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Announcements

Announcements. Course Evaluations Final is Wednesday Afternoon on May 9 th Homework 14-15, 14-26, 15-6, 16-3 , 16-6,. Course Evaluations. Request from Thelmo

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Announcements

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  1. Announcements • Course Evaluations • Final is Wednesday Afternoon on May 9th • Homework • 14-15, 14-26, 15-6, 16-3, 16-6,

  2. Course Evaluations • Request from Thelmo • “Teaching is a complex endeavor, capable of an almost infinite variety of successful expressions, and thus, success as a teacher cannot be judged by any one criterion or through one single mechanism.” • Consider the many facets of the learning environment over the course of whole semester • How will your feedback be used? • Read by me to evaluate which aspects of the course most contribute to student learning • formative • Read by PSC as part of faculty’s permanent file to evaluate faculty for promotion and tenure • summative

  3. Concentration dependency of E

  4. Concentration Dependency of E • Eo values are based on standard conditions. • The E value will vary if any of the concentration vary from standard conditions • Theoretically • Predicted by the Nernst Equation

  5. The Nernst Equation • The Nernst Equation For aA + ne- bB

  6. Example Equilibrium constant and Eo • Find the equilibrium constant for the reaction Cu (s) + 2Fe3+2Fe2+ + Cu2+

  7. Example • Find the voltage of the cell • Half reaction • Ag (s) into a solution of 0.50 M AgNO3 (aq) • The other half-reaction • Cd (s) is immersed into a 0.010 M Cd(NO3)2 (aq) • Metals are connected by wires • Solution connected with salt bridge

  8. Find the voltage of the cell 0.010 M Cd(NO3)2 0.50 M AgNO3

  9. Potentiometric Methods

  10. Potentiometric Methods • Basis of Method • The difference b/w the E (not Eo) values for two halves of a cell give rise to Eoverall., • If one half reaction is known and held constant, we can measure the concentration of species on the other side!!!

  11. Indicating electrode – The part of the cell that contains the solutions we are interested in measuring Reference

  12. Electrodes • The previous cell would be difficult to use for many systems. • We would like something that can be placed in the solution we wish to measure • The electrodes in the following slides have that goal in mind but • THEY STILL represent a complete electrochemical cell when used

  13. Reference Electrodes • Ag/AgCl

  14. Reference Electrodes • Calomel Electrode (SCE) • Very Common • Hg|Hg2Cl2 (sat), KCl|| • Chloride is used to maintain constant ionic strength

  15. Reference Electrodes • The SCE (Saturated Calomel Electrode) • Different KCl concentrations can (and are used) • 0.1 M – least temperature sensitive • Saturated – easier to make and maintain. Eref = 0.244 V @STP • Reaction Eo (V) • Hg2Cl2 + 2e- ->2Hg(l) + 2Cl- 0.244 V

  16. Hg2Cl2 + 2e- ->2Hg(l) + 2Cl- 0.241 AgCl (s) + e- ->Ag(s) + Cl- 0.197

  17. Sensing electrodes • Several types • Simple Metal • Solid State Electrodes • Glass Membrane • Etc. Let’s look at some examples.

  18. Sensing electrodes • Several types • Simple Metal • Solid State Electrodes • Glass Membrane • Etc. Let’s look at some examples.

  19. Simple metal electrodes • A bare metal in contact with a solution. • General Form: • Mn+ + ne- -> M(s)

  20. Simple Metal Electrodes • A bare metal in contact with a solution of its cation. • Ag+ + 1e- -> Ag(s) 0.799 V • General form Mn+ + ne- -> M(s)

  21. Example • A potential of 0.5000V was measured vs. SCE. What is the concentration of Ag+? Hg2Cl2 + 2e- 2Hg(l) + 2 Cl- Eo = 0.241V Ag+ + e-  Ag(s) Eo = 0.799V Using a simple metal electrode (Ag) and a reference electrode (Calomel), the voltage determined from this potentiometric set-up provides us with a direct measure of concentration no calibration plot required!!

  22. Simple Metal Electrodes • Example • Silver sensing electrode

  23. Example (cont’d)

  24. Simple Metal Electrodes • For some metals, a good electrode can’t be made or no metals are involved – just ions or gas! An inert indicating electrode is used (graphite or Pt). • This type only measures the ratios of ions. • No quantitation but suitable for titrations!

