Unit 10 Redox Titrations

5

Redox Titrations UNIT 10 REDOX TITRATIONS

Structure

10.1 Introduction Objectives

10.2 Redox Reactions and Redox Potential Redox Reaction as Two Half Reactions

Redox Potential

10.3 Electrochemical Cells Nernst Equation

Cell Potential

Redox Equilibrium Constant

10.4 Redox Titration Curves

10.5 Redox Indicators

10.6 Redox Titrations in Nonaqueous Solvents Criteria for Solvent Selection

Characteristics of Common Nonaqueous Solvents

10.7 Applications of Redox Titrations Oxidimetric Reagents

Reductimetric Reagents

10.8 Summary

10.9 Terminal Questions

10.10 Answers

10.1 INTRODUCTIOIN

You have learnt about the theory and applications of acid-base titrations in Units 7, 8

and 9 of Block 3. A large number of analytical determinations make use of another

important kind of titration, namely, redox titration. As the name suggests these

titrations are based on oxidation-reduction reactions. In contrast to acid-base titrations

in which the titration reaction involves the formation of undissociated molecules of a

weak electrolyte (water or a weak acid), a redox titration reaction is associated with

the transfer of electrons. The electrons are transferred from a reducing agent to an

oxidising agent. Oxidation-reduction or redox reactions are widely used in the

titrimetric determination of both inorganic and organic compounds.

In this unit we begin with a brief review of the concept of redox reactions and redox

potential. Then we shall take up the measurement and significance of standard

reduction potentials. It will be followed by the calculation of the cell potential in the

course of a titration. The theory of redox indicators and the redox titrations in the

nonaqueous solvents will be followed by a few important applications of redox

titrations in analytical determinations. In the next unit you would learn about

complexometric titrations having extensive analytical applications.

Objectives

After studying this unit, you will be able to:

• define oxidation and reduction,

• explain the terms like electrode potential, cell potential and redox potential,

• calculate the cell potential for a given cell,

• calculate the equilibrium constant for a redox reaction and predict the direction

of the reaction,

• construct a titration curve for a redox reaction and compute the equivalence

point,

6

Estimations Based on

Redox and

Complexation

Equilibria Studies

• select a suitable indicator for a given redox titration on the basis of potential at

equivalence point, and

• carry out the determination of different inorganic and organic substances by

performing their redox titrations with suitable oxidising or reducing agents.

10.2 REDOX REACTIONS AND REDOX POTENTIAL

You would recall from your earlier studies that the oxidation process results in the loss

of one or more electrons by an atom or an ion, e.g. in case of iron we can write the

process as follows.

eFeFe 32+→

++

Here, the ferrous ions are getting oxidised to ferric ions by the loss of one electron.

Reduction, on the other hand, is the process which results in the gain of one or more

electrons by an atom or an ion, e.g., the chlorine gets reduced to chloride ions by gaining electrons, as shown below.

−→+ Cl22eCl2

Further, oxidation and reduction are always found to occur together in a reaction. In

other words, electron loss of one substance is always accompanied by an electron gain

on the part of the other and consequently there are no free electrons in a chemical system. Such reactions are termed redox reactions and are also termed electron

transfer reactions. Let us take the example of the reaction between the Ce4+

and Fe2+

ions as given below.

Ce4+

+ Fe2+ � Ce

3+ + Fe

3+

Here, the electron lost by ferrous ion are picked up by Ce4+

ions and as you can see

that there are no free electrons in the reaction. The transfer of electrons from one species of the reactants to another can be rationalised in terms of changes in the oxidation number. An oxidation-reduction reaction is accompanied by an increase in

the oxidation number of a given species and a complementary decrease in oxidation number of another species. In the example given above, the oxidation number of iron

increases from 2 to 3 whereas that of cerium decreases from 4 to 3.

10.2.1 Redox Reaction as Two Half Reactions

As a redox reaction involves both oxidation and reduction components, it is possible

to divide the total reaction into two half reactions. These clearly demonstrate which

species gains electrons and which loses them. This can be suitably illustrated by the reaction between iron (II) and cerium (IV) given above. This reaction can be divided into two half reactions as shown below.

The first reaction shows oxidation of iron (II), the second reaction depicts the

reduction of cerium (IV). On adding the half reactions, the electrons cancel out and we get the overall equation. This method of expressing a redox reaction as two half

reactions provide a kind of flexibility to the concept of redox reactions.

The species possessing strong affinity for electrons cause the oxidation of other substances by abstracting electrons from them. These species are called oxidising

agents. In the process, the oxidising agents are themselves reduced. In the above

reaction, for example, cerium (IV) ions act as oxidising agents and are consequently reduced to cerium (III) ion. On the other hand the species that readily give up

( Oxidation) Fe2+ Fe3+ e+

(Reduction)Ce3+Ce4+ e+

You would recall that the

oxidation is defined as an

increase in the oxidation

number of an element.

Reduction is defined analogously as a decrease

in the oxidation number.

7

Redox Titrations electrons cause the reduction of other substances. These are called reducing agents and in this process, a reducing agent itself gets oxidised. Thus, in the above example, iron (II) ions act as reducing agents and are consequently oxidised to iron (III) ions.

10.2.2 Redox Potential

If a system contains both an oxidising agent and its reduction product, there will be equilibrium between them and the electrons. If an inert electrode, such as platinum, is

placed in such a redox system, for example, one containing Fe (III) and Fe (II) ions, it

will assume a definite potential indicative of the position of the equilibrium. If the system tends to act as an oxidising agent i.e., it will take electrons from platinum,

leaving the latter positively charged. On the other hand, if the system has reducing

properties i.e., electrons will be given to the metal, which will acquire a negative

charge. The magnitude and sign of the potential on the platinum electrode will thus be a measure of the oxidising or reducing properties of the system.

Different oxidising / reducing agents differ from one another in their strength. An oxidising agent can behave as a reducing agent in the presence of a stronger oxidising

agent. The reaction can be depicted as under.

(II)agentOxidisingagent(I)Reducing(II)agent Reducing(I)agentOxidising +→+

For a reaction of the above type, where we have a pair of oxidising-reducing agents

such that either of the species can act as an oxidising or a reducing agent, the direction of the reaction is determined by comparing the redox potential of the oxidising/reducing agents. The redox potential is a quantitative characteristic of the

oxidising/reducing power of a reagent. Let us try to understand the meaning and the

significance of redox potential.

Significance of Redox Potentials

In order to obtain the comparative strengths of oxidising agents, it would be necessary

to measure the potential difference between the platinum and the solution relative to a standard or a reference under standard experimental conditions. The reference

electrode used is the standard or normal hydrogen electrode whose potential is

taken as zero at all temperatures. The standard experimental conditions for redox systems are those in which the ratio of the activity of the oxidant to that of the

reductant is unity. Hence for the following electrode,

Fe3+ Fe2+

the standard redox cell would be written as follows:

Pt,Fe2+

Fe3+ (a = 1)

(a = 1)H+ (a = 1) PtH2

The potential measured in this way is called the standard redox potential or

standard reduction potential. The standard reduction potentials for a number of redox systems are given in Table 10.1.

Table 10.1: Standard reduction potentials (E0)

* for some redox systems at 298 K

Half reaction E0, Volts

F2 + 2e 2F

+ 2.65

+ 2e S2O8

2-2SO4

2-

+ 2.01

Co3+ + e Co2+

+ 1.82

Fe3+ Fe2+

Fe3+Fe2+

You would recall that in a

cell notation a single

vertical line indicates a

solid liquid interface

whereas a double vertical

line represents a salt

bridge.

