Redox titrations in pharmaceutical analysis

Redox titrations introduction, Oxidation , reduction, oxidation number , oxidation number calculations, Balance redox reactions, Oxidizing agents, reducing agents,Cerimetry Iodometry, iodimetry, Bromatometry, dichromatometry

1. Redox Titrations Archana.K

2. Introduction Oxidation-reduction reactions are reactions in which electrons are transferred from one reactant to another reactant. Oxidation is defined as the loss of electrons Reduction as the gain of electrons. They must occur simultaneously, when a substance gives up electrons, There must be another substance to receive them The first substance is oxidized and the other is reduced Substances which lose electrons are reducing agents or reductants Those which gain electrons are oxidising agents or oxidants. LEO the lion says GER

3. Electrons carry negative charge So more numbers of electrons more negative charge Example: Na + cl = Nacl (no charge) (no charge) (Na ⁺ cl⁻) Na donates electron Cl Receives electrons Sodium is reducing and cl is oxidising As Na is losing electrons the negative charge decreases and converts to positive + Cl is receiving electrons so more negative charge as more electrons are adding. How?

4. No reaction can happen in half If there is reduction there should be oxidation. Half reactions: Na + cl = Nacl How to write half reactions…let’s see.. For chloride Cl + e ⁻ = cl⁻ For sodium can we write like this. Let’s see.. Na - e⁻ = Na⁺ Something is not right here…can you guess?

5. Exactly……. Is there any “-” symbol in a reaction phase… No….. So how to write the half reaction then? Ok here we go….. Na Na ⁺ + e⁻ Here the reaction conveys the same information as it is donating electrons. Na Na ⁺ + e⁻ cl + e⁻ cl⁻ Na + cl Na ⁺cl⁻ (Nacl)

6. Trick question K ⁺ + cl⁻ = Kcl IS THIS A REDOX REACTION??? YES / NO?

7. Oxidation Number The oxidation number is basically the count of electrons that atoms in a molecule can share Rules: Elements by itself = 0 eg : Ag Group 1A = always +1 Group 2A = always +2 Halogens = usually -1, positive with oxygen Monoatomic ion = ion charge is the oxidation number (cl ⁻, cu²⁺) H = +1 with nonmetals (oxygen, carbon) -1 with metals (cu, iron) o = usually -2 -1 in peroxide F = always -1 sum of oxidation number for a neutral compound = 0 sum of oxidation number for a polyatomic ion = ion charge

8. Oxidising agent A substance that tends to bring about oxidation by being reduced and gaining electrons. Na + cl = Nacl Chlorine is reducing by gaining electrons but it is an oxidising agent Cl + e ⁻ = cl⁻ Reducing agent A substance that tends to bring about reduction by being oxidized and losing electrons. Sodium is oxidising by loosing electrons but is a reducing agent Na Na ⁺ + e⁻

9. Equivalent weight Equivalent weight of a substance (oxidant or reductant) is equal to molecular weight divided by number of electrons lost or gained by one molecule of the substance. It is not a constant quantity but depends up on the reaction it is taking place. Equivalent weight of oxidising agents = Molecular weight No.of electrons gained by one molecule Equivalent weight of reducing agents = Molecular weight No. of electrons lost by one molecule.

10. Example : KMno ₄ K⁺ + Mno₄⁻ Basic medium Mno₄⁻ + e⁻ Mno₄⁻ ² Oxygen oxidation number = -2 o₄⁻ = -2x4 = -8 Mn = ? As it is a polyatomic ion oxidation number = ion charge Mno₄⁻ + e⁻ Mno₄⁻ ² (+7) (-8) = -1 (+6) (-8) = -2 1 electron Equivalent weight = Molecular weight No of electrons gained 158 1