  25. Simple Metal Electrodes • For some metals, a good electrode can’t be made or no metals are involved – just ions or gas! An inert indicating electrode is used (graphite or Pt). • This type only measures the ratios of ions. • No quantitation but suitable for titrations! Calomel (Hg2Cl2)

  26. Constant = -0.241 V Simple Metal Electrodes • Eoverall = Eox + Ered Reduction at Platinum Electrode: Reaction Eo Fe3+ + 1e- -> Fe2+ 0.771 V Calomel (Hg2Cl2)

  27. Simple Metal Electrodes Ce4+ • For some metals, a good electrode can’t be made or no metals are involved – just ions or gas! • This type only measures the ratios of ions. • No direct quantitation but suitable for titrations! Calomel (Hg2Cl2)

  28. REDOX titrations • “Your titrant is commonly an oxidizing agent although reducing titrants can be used.” Consider: Ce4+ + Fe2+ Ce3+ + Fe3+ General form: Aox + Bred Ared + Box

  29. Determination of the Equivalence Point • The equivalence point is based on the concentration of the oxidized and reduced form of all species involved • Use Nernst Equation to find Eeq.

  30. Equivalence Point Nernst Equation for A Nernst Equation for B Since at equilibrium, [Ared] = [Box] and [Bred] = [Aox] we massage the two general equations to yield:

  31. Equivalence Point Note: This expression only works for simple REDOX TITRATIONS: Simple redox titrations: Only Aox, Box, Ared, Bred are involved in the reaction …

  32. Two examples • Determine Eeq for the following reactions: • Fe2+ + Ce4+ -> Fe3+ + Ce3+ • Sn2+ + 2Ce4+ -> Sn4+ + 2Ce3+

  33. Titration curves • What does a titration curve look like for an acid/base titration?

  34. REDOX Titrations OverTitration Ecell

  35. Just like Acid/Base Titrations There are four significant regions, • The Start • The Buffer Region • The equivalence Point • Overtitration Let’s Use our simple example: • Fe2+ + Ce4+ Fe3+ + Ce3+

  36. Our simple example Let’s Use our simple example: • Fe2+ + Ce4+ Fe3+ + Ce3+ Titrate 50 mL of 0.05 M Fe2+ with 0.10 M Ce4+

  37. 0% Titration Unlike acid/base titrations, we can’t find this point exactly. • While some Fe3+ must be present, we can only guess what the concentration is. • No Ce4+ or Ce3+ present, so we don’t have a complete reaction

  38. 0% Titration NO … some of the iron is oxidized by air to give some Fe3+ … how much ? We generally estimate that less Than one in 1000 are oxidized.

  39. “Buffer Region” Fe2+ + Ce4+ Fe3+ + Ce3+ 10 ml of Ce4+ is added Goes to completion … Excess Fe2+ pushes equilibrium to the right. Thus Eis not dependent on Ce3+/Ce4+, but only on Iron.

  40. “Buffer Region”

  41. Closer look at the “buffer” region

  42. Equivalence Point • From Before • Eeq = 1.24 V What volume? 25 ml

  43. Excess Ce4+ (post titration) • Fe2+ + Ce4+ Fe3+ + Ce3+ • The predominate change is that Ce4+ is being added and diluted into a solution of Ce3+. • All Fe2+ has been converted to Fe3+ and no longer figures into the calculations • We just need to keep track of the amounts of Ce3+ and Ce4+ as well as the VOLUME of the system.

  44. Excess Ce4+ (post titration) • At 30.0 mL Ce4+ Vt = 30.0 mL+ 50.0 mL

  45. Excess Ce4+ (post titration) • Fe2+ + Ce4+ Fe3+ + Ce3+ Ce3+/Ce4+

  46. Excess Ce4+ (post titration)

  47. Redox Indicators General Specific

  48. General Redox Indicators • Varies as a function of Ecell • Rely on a color change with Indox and Indred being different colors. Indox + ne- Indred

  49. General Redox Indicators • In order to see a color change, you typically need approximately a 10% conversion from one form to another.

  50. General Redox Indicators • Examples • Consider 1,10 phenanthrolene-Fe + e- BLUE RED Eo = 1.06 V

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