8


Redox and

Complexation

Equilibria Studies

Pb4+ + 2e Pb2+

+ 1.70

MnO4

-+ 4H+ + 3e MnO2 + 2H2O

+ 1.69

BrO3 + 6H+ + 5e 1/2 Br2 + 2H2O-

+ 1.52

MnO4 + 8H+ + 5e Mn2+ + 4H2O-

+ 1.52

Ce4+ + e Ce3+

+ 1.44

Cl2 + 2e 2Cl-

+ 1.36

2-Cr

2O

7 + 14H+ + 6e 2Cr3+ + 7H

2O

+ 1.33

Tl3+ + 2e Tl+

+ 1.25

MnO2 + 4H + + 2e Mn2+ + 2H2O + 1.23

O2 + 4H+ + 4e 2H2O + 1.23

IO3 + 6H+ + 5e 1/2 I2 + 3H2O-

+ 1.20

Br2 + 2e 2Br

-

+ 1.07

HNO2 + H+ + e NO + H2O + 1.00

2Hg2+ + 2e Hg22+

+ 0.92

Cu2+ + I + e CuI-

+ 0.86

Ag+ + e Ag+ (s)

+ 0.80

_Hg2 + 2e 2Br

2+

+ 0.79

Fe3+ + e Fe2+

+ 0.77

BrO3 + 3H2O + 6e Br + 6OH---

+ 0.61

-MnO4 + 2H2O + 2e MnO2 + 4OH-

+ 0.60

2-MnO4 + e MnO4

-

+ 0.56

H3AsO4 + 2H+ + 2e H3AsO3 + H2O

+ 0.56

Cu2+ + Cl + e CuCl-

+ 0.54

I2 + 2e 2I-

+ 0.54

[Fe(CN)6]3+ + e [Fe(CN)6]

4-

+ 0.36

IO3 + 3H2O + 6e I + 6OH---

+ 0.26

Cu2+ + e Cu+

+ 0.15

Sn4+ + 2e Sn2+

+ 0.15

S4O6 + 2e 2S2O3

2- 2-

+ 0.08

2H+ + 2e H2 + 0.00

Ni2+ + 2e Ni (s)

- 0.25

V3+ + e V2+

- 0.26

Cr3+ + e Cr2+

- 0.41

9

Redox Titrations U4+ + e U3+

- 0.61

AsO4 + 3H2O + 2e H2AsO3 + 4OH---

- 0.67

Zn + 2e Zn (s) -0.76

Al3+ + 3e Al (s)

-1.66

Mg2+ + 2e Mg(s) -2.38

Ca2+ + 2e Ca (s)

- 2.87

K+ + e K (s)

- 2.92

Li+ + e Li (s) - 3.04

*A negative E

0 means that the redox couple is a stronger reducing agent than H

+ / H2 couple

while a positive value means that the redox couple is a weaker reducing agent than hydrogen

couple.

These redox potentials enable us to predict which ion will oxidise or reduce other ions

at unit activity (or molar concentration). The most powerful oxidising agents are those

which are at the upper end of Table 10.1 and most powerful reducing agents are at

bottom end of the table. For example, the standard reduction potential for the

reduction of Ce (IV) ions is highly positive, that is the reduction of cerium (IV) ions is

quite spontaneous, so Ce (IV) acts as a strong oxidising agent. On the other hand, the

standard reduction potential for the Zn (II) ions is negative, meaning thereby that the

reduction of zinc ions is not favourable. Therefore, Zn (II) ion is a weak oxidising

agent, while metallic zinc is a strong reducing agent.

It is important to mention here that for many oxidants, the pH of the medium is of

great importance since they are generally used in acidic media. For example, in

measuring the standard redox potential of the−

4MnO , Mn2+

system, i.e.,

OH4Mn5e8HMnO 22

4 +→++++−

it is necessary to state that hydrogen activity is unity. It needs to be emphasised that

standard redox potentials do not give any information about the speed of the reaction.

In some cases, a catalyst has to be added to enable the reaction to proceed with

reasonable velocity.

SAQ 1

Using Table 10.1 determine which of the two ions, Fe (II) or Cr (II) would be easily

oxidised.

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

10.3 ELECTROCHEMICAL CELLS

The direct transfer of electrons from an electron donor to the electron acceptor in a

redox reaction can be best illustrated by immersing a piece of zinc in a copper (II)

sulphate solution. It is observed that copper is deposited at the surface of the metal.

The copper ions migrate to the zinc electrode and get reduced there by extracting

electrons from it; the zinc metal dissolves and generates Zn2+

ions in the solution. It is

important to mention here that such an transfer of electrons can still be accomplished

Ce4+ + e Ce3+

Zn2+ + 2e Zn

10


Redox and

Complexation

Equilibria Studies

even when the electron donor and accepter are kept away from each other. This can be achieved in an electrochemical cell as shown in Fig.10.1.

Fig. 10.1: Schematic diagram of a simple electrochemical cell

The electrochemical cell shown in Fig.10.1 is called Daniel cell. In this set up the

direct contact between metallic zinc and Cu2+

ions is prevented by connecting the two

electrodes with the help of a salt bridge. The salt bridge contains a saturated solution

of an inert electrolyte like KNO3 in agar agar gel. In such a situation, the electrons are transferred by means of the external circuit and within the cell the current is carried by

the ions. This electron transfer continues until the Cu2+

and Zn2+

concentrations achieve levels corresponding to the equilibrium for the reaction.

Zn + Cu2+ Zn2+ + Cu

When this stage is reached, there will be no further flow of electrons. It is important to

note that the overall reaction and the position of its equilibrium is the same regardless

of the manner in which the reaction is carried out.

An electrochemical cell is capable of producing electrical energy because of the

tendency of reacting species to transfer electrons and thus achieve the condition of equilibrium. A cell operated in such a way is called a galvanic cell and is used to produce electric energy. The voltage produced by an electrochemical cell is directly

related to the equilibrium constant for the particular oxidation-reduction process

involved. The measurement of these potentials constitutes an important source of numerical values for these constants.

10.3.1 Nernst Equation

In our arguments above, we mentioned that the activity of different species in an electrode reaction is unity. However, this need not be always true. Nernst showed that for the electrode reaction:

)s(Me)aq(M n→+

+ n

the electrode potential at any concentration measured with respect to standard hydrogen electrode is given by the following equation called Nernst

equation.

]M[

]M[ln

n

0

M/MM/M nn+

−= ++

nF

RTEE

where E0 is the standard electrode potential (characteristic of the particular half

reaction), R is the gas constant, T is the absolute temperature, n is the number of electrons participating in the reaction and F is the Faraday’s constant. The logarithm

term is the ratio of the concentrations of the reduced and oxidised forms of the species.

It is customary to employ concentrations rather than activities of reactants and products and the same are given in molarities (for ionic and molecular species). For

The electrical potential

existing between zinc and

copper electrodes is a

measure of this driving

force and can be easily

measured by a voltmeter,

V, placed in circuit.

The electrode potential,

which is a measure of the

chemical driving force of

a half-reaction, is affected

by the concentration. The

concentrations of

reactants and products in

a half-reaction have a marked effect on

electrode potential.

11

Redox Titrations gaseous species, the pressure in atmosphere is used. The concentration of pure substances such as solids (precipitates) and liquids (water) is taken to be unity. As the numerator in the logarithm term in equation is the concentration of metal, a pure

substance, it is taken as 1 and the equation becomes,

]M[

1ln

n

0

M/MM/M nn+

−= ++

nF

RTEE

After substituting common logarithm for natural logarithms and inserting numerical

values for the constants and assuming the temperature as 298K, we can write the equation as given below.

][M

1log

0.059n

0

/MM/MM nn+

−= ++

nEE

Let us write the Nernst equation for some half reactions.

Cu2+ + 2e CuHalf reactionI

Nernst equation ][Cu

1log

2

0.0592

0

/CuCu/CuCu 22+

−= ++ EE

The electrode potential varies with the logarithm of the reciprocal of the molar

Cu2+

ion concentration in the solution (the activity of metallic copper is taken as unity).