11. Theory of redox Titrations Redox titration consists of two different types of electrodes. 1. Indicator Electrode 2. Reference Electrode Indicator Electrode: Used to sense the presence or change in concentration of the oxidized and reduced forms of a redox couple Usually, the indicator electrode is an inert noble metal, such as Pt Pt half reactions at the electrode: Fe ³⁺ + e⁻ Fe²⁺ Eº = 0.767 V Reference Electrode: Standard hydrogen electrode and standard calomel electrode used as reference electrode. It has accurately maintained potential Redox potential (also known as oxidation / reduction potential 'ORP', pe, E ', or. ) Is a measure of the tendency of a chemical species to acquire electrons from or lose electrons to an electrode and thereby be reduced or oxidized respectively

12. Redox Indicators A redox indicator is an indicator compound that changes color at specific potential differences A redox indicator compound must have a reduced and oxidized form with different colors and the redox process must be reversible. In(oxidation) + ne ⁻ = In(red) Types of Indicators: Self Indicator : Potassium permanganate is a good example for the self indicator. Cerric sulphate and Iodine are other examples After equivalence point, the titrant will impart a definite pink color at the end of the titration. External Indicator : Based on some visible reactions of the titrated substances with suitable reagent. End point is marked by failure to elicit reaction

13. Eg: potassium ferricyanide Titration of ferrous ions with potassium dichromate. Drops removed during titration on to a tile gives Prussian blue colour because ferrous ions still present. At the end point ferric ions are present and does not give colour Internal or redox Indicators: These Indicators have different colours in oxidised or reduced form Most of these are dyes. Eg: Diphenylamine, Diphenyl Benzidine Potentiometric method: This method is useful when suitable indicators are not available and also when visual indicators fail or have limited accuracy.

14. Cell Representation : Cu(s) cuso ₄ (0.100M) Zncl₂ (0.200M) Zn Copper electrode immersed in 0.100 M cuso₄ ( First half cell electrode) Zinc electrode immersed in 0.200 M zncl₂ (second half cell cathode) If Eº is positive it is spontaneous reaction If Eº is negative it is non spontaneous reaction which has to be reversed for spontaneous reaction

15. Measurement of electrode potential Nernst equation More positive half cell reaction is by oxidizing agent (anode) Less positive half cell reaction is by reducing agent ( cathode) Nernst Equation is the relationship represented between the concentration and electrode potential for the half cell reaction. E = Reduction potential E º = Standard potential

16. Cerimetric Titration It is a redox titration in which an iron color change indicates the end point.

17. The potential difference is caused by the ability of electrons to flow from one half cell to the other.

18. Iodimetry and Iodometry Iodimetry: Principle: Standard Iodine solution is used as standard. Iodine is a weak oxidant and it can be reduced by reductants to Iodide ions I ₂ ↔ 2I⁻ Strong reducing agents examples Stannous chloride, sodium thiosulphate Sn²⁺ +I₂ Sn⁴⁺ + 2I⁻ (stannous) 2S₂O₃ +I₂ S₄O₆²⁻ + 2I⁻ (sodium thiosulphate) Weak reducing agents Arsenic As³⁺ + I₂ As⁵⁺ +2I⁻ This method is used to quantify oxidising agnets

19. Steps involved: 1. Take a standard solution of Iodine in the Iodine flask. 2. Add 1ml of Indicator solution Eg: Starch or sodium starch glycolate I ₂ + Indicator Blue color 3. Titrate the above solution using analyte solution in burette Eg : sodium thio sulphate 4. At the equivalence point all the I₂ will react with the sodium thio sulphate The solution present in Iodine flask is colorless In + 2I⁻ No reaction (colorless) As the indicator does not react with the Iodide ions there is no reaction and it will be colourless.

