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Dr. Marc Madou

BIOMEMS. Class III. Electrochemistry Background (II) Winter 2011. Dr. Marc Madou. Oxidants and reductants Battery Reference Electrodes Standard Reduction Potentials Thermodynamic Significance of Potentials How do Cell Potentials Change if We are Not at Standard State? Nernst-Equation

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Dr. Marc Madou

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  1. BIOMEMS Class III. Electrochemistry Background (II) Winter 2011 Dr. Marc Madou

  2. Oxidants and reductants Battery Reference Electrodes Standard Reduction Potentials Thermodynamic Significance of Potentials How do Cell Potentials Change if We are Not at Standard State? Nernst-Equation Cyclic voltammetry Potentiometric sensors Amperometric sensors Contents

  3. Oxidants and Reductants • oxidant = oxidizing agent • reactant which oxidizes another reactant and which is itself reduced • reductant = reducing agent • reactant which reduces another reactant and which is itself oxidized

  4. Oxidants and Reductants • Identify the oxidant and reductant in each of the following reactions:a) Karl Fischer reaction – for quantitation of moisture:I2 + SO2 + H2O = 2HI + SO3b) Hall Heroult process – production of Al:2Al2O3 + 3C = 4Al + 3CO2c) the Thermite reaction – used to produce liquid iron for welding2Al + Fe2O3 = 2Fel + Al2O3

  5. Zn strip CuSO4 Oxidants and Reductants • Reactions occur pair wise: One cannot have oxidation without reduction • Charge must be conserved: Number of electrons lost in oxidation must equal number of electrons gained in reduction • Suppose we add a strip of Zinc metal to a solution of CuSO4 • Zn - 2e- = Zn2+ • Cu2+ + 2e- = Cu

  6. 1.1 V Zn strip salt bridge Cu Zn CuSO4 ZnSO4 CuSO4 Oxidants and Reductants • It is the relative tendencies of oxidants and reductants to gain/lose electrons that determines the extent of a redox reaction • Strong oxidant + strong reductant  completion • What if we could separate the oxidant from the reductant? • We would have set up a constant flow of electrons = current = electricity! 1836 The Daniell Cell

  7. Battery • Electrode • anode = electrode at which oxidation occurs • cathode = electrode at which reduction occurs • Salt bridge = completes the electrical circuit • allows ion movement but doesn’t allow solutions to mix • salt in glass tube with vycor frits at both ends • Since electrons flow from one electrode to the other in one direction, there is a potential difference between the electrodes • This difference is called • The electromotive force (EMF) • Cell voltage • Cell potential

  8. Battery Problem: True or False • In the Daniell cell, zinc metal is reduced to zinc(II) at the cathode and copper is oxidized to copper(II) at the anode • In the Daniell cell, zinc is the oxidant and copper is the reductant • Since all redox reactions occur pair wise, i.e., reduction and oxidation always occur at the same time we cannot measure the cell potential for just one half cell reactionand this means we must establish a RELATIVE scale for cell potentials

  9. Reference Electrodes • Electrodes with a potential independent of solution composition • Standard hydrogen electrode (SHE) • 1 M H+(aq)+ 2e- = H2(g) (1 atm) • We define E0 0 V for this electrode • where 0 stands for standard state: • 1 M all solutes • 1 atm all gases • 250C (298 K) H2(gas) HCl Pt black

  10. Reference Electrodes

  11. Reference Electrodes 2H+(1M) + 2e-  H2(g,1atm) Eoredn = 0.0V

  12. Reference Electrodes

  13. Reference Electrodes 0.244 V v. SHE

  14. Reference Electrodes

  15. Reference Electrodes

  16. Standard Reduction Potentials • Li+ + e- = Li -3.0 V • 2H2O + 2e- = H2 + 2OH- -0.83 V • Zn2+ + 2e- = Zn -0.76 V • 2H+ + 2e- = H2 0 V (SHE) • Cu2+ + 2e- = Cu 0.34 V • MnO4- +8H+ +5e- = Mn2+ 1.51 V

  17. Standard Reduction Potentials • Always write the redox ractions as shown :

  18. Standard Reduction Potentials • Half cell reactions are reversible, i.e., depending on the experimental conditions any half reaction can be either an anode or a cathode reaction • Changing the stoichiometry does NOT change the reduction potential (intensive property) • Oxidation potentials can be obtained from reduction potentials by changing the signEcell = Eanode + Ecathode

  19. Li+ + e- = Li -3.0 V • 2H2O + 2e- = H2 + 2OH- -0.83 V • Zn2+ + 2e- = Zn -0.76 V • 2H+ + 2e- = H2 0 V (SHE) • Cu2+ + 2e- = Cu 0.34 V • MnO4- +8H+ +5e- = Mn2+ 1.51 V Problem: Standard Reduction Potentials • Calculate the cell potential for the Daniell cell.