2H+ + 2e H2II Half reaction

Nernst equation 2

20

2/ H2H/ H2H][H

pHlog

2

0.059

2 ++−=+ EE

pH2 represents the partial pressure of hydrogen (expressed in atmospheres).

Ordinarily, it will be very close to atmospheric pressure.

Cr2O7 + 14H+ + 6e 2Cr3+ + 7H2O-III Half reaction

Nernst Equation

14272

23

14272

230

/CrOCr/CrOCr

]H[]OCr[

][Crlog

6

0.05933.1

]H[]O[Cr

][Crlog

0.05932

7232

72

+−

+

+

+

−=

−= +−+−

nEE

In this case, the potential depends not only on the concentration of Cr3+ and

Cr2O72 −

ions but also on the pH of the solution.−

10.3.2 Cell Potential

When two electrodes combine to give a cell, the potential of the cell depends on the

concentrations of the species involved in both the electrodes or the half cells. We can

develop a suitable Nernst equation for the same also. Let us do this for the Daniel cell

described above. The cell potential of the Daniel cell can be determined by taking the

electrode potentials of the two electrodes for any given concentration of Cu2+

and Zn2+

ions. Let us write the expressions for the reactions at cathode and anode,

Reaction at Cathode: )]aq(Cu[

1ln

2

0

/CuCu/CuCu 22+

−= ++

2F

RTEE

12


Redox and

Complexation

Equilibria Studies

Reaction at Anode:

)]aq(Zn[

1ln

2 2

0

Zn/ZnZn/Zn 22+

−= ++

F

RTEE

The cell potential, Zn/ZnCu/Cucell 22 ++ −= EEE

)]aq(Zn[

1ln

2)]aq(Cu[

1ln

2 2

0

Zn/Zn2

0

Cu/Cu 22++

+−−= ++

F

RTE

F

RTE

)]aq(Zn[

1ln

2)]aq(Cu[

1ln

2)(

22

0

Zn/Zn

0

Cu/Cu 22++

−−−= ++

F

RT

F

RTEE

)]aq(Cu[

)]aq(Zn[

2 2

20

)cell(cell +

+

−= lnF

RTEE

Taking clue from here, we can write a generalised Nernst equation for any cell. For a

cell with the following general equation

aA + bB + . . . + ne cC + dD + . . .

The Nernst equation would be

.....[B][A]

.....[D][C]ln

ba

dc0

n

RTEE −=

Converting to the logarithm to the base ten and substituting the values of constants and

the temperature as 298 K, we get the following.

......[B][A]

......[D][C]log

0.059ba

dc0

nEE −=

Let us take an example to learn how to write the Nernst equation for a given cell

reaction.

Example 10.1

Compute the standard electrode potential of the cell in which the following reaction

takes place and write the Nernst equation for the cell. (You may use Table 10.1 for the

data required.)

)s(Ag2)aq(Ni)aq(Ag2)s(Ni2

+→+++

Solution

In such a situation we need to identify the species getting oxidised and the one getting

reduced. In other words we need to identify the oxidation and reduction half reactions.

In this case as you can see that the nickel is getting oxidised to nickel ions while silver

ions are getting reduced to metallic silver. The reactions being,

2e(aq)Ni(aq)Ni(s) 2+→

+

(s)Ag22e(aq)2Ag →++

The standard reduction potentials for the two redox systems are – 0.25 V and 0.80 V

respectively. The standard cell potential would be [0.80 – (– 0.25) = 1.05 V]

The Nernst equation for the cell with given cell reaction would be as follows.

13

Redox Titrations = V

)]aq(Ag[

)]aq(Ni[ln

2 2

2

cell +

+

−=F

RT1.05E

V(aq)][Ag

(aq)][Niln

2

0.05921.05

2

2

+

+

−=

10.3.3 Redox Equilibrium Constant

The equilibrium constant of a redox reaction can be obtained by the application of

Nernst equation. It provides information regarding the direction and extent of the reaction.

Let us consider two redox systems that involve the same number of electrons.

Ox1 + ne Red1

Red2Ox2 + ne

Ox1

+ Red2 Red

1 + Ox

2

The equilibrium constant for the reaction can be written as follows.

][Red][Ox

][Ox][Red

21

21=K

For the above two systems, Nernst equations for the reduction potentials can be

written as follows.

][Ox

][Redlog

0.059

][Ox

][Redlog

0.059

2

2022

1

1011

nEE

nEE

−=

−=

When the systems are at equilibrium i.e., E1 = E2, then we have the following

expressions.

][Ox

][Redlog

0.059

][Ox

][Redlog

0.059

2

202

1

101

nE

nE −=−

Rearranging the equation and substituting the equilibrium constant into the expression,

we get the following relation

0.059

)( log

log0.059

][Red

][Ox

][Ox

][Redlog(

n

059.0

)][Ox

][Red

][Ox

][Red(log

0.059

)][Ox

][Redlog

][Ox

][Redlog(

0.059

][Ox

][Redlog

0.059

][Ox

][Redlog

0.059

o2

o1

2

2

1

1

2

2

1

1

2

2

1

1

2

2

1

102

01

EEnK

Kn

n

n

nnEE

−=

=

=

−=

−=

−=−

If the number of electrons interchanged is not the same (n1 ≠ n2), the following equation can be derived.

14


Redox and

Complexation

Equilibria Studies

( )0.059

log

02

01

21

EEnnK

−=

Let us take an example to learn about the application of Nernst equation in calculating the equilibrium constant of a reaction.

Example 10.2

Calculate the equilibrium constant of the following reaction and predict the direction of the reaction

Sn4+ (aq) + 2Fe2+(aq) Sn2+(aq) + 2Fe3+(aq)

V0.77;V0.14 Given, 0

Fe/Fe

0

Sn/,Sn2324 == ++++ EE

Solution

Substituting the given values of the standard electrode potentials, we get

21.310

21.3)0.059

0.770.14(21log

−=

−=−

×=

K

K

The extremely low value of equilibrium constant suggests that the equilibrium lies far to the left.

SAQ 2

Calculate the electrode potential of a half cell containing an aqueous solutions of

0.100 M KMnO4 and 0.100 M MnCl2 having a pH = 1.000. (Note: You may use Table 10.1.)

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

10.4 REDOX TITRATION CURVES

In redox titrations, the concentration of the substances or ions involved in the reaction continuously keeps changing in the course of the titration. Hence, the redox potential

of the solution must also change (the phenomenon may be compared to the change of

the pH of solution during acid-base titrations). By plotting the redox potential

corresponding to different points in the titration, a titration curve similar to the curve obtained in an acid-base method is obtained. This can be illustrated by computing the titration curve for a titration of 100 cm3 of 0.1M iron (II) with 0.1 M cerium (IV) in

the presence of dilute sulphuric acid.

Fe2+ + Ce4+ Fe3+ + Ce3+

This titration involves Fe3+/Fe2+ and Ce4+/Ce3+ ion electrode systems (let us call them as 1 and 2). These can be represented by Nernst equation as follows.

15

Redox Titrations

][Fe

][Felog0.0590.75

][Ce

][Celog

1

0.059

;][Fe

][Felog

1

0.059

3

2

1

4

3022

3

20

11

+

+

+

+

+

+

−=

−=

−=

E

EE

EE

][Ce

][Celog0.0591.45

4

3

+

+

−=2E

As you know, the equilibrium constant of the above reaction can be written as follows.

11

o

2

o

1

24

33

107

11.84

0.75)(1.450.059

10.059

)E(En log also,

][Fe][Ce

][Fe][Celoglog

×=

=

−=

−=

×

×=

++

++

K

K

K

As the value of the equilibrium constant is quite large, the above reaction goes almost

to completion.

