Redox reaction is the sum of half-reactions 1 and 2 and involves the simultaneous loss and gain of electrons:

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1 Chapter 4: - (Redox) Titrations A redox titration is a type of titration based on a redox reaction between the analyte and titrant. Redox titration may involve the use of a redox indicator and/or a potentiometer. 1. Redox Reaction Whereas acid-base reactions can be characterized as proton-transfer processes, the class of reactions called oxidation-reduction, or redox, reactions are considered electron transfer reactions. The term oxidation reaction refers to the half-reaction that involves loss of electrons. A reduction reaction is a half-reaction that involves gain of electrons. Thus an oxidizing agent (or oxidant) is one which accepts electrons while a reducing agent (or reductant) is one which loses the electrons. An oxidizing agent oxidizes the other substance by stripping off electrons from it. A reducing agent reduces the other substance by donating electrons to it. and reduction reactions always occur simultaneously. One cannot take place in isolation from the other. During a redox reaction the oxidizing agent itself undergoes reduction while the reducing agent undergoes its oxidation. Thus, reaction is represented as: Reductant1 (reducing agent) Oxidant1 (oxidizing agent) + ne; half-reaction 1 reaction is represented as: Oxidant2 (oxidizing agent) + ne Reductant2 (reducing agent); half-reaction 2 Redox reaction is the sum of half-reactions 1 and 2 and involves the simultaneous loss and gain of electrons: Example: Therefore, Reductant1 + Oxidant2 Oxidant1 + Reductant2; Fe 2+ Fe 3+ + e Ce 4+ + e Ce 3+ Fe 2+ + Ce 4+ Fe 3+ + Ce 3+ () () (Redox reaction) redox reaction of a substance leads to increase in its oxidation number while of a substance lead to decrease in the oxidation number. The relative tendency to accept or lose electrons by any reagent is measured in terms of their standard reduction potential values. 2. Redox Titrations (or Indicators) In the redox titrations, we need a chemical species that can change color in the potential range corresponding to the sharp change at the end point. A chemical substance, which changes color when the potential of the solution reaches a definite value, is termed as an oxidation-reduction or redox indicator. A half-reaction can be written for the indicator: In Ox + ne In Red Color A Color B A redox indicator may be defined as a substance whose oxidized form is of different color from that of its reduced form. The oxidation and reduction of the indicator is readily reversible. So the ratio [InRed]/[InOx], and therefore the color will change as the potential of the solution changes. The redox indicator reaction must be rapid and eversible, if the reaction is slow or is irreversible, the color change will be gradual and a sharp end point will be not detected

2 The main redox titration (or indicators) types are: Redox Titration (or Indicators) Iodometry Permanganometry Chromotometry Cerimetry Bromatometry Titrant Iodine (I2) Potassium permanganate (KMnO4) Potassium dichromate (K2Cr2O7) Cerium (IV) salts Bromine (Br2) 2.1. Iodometry The titration of iodine against sodium thiosulfate, using starch as the indicator of color change, is one of the most accurate volumetric redox processes. The descriptive term for the titration procedure depends on which reagent is used as the titrant. If iodine, I2, is used as the titrant, then the process is termed an iodimetric titration (Iodimetry: a direct titration with only one reaction). Analyte (unknown) + Titrant (Iodine: known) Product If on the other hand thiosulfate, S2O3 2-, is used as the titrant, then this type of titration is termed an iodometric titration (Iodometry: Not a direct titration because there are two reactions). Analyte (unknown) + I - I2 I2 + Titrant (standard thiosulfate: known) Product In either iodimetric and iodometric titration, the principal reaction is the oxidation of thiosulfate by iodine to produce iodide ion, I -, and the tetrathionate ion, S4O6 2-. This process is showed in the following reaction: I 2 (aq) + 2S 2O 3 (aq) 2I (aq) + 2S 4 O 6 (aq) The brown color of molecular iodine in an aqueous solution is sufficiently intense to serve as an indicator of color change, because the brown color will begin disappearing at the same time as I2 is consumed, but this color change is possible only if there are no other colored substances present to interfere. Usually though, an indicator is preferred, and starch is commonly used for this purpose. "Soluble" starch forms an intensely blue-colored complex with molecular iodine. Even traces of iodine produce a visible color, making an indicator blank unnecessary. The blue color of the complex disappears if the solution is heated, but returns again with cooling. When iodine is titrated with thiosulfate (an iodometric titration), starch should be added only after most of the iodine has been consumed; otherwise, the disappearance of the blue color at the end point is sluggish. Example: Iodometric Titration of Copper (Cu) An excess of iodide ions react with copper(ii) ions thus: The reaction of iodine with thiosulphate ions is: 2Cu 2+ (aq) + 4I excess(aq) 2CuI (s) + I 2 (aq) I 2 (aq) + 2S 2O 3 (aq) 2I (aq) + 2S 4 O 6 (aq) 2.2. Permanganometry Permanganatometry is a method based on a redox reaction, which is often used in analytical chemistry for determining the concentration of various compounds. A useful property of potassium permanganate solution is its intense purple color, which to serve as an indicator for most titrations

