Redox Titration

Redox Titration - Permanganate, Dichromate and Idometric Titration

A redox titration is a form of titration established on a redox reaction between the titrant and analyte. Redox titration can involve the use of a redox indicator or a potentiometer. A general example of a redox titration is treating a solution of iodine with a reducing agent to create iodide using a starch indicator to help spot the endpoint. Iodine (I2) can be reduced to iodide (I−) example. thiosulfate (S2O32−), when all iodine is consumed the blue color vanishes. This is process is known as an iodometric titration.




The reduction of iodine to iodide is the final step in a chain of reactions where the first reactions are used to transform an unknown amount of the solute (the material being analyzed) to an equivalent amount of iodine, which can then be titrated. Sometimes other halogens (or haloalkanes) than iodine are used in the in-between reactions as they are available in better measurable standard solutions and react more quickly with the solute. The additional phases in iodometric titration might be valuable due to the equivalence point, where the blue turns slightly colorless, is more distinct than few other analytical or can be by volumetric methods.

Sub-Divisions of Redox Titrations


  • • Permanganate Titrations

  • • Dichromate Titrations

  • • Iodimetric or Iodometric Titrations

  • 1) Permanganate Titrations


    It is an oxidizing agent in this kind of redox titration method. Maintenance of the mixture is done with the help of dilute sulphuric acid. Moreover, the addition of Sulphuric acid also helps to raise the hydrogen ions present in the solution.
    In this method, the reagent has strong color because of the permanganate ion MnO4. In this case, the permanganate ion behaves as a self-indicator in this method. The solution remains colorless earlier to the endpoint. The equation of redox reaction



    The outcome of the endpoint is visible when oxidation of the last of the reductant like Fe2+ or C2O42– happens. At this point, the solution recalls the first lasting tinge of the MnO4 (pink color) appears. The concentration can be a minimum of 10–6 mol L–1. This allows the slight “overshoot” of the pink color after the equivalence point.

    The same point is where oxidant and reductant are the same with respect to the mole stoichiometry or the total number of electrons lost and electron increased in oxidation and reduction reaction will be similar. The potassium permanganate titration helps in the estimation of ferrous salts, hydrogen peroxide, oxalic acid, oxalates and more. Still, it is very vital to always standardize the solution prior to use.

    2) Dichromate Titrations


    In this technique, potassium dichromate behaves as the oxidant in the acidic medium. It is essential to maintain the acidity of the medium by addition of dilute sulphuric acid. The equation of the reaction is




    In this condition, there is no important auto color change as observed in the MnO4titration. Cr2O2– is not a self-indicator. Still, Cr2O2– oxidizes the indicator material diphenylamine soon after achieving the equivalence point thereby making an intense blue color. The variation in signals the end point of the titration. We can use the potassium dichromate mixture in titrations directly. This technique helps in the estimation of ferrous salts and iodides.

    3) Iodometric Titrations


    This is an exciting but simple method. In this case, free iodine reduction to iodide ions happens as well as iodine ion oxidation to free iodine occurs. The reduction and oxidation reactions are


    The solution behaves as an indicator. The use of this technique is limited to the reagents capable of oxidizing I ions. Examples of such reaction are of Cu(II)

    The ability of iodine to yield an intense blue color with starch as the material and its capacity to react with thiosulphate ions (S2O32–) forms the basis of this technique. The definite reaction with (S2O32–) is also a redox reaction


    I2 is insoluble in water but it remains in the solution in the form of KI3 that has KI. After the adding of starch, the iodine in the reaction releases as iodide ions generating intense blue color. However, the color vanishes on the consumption of iodine by the thiosulphate ions. The endpoint of the reaction is easily visible. Hence, it is easy to conclude the concentration of the unknown mixture by stoichiometric calculation.

    Finding the Endpoint with an Indicator


    Three kinds of indicators are needed to signal a redox titration’s end point. The reduced and oxidized forms of certain titrants, such as MnO4–, have different colors. A solution of MnO4– is strong purple. In an acidic solution, still, permanganate’s reduced form, Mn2+, is almost colorless. When using MnO4– as a titrant, the titrand’s mixture remains colorless till the equivalence point. The first drop of excess MnO4– produces a permanent tinge of purple, indicating the endpoint.

