Reduction- Oxidation
Reactions
5thlecture
Ceric as titrant:
4+
Ce
Properties
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
(1) Against arsenious trioxide:
2Ce4+ + H3 AsO3 + H2O  2Ce3+ + H3AsO4+ 2H+
(2) Against oxalate
2Ce4++ H2C2O4 ↔ 2Ce3+ + 2CO2 + 2H+
In both cases , the reaction is slow it requires heat
to 50°C, using ICl as catalyst and ferroin indicator
Applications
Ceric as titrant: Ce4+
(a)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]3H2O2 + 2Ce4+  2Ce3+ + 2H+ + O2
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.
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
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
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
Applications
1-Determination of Fe2+ Iron (internal indicator)
E° Fe3+/Fe2+ 0.77
E° ferroin 1.06
E° diphenylamine 0.76
Is there need for
H3PO4
H3PO4??
diphenyl Titrate with
2Cr
O
2
7
amine
0 = 1.33 v
E
H2SO4
3+
Fe
Fe2+
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
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.
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.
Applications
3-reducing agents
Na2SO3
Potassium dichromate as titrant
4-Organic compd
glycerol
5-Pb2+
PbO
Add measured 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
Properties:
Iodine as oxidant
The iodine/iodide half reaction is
I2 + 2e  2I↑ E°
↓ E°
(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-/S2O32ClO3-/ClS/S2Iodometric 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
Properties:
Iodine as oxidant
Systems having oxidation potentials near to that
of iodine/iodide e.g AsO43-/AsO33-, Fe3+/Fe2+
Their reactions with 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:
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+
The potential of:
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  2IE= = Eo -
0.059 Log [I-]2 / [ I ]
2
2
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
How to determine Ferrous
E° Fe3+/Fe2+= 0.77V
salts using Iodine?
E° I2/ 2I- = 0.54V
Fe3+ + e  Fe2+
E Fe2+ / Fe3+ = 0.559 -
0.059
[ Fe2 ]
log
1
[ Fe3 ]
When pyrophosphate, EDTA or F- is added to the
Fe3+/Fe2+ system it form [FeF6]3- or [Fe(PO4)6]3Thus:
 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-
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
4[
Fe
(CN)
]
0
.
059
6
E = Eo log
1
[Fe(CN)6 ]3-
 minimizing conc of ferrocyanide
 increasing ferri/ferro potential
 So Ferri/Ferro system can oxidize I- to I2
E° Cu2+/Cu+ = 0.46
E° I2/2I- = 0.54
Iodine as oxidant
How to determine Cu2+
salts using KI ?
 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 S2O322Cu2+ + 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
oxidation
Direct
Indirect
Iodimetric method Iodometric method
Iodine
I- is added to oxidizing
Titrating agent for determination
agents,the librated I2
of reducing agents
is titr. with Na2S2O3
Added near the
Added at the
Indicator
end of titr (when the
beginning of titr.
(Starch)
brown color of I2
becomes pale)
E.P.
permanent blue
color
disappearance of
blue color
Iodine as oxidant
Iodine
Na2S2O3
Na2S2O3
oxidant
+
KI→I2
Reductant
+
starch
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
added
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 S2O32to react with I2 in CHCl3
Sources of error in iodimetry
Iodine as oxidant
A- Error due to I2:
(1)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 II-+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 II2 + 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.
4H+ + 4I- + O2  2I2 + 2H2O
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.
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 I2
to 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 I2 to 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.