3 Potassium permanganate is a very strong oxidizing agent and is employed in the estimation of reducing agents like ferrous salts, oxalic acid, arsenious oxide, etc. The permanganate ion, MnO4 -, gets reduced to Mn 2+ ion in acidic medium according to following equation: MnO 4 (intense purple color) + 8H + + 5e Mn 2+ (colourless) + 4H 2 O +VII Titrations involving potassium permanganate are usually carried out in acidic medium. This is due to higher oxidizing power of permanganate ion in acidic medium than in neutral or alkaline medium; secondly, the formation of brown colored, MnO2 in alkaline medium interferes with the detection of the end point. For acidification of KMnO4 solution, only H2SO4 is suitable whereas the other mineral acids like HCl and HNO3 are not. HCl is not used because some of the KMnO4 will oxidize Cl - ions to chlorine gas according to the following equation and thus interferes in the quantitative estimations of reducing agents: 2MnO H Cl 2 Mn H 2 O + 5Cl 2 Nitric acid (HNO3) cannot be used because it is itself a strong oxidizing agent and may oxidize the reducing agent, thereby introducing error. Example: Permanganometry Titration of Iron (Fe) In this redox titration you will use potassium permanganate, KMnO4, as the titrant in the analysis of an unknown sample containing iron. In acidic solution, potassium permanganate rapidly and quantitatively oxidizes iron(ii) to iron(iii), while itself being reduced to manganese(ii). The half reactions for the process are: MnO 4 + 8H + + 5e Mn H 2 O; 5 x (Fe 2+ Fe 3+ + e ); When these half-reactions are combined to give the overall balanced chemical reaction equation, a factor of five must be used with the iron half-reaction so that the number of electrons lost in the overall oxidation will equal the number of electrons gained in the reduction: MnO 4 + 8H + + 5Fe 2+ Mn H 2 O + 5Fe 3+ Potassium permanganate is one of the most commonly used oxidizing agents because it is extremely powerful, inexpensive, and readily available. It does have some drawbacks, however. Because KMnO4 is such a strong oxidizing agent, it reacts with practically anything that can be oxidized. This tends to make solutions of KMnO4 difficult to store without decomposition or a change in concentration. Because of this limitation, it is common to prepare, standardize, and then use KMnO4 solutions for an analysis all on the same day. It is not possible to prepare KMnO4 standard solutions determinated directly by mass: solid potassium permanganate cannot by obtained in a completely pure state due to the high reactivity mentioned above. Rather, potassium permanganate solutions are prepared to be an approximate concentration, and are then standardized against a known primary standard sample of the same substance which is to be analyzed in the unknown sample. Potassium permanganate is particularly useful among titrants since it requires no indicator to signal the endpoint of a titration. Potassium permanganate solutions even at fairly dilute concentrations are intensely colored purple. The product of the permanganate reduction half-reaction, manganese(ii), in dilute solution shows almost no color. Therefore, during a titration using KMnO4, when one drop excess of potassium permanganate has been added to the sample, the sample will take on a pale red/pink color (since there are no more analyte ions remaining to convert the purple MnO - 4 ions to the colorless Mn 2+ ions) Chromimetry Chromimetry is a method based on a redox reaction, which is often used in analytical chemistry for determining the concentration of various compounds. Potassium dichromate is also a very strong oxidizing agent. However it is not as strong oxidizing agent as permanganate. Still it is widely used in +II

