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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

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;

T 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

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 E⁰are 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

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). I₂in 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 KMnO_4;
Not very selective where reducing substances are concerned.