It may be noted that upto the equivalence point of the titration, cerium(IV) solution

added will only oxidise iron(II) to iron(III); since K is large all the cerium ions

would be consumed and the ratio [Fe2+

] / [ Fe3+

] would change. Therefore, we can

compute the cell potential by using the Nernst equation for the Fe3+ / Fe2+ half cell.

Let us compute the cell potential as a function of the progress of the reaction.

a) Potential before starting the titration

In the beginning of the titration, the solution would contain only Fe2+ ions and may

have traces of Fe3+

due the aerial oxidation of the solution. As the concentration of the

tripositive ion is too small, the calculation of the potential has no meaningful

significance.

b) Potential after the addition of 10 cm3 of 0.1 M cerium (IV)

As the titrand and titrant have same concentration, on addition of 10 cm3 of the titrant

10% of the ferrous ions would have got oxidised to ferric ions and the relative

concentrations of the two ions would be Fe2+

: Fe3+

:: 90:10. In other words,

( )approx.10

90

][Fe

][Fe3

2

=+

+

We can compute the cell potential as following.

( )approx.V0.690.060.75

10

90log0.0590.75

=−=

−=E

16


Redox and

Complexation

Equilibria Studies

c) Potential after the addition of 50 cm3of 0.1 M cerium (IV)

It can be calculated as follows.

V0.7500.75

50

50log0.0590.75

50

50

][Fe

][Fe3

2

=−=

−=

=+

+

E

d) Potential after the addition of 90 cm3 of 0.1 M cerium (IV)

(approx.)V0.810.060.75

90

10log0.0590.75

(approx.)90

10

][Fe

][Fe3

2

=+=

−=

=+

+

E

e) Potential after the addition of 99 cm3 of 0.1 M cerium (IV)

(approx.)V0.870.120.75

99

1log0.0590.75

(approx.)99

1

][Fe

][Fe3

2

=+=

−=

=+

+

E

f) Potential after the addition of 99.9 cm3 of 0.1 M cerium (IV)

Preequivalence point potential

(approx.)V0.930.180.75

99.9

0.1log0.0590.75

(approx.)99.9

0.1

][Fe

][Fe3

2

=+=

−=

=+

+

E

g) Potential after the addition of exactly 100 cm3 of 0.1 M cerium (IV)

Equivalence point potential

V1.102

1.450.75

2

][Ce][Ce;][Fe][Fe

21

4332

=+

=+

==++++

EE

After this stage there won’t be any free ferrous ions, with further addition of the

reagent, it is the ratio [Ce3+

] / [Ce4+

] which will change. Therefore, for the

further computation of the titration curve we need to use the ratio of the

concentrations of these ions and the Nernst equation for the Ce4+

/Ce3+

half cell

will be used.

h) Potential after the addition of 100.1 cm3 of 0.1 M cerium (IV)

Postequivalence point potential.

The added excess (0.1 cm3) of the titrant will be as Ce

4+ while the earlier 100

cm3 of the titrant would have got reduced to Ce

3+ ions.

17

Redox Titrations

(approx.)V1.20.181.45

0.1

100log0.0590.145

(approx.)0.1

100

][Ce

][Ce4

3

=−=

−=

=+

+

E

i) Potential after the addition of 101 cm3 of 0.1 M cerium (IV)

(approx.)1

100

][Ce

][Ce4

3

=+

+

(approx.)V1.330.121.45

1

100log0.0590.145

=−=

−=E

j) Potential after the addition of 110 cm3 of 0.1 M cerium (IV)

(approx.)V1.390.061.45

10

100log0.0590.145

(approx.)10

100

][Ce

][Ce4

3

=−=

−=

=+

+

E

k) Potential after the addition of 190 cm3 of 0.1 M cerium (IV)

(approx.)V1.4501.45

90

100log0.0590.145

(approx.)90

100

][Ce

][Ce4

3

=−=

−=

=+

+

E

The above results of the calculations can be plotted as a titration curve as shown

in Fig. 10.2.

Fig. 10.2: Calculated titration curve of 100 cm3 of 0.1 M iron (II) with 0.1 M cerium (IV)

solution

18


Redox and

Complexation

Equilibria Studies

The general appearance of the curve is similar to that for an acid base titration. It may

be noted that the potential, which changes slowly in the early stages of the titrations,

changes very abruptly as the titration passes through the equivalence point. The

magnitude of potential change evidently depends on the difference between the

standard potentials of the two redox systems that are involved; the greater the

difference, the greater is the change of potential at the equivalence point. Further, the

curve is symmetric about the equivalence point when oxidants and reductants react in

equimolar amounts or exchange same number of electrons per molecule (as in the

above titration). The curve would, however, would not be symmetric about the

equivalence point when the two redox systems exchange different number of electrons

per molecule.

Redox titration curves are usually independent of dilution because Nernst equation

contains the ratio of the concentrations of reduced and oxidised forms, which does not

alter with dilution. It is worth mentioning here, that the break in the potential on the

titration curve can be used in the selection of an indicator for the redox titration.

SAQ 3

The solution potential at the equivalence point in a titration of 0.1 M Ce (IV) solution

with 0.1 M Fe (II) solution has been worked out to be 1.10 V. What will be the

potential at equivalence point if the concentration of titrant and the titrand were

0.01 M each?

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

10.5 REDOX INDICATORS

A redox indicator is a compound which can undergo a redox reaction and change

colour depending on the potential of the solution. The oxidised and reduced forms of

the indicator have different colours. The indicator reaction may be represented as

follows.

Inox + ne Inred

(colour X) (colour Y)

where, Inox is the oxidised form of the indicator having colour X and Inred is the

reduced form of the indicator having a different colour, Y. The potential of the above

half reaction can be represented in terms of Nernst equation as given below.

][In

][Inlog

0.059

ox

red0red/ox

nEE −=

E0

Ox, Red is the standard reduction potential for the indicator half reaction. As the colour

intensity is proportional to the concentration of the species causing it, the Nernst

equation for the indicator half cell may be rewritten as given below.

Xof colour Intensity

Yof colour Intensity log

0.0590red/ox

nEE −=

Thus, the redox indicator will show colour X if the intensity of this colour (oxidised

form) is at least 10 times greater than the intensity of the colour Y. Similarly, the

indicator will show colour Y if the intensity of the colour is 10 times more than that of

the colour X. In other words colour X will be seen when the solution potential is as

follows.

19

Redox Titrations

nEE

nEE

0.059

10

1log

0.059

0red/ox

0red/ox

+=

−=

And colour Y will be seen when the solution potential is as follows.

10log0.0590

red/oxn

EE −=

nEE

0.0590red/ox −=

At intermediate potential values of the solution, the indicator will show a mixture of

colour X and Y. From these expressions we can conclude that a redox indicator

involving 1 electron in the half reaction requires a potential change of 2 × 0.059/1

(~0.12 V) to change colour from one form to the other. For indicators involving 2

electrons in their half-reactions, a potential change of 2 × 0.059/2 (~0.060 V) will be

sufficient to cause the colour change. The potential range of the colour change for a

redox indicator can be calculated theoretically on the basis of their standard potential

values as, 0red/oxE ± 0.059/n V.

Let us take an example of the titration of iron (II) with cerium (IV) in acidic medium

having 1M sulphuric acid. In this titration 1, 10-phenanthroline iron (II) sulphate is

used as an indicator. The indicator is called Ferroin and is prepared by dissolving

1,10 phenanthroline and iron(II) sulphate in 3:1 molar ratio. Three molecules of 1,10

phenanthroline are coordinated to an iron(II) cation to form red octahedral indicator

complex which is frequently written as [ Fe (Phen)3]2+

.

N

N

N

N

3 + Fe2+

3

Fe

2+

The colour change of the indicator is due to the oxidation of Fe (II) to Fe (III) in the

indicator. The reduced iron (II) form of the indicator has an intense red colour

whereas the oxidised iron (III) indicator complex is pale blue. The indicator half-reaction can be written as given below.

