Cr 2 O H +  14Cr CO H 2 O 2Pb 2+ + Cr 2 O H 2 O  2PbCrO 4 ↓ (ppt)+ 2H +

Iodine as oxidant

The iodine/iodide half reaction is I 2 + 2e  2I - (E o = V) I - can be oxidized by systems I 2 can oxidize systems of of higher oxidation potential lower oxidation potential MnO 4 - /Mn 2+ Sn 4+ /Sn 2+ Cr 2 O 7 2- /Cr 3+ S 4 O 6 2- /S 2 O 3 2- ClO 3 - /Cl - S/S 2- Iodine as oxidant Properties: ↑ E°↓ E° Iodometric method Indirect titration Add KI to oxidizing agents, equivalent I 2 is libarated and titr with Na 2 S 2 O 3 To determine oxidizing agents Iodimetric method Direct titration with I 2 To determine reducing agents

 Systems having oxidation potentials near to that of iodine/iodide e.g AsO 4 3- /AsO 3 3-, Fe 3+ /Fe 2+ Their reactions with Iodine is directed forward or backword by control of experimental conditions. i.e. Change in oxidation potential Iodine as oxidant Properties: 1-the pH of the medium 2-addition of complexing agents 3-addition of precipitating agents

1-Effect of pH: The potential of: AsO 4 3- /AsO 3 3- = I 2 /2I - = To determine arsenite sample using Iodine the pH of the solution should be adjusted to 8.3 by adding NaHCO 3 I 2 + AsO H 2 O  2I - + AsO H + E AsO4 3- / AsO3 3- =E o – / 2 log [ AsO 3 3- ] / [ AsO 4 3- ][H + ] 2  ↓ [H + ] by addition of NaHCO 3 ↓ the oxidation potential of AsO 4 3- / AsO 3 3- system.  NaHCO 3 reacts with H + giving CO 2 and H 2 O shifting the reaction to the right and prevent reversibility.  At higher pH if using NaOH, I 2 reacts with OH - producing OI - so consuming more I 2. Also OI - has oxidizing properties which differ than I 2. Factors affecting the potential of I 2 /I - system:

2-Effect of Complexing agents: Iodine as oxidant When HgCl 2 is added to the I 2 /I - system it forms [HgI 4 ] 2- Thus:  removing the I - ions from the share of the reaction,  minimizing its concentration,  increasing the ratio of I 2 / [I - ] 2  increasing the oxidation potential of I 2 /2I - system  So I 2 could determine AsO E= = E o - Log [I - ] 2 / [ I 2 ] I e  2I -

E° Fe 3+ /Fe 2+ = 0.77V E° I 2 / 2I - = 0.54V Fe 3+ + e  Fe 2+ E Fe2+ / Fe3+ = When pyrophosphate, EDTA or F - is added to the Fe 3+ /Fe 2+ system it form [FeF 6 ] 3- or [Fe(PO 4 ) 6 ] 3- Thus:  removing the Fe 3+ ions from the share of the reaction,  minimizing its concentration,  decreasing the ratio of Fe 3+ / Fe 2+  lowering the oxidation potential of Fe 3+ / Fe 2+ below that of I 2 /2I - system. Another example How to determine Ferrous salts using Iodine? Iodine as oxidant

Fe(CN) e  Fe (CN) 6 4-  minimizing conc of ferrocyanide  increasing ferri/ferro potential  So Ferri/Ferro system can oxidize I - to I 2 E = E o - 3- Effect of precipitating agents E° Ferri/Ferro= 0.36V E° I 2 / 2I - = 0.54V To determine [Fe(CN) 6 ] 3- ion iodometrically; Zn 2+ should be present: it precipitate Zn 2 [ Fe(CN) 6 ] ion Iodine as oxidant

E° Cu 2+ /Cu + = 0.46 E° I 2 /2I - = 0.54  It is expected that I 2 oxidizes Cu + (cuprous), however, Cu 2+ (cupric) oxidizes I -  Procedure: Cu 2+ is treated with KI and the liberated I 2 is titrated with S 2 O Cu I -  I 2 + Cu 2 I 2 ↓  The precipitation of Cu 2 I 2 increases the oxidation potential of Cu 2+ /Cu +  E = E log [Cu + ] 1 [Cu 2+ ] So Cu 2+ oxidizes I - to I 2  I 2 tends to be absorbed on Cu 2 I 2 so the reaction with S 2 O 3 2- is incompltete so add SCN - near the end point to form Cu 2 (SCN) 2 which has no tendency to adsorb I 2. How to determine Cu 2+ salts using KI ? Iodine as oxidant

