2011-2012 chem 27.1 WEJ2

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chem 27.1 experiment 10 potentiometric redox titration

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2011-2012 chem 27.1 WEJ2

  1. 1. Potentiometricredox titration<br />Czarina Charmaine S. Diwa<br />Ralph Marco S. Flores<br />
  2. 2. Objectives:<br />To construct a potentiometricredox titration curve<br />Determine percent weight of iron in the sample<br />
  3. 3. Potentiometricredox titration<br />Introduction: Terms defined<br />
  4. 4. Potentiometry<br />A method of analysis based on measuring the potential of the electrochemical cells with a production minimal amount of current.<br />One of the most used measurements because of its convenience.<br />Aside from its applications on determining analyte concentrations, it is also used to determine thermodynamic equilibrium constants.<br />
  5. 5. Redox (as in Redox reaction!)<br />Reactions that require electrons to be transferred from one reactant to the other<br />It is performed by bringing an oxidizing and reducing agents into direct contact in a container <br />Or it is carried out in an electrochemical cell where the components are not in direct contact with each other<br />
  6. 6. Titration<br />A process wherein a standard reagent is added to a solution of an analyte until the reaction between the analyte and reagent is judged complete<br />Completeness of the titration is based on the equivalence point, wherein the reagent is approximately equal to the analyte<br />
  7. 7. Equipment and procedure<br />
  8. 8. The equipments and important parts<br />Multimeter/voltmeter – a device that measures the cell potential<br />Reference electrode – a half-cell having a known electrode potential that remains constant and is independent of the analyte solution (Mohr’s salt solution)<br />Indicator electrode – an electrode that varies as a result of the change in analyte concentration<br />*together with the salt bridge and analyte solution, they constitute the electrochemical cell*<br />
  9. 9. Methodology: preparation of the iron (III) solution<br />Plug the end of a Pasteur pipet with cotton and soak it in a saturated solution of KCl or KNO3<br />Fill the pipet with 0.1 M CuSO4 and insert a copper wire.<br />Dissolve pre-measured KMnO4 in 100mL dH2O and add enough dH2O to make 300 mL. Store in a colored container.<br />Weigh 0.3 to 0.4 g Fe(NH4)2(SO4)2·6H2O and dissolve with 20 mL dH2O in a 100mL beaker. Add 5mL 2M H2SO4.<br />
  10. 10. Methodology: Standardization of KMnO4<br />Dip the a carbon rod and the Cu/Cu2+ electrode in the prepared solution. Connect both electrodes to the multimeter.<br />Add 2.0 mL of the KMnO4 solution into the Fe(NH4)2(SO4)2·6H2O + H2SO4 solution and measure the reduction potential (reading on the multitester). Repeat until 20 mL of KMnO4 solution has been added. Do not forget to stir after each addition.<br />Add 1.0 mL KMnO4 solution until 30 mL KMnO4 solution has been added. Stir after addition.<br />Plot the volume used versus the observed potential with the highest peak (in the first derivative plot) as the equivalence point.<br />
  11. 11. Methodology: analysis of the unknown<br />Pipet10mL of the unknown sample in a 100mL beaker. (pre-mixed with H2SO4)<br />Add 2.0 mL of the KMnO4 solution into the unknown solution and measure the reduction potential (reading on the multitester). Repeat until 20 mL of KMnO4 solution has been added. Stir after each addition.<br />Add 1.0 mL KMnO4 solution until 30 mL KMnO4 solution has been added. Stir after each addition.<br />Plot the volume used versus the observed potential with the highest peak (in the first derivative plot) as the equivalence point. Calculate weight of Fe in the sample.<br />
  12. 12. The electrochemical cell<br />Reference electrode | salt bridge| analyte solution| indicator electrode<br />
  13. 13. Results and discussion<br />
  14. 14. Reactions involved:<br />Standardization of potassium permanganate:<br />Fe3+ + e- Fe2+<br />MnO4- + 5e- + 8H+  Mn2+ + 4H2O<br />MnO4-+ 5Fe2+ + 8H+  Mn2++ 4H2O + 5Fe3+<br />Analysis of the unkown:<br /> MnO4-+ 5Fe2+ + 8H+  Mn2++ 4H2O + 5Fe3<br />Fe2+  Fe3+ + e-<br />5Fe2+  5Fe3+ + 5e-<br />
  15. 15. Results: dE/dV titration curve<br />
  16. 16. Results: 1st derivative plot<br />
  17. 17. Results: 2nd derivative plot<br />
  18. 18. Sample computation<br />equivalence point: 20.5 mL<br />weight of Fe(NH4)2(SO4)2.6H2O: 0.4g <br />computation: <br />M KMnO4 = (mass Fe)(molar mass Fe)(ratio)/(volume titrant)<br />M KMnO4 = {[(0.4g Fe)/(392.14g Fe/mol Fe)][mol KMnO45 mol Fe]}/ .0205 L <br />=9.95x10-3M KMnO4<br />
  19. 19. Analysis of the unkown<br />
  20. 20. Results: dE/dV titration curve<br />
  21. 21. Results: 1st derivative plot<br />
  22. 22. Results: 2nd derivative plot<br />
  23. 23. Sample computation<br />equivalence point: 9.5mL<br />weight of Fe(NH4)2(SO4)2.6H2O: unkown<br />computation: <br />Xg Fe ={[9.5mL KMnO4][9.95-3M KMnO4][5 mol Femol KMnO4][(55.85g Fe/mol Fe)]}<br />0.0.02640 g<br />=0.26% w/v <br />
  24. 24. Some notes on:<br />Indicator electrode: Cu2+/Cu<br />Analyte: mohr’s salt<br />Titrant: potassium permanganate<br />
  25. 25. Sources of systematic errors<br />Instrument error<br />Multimeter<br />Method error<br />Photodecomposition of permanganate<br />Acid error<br />Personal error<br />Detection of end point<br />
  26. 26. Conclusion: Potentiometricredox titration<br />An analytical process wherein a standard reagent is added to the analyte, one of which behaves as an electron donor and the other electron acceptor, until all analyte has been consumed, as judged by the change in the electrode potential.<br />
  27. 27. End<br />

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