Standard potentials Eo are evaluated with full regard to
activity effects and with all ions present in simple form: they
are really limiting or ideal values and are rarely observed
in a potentiometric measurement. In practice, the solutions
may be quite concentrated and frequently contain other
electrolytes; under these conditions the activities of the
pertinent species are much smaller than the concentrations,
and consequently the use of the latter may lead to unreliable
Also, the actual active species present (see example below)
may differ from those to which the ideal standard potentials
apply. For these reasons 'formal potentials' have been
proposed to supplement standard potentials
The concept of formal potential was introduced by Swift and he
defined it as the experimentally determined potential of a
solution containing both the oxidised and reduced species
each at a concentration of 1 formal together with other
substances at a given concentration.
1F ≡ 1g formula weight per litre solution
The std. Potential or E0 (normal potential) is a limiting or ideal value
assuming that the reacting species are in their simple forms. Since in
practice such ideal solutions are rarely found. The term formal
potential has been proposed and found to be more practical one.
[ Fe +3 ] Where the ionic sp. are aqueous forms
E = E0 log
[ Fe +2 ]
Fe+3 aq or Fe+3(H2O)6 ; Fe+2 aq or Fe+2 (H2O) simple form of Fe+3 or Fe+2
In IM HCl soln, however part of Fe+3 exists as various sp.like
FeCl+2, FeCl2+, FeCl4-2 (all in Feric state).
Similarly Fe+2 as FeCl+, FeCl4-2 etc.
Let, α be the fraction of total Fe+3 irons, β is the fraction of
total Fe+2 iron existing in simple aqueous form
E = E 0 - 0.0591log
α . CFe III .f Fe III aq
β .CFe II .f Fe + 2aq
= E - 0.0591 log - 0.0591log
= E - 0.0591 log
Fe III aq
E : normal potential i.e., std for
ionic strength of the soln,
where CFeIII and CFe+2 are the
molar concentration of total
Fe+3 and Fe+2 iron present in
the soln.,f: activity coefficient
− 0.0591 log
Fe + 2 aq
E 0/ = Formal
Where Fe+3 forms a more stable complex in comparison to
the Fe+2 by interaction with the anions or any comlexing
reagent present in the soln. The value of the ratio α/β is less
than 1 and in that case the value of Formal potential
becomes more +ve which means the Fe+2 behaves as stronger
reducing agent. PO4-3 is such an anion in presence of which
Fe+2 is a stronger reducing agent.
E0 : Fe+2/Fe+3 = -0.77V
0.5 M H3PO4 + 1M H2SO4
E0/ Diphenyl amine = -0.76V in (0.5M H3PO4 + 1M H2SO4)
Minm 0.15 V difft of the redox potential value for making a titration
Fe(CN)6 ⇋ Fe(CN)6-3 + e4-
2I- = I2 + 2e-
E0 = -0.36 V
E0= -0.54 Volt
I2 + 2Fe(CN)6-4 ⇋ 2I- +2Fe(CN)6-3
It can be seen that I2 should normally oxidise ferrocyanide to
Ferricyanide. But in 1M HCl medium or 1M H2SO4 medium
E0/ [(H2Fe(CN)6-2 /H2Fe(CN)6-1 ] = -0.71 V
α/β > 1
+ 2I- ⇋ I2 + 2Fe(CN)6-4
Hence in 1M HCl medium Ferricyanide liberates I2 from I-. Similarly
when Zn+2 is present in the soln it reacts with Ferrocyanide to form
2K + 3Zn + [Fe(CN)6] → K2Zn3[Fe(CN)6]2
And so in presence of Zn+2 sufficient K+Ferricyanide liberates I2 from KI
and the rn can be used for indirect iodometric determination
determination of Zn+2 in soln.
Red Ox indicators
InRed ⇋ InOx +ne-
InRed and InOx must have different colours
= E In log
[ In Red ]
In order that indicator can be used in a redox titration the
potential of the analyte soln must lies near
Criteria for using a redox indicator
1. The indicator potential must be intermediate between that
of the analyte and titrant system
2. At eq. point redox potential of the soln should change from
0.0591 to 0
0.0591 with sharp detectable
E 0 In +
E In −
3.E0In or better Eo/In should differ by at least 0.15 V from the
formal potential of the redox couple titrated
Classification of RedOx Indicator
[ MLx ]+p ⇋ [MLx ]+(p+n)
Where L is a ligand, M= Metal ,
the two forms must have difft colour. e.g.
pale blue less stable
E0In = -1.14 V
E0/In (1MHCl) = -1.06 V
As the red colour is more intense the colour change is actually
observed at E0In value of -1.12 V
Introducing various substitutes in the phenanthroline ring
indicator potential value can be changed.
