Oxidation and Reduction Titration, Slides of Pharmaceutical Analysis

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Oxidation-Reduction

Titration

Reference Book:

Quantitative Analysis by V. Alexeyev

Oxidation-reduction

titration

Oxidation-reduction titration is the

titrimetric method in which the

analyte is quantified via a redox

reaction with the titrant.

Basic terminologies

Reduction and reducing agent:

• Reduction is the gaining of electron of a substrate.
• Reducing agent is the substance that causes the gaining of

electron of a substance. The reducing agent itself is oxidized in

the process.

• If an agent has high oxidizing capacity, the other part of the

couple will have lower reducing capacity.

• Reducing agent/reductant: loses electron

Fe

3+

• e

• Fe

2+

M M

**+

• e- (Release (Left in solution) in metal)**

M+^ + e- M

(From (From metal) Solution)

If a metal is dipped into water or an aqueous

solution, one of the following will happen -------

Reduction potential :

The tendency of a chemical species to be reduced by acquiring electrons. It is

measured in volts.

Oxidation potential :

The tendency of a chemical species to be oxidized by releasing electrons. It is also

measured in volts.

Redox potential:

The Metal atom leaves the metal as ion and electrons are left in the metal.

The metal ion from the solution enters the metal by

gaining electron.

Standard oxidation potential

• When the concentrations of the oxidizing and reducing agents are equal

( 1 M) and the temperature is 25 ⁰C, the oxidation potential is called

standard oxidation potential.

• There are a few rules that help along with standard oxidation potential

value to determine the oxidation-reduction reaction. These are-

1. The higher the standard oxidation potential of a given system the

stronger the oxidizing power of its oxidized form and the weaker the

reducing power of its reduced form and vice versa.

2. When two redox systems are combined, the stronger oxidizing agent

will receive electrons from stronger reducing agent, with the formation of

weaker reducing and oxidizing agents.

Potentiometer

To obtain comparable results when determining standard

oxidation potential, different oxidation-reduction couples

should be paired with the same standard couple (called

standard hydrogen electrode, 2 H+/H 2 , 1 g-ion per liter/ 1 atm

pressure).

Electrolytic bridge

KCl

H 2 - 2e = 2H+^ 2Fe

++++ 2e = 2Fe++

• Let us try to explain the rules with the following example-
• Now in the above, the oxidation capacity of Cl 2 (oxidized form) is higher

thus the reducing capacity of Cl─ (reduced form) is low. Similarly, since

the oxidation capacity of Fe^3 + (oxidized form) is low the reduction

capacity of Fe^2 + (reduced form) is high. So according to the above rules,

Cl 2 reacts with Fe^2 +^ and Cl─^ will react with Fe^3 +. We need to know the

standard redox potential cause-

1. To predict the direction of any oxidation-reduction reaction
2. Choose suitable oxidizing or reducing agent
3. Possible to solve different problems in analytical practice.

Cl 2 2Cl-^ E 0 = + 1.36V

Fe3+^ Fe2+^ E 0 = +0.77V

If the hydrogen electrode is attached with a Zn++/Zn system-

Zn - 2e = Zn++^ (reaction at cathode)

2H+^ + 2e = H 2 (reaction at anode)

The standard oxidation potential of this system is found to be

• 0.76 V

Standard Reduction Potentials at

(^25)

Co

 - F2+2e−→F− +2. Half-Reaction E 0 (V) 

• PbO2+4H++SO2−4+2e−→PbSO4+2H2O +1. - MnO−4+8H++5e−→Mn2++4H2O +1. - Au3++3e−→Au +1. - Cl2+2e−→2Cl− +1.
  • Cr2O2−7+14H++6e−→2Cr3++7H2O +1. - O2+4H++4e−→2H2O +1. - Br2+2e−→2Br− +1. - NO−3+4H++3e−→NO+2H2O +0. - 2Hg2++2e−→Hg2+2 +0. - Hg2++2e−→Hg +0. - Ag++e−→Ag +0. - Fe3++e−→Fe2+ +0. - I2+2e−→2I− +0. - Cu++e−→Cu +0. - O2+2H2O+4e−→4OH− +0. - Cu2++2e−→Cu +0. - Sn4++2e−→Sn2+ +0. - 2H++2e−→H2 0. Half-Reaction E 0 (V) - Pb2++2e−→Pb - 0. - Sn2++2e−→Sn - 0. - Ni2++2e−→Ni - 0. - Co2++2e−→Co - 0.
• PbSO4+2e−→Pb+SO2− 4 - 0. - Cd2++2e−→Cd - 0. - Fe2++2e−→Fe - 0. - Cr3++3e−→Cr - 0. - Zn2++2e−→Zn - 0.
  • 2H2O+2e−→H2+2OH− - 0.
    • Mn2++2e−→Mn - 1. - Al3++3e−→Al - 1. - Be2++2e−→Be - 1.
    • Mg2++2e−→Mg - 2. - Na++2−→Na - 2.
      • Ca2++2e−→Ca - 2. - Sr2++2e−→Sr - 2.
        • Ba2++2e−→Ba - 2. - Rb++e−→Rb - 2. - K++e−→K - 2. - Cs++e−→Cs - 2. - Li++e−→Li - 3.

Influence of concentrations on redox potential

The relationship between the oxidation potential of any given

system and the concentration of oxidized and reduced form

is given by Nernst equation-

E = E 0 +

RT

nF

ln

[Ox.]

[Red.]

E = E 0 +

n

log

[Ox.]

[Red.]

Fe++^ - e Fe+++

E = Oxidation potential of the system

E 0 = Standard oxidation potential of the system

The reduction potential can be determined using following

equation-

[Oxidized form]

[Reduced form] log n

E E

[Oxidized form]

[Reduced form] ln nF

RT

E E

red (^0) red

red (^0) red

For the oxidation-reduction couple Br 2 /2Br

--

E (^) Br 2 /2Br^ --^

log

[Br 2 ]

[Br - ]^2

Br 2 + 2 e 2 Br

2 Br - 2e Br 2

ox (^0) ox 2 [Br ]

[ ] log 1

E E^2

Br   

[Br ]

[ ] log 1

E E 2

2

red (^0) red

Br  

Here the following reaction occurs-

Above is a reduction reaction. To determine the oxidation potential the

reaction must be converted to oxidation form.

Now, the oxidation potential is-

If the reduction potential is to be determined then-

It should be noted that the product is always placed in the

numerator and reactant is always placed in the

denominator. With that in mind, the (+) sign is placed

when oxidation potential is determined and (-) is placed

when reduction potential is determined.

Influence of conditions of the reaction

medium on redox potential

• The oxidation (and the reduction potential) strongly depends upon the pH of the

medium (i.e. the concentration of H+ ). This can be explained using the following equations-

In this reaction, the permanganate ion is reduced to Mn^2 +. So, this is the reduction

reaction. When we write it in the oxidized form:

As we can see from the reaction representation (both the reduced form and the

oxidized form), that proton actively participate in the reaction by gaining electron

(reduced form) or by donating electrons (oxidized form). So, the oxidation potential

is determined using the following equation-

MnO 4 8 H+^ 5 e Mn2+^ 4 H + (^) + (^) + 2 O

Mn2+^ + 4 H 2 O 5 e MnO 4 + 8 H+

[Mn ]

[MnO ][H ] log 5

E E

2

8 4 ox (^0) ox 

 

 