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### redoxtitrations-210614081400.pdf

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 :  KMno₄ 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:  Fe³⁺ + 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, E0', 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) cuso₄ (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 Eº= 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  I₂ ↔ 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  I₂ + 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 CuSO4, 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 Cu²⁺ + I⁻ ↔ 2CuI + I₂ ↑ (Iodine is liberated from Iodide) Step 2: I₂ + 2S₂O₃ ↔ S₄O₆²⁻ + 2I⁻ (Titrated with sodium thio sulphate) 2 Cu²⁺ ≡ I₂ ≡ 2S₂O₃
21.  Steps involved in Iodometry:  1. Take the analyte solution in the Iodine flask - Cu²⁺ solution  2. Add excess of KI solution so reaction between Cu²⁺ 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 Bro₃⁻(bromate ions) reacts with analyte directly which is a direct titration.  KBro₃ ↔ K⁺ + Bro₃⁻  Arsenite  3AsO₃³⁻ ( Analyte example) + Bro₃⁻ 3AsO₄³⁻ + Br⁻  This liberated Br⁻ (bromine ion) further reacts with Bro₃⁻ in presence of acidic medium H⁺ ions.  5 Br⁻ + Bro₃⁻ + 6H⁺ 3 Br₂ + 3 H2O  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 K2Cr2O7  It is an Oxidizing agent in presence of acidic medium and used as primary standard  K2Cr2O₇ is used only in acidic medium Cr₂O₇²⁻(dichromate) is rapidly reduced to Cr³⁺(chromium) which is green in colour.  K ₂Cr₂ O₇ 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 Fe²⁺  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.
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