20. Iodometry : principle: A redox titration where the appearance or disappearance of elementary iodine indicates the end point. Liberated Iodine from Iodide is used for titration and the method is considered as Indirect titration. KI ↔ K⁺ + I⁻ 2I⁻↔ I₂ ↑ + 2e⁻ When we have solutions of strong oxidant CuSO 4, KMnO₄ add excess KI solution in acidic medium so that Iodide ions are oxidized to Iodine ions. Reaction with Iodide ions with analyte as follows: (cupric ions) (Iodide) (copper Iodide) (Iodine) Step 1: 2 Cu ²⁺ + I⁻ ↔ 2CuI + I₂ ↑ (Iodine is liberated from Iodide) Step 2: I₂ + 2S₂O₃ ↔ S₄O₆²⁻ + 2I⁻ (Titrated with sodium thio sulphate) 2 Cu ²⁺ ≡ I₂ ≡ 2S₂O₃

21. Steps involved in Iodometry: 1. Take the analyte solution in the Iodine flask - Cu ²⁺ solution 2. Add excess of KI solution so reaction between Cu ²⁺ and KI takes place and it will liberate I₂ 3. Add indicator solution in to the Iodine flask which gives blue colour Starch or Sodium starch glycolate 4. Titrate the above solution by using standard sodium thio sulphate till the appearance of colourless solution. I₂ + 2S₂O₃ ↔ S₄O₆²⁻ + 2I⁻ I⁻ + Indicator No reaction Applications: Iodometry in its many variations is extremely useful in volumetric analysis. Examples include the determination of copper(II), chlorate, hydrogen peroxide, and dissolved oxygen

22. Bromatometry Principle: Potassium bromate is a strong oxidizing agent in acidic medium. Bromatometry is a titration process in which the bromination of a chemical indicator is observed. Reaction takes place generally in presence of acidic medium 1M Hcl. The liberated Bro ₃⁻(bromate ions) reacts with analyte directly which is a direct titration. KBro ₃ ↔ K⁺ + Bro ₃⁻ Arsenite 3AsO ₃³⁻ ( Analyte example) + Bro ₃⁻ 3AsO ₄³⁻ + Br⁻ This liberated Br⁻ (bromine ion) further reacts with Bro ₃⁻ in presence of acidic medium H⁺ ions. 5 Br⁻ + Bro ₃⁻ + 6H⁺ 3 Br₂ + 3 H 2 O Br₂ reacts with indicator which is methyl orange, methyl red This Br₂ oxidizes indicator and solution is colorless.

23. Steps: 1. Take analyte in stoppered conical flask at low temp 2. Add 1M Hcl to analyte to make it acidic medium. 3. Two to three drops of Indicator solution- color becomes red because of acidic medium 4. Titrate the analyte with standard in the burette 5. End point is red color to colorless Applications: Bromination of Indicators can be analysed It is used to determine arsenic, antimony ,iodide compunds

24. Dichromatometry Potassium Dichromate is used as standard K 2 Cr 2 O 7 It is an Oxidizing agent in presence of acidic medium and used as primary standard K 2 Cr 2 O ₇ is used only in acidic medium Cr₂O₇²⁻(dichromate) is rapidly reduced to Cr³⁺(chromium) which is green in colour. K ₂ Cr ₂ O ₇ 2K⁺ + Cr₂O₇²⁻ Cr₂O₇²⁻ +14 H⁺ 6e⁻ 2 Cr³⁺ +7H₂O Iron II salt is used as analyte 6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ 2 Cr³⁺ + 6Fe³⁺ + 7H₂O Cr³⁺ is green in color after reduction of Cr₂O₇²⁻ ions By using simple indicator method end point can’t be determined so we need to use external indicator method. Eg: potassium ferricyanide

25. Steps involved in dichromatometry: 1.Take sample solution in conical flask Fe ²⁺ 2. Add sulphuric acid for acidic medium Titrate with the potassium dichromate in burette. Take one drop of solution from the conical flask near the end point and put it in the external indicator Before the ed point the ferrous ions reacts with potassium ferricyanide and converts to ferric so Prussian blue colour will appear At the end point this reaction will not occur so the colour will not change. Applications: Used to determine Iron II salt.