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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 .

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Ceric as titrant: Ce4+


  • 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-
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+

Ceric as titrant: Ce4+


  • 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

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.

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 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

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.
Potassium dichromate as titrant


1-Determination of Fe2+ Iron (internal indicator)

E° Fe3+/Fe2+ 0.77

E° diphenylamine 0.76

E° ferroin 1.06

Is there need for



Titrate with


E0 = 1.33 v






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.
Potassium dichromate as titrant


  • 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.

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.

Potassium dichromate as titrant


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+


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

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

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.
Iodine as oxidant



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-.
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.
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 -

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.
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

Iodine as oxidant

Titration methods:

Since iodine may be either reduced or produced by



Iodimetric method


Iodometric method


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.



permanent blue


disappearance of

blue color


Iodine as oxidant











Add starch

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 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.

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


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

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.
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

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.
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.

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.