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 Cu 2+ /Cu + E = E log Cu + 1 Cu +2

Titration methods: Since iodine may be either reduced or produced by oxidation Direct Iodimetric method Indirect Iodometric method Titrating agent Iodine for determination of reducing agents I - is added to oxidizing agents,the librated I 2 is titr. with Na 2 S 2 O 3 Indicator (Starch) Added at the beginning of titr. Added near the end of titr (when the brown color of I 2 becomes pale) E.P. permanent blue color disappearance of blue color Iodine as oxidant

Reductant + starch Iodine E.P. oxidant + KI→I 2 Na 2 S 2 O 3 Add starch Na 2 S 2 O 3 Colorless E.P. 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 I 2  In exx I 2 an irreversible blue adsorption complex is formed which is not changed  Starch consists of  amylase and amylopectin I 2 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. Iodine as oxidant

Detection of the end point in iodine titrations: 1- The use of starch:  Starch indicator solution must be freshly prepared when it stands decomposition takes place and its sensitivity is decreased. A preservative can be added  Starch can not be used in alcoholic solu.because alcohol hinders the adsorption of I 2 on starch  The sensitivity of the blue color decreases with temperature due to gelatinization of starch and volatility of Iodine Iodine as oxidant

Detection of the end point in iodine titrations: 2- Use of organic solvent (CHCl 3 or CCl 4 )  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 I 2 gives an intense colour, and the end point will be the appearance Or disappearance of the colour in the organic solvent layer.  I 2 is soluble in CHCl 3 or CCl 4 90 times more than in H 2 O  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 S 2 O 3 2- to react with I 2 in CHCl 3

A- Error due to I 2 : (1)I 2 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 - +I 2 → I 3 - (triiodide) less volatile and more stable (2) I 2 conc is changed if the solution gets in contact with rubber, organic matter, dust, SO 2, H 2 S (3) I 2 may undergo disproportionation into HOI and I - I 2 + H 2 O  HOI + I - + H + To overcome this difficulty the solution may be acidified to shift the reaction to the left. Sources of error in iodimetry Iodine as oxidant

Sources of error in iodimetry B- Error due to I - ion: I - ion is liable to atmospheric oxidation. This is catalysed by light, heat, Cu 2+, NO gas  The medium must be completely free from O 2 so introduce CO 2 (add little NaHCO 3 ).  In titration which needs standing for time, standing should be away from light.  If we need acid medium, never use HNO 3, it contains nitrous oxide. Iodine as oxidant 4H + + 4I - + O 2  2I 2 + 2H 2 O

 Thiosulphate is affected by pH, the most favourable pH is 7 till pH 9  Under these conditions: S 2 O 3 2- is oxidized to S 4 O 6 2-, where every 2 S 2 O 3 2- is oxidized by 1 I 2 to S 4 O 6 2- (tetrathionate) 2S 2 O I 2  S 4 O I -  Under acidic conditions: thiosulphate is changed to bisulphite (HSO 3 - ) with the precipitation of S. Every 2 HSO 3 - is oxidized by 2 I 2 to 2HSO 4 -  Therefore, The consumed I 2 in acid medium is double that consumed in neutral medium. Sources of error in iodimetry Iodine as oxidant C- Error due to S 2 O 3 2- ion:

Sources of error in iodimetry Iodine as oxidant C- Error due to S 2 O 3 2- ion:  In pH>9: I - changes to IO - (hypoiodite) oxidizing S 2 O 3 2- to SO 4 2- 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, CHCl 3, or HgI 2 may be added. ●The pH is adjusted by adding borax, Na 2 CO 3 or NaHCO 3 to about pH 9 which inhibits bacterial action.

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 I 2 which masks the true end point. To avoid this, preservatives e.g. H 3 BO 3 and formamide are added. Iodine as oxidant Sources of error in iodimetry