4 redox titrations because of several advantages over permanganate. Unlike potassium permanganate, potassium dichromate is available in high purity and is highly stable up to its melting point. Its aqueous solutions are not attacked by oxidisable impurities and thus composition of aqueous solution does not change on keeping. The aqueous solutions are quite stable towards light. It is thus an excellent primary standard and its standard solutions can be prepared by direct weighing of an amount of it and dissolving in a known volume of distilled water. Potassium dichromate acts as oxidizing agent in acidic medium only. The neutral aqueous solution of Potassium dichromate is 1:1 equilibrium mixture of dichromate and chromate, a consequence of hydrolysis of dichromate ions according following equation: Cr 2 O 7 + H 2 O 2CrO 4 + 2H + Chromate ions are weaker oxidizing agent than dichromate. Thus oxidizing strength of dichromate is reduced in neutral solution. The above hydrolysis reaction however can be reversed by adding acid to the solution and this explains the necessity of acidic medium for the reaction. Also the reduction reaction of dichromate can be represented as: Cr 2 O 7 (orange) + 14H + + 6e 2Cr 3+ (green) + 7H 2 O +VI +III The medium is generally acidified with dilute H2SO4. Unlike permanganate case, cold HCl can be used here for acidifying the reaction mixture provided the acid conc. Does not increase beyond 1-2 M. Though the dichromate solutions are intensely orange colored solutions and a single drop of it imparts yellow color to a colorless solution, it can t be used as a self-indicator like KMnO4. This is because its reduction product (Cr 3+ ) is green which hinders in the visual detection of end point by observing dichromate color. Thus an indicator is must in this titration. The indicator should be redox active and must be properly chosen keeping in mind the electrode potential values of the reducing agent being titrated with dichromate. Suitable indicators for dichromate titrations are Diphenylamine (specifically sodium diphenylamine sulphonate) in presence of orthophosphoric acid, and N- phenylanthranilic acid. Example: Chromimetry Titration of Iron (Fe) Based on the same principle of permanganometry titration of Iron (Fe). But the endpoint of this titration as indicated earlier has to be defined with the help of an indicator. Cr 2 O H + + 6e 2Cr H 2 O; 6 x (Fe 2+ Fe 3+ + e ); When these half-reactions are combined to give the overall balanced chemical reaction equation, a factor of six must be used with the iron half-reaction so that the number of electrons lost in the overall oxidation will equal the number of electrons gained in the reduction: Cr 2 O H + + 6Fe 2+ 2Cr H 2 O + 6Fe 3+ Diphenylamine is one such indicator (internal indicator as it is added to the reaction mixture). The endpoint is marked with an intense blue violet coloration Cerimetry - Titrations involving ceric sulfate, symbolized here by Ce 4+, as an oxidizing agent are sometimes grouped under the generic term cerimetry. Most of the time, ceric sulfate is used. Ceric sulfate is a powerful oxidant and can be used only in acidic solution. In neutral solution ceric hydroxide or basic salts precipitate Ce(OH)4. Ceric sulfate have intense yellow color and endpoint detection can be possible without any indicator in hot solutions. The advantage of ceric salts (sulfate) over permanganate and dichromate as a standard oxidizing agent are:

5 - Ceric sulfate solution is indefinitely stable; the concentration does not vary in sunlight or even on boiling. - It may be employed in the titration of reducing agent in the presence of a high concentration of HCl. - In acid solution with reducing agent, the simple valence change Ce 4+ + e Ce 3+ is assumed to take place. No intermediate products are formed. - Cerous Ce(III) salt, which is the reduction product in ceric titrations, is practically colorless and hence allows a more effective use of indicators. Example: Cerimetry Titration of Oxalic Acid (H2C2O4) Ceric sulfate solution reacts with hot solution of oxalic acid acidified with dilute sulfuric acid by the following reaction: H 2 C 2 O 4 + 2Ce 4+ 2Ce CO 2 + 2H + In this reaction, an electron is transferred from H 2 C 2 O 4 to Ce 4+ to form Ce 3+ and (CO2+H + ). A substance that has a strong affinity for electrons, such as Ce 4+, is called an oxidizing agent, or an oxidant. A reducing agent, or reductant, is a species, such as H 2 C 2 O 4, that easily donates electrons to another species. We can split any oxidation-reduction equation into two half-reactions that show which species gains electrons and which loses them. H 2 C 2 O 4 2CO 2 + 2H + + 2e ; 2(Ce 4+ + e Ce 3+ ); The titration is carried out until the reaction mixture gets a pale yellow color. The ceric sulfate solution acts as a self-indicator