Thus, with 4,7 dimethyl Ferroin , E0/In (1MHCl) = -0.88 V
which is very suitable for titration of chromate or dichromate by
Fe+2 soln. (Mohr’s titrant soln)
when NO2- group is introduced RedOx potential value is more
–ve -1.25V (Nitroferroin)
II . Organic molecules forming a Redox Couple,
Diphenyl amine (knop, 1924) in H2SO4 soln (DPA, base and
insoluble in water)
Red-Ox potential of In (0.5-1M H2SO4 + 0.5M H3PO4 ) medium,
Eo/In = - 0.76 V
Suitable for titration of Fe+2 by dichromate soln, in presence of
Eo/In = - 0.61 V
Due to low solubility of free DPA it is replaced by its sulfonic
Diphenylamine 4-sulphonic acid
The sodium or Ba salt of which is used as indicator in the form
of 0.2% aq soln. The formal potential o/
E In = - 0.85 V
So, even without using H3PO4, DPA sulfonate indicator can
be used safely in the titration of Fe+2 by dichromate.
The violet colour Diphenyl benzidine is oxidised further to a yellow
or reddish yellow product and hence in the reverse titration of
chromate or dichromate by Mohr’s salt soln, DPA is not convenient,
N-methyl DPA sulfonic acid is ecellent for reverse titration
Eo/In = - 0.90 V
N-PHENYL ANTHRANILIC ACID is best for reverse titration
Another group of Organic compounds are used as redox
indicator of which “ Variamine blue” is a versatile indicator
In the pH range 1.5-6.3 it is colourless in reducing medium and
blue in Oxidising medium.
Eo/In = - 0.6 V (pH= 1.5) to 0.36 V (pH =6.3)
It has been used in the titration using Ascorbic acid(Vitamin C)
BTB can be used as a Red-Ox indicator in the titration of AsIII,
NH3, I- etc. With ClO- hypochlorite soln,
Deep blue (Reducing) – greenish yellow (Oxidising medium)
EoIn = - 0.52 V
EoIn = - 0.53 V
Oxidation with permanganate : Reduction of
Powerful oxidant that the most widely used.
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.
In feebly acidic, neutral, or alkaline solutions
MnO4– + 4H+ + 3e = MnO2 (s) + 2H2 O
Eo = 1.695 V
brown manganese dioxide solid
In very strongly alkaline solution (2M NaOH)
MnO4– + e = MnO42 –
Eo = 0.558 V
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
158.03 g / 5
5 Na2(COO)2 ≡ 10 Equivalent
≡ 134.01 g / 2
≡ 1 Eq.
1N × 1000 ml
x N × V ml
x N = ( a g × 1N × 1000 ml) / (67.005 g × V ml)
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.
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
2) Filter the liquid through a sintered-glass filter crucible to remove solid
3) Transfer the filtrate to a clean stoppered bottle freed from grease with
4) Protect the solution from evaporation, dust, and reducing vapors, and
keep it in the dark or in diffuse light. Preserve it in amber –coloured
5) Standardise from time to time. If in time managanese dioxide settles out,
refilter the solution and restandardize it.
The concentrations of the various species must be taken into consideration
especially if combined redox and acid-base systems are involved. From the
data, taken from Table 1.17 and I.l8 for example:
one could draw the conclusion that arsenate ions will oxidize iodide:
but the reaction cannot go in the opposite direction. This in fact is true only
if the solution is strongly acid (pH ~ 0). The oxidation-reduction potential of
the arsenate-arsenite system depends on the pH:
At pH = 6 the potential of a solution containing arsenate and arsenite ions at
equal concentrations decreases to +0'20 V. Under such circumstances therefore
the opposite reaction will occur:
Starch solution(05~ 1%) is not redox indicator.
The active fraction of starch is amylose, a polymer of the sugar α-D-glucose
( 1,4 bond).
The polymer exists as a coiled helix into which small molecules can fit.
In the presence of starch and I–, iodine molecules form long chains of I5– ions
that occupy the center of the amylose helix.
••••[I I I I I]– ••••[I I I I I]– ••••
Visible absorption by the I5– chain bound within the helix gives rise to the
characteristic starch-iodine color.
Structure of the repeating unit of the sugar
Schematic structure of the starch-iodine
complex. The amylose chain forms a
helix around I6 unit.
View down the starch helix, showing
iodine, inside the helix.
Starch is the indicator of choice for those procedures
involving iodine because it forms an intense blue
complex with iodine. Starch is not a redox indicator;
it responds specifically to the presence of I2, not to a
change in redox potential.
The active fraction of starch is amylose, a polymer of
the sugar α-d-glucose.
In the presence of starch, iodine forms I6 chains
inside the amylose helix and the color turns dark
16-7 Methods Involving Iodine
• 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
• Iodine only dissolves slightly in water. Its solubility is
enhanced by interacting with I-
• A typical 0.05 M solution of I2 for titrations is prepared by
dissolving 0.12 mol of KI plus 0.05 mol of I2 in 1 L of water.
When we speak of using iodine as a titrant, we almost always
mean that we are using a solution of I2 plus excess I−.
Preparation and Standardization of Solutions
• 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
• 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 I2 by quantitative