Redox indicator is an organic compound that exhibit a change in color in both the reduced and oxidized states, or at various potential values. Depending on oxidation and reduction potential, these organic molecules undergo structural change during redox titration.
In oxidation + n e- ⇌ In reduction
and E = E + In + RT/NF In [In oxd] / [In red]
E + In = Formal potential of indicator
E = Potential of solution
[In oxd] = Concentration of oxidized indicator
[In red] = Concentration of reduced indicator
The formal potential is the potential observed in the solution containing one mole of each of the oxidized and reduced substances along with other specified substances with a specific concentration.
Formal potential depends on the nature and concentration of the acid.
Example: the formal potential of Fe (II) – Fe (III)
E o = +0.77 V, But
E o = +0.73 V, in 1 M HClO4
E o = +0.70 V, in 1 M HCl
E o = +0.73 V, in 1 M H2SO4
The standard potential is measured under the standard potential i.e., at 25oC 1M concentration for each ion participating in the reaction with a partial pressure of 1 atm for each gas that is part of the reaction of metals in their pure state.
Characteristic of the redox indicators
- Near the equivalent point in the redox titration, a redox indicator should mark the sudden change in oxidation potential.
- At the endpoint, there should be a distinct color change.
- The oxidation-reduction process must be quick and reversible.
- To achieve a sharp color change at the endpoint, the formal potential of redox indicator should be at least 0.15 volts different from the standard potential of the other systems involved in the reaction.
- To prevent the reactant from reacting with an indicator before the endpoint the indicator potential must always be larger than that of the reactant.
- The color of the indicator must vary within a specific range of potential.
Types of the redox indicator
A colored substance may act as the self-indicator. For example, KMnO4 during the reaction between KMnO4 and Fe (II) solution. In this reaction, Fe (II) ions are oxidized to Fe (III) by MNO4- and KMnO4 itself gets reduced. The solution remains colorless before the endpoint when Fe (II) ions are completely oxidized by KMnO4. The added KMnO4 in slight excess imparts pink color to the solution.
It is a type of redox indicator that reacts with one of the reagents in a specific manner to give color. For example, starch reacts with iodine to give deep blue color.
It is not added to the solution being titrated but used outside the titrating system. For example, the ferrocyanide ion was used to detect Fe (II) ion by the formation of Fe (II) ferricyanide on a spot plate outside the titration vessel.
The redox potential can be measured during redox titration. The equivalent point is detected from the large change in potential in the titration curve. This is called potentiometric titration.
True-redox indicators (internal indicator)
These are the indicators that themselves undergoes oxidation-reduction and exhibit different color in oxidized and reduced forms. E.g., Diphenylamine, 1,10-phenanthroline
Examples of redox indicators
1. Diphenyl amine (0.76 V)
2. Diphenyl amino sulphuric acid
3. Methylene blue (0.52)
4. Ferroin (i.e., 1, 10-Phenanthroline iron (II) sulfate) (1.06 V)
It is the internal redox indicator, widely used in the titrimetric analysis. It is soluble in conc. H2SO4.
- Diphenyl amine is used for the titration between Fe (II) and K2Cr2O7.
- As the formal potential of the indicator is 0.76 V while that of Fe (III)- Fe (II) is 0.77 V, phosphoric acid is used to decrease the formal potential of the Fe (III)- Fe (II) solution.
- K2Cr2O7 is a strong oxidizing agent, so diphenylamine is oxidized first to diphenyl benzidine. Diphenyl benzidine is colorless which further oxidizes to diphenyl benzidine violet. Due to the long conjugate system in diphenyl benzidine, which causes light to be absorbed in the visible spectrum, the substance is violet in color.
Ferroin i.e., 1, 10-Phenanthroline iron (II) sulfate)
It is one of the best oxidation-reduction indicators. During the titration 1, 10-Phenanthroline and Fe (II) salt solution in a ratio of 3; 1 gives red 1, 10-Phenanthroline Fe (II) complex ion which on oxidation produced produces a pale blue color.
[ Fe ( C12 H8 N2 )3 ] 3+ + e- ⇌ [ Fe ( C12 H8 N2 )3 ] 2+
Pale blue deep red
Potassium bromate (KBrO3)
It is a powerful oxidizing agent that is commonly employed as the primary standard solution in redox titration. KBrO3 is converted to bromine throughout the titration. Excess bromine is present at the endpoint, as seen by the pale color. For better detection, methyl red and orange indicators are occasionally employed.
Bro3- + 6 H+ + 6 e- → Br- + 3 H2O
At end point,
Bro3- + 5 Br- + 6 H+ → 3 Br2 + 3 H2O
KBrO3 is commonly used for the determination of metals with 8- hydroxy quinoline.
Titration of potassium bromate can be carried out in two ways
I. Direct titration
With the KBrO3 solution, several reducing substances can be titrated directly, including arsenic (III), antimony (III), iron (III), and organic sulfides and disulfides. The presence of bromine marks the endpoint of the titration.
BrO3- + 3 HAsO2 → Br- + 3 HAsO3
BrO3- + NH2OH → Br- + N03– + H+ + H2O
II. Indirect method
Numerous compounds that cannot be directly oxidized by KBrO3 but can quantitatively react with excess bromine are indirectly titrated with KBrO3. In this titration, known amounts of bromine are produced in the presence of Br- and acid using a standard solution of KBrO3.
Cerium can exist in two oxidation states, +3 and +4, in the solution. Ce (IV) is a strong oxidizing agent and its oxidizing power varied with the pH of the solution. It serves as a strong oxidizing agent for quantitative analysis. As hydrolysis causes Ce (IV) hydroxide to precipitate at low [H+] concentrations, this can only be employed in acidic solutions. It is commonly used in the analysis of hydrogen peroxide.
2 Ce 4+ + H2O2 → Ce 3+ + O2 + 2 H+
- A.I. Vogel, qualitative analysis, E.L,B.S. , 1994.
- D.A. Skoog, Principle of Instrumental Analysis, 3rd Edition, saunders college publishing, 1985.