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Daniele Naviglio » 16.Oxidation reduction titration


Oxidation reduction titration

Volumetric analysis through oxidation and reduction, is based on reactions which involve the transfer of electrons from a molecule, atom or ion to another chemical species.

Redox reactions

These are chemical reactions whereby a reducing agent oxidizes an oxidizing agent, which is itself then reduced. An example is that of the oxidation of iron (II) ions, with cerium (IV) ions. The reaction is described in the following equation:

Ce4+ + Fe2+ ↔ Ce3+ + Fe3+

The  cerium removes electrons from the iron, and is termed the oxidizing agent as it is reduced in the process and its oxidation number decreases.

The iron, which donates electrons to the other chemical, is termed the reducing agent because it is oxidized, and thus loses electrons thereby increasing its oxidation number.

Thus we get the following definitions:

Oxidation: the loss of electrons from a molecule, atom or iron;

Reduction: the gain of electrons by a molecule, atom, or ion.

Redox reaction (cont.)

In oxidation reduction reactions, the titrant is usually a strong oxidant because if we use reducing agents there can be too much interference from the o2 present in the atmosphere.

In an oxidation reduction titration we need to know:

  • the oxidimetric equivalent weight (E.W.), i.e. the number of electrons transferred;
  • the strength or tendency of the electrons to be transferred, i.e. the intensity of the oxidization or reduction (oxidation reduction potential expressed using the Nernst equation).
Diagram showing oxidation reduction reactions

Diagram showing oxidation reduction reactions


Nernst equation

The Nernst equation expresses the electrode potential of a pair of electrodes, or of a semielement of a battery compared to standard electrode potential. In other words, the equation enables us to calculate electrode potential in non-standard conditions.

Where:

R is the universal constant of gases, equal to 8.314472 J K-1 mol-1 o 0.082057 L atm mol-1 K-1;

is the absolute temperature;

a is the chemical activity;

F is Faraday’s constant equal to 9.6485309*104 C mol-1;

n is the number of electrons transferred in the half-reaction;

[red] is the concentration of the oxidizing agent (or reduced species)

[Ox] is the concentration of the reducing agent (or oxidized species). The concentration of the oxidized and reduced species is the equivalent of the activities in dilute solutions.

Nernst equation

Nernst equation


Titrant choice

The titrant is chosen on the basis of its standard reduction potential E0 and in particular:

  • if the oxidant has an E0 value that is quite a few decivolts above that of the reducing agent: the reaction proceeds spontaneously until it reaches completion;
  • if the two E0 are very similar: the two substances react until the two reagents and the products of the reaction are in equilibrium; in other words the reaction is incomplete;
  • if the reducing agent has an E0 which is clearly higher than that of the oxidant, the two substances do not react.

Some of the titrants used


Titration curves

Titration curve for generic oxidation reduction titration. Source: ExpoMix Forum Italia

Titration curve for generic oxidation reduction titration. Source: ExpoMix Forum Italia


Oxidation reduction indicators

There is a redox system whereby the oxidized form (Ox) is a different colour from the reduced form (Red). As with acid-base indicators, the colour change occurs within a specific interval (turning interval).
In(Ox) + e- = In(Red)

The colour of In (Ox) dominates when [In(Ox)] / [In(Red)] >10

The colour of In (Red) dominates when [In(Ox)] / [In(Red)] <0.1 and it is only in this condition that the colour change can be noted.

Some reactives are highly coloured, so if they are used to excess, or if they disappear, there will be an obvious colour change. For example, the oxdized form of permanganate is an intensive violet (even for concentrations of < 10-5). I2 in the presence of a starch indicator turns an intense blue.

Examples of indicators


Permangonometry

Permanganate is a strong oxidant, and this means that it should not be used in chemical analysis because its organic substrates are not only titrated but also degraded.

In an acid environment it is a strong oxidant

MnO4- + 8H+ +5e-→ Mn2+ + 4H2O                    E0= +1.51 volt

In a neutral or slightly alkaline environment it is an even stronger oxidant

MnO4- + 2H2O +3e- → MnO2 + 4OH- E0= +1.695 volt

In permanganometry no indicator is used. The end point of titration is clearly visible because the solution turns an intense purple colour when the permanganate is in excess.

Disadvantages of permanganometry

  • Permanganate solution in a sulfuric acid solution will slowly decompose, and may eventually need restandardisation;
  • The reaction needs to be carried out at quite a high temperature to speed up the reaction, resulting in decomposition of some of the KMnO4;
  • Not very selective where reducing substances are concerned.

Advantages of permanganometry

  • Permanganate solutions are sufficiently intense in colour to serve as indicators in oxidation reduction titration;
  • The standard is cheap.

Application of permanganometry

  • Determination of oxygen in oxygenated water;
    2MnO4- + 5H2O2 + 6H+ = 2Mn2+ + 5O2 + 8H2O
  • Determination of nitrites;
    2MnO4- + 5NO2- + 6H+ = 2Mn2+ 5NO3- + 3H2O
  • Determination of bromides and iodides;
    2MnO4- + 10Br- + 16H+ = 2Mn2+ + 5Br2 + 8 H2

Applications of permanganometry

  • Dosage of Fe2+;
    MnO4-+ 5Fe2+ + 8H+ = Mn2+ + 5Fe3+ + 4H2O
  • Dosage of C2O42- (oxalate ion);
    2MnO4- + 5H2C2O4 + 6H+ = 2Mn2+ + 10CO2 + 8H2
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