    Certain indicators form a colored compound with an exact reduced or oxidized form of the titrant. for instance, starch forms a dark blue complex with I3–. We can use this distinct color to indicate the presence of excess I3– as a titrant—a shift in color from colorless to blue—or the completion of a reaction using I3– as the titrand—a shift in color from blue to colorless. Another example of an exact indicator is thiocyanate, SCN–, which create a soluble red-colored complex of Fe(SCN)2+ with Fe3+.
    The most vital class of indicators are materials that do not contribute in the redox titration, but whose reduced and oxidized forms differ in color. When adding a redox indicator to the titrant, the indicator imparts a color that depends on the solution’s potential. As the mixture’s potential changes with the addition of titrant, the indicator alters oxidation state and changes color, signaling the endpoint.

    Inox+ne−⇌Inred(9.4.39)(9.4.39)Inox+ne−⇌Inred

    where Inox and Inred are, respectively, the indicator’s reduced and oxidized and forms.



    Redox Titration Curves

    To estimate a redox titration we need to know the shape of its titration curve. In and a complexation titration or acid-base titration the titration curve displays how the concentration of H3O+ (as pH) or Mn+ (as pM) changes on adding titrant. For a redox titration, it is suitable to monitor the titration reaction’s potential instead of the concentration of one species.
    You may recall from earlier chapters that the Nernst equation relates a solution’s potential to the concentrations of reactants and products contributing to the redox reaction. Consider, for instance, a titration in which a titrant in a reduced state, Ared, reacts with a titrant in an oxidized state, Box.

    Ared+Box⇌Bred+Aox


    where Aox is the titrand’s oxidized state, and Bred is the titrant’s reduced state. The reaction’s potential, Erxn, is the difference among the reduction potentials for each half-reaction.

              Erxn=EBox/Bred−EAox/Ared


    After every addition of titrant, the reaction between the titrand and the titrant touches a state of equilibrium. Due to the potential at equilibrium is zero, the titrand’s and the titrant’s reduction potentials are similar.

    Previously the equivalence points the titration mixture have an appreciable quantity of the titrand’s oxidized and reduced forms. The unreacted titrant concentration is very small. Therefore, is easier to measure if we use the Nernst equation for the titrand’s half-reaction

    Erxn=EoAox/Ared−RT/nF ln [Ared]/[Aox]

    After the correspondence or equivalence point it is easier to measure the potential using the Nernst equation for the titrant’s half-reaction.

    Erxn=EoBox/Bred−RT/nFln[Bred]/[Box]

    Selecting and Evaluating the Endpoint

    A redox titration’s equivalence point happens when we react stoichiometrically equivalent quantities of titrant and titrand. As is the case with acid-base and complexation titrations, we evaluate the equivalence point of a complexation titration with the help of an experimental endpoint. A variety of techniques are available for pinpointing the end point, containing indicators and sensors that react to a change in the solution conditions. 

    Other Methods for Finding the Endpoint


    Another technique for finding a redox titration’s endpoint is with the help of potentiometric titration in which we observe the change in potential while adding the titrant to the titrant. The endpoint is seen by visually examining the titration curve. The common experimental design for a potentiometric titration contains a Pt indicator probe whose potential is controlled by the titrand’s or titrant’s redox half-reaction and a reference probe or electrode that has a permanent potential
    Other methods for determining the titration’s endpoint contain thermometric titrations and spectrophotometric titrations.

    Quantitative Applications


    Although several quantitative applications of redox titration have been substituted by other analytical methods, some important applications continue to be applicable. The general application of redox titration with an impact on environmental, pharmaceutical, and industrial applications.


    Selecting and Standardizing a Titrant
    If it is to be used quantitative purpose, the titrant’s concentration must remain stable during the analysis. Because a titrant in a reduced state is prone to air oxidation, mostly redox titrations use an oxidizing agent as the titrant. There are some common oxidizing titrants, containing Ce4+, Cr2O72–, MnO4–, and I3–. Which titrant is used often rest on how easy it is to oxidize the titrant. A titrand which is a weak reducing agent requires a strong oxidizing titrant if the titration reaction is to have an appropriate endpoint.