[ Fe (Phen)3 ]3+ + e [ Fe (Phen)3 ]

2+ ;

Why don’t you calculate the range of solution potential at which the ferroin indicator

would change colour.

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

V1.060

=E

20


Redox and

Complexation

Equilibria Studies

You seem to have calculated correct, it is 1.00 to 1.12 V, as the indication involves only one electron, it will show one colour at E = E0 – 0.06 V (1.06 – 0.06 = 1.00 V) and other at E = E

0 + 0.06 V (1.06 + 0.06 = 1.12 V ).

As has been discussed earlier, the potential at the equivalence point is the mean of the

two standard redox potentials. The potential at equivalent point for the titration of 0.1M Fe(II) with 0.1 M cerium(IV) solution will be 1.10 V. Ferroin is a suitable indicator for the titration as it changes from deep red to pale blue at a redox potential

of 1.12 V. The indicator will therefore be present in the red form after the addition of just 0.1 percent excess of cerium (IV) solution, the potential rises to 1.27 V and the

indicator is oxidised to pale blue form. The titration error is evidently negligibly small.

It is important to mention here that diphenylamine is one of the earliest redox indicators. It was introduced by Knoop for the titration of iron (II) with potassium

dichromate. The reduced form of the indicator is colourless and the oxidised form has

a deep violet colour. In the presence of a strong oxidising agent, diphenylamine undergoes the following reactions:

N N

N

H

N

H

N

H

2

Phenylbenzidine

(colourless)

+ 2H+ + 2 e

E0 = 0.76'

(colourless)

Diphenylbenzidine

(violet)

Diphenylamine

+ 2H+ + 2 e

The first reaction involving the formation of colourless diphenylbenzidine is irreversible, the second, however, giving a violet product, is reversible and is the true

indicator reaction. However, this indicator suffers from the drawback that its solution

must be prepared in concentrated sulphuric acid. The sulphonic acid derivatives of diphenylamine are frequently used instead.

HO3S N

H

Diphenylamine sulphonic acid

The barium or sodium salt of this acid are soluble in water and may be used to prepare

aqueous indicator solutions. The standard potentials values of some redox indicators along with the colour of the oxidised and reduced form are compiled in Table 10.2.

Table 10.2: The standard transition (reduction) potential values and the colour

of the oxidised and reduced form of some redox indicators

Indicator Colour of the reduced form

Colour of the oxidised form

Std. transition potential, E0

Indigo tetrasulfate Colourless Blue 0.36

Methylene blue Colourless Blue 0.53

Diphenylamine Colourless Violet 0.76

Barium diphenylamine Colourless Red-violet 0.84

21

Redox Titrations sulfonate

Erioglaucine A Yellow- green Blue- red 1.00

Tris (5-methyl-1,10-phenanthroline) iron (II)

sulfate

Red Pale-blue 1.02

Tris (1, 10-phenanthroline)

iron (II) sulfate

Red Pale-blue 1.11

Tris (5-nitro-1,10-phenana-

throline) iron (II) sulfate (nitroferroin)

Red Pale-blue 1.25

SAQ 4

Compute the potential at which the indicator 5 – nitro ‒ 1, 10 phenanthroline would acquire the colour of the oxidised form of the indicator. The indicator reaction can be

given as under.

[ Fe (5- Nitro-Phen)3 ]3+ + e [ Fe (5- Nitro-Phen)3 ]

2+ ; E0 = 1.19 V

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

10.6 REDOX TITRATIONS IN NONAQUEOUS SOLVENTS

The discussion so far on the redox reaction was confined to aqueous solution. However, as a number of analytical determinations are carried out in nonaqueous

media, it would be worth while to learn about the redox titration in nonaqueous solutions. The scope of redox reactions can be extended by using nonaqueous solvents as media of titrations in place of water. Redox titrimetry in nonaqueous media has a

wide scope for the determination of compounds which have the following

characteristics.

i) The compounds are insoluble in water.

ii) The compounds react with water through hydrolysis or oxidation.

iii) The compounds are decomposed by the medium of aqueous redox titrations i.e.,

acids, bases, etc.

10.6.1 Criteria for Solvent Selection

The following are the properties that are to be considered while selecting a solvent for

use in nonaqueous redox studies.

i) The solvent should be resistant to the attack by the reagent. This permits

reactions to be carried out over a wide range of potentials. This is so because some of the nonaqueous solvents may extend the range towards reducing

conditions while others may do so towards oxidising conditions. For example,

dimethyl formamide can be used as a solvent for redox titrations involving powerful reductants, such as Cr(II) but is oxidised even by mild oxidising agents. On the other hand, nitrobenzene is resistant to oxidation but susceptible

to reduction whereas, acetonitrile is resistant to both oxidation and reduction.

This makes it a good choice as a solvent for redox determinations in nonaqueous media.

22


Redox and

Complexation

Equilibria Studies

ii) The solvent should be capable of keeping reactants and, if possible, products as well in solution. As you know electron transfer reactions tend to be slow unless appreciable portions of the reactants are present in the ionic form. The dielectric

constant of the solvent must therefore be kept very high so as to minimise ion

pairing that slows down the reaction rates. Acetic acid (dielectric constant 6.2)

has good stability but is poor in providing sufficient salt dissociation. Acetonitrile is resistant to oxidation or reduction, on the other hand the high dielectric constant (36) of acetonitrile encourages salt ionisation.

iii) In addition to the above considerations the solvent should be easy to handle. Let us take up some such issues related to the ease of handling.

a) Some solvents may have favourable resistance and dielectric characteristics but are sometimes difficult to purify. A significant source of error and hence such solvents are ignored.

b) Some solvents tend to react with or absorb atmospheric components and need to be protected from exposure to air, e.g. acetonitrile and dimethylformamide take up water even on brief exposure to atmosphere

and ethylenediamine tends to absorb carbon dioxide.

c) Many solvents like ammonia, sulphur dioxide, hydrogen sulphide,

nitrosyl chloride etc. have low boiling points and require low temperature conditions. As a solvent working at room temperature is an

ideal choice, such solvents with low boiling points are not desirable.

Similarly, the high melting points of a number of solvents likes N-methylacetamide makes them unfit for the purpose.

d) Some solvents are toxic and hence are discouraged for use in analytical

applications. For examples, hydrogen cyanide (b.p. 25.7o C, dielectric

constant 106.8) would be very useful solvent but the problems associated

with its handling and disposal due to its toxicity make it generally undesirable medium.

10.6.2 Characteristics of Common Nonaqueous Solvents

The above discussion has raised issues regarding the desirable attributes of a solvent to

be employed for redox determination. Let us now learn about the characteristics of the commonly employed nonaqueous solvents. Some of the commonly used solvents for

redox determination are given below.

Acetic acid: It has a low dielectric constant, is difficult to purify and is not convenient to handle. That is, it fails to meet the criteria mentioned in the previous subsection.

Yet, it has been more widely used as an analytical medium than any other solvent; this

probably is because of its availability at low cost and its history of analytical use as solvent for acid-base titrations. The most common oxidation-reduction titrant used in

this solvent is lead (IV) acetate. The oxidations by lead (IV) acetate are slow and require the presence of an acid or base during the titration to increase the rate of the

reaction. Perchloric acid and Sodium acetate (strong base in acetic acid) have been found satisfactory for this purpose. These compounds help in the formation of ionic

species in acetic acid, which are able to undergo rapid reactions.

Acetonitrile: It is a versatile solvent for a large number of interesting analytical

electron transfer reactions. The reasons for its popularity include a convenient liquid range, ready availability (as a by product of acrylonitrile synthesis), wide electrochemical range, moderate dielectric constant and low toxicity. However, it has

the disadvantage of being difficult to purify and has a tendency to polymerise in the

presence of acids and bases. It is a powerful solvent for a wide range of organic

compounds and a number of inorganic salts, though it solvates anions of low

23

Redox Titrations polarisability relatively poorly with the result that many anionic metal chlorides for example, are slightly soluble. On the other hand, salts with higher polarisable anions such as iodide and thiocyanate, and predominantly covalent chlorides such as zinc

chloride typically are quite soluble. The purification of acetonitrile is difficult and

time consuming. The available potential range in acetonitrile extends from about

+ 2.2 V to – 3.1 V versus Ag/Ag+ (0.01 M) as reference couple.

Dimethylformamide: It is a solvent for which only a few applications as an analytical

redox solvent have been reported but is of considerable promise. It has a dielectric constant close to that of acetonitrile, nitromethane and dimethyl acetamide. It

dissolves many organic and inorganic compounds, has a wide liquid range and is

thermally stable. In addition, it is readily available; and it has low toxicity upon inhalation or on contact with skin. The useful electrochemical range extends past – 2.5 V versus the aqueous saturated calomel electrode, making it useful for a variety of

studies. However, susceptibility to oxidation limits its utility in the anodic direction.

The calomel electrode is not stable in this solvent and an aqueous saturated calomel electrode with a suitable bridge or a silver-silver ion electrode is the best choice in

electrochemical work. Purification of dimethylformamide is however fairly

straightforward.

Dimethylsulphoxide: It has received considerable attention as an electrochemical

solvent and has some promise for homogenous analytical oxidations/reductions. Its

excellent solvent properties for organic compounds and ability to act as a mild oxidant in conjunction with a number of reagents have made it of interest recently in organic

synthesis. It is also a good solvent for many organic compounds and has a potential range of over 4.0 V. Dimethylsulphoxide, though not very toxic, still careful handling of toxic substances in conjunction with this solvent is desirable.

Many other solvents such as acetone, lower aliphatic alcohols tetrahydrofuran, dioxane, morpholine, nitromethane and nitrobenzene have been used in nonaqueous

redox titrimetry, but they only find limited applications.

SAQ 5

Enlist the conditions under which it becomes pertinent to use nonaqueous medium for the titrimetric determination of oxidising/reducing agents.

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

…………………………………………………………………………………………...

10.7 APPLICATIONS OF REDOX TITRATIONS

Having learnt about the basics aspects of redox reaction, the parameter characterising

these reactions and their determination, the meaning and significance of redox titration, we are now equipped to learn about the applications of redox titrations. Let

us learn about some common oxidimetric and reductimetric applications of redox titrations. In oxidimetric titrations the oxidising agents are used to determine the

reducing and while is reductimetric agents are employed for the determination of

oxidising agents.

24


Redox and

Complexation

Equilibria Studies

10.7.1 Oxidimetric Reagents

i) Potassium permanganate

Potassium permanganate is a powerful and versatile oxidising agent as it can be

used in acidic, neutral or alkaline medium. It has specific applications in

different media. In acidic medium, KMNO4 is a strong oxidising agent as it gets reduced to Mn

2+ ions as per the following reaction.

V1.51O4HMn5e8HMnO 02

24 =+→++

++− E

In neutral medium potassium permanganate is reduced to manganese dioxide

and again acts as an oxidising agent.

V1.68OH2MnOe3H4MnO 0224 =+→++

+− E

On the other hand in strongly alkaline medium, the oxidising property of

permanganate is attributed to its reduction to mangnate ion.

V.670MnOeMnOo-2

4 4=→+

−E

In typical determination of reducing agents using KMnO4, a standard solution of

KMnO4 is prepared by dissolving a slightly more than the calculated amount in

required volume of water, boiling for half an hour or so and filtering off the precipitated MnO2. The solution so obtained is then standardised against sodium oxalate (a primary standard). KMnO4 not being a primary standard is a big

disadvantage. Besides, while preparing, its solution needs to be filtered off MnO2 after prolonged period of storage also. More so its solution needs to be

stored in dark so as to check formation of MnO2 catalysed by light. The

standardization of KMnO4 by oxalic acid can be represented by the following

equation,

O8H2Mn10CO16H2MnOO5C 22

24242 ++→++

++−−

The oxidation of oxalate with permanganate at room temperature is slow and

hence heating is required. The end point is marked by pink colour imparted to

the solution by the first drop of oxidant added in excess. In this way KMnO4 acts as a self indicator. This property of KMnO4 is an added advantage. The

oxidant is commonly used for the analysis of iron(II), As(III), H2O2 and −

2NO

ion, the corresponding equations being as follows.

O4HMn5Fe8HMnO5Fe 223

42

++→+++++−+

OH6Mn4O5AsH12MnO4O5As 22

52432 ++→++++−

OH8Mn25OH6MnO2O5H 22

2422 ++→++++−

OH3Mn25NOH6MnO25NO 22-

34-2 ++→++

++−

It is important to mention that a difficulty arises when iron (II) is titrated with permanganate in the presence of hydrochloric acid. High results are obtained

due to the oxidation of chloride ions by the oxidant.

O8H5Cl5Mn16H10Cl2MnO 222

4 ++→++++−−

This reaction which normally does not occur rapidly enough to cause serious

error, is induced by the presence of iron (II). Its effect can be avoided by preliminary removal of chloride. In alkaline medium, permanganate easily

oxidises iodides, cyanides, thiocyanates and many organic substances.

25

Redox Titrations ii) Cerium (IV) salts

Cerium (IV) sulphate is a powerful oxidising agent in acidic medium.

++→+

34 CeeCe

The E0 values are 1.70 V in HClO4 ; 1.60 V in HNO3 and 1.42 V in H2SO4

solutions. Ceric sulphate is a primary standard and solutions are remarkably

stable. It can be stored almost indefinitely and may be employed in the

determination of reducing agents in the presence of high concentrations of hydrochloric acid, in contrast to potassium permanganate. The reagent has an intense yellow colour and if the solutions are not too dilute, the end point may

be detected without an indicator. However, it is not used as a self indicator and it is preferable to use a suitable indicator such as ferroin, N-phenylanthranilic

acid or 5,6- dimethylferroin. Cerium(IV) solutions are standardised with

arsenic(III) oxide or with sodium oxalate solution. Cerium(IV) solution may be prepared by dissolving Cerium(IV) sulphate but usually ammonium hexanitratocerate (NH4)2[Ce(NO3)6] is used for this purpose.

iii) Potassium dichromate

It is a strong oxidant in acidic medium; the reduction equation being the following.

Cr2O7 + 14H+ + 6e 2Cr3+ + 7H2O2- ; Eo = 1.33 V

Potassium dichromate is a primary standard and is obtainable in a high state of

purity. Its standard solution can, therefore, be prepared by direct weighing. The end point in titrations with dichromate ion is generally detected by using

suitable redox indicators such as diphenylamine, sodium diphenylamine-

sulphonate and N-phenylanthranilic acid, etc.

Although, potassium dichromate is a weaker oxidising agent than KMnO4, it

can be used in acidic medium to determine most compounds that can be titrated

with KMnO4 . Potassium dichromate possesses an advantage over permanganate that it does not oxidise chloride ion and consequently can be used to determine

iron(II) in hydrochloric acid medium.

O7H6Fe2Cr14HOCr6Fe 2332

722

++→+++++−+

The determination of Fe(II) is probably the most important application of K2Cr2O7.

iv) Potassium bromate

It is also a strong oxidant in acidic medium.

V1.440

=EBrO3 + 6H+ + 6e Br + 3H2O- - ;

Potassium bromate is a primary standard as the reagent can be obtained in a pure state. Its standard solution can be prepared by direct weighing. The end

point in titrations with bromate is indicated by the yellowish colour due to

bromine released from the slight excess of reagent added.

BrO3 + 6H+ + 5Br 3Br2 + 3H2O- --

The end point may also be detected by irreversible bleaching of azo indicators such as methyl orange, methyl red etc. by bromine. The reagent is used in hydrochloric acid for the direct determination of arsenic(III), antimony(III),

tin(II), hydrazine, etc.

26


Redox and

Complexation

Equilibria Studies

O3HBr)(Sb3As6HBrO)(Sb3As 255

333

++→++−+++−++

O3HBr3Sn6HBrO3Sn 24

32

++→++−++−+

OH6Br23NBrO2H3N 22342 ++→+−−

There are many compounds, which do not react with bromate as such but with bromine liberated from reaction between bromate and bromide, and to accomplish this, an excess of bromide is added to the solution. Bromine thus

liberated is used for quantitatively saturating the double bond and in substitution

reactions.

C C

Br

Br

OH OH

BrBr

Br

+ Br2

+ 3 Br2+ 3 HBr

C C

The above reactions are used in the determination of unsaturated organic

compounds, phenols, aniline, salicylic acid, etc. Particular mention may however be made of the substitution reaction with 8-hydroxyquinoline (oxine)

for the indirect determination of various metals. The metal is precipitated as sparingly soluble oxinate, which is filtered off and heated with potassium bromate and bromide, resulting in the bromination of oxine.

N

OH

N

OH

Br

Br

+ 2 Br2+ 2 HBr

The unreacted bromine is determined by adding a solution of potassium iodide

and titrating the liberated iodine with thiosulphate.

v) Potassium iodate

The compound is available in a high state of purity and consequently serves as a primary standard as its standard solution can be prepared by direct weighing.

The reaction of iodate with iodide is analogous to bromate-bromide reaction.

This reaction is very useful for the generation of known amount of iodine.

O3H3l6H5lIO 223 +→+++−−

In hydrochloric acid solutions (3-6 M), the reduction of iodate occurs to iodine

monochloride and it is under these conditions that it finds its maximum use.

IO3 + 6H+ + CI + 4e ICI + 3H2O- -

27

Redox Titrations Oxidation of iodate in strong hydrochloric acid medium infact proceeds through the following stages.

IO3 + 6H+ + 6e-

I + 3H2O-

IO3 + 5I + 6H+ 3I2 + 3H2O- -

IO3 + 2I2 + 6H+ 5I+ + 3H2O-

Free iodine is liberated in the initial stages but as more iodate is added, iodine is

oxidised to iodine monochloride and the dark colour (due to iodine) of the

solution gradually disappears. The reaction has been used for the determination of many reducing agents. Starch cannot be used as an indicator as the blue

starch-iodine complex is not formed at high concentration of the acid. Carbon

tetrachloride or chloroform is added to the solution being titrated in a glass

stopperred conical flask. The end point is marked by the disappearance of the last trace of violet colour (due to iodine) from the organic solvent. The method

suffers from the drawback that vigorous shaking is required after addition of each instalment of the oxidant.

vi) Iodine

The I2/2I‒ couple is of medium oxidising power. Whereas iodine is a

weak oxidant, iodide is a relatively weak reductant. Hence, the I2/I– redox

couple can be used for the determination of both reductant as well as oxidants.

0.540=EI2 + 2e 2I -

I2 + red 2I + Ox-

If the couple to be determined has a standard potential lower than 0.54 V, the

reaction would proceed to the right (or iodine would oxidise the reduced form of the compound). In case the couple to be determined has a potential higher than

0.54 V, the reaction would proceed to the left (or iodide will be oxidised by the

oxidised form of the couple).

Iodine is sparingly soluble in water and its aqueous solution is prepared by adding potassium iodide. Though in such a solution, iodine is in the form of a

tri-iodide ion I3- , but for sake of simplicity, iodine is usually expressed as I2 in

its reactions.

2I3

-I2 + 2l -

Iodine solutions are standardised by means of arsenous oxide (a primary standard) or sodium thiosulphate (to be standardized first). Starch is generally used as an indicator; the solution turns blue at the end point.

Among the substances which can be titrated directly with a standard iodine solution are hydrogen sulphide, stannous, arsenite and sulphurous acid.

I2 + H2S S + 2I- + 2H+

Sn2+ + I2 Sn4+ + 2I-

H3AsO3 + I2 + H2O H3AsO4 + 2I- + 2H+

28


Redox and

Complexation

Equilibria Studies

H2SO3 + I2 + H2O SO4 + 2I- + 4H+2-

Many oxidants are capable of converting iodide quantitatively to free iodine

which can be titrated with a standard solution of sodium thiosulphate.

Ox + I I2 + Red-

I2 + 2S2O3

2-2I- + S4O6

2-

The end point can be detected by adding starch. The blue colour disappears at the end point. In this manner, potassium iodide can be used for the

determination of copper (II), MnO2, Cl2, Br2, −

3BrO , −272OCr , etc.

2 Cu2+

+ 4I‒ � I2 + 2CuI

MnO2 + 4I‒ + 4H

+ � Mn

2+ + 2I2 + 2H2O

Cl2 (or Br2) + 2I‒ � I2 + 2Cl

‒ (or 2 Br

− )

−

3BrO + 6I‒

+ 6H+ � 3I2 + Br −

+ 3H2O

−272OCr

+ 6I

‒ + 14H

+ � 3I2 + 3Cr

3+ + 7H2O

The method in which iodine is used for the determination of reductants is called iodimetry. The method in which iodide is used for the determination of oxidants is called iodometry.

10.7.2 Reductimetric Reagents

Standard solutions of reducing agents are not used as widely as oxidising agents,

because most of them are oxidised by dissolved and atmospheric oxygen. They are,

therefore, less convenient to prepare and use. Thiosulphate is the only common reducing agent that is stable to air oxidation and can be kept for long periods of time.

This is the reason that Iodometric titrations are so popular for determining oxidising agents. However, stronger reducing agents than iodide are sometimes required.

i) Iron

It is a weak reductant. Iron(II) is slowly oxidised by air in sulphuric acid solution and finds common use.

Its standard solution is prepared from Mohr’s salt FeSO4. (NH4)2SO4.6H2O and standardised against primary standard K2Cr2O7, or by titration with a standardised KMnO4 or cerium(IV) solution. The end point in titrations

involving iron(II) is detected by using redox indicators such as ferroin, N-

phenyl anthranilic acid etc. In the presence of concentrated sulphuric acid at 0

oC, iron(II) is used for the determination of nitrates, nitro compounds,

indophenols, etc.

ii) Chromium (II) and Titanium(III)

These two are very powerful reducing agents, but they are readily oxidised by

air and consequently are difficult to handle. The reduction equations and the

standard potential are as follows.

V0.770=EFe3+ + e Fe2+

29

Redox Titrations ++→+

23 CreCr E0 = – 0.41 V

OHTie2HTiO 2

32

2 +→+++++

E0 = 0.04 V

Chromium (II) has been used for the titrations of oxidised forms of copper, iron, silver, gold, bismuth, uranium, tungsten, etc. Titanium (III) has, however, been

used for the titrations of iron(III), copper(II), tin(IV), chromate, vandate and

chlorate.

SAQ 6

Iodine can be used for the determination of oxidising as well as reducing agents.

Illustrate with the help of examples.

…………………………………………………………………………………………...

…………………………………………………………………………………………...

10.8 SUMMARY

Chemical reactions involving oxidation-reduction are widely used in volumetric

analysis. Redox reactions are also called electron transfer reactions since oxidation proceeds with loss of electrons while reduction is accompanied by gain of electrons

and both the processes must occur together and must compensate for each other. The separation of redox reactions into its components i.e. half-reactions, is a suitable way of indicating species that gains electrons and the one that looses electrons.

Oxidising and reducing agents may differ among themselves in strengths. For

obtaining comparative values of their strengths, it is essential to measure under standard experimental conditions, the potential difference between the platinum and their solutions relative to a standard of reference. The reference electrode is the

standard hydrogen electrode and its potential is taken as zero. The standard experimental conditions for redox systems are those in which the ratio of the activity

(or the molar concentration) of the oxidant to that of the reductant is unity.

An understanding of the oxidising and reducing tendency of a substance can be obtained from an electrochemical cell, which develops an electrical potential from the

tendency of the reacting species (to transfer electrons) and thereby approach the

condition of equilibrium. The voltage produced by an electrochemical cell is directly related to the equilibrium constant for the particular oxidation- reduction process

involved as well as the extent to which the existing concentration of the participants

differ from their equilibrium values. Measurement of such potential, in fact constitutes an important source of numerical values for equilibrium constants for redox reactions.

The electrode potential which is a measure of the chemical driving force of a half-

reaction is affected by the concentration. The dependence of the redox potential of a

system on the ion concentration of the reduced and oxidised forms is expressed by the Nernst equation.

As regards the suitability of a redox reaction to volumetric analysis, it is desirable that the equilibrium is reached very rapidly following each addition of the titrant and the

indicator is available for locating the end-point with reasonable accuracy. The

equilibrium constant of a redox reaction which provides information regarding the direction and extent of the reaction, can be derived by the application of the Nernst

equation with the standard potential data serving as a guide.

The redox indicators are used to determine the end point in a redox titration. These are the substances which undergo redox reactions and their oxidised and reduced forms

differ in colour from each other. The colour-changes of redox indicators can be

30


Redox and

Complexation

Equilibria Studies

calculated theoretically on the basis of its standard potential values. A good redox indicator should have a sharp colour change, preferably reversible, at the equivalence point. Hence it is extremely important to examine the changes in the potential that

occur in the course of a redox titration and pay particular attention to those changes

that are more pronounced in the region of the equivalence point. In brief, the potential

of the redox system at the equivalence point is of a particular importance from the standpoint of indicator selection.

The scope of the redox reactions can be extended by using organic solvents as media of titration in place of water. Redox titrimetry in nonaqueous media has a wide scope

for the determination of compounds which are (i) insoluble in water, or (ii) react with

water through hydrolysis or oxidation, or (iii) are decomposed by the media i.e. acids, bases, etc. of aqueous redox reactions.

Potassium permanganate, cerium(IV), potassium dichromate, potassium bromate,

potassium iodate and iodine are commonly used oxidimetric agents. These are used to determine the oxidising agents on the other hand iodide, iron (II), chromium(II) and

titanium(III)) are commonly used agents for determining the reducing agents that is

theses are reductimetric agents.

10.9 TERMINAL QUESTIONS

1. Describe overall redox reaction and then half- reactions involved when a piece

of copper metal is placed in a solution of silver ion and the copper is coated

with metallic silver and the solution acquires a characteristic blue colour.

2. A solution is 310− M in −272OCr and 210− M in Cr

3+. If the pH is 2.0, what is the

potential of the half reaction? You may consult Table 10. 1 for the reaction

involved.

3. Calculate the potential for a platinum electrode immersed in a solution that is 210− M in KBr and 310− M in Br2. the electrode reaction is as follows.

4.

5. Calculate the redox potential of Sn4+

, Sn2+

system, if the Sn 4+

ion concentration is 0.1 g ion/dm3 and the Sn2+ ion concentration is 0.0001 g ion/dm3.

6. Consult Table 10.1 on the standard reduction potentials for the following

substances and list the oxidising agents in the decreasing order of oxidising capability and reducing agents in the decreasing order of reducing capability.

−

4MnO , Ce3+

, Cr3+

, IO −

3 , Fe3+

, I‒, H

+, Zn

2+.

7. Calculate the equilibrium constant for the reaction:

MnO4 + 5Fe2+ + 8H+ Mn2+ + 5Fe3+ + 4H2O-

Given :

V0.77

V1.52

0

2Fe/3Fe

0

2/Mn4MnO

=

=

++

+−

E

E

8. Calculate the potential at the equivalence point in the titration of tin(II) with

cerium(IV):

-Br2 (aq) + 2e 2Br V1.0870

=E

31

Redox Titrations Sn2+ + 2Ce4+ Sn4+ 2Ce2+

Given:

V1.44E

V0.15E

0

Ce/Ce

0

/SnSn

34

24

=

=

++

++

9. A solution of Fe2+

is titrated with an oxidising agent. Calculate the potential of

the Fe3+

, Fe2+

couple when the following percentages of Fe2+

have been oxidised.

a) 10% (b) 25% (c) 33% (d) 50% (e) 99% (f) 99.9%

Given:

V0.7723 Fe/Fe

0=++E

10.10 ANSWERS

Self Assessment Questions

1. The reduction potential for Fe (III) is more positive as compared to that of Cr

(III) which means that Fe (III) is easier to be reduced than Cr (III). This in turn

means that Cr (II) would be easier to be oxidised.

2. From Table 10.1 the reduction of MnO4 – to Mn2+ is given by the following

equation

MnO4 + 8H+ + 5e Mn2+ + 4H2O-

EO = 1.52 V

The Nernst equation would be

84

2

)cell(cell]H[]MnO[

]Mn[log

5

0.059+−

+θ

−= EE

As the pH is 1.00, the concentration of H+

ions would be 1 × 10–1

Substituting the values in the equation

VE 43.1]1.0[]100.0[

]100.0[log

5

0.0591.52

8cell =−=

3. The potential at the equivalence point is given as 2

21 EE + and it does not

depend on the concentration of the reactants. Therefore the potential would be

1.10 V.

4. As the potential range of the colour change for a redox indicator is given as, 0

red/oxE ± 0.059/n V; where, n is the number of electrons involved in the

indicator redox reaction. As only one electron is involved in the redox reaction

of the given indicator, the potential range would be

0.1 ± 0.059/n = 0.1 ± 0.059/1 = 0.1± 0.059 = 1.18 to 1.25

� the indicator range would be 1.18 – 1.25 V

32


Redox and

Complexation

Equilibria Studies

5. It is advisable to use nonaqueous medium of titration for the determination of

compounds which

i) are insoluble in water.

ii) react with water through hydrolysis or oxidation.

iii) are decomposed by the medium of aqueous redox titrations i.e., acids,

bases, etc.

6. The I2/2I– couple is of medium oxidising power. Therefore, iodine acts a weak

oxidant and gets reduced to I– ions. This can be used for the determination of

weak reducing agents. For example,

Sn2+

+ I2 � Sn4+ + 2I‒

The method in which iodine is used for the determination of reductants is

called iodimetry. Similarly, I– ions can get oxidised to I2 and can be used for

the determination of weak oxidising agents. In such cases, the determination is

performed indirectly. The oxidants quantitatively convert iodide ions to free

iodine which is then titrated with a standard solution of sodium thiosulphate.

For example,

2 Cu2+

+ 4 I‒ � I2 + 2 CuI

−+

2322 O2SI � −

+264OS2I

The method in which iodide is used for the determination of oxidants is called

iodometry.

Terminal Questions

1. The overall reaction would be as follows.

While the half reactions would be as follows.

2. 1.06 V

3. 1.117 V

4. 0.24 V

5. The decreasing order of the oxidising capability of the oxidising agents

is: −

4MnO , −

3IO , Fe3+

, H+, Zn

2+.

The decreasing order of the reducing capability of the reducing agents is as

follows.

I‒, Cr

3+, Ce

3+

6. K = 3 × 1063

7. E = 0.5 8 V

8. (a) 0.71 V (b) 0.74 V (c) 0.75 V (d) 0.77 V (e) 0.89 V (f) 0.95 V

Ag(s)(aq.)Cu(aq)2AgCu(s) 2+→+

++

(s) Ag2e2(aq)Ag2

2e(aq)CuCu(s)2

→+

+→

+

+

Documents

Unit 10 Redox Titrations