  20. Standard Reduction Potentials

  21. Standard Reduction Potentials Zn --> Zn2+ + 2e- oxidation Cu2+ + 2e- -->Cu reduction

  22. Standard Reduction Potentials • Anode reaction appears leftmost while cathode reaction appears rightmost • All redox forms of reagents present should be listed. Phase and concentration specified in brackets, e.g., ZnSO4(aq, 1 M) • A single vertical line (|) is used to indicate a change of phase (s to l to g) • A double vertical line (||) indicates a salt bridge • A comma should be used to separate 2 components in the same phase

  23. Thermodynamic Significance of Potentials • We usually operate electrochemical cells at constant P and T • Recall, • G = H - T S (change in Gibbs free energy) • H = E + (PV) • So, GT,P=welec = -QE = -(nF)E • since Q = n F • Recall, F is Faraday’s constant 96,485 C/mole

  24. Thermodynamic Significance of Potentials • The maximum electrical work done by an electrochemical cell equals the product of the charge flowing and the potential difference across which it flows. The work done on the cell is: • W = -E x Q, where E is the Electromotive Force of the Cell (EMF), and Q is the charge flowing: Q = n x NA x e • where n is the number of moles of electrons transferred per mole of reaction, NA is Avogadro's Number (6.02 x 1023), and e is the charge on an electron (-1.6 x 10-19 C). • Note: NA x e = F (one Faraday). Thus: W = -nFE and: W = DG = -nFE

  25. Thermodynamic Significance of Potentials • Recall sign of G provides information on spontaneity:G negative  spontaneous reactionG positive  non-spontaneous reaction • So, since G = - nFEE positive  spontaneous reactionE negative  non-spontaneous reaction

  26. Thermodynamic Significance of Potentials • Since half-cell potentials are measured relative to SHE, they reflect spontaneity of redox reactions relative to SHE • More positive potentials  more potent oxidants (oxidants want to be reduced) • More negative potentials  more potent reductants (reductants don’t want to be reduced; they spontaneously oxidize)

  27. Thermodynamic Significance of Potentials • Galvanic • Chemical energy  electrical energy • Spontaneous(so Ecell is positive)EXAMPLES: • Primary (non-rechargeable) • Le Clanche (dry cell) • Secondary (rechargeable) • Lead storage battery • Hydrogen-Oxygen Fuel Cell

  28. Thermodynamic Significance of Potentials • Electrolytic • Electrical energy  chemical energy • Non-spontaneous(Ecell is negative)EXAMPLE: • Lead storage battery when recharging • Electrolysis of water

  29. Thermodynamic Significance of Potentials

  30. Thermodynamic Significance of Potentials

  31. Thermodynamic Significance of Potentials

  32. Thermodynamic Significance of Potentials

  33. Thermodynamic Significance of Potentials

  34. Thermodynamic Significance of Potentials-Problems • Arrange the following in order of increasing oxidizing strength: • MnO4- in acidic media • Sn2+ • Co3+ • Co3+ + e- = Co2+ 1.82 V • MnO4- + 4H+ + 3e- = MnO2 + 2H2O 1.70 V • MnO4- + 8H+ + 5e- = Mn2+ + 4H2O 1.51 V • Sn2+ + 2e- = Sn -0.14 V • So, Co3+ > MnO4- > Sn2+

  35. Thermodynamic Significance of Potentials-Problems • A galvanic cell consists of a Mg electrode in a 1.0 M Mg(NO3)2 solution and a Ag electrode in a 1.0 M AgNO3 solution. Calculate the standard state cell potential and diagram the cell. • Consider the following cell:Ag(s)/AgNO3(aq, 1 M)//CuSO4(aq, 1 M)/Cu(s)a) what is the anode reaction?b) what is the cathode reaction?c) what is the net number of electrons involved?d) what is the net reaction?e) what is the cell potential at standard state?f) is the cell galvanic or electrolytic?

  36. Thermodynamic Significance of Potentials -Problems • Is the following redox reaction spontaneous?Mg2+ + 2Ag = Mg + 2Ag+ given:Ag+ + e- = Ag +0.80 VMg2+ + 2e- = Mg -2.37 V

  37. Thermodynamic Significance of Potentials • Using a table of standard reduction potentials, any species on the left of a given half reaction will react spontaneously with any species appearing on the right of any half reaction that appears below it when reduction potentials are listed from highest and most positive to lowest and most negative.

  38. Thermodynamic Significance of Potentials -Problems • What would the cell potential be for the following cell?Ag(s)/AgNO3(aq, 1 M)//CuSO4(aq, 0.5 M)/Cu(s) • This represents a set of non-standard state conditions so we need derive an equation relating the standard state to the non-standard state or the Nernst Equation • Standard state: • Temperature 250C (K = 273.15 + 0C) • Pressure 1 atm • Concentrations of all solutes 1 M • 0 (not) is used to indicate at standard state • Example: E0 = cell potential at standard state

  39. How do Cell Potentials Change if We are Not at Standard State? • For the reaction:aA + bB = cC + dD • G = DG0 + 2.303 RT log Qwhere Q is the reaction quotient: • Where c is the activity for product C

  40. How do Cell Potentials Change if We are Not at Standard State? • Since G = - nFE thenE = E0 - 2.303 (RT/nF) log Q • At standard state,E = E0 - (0.0591 /n) log QThis is called the Nernst equation • Apply the Nernst Equation to a pH sensor: pH=-log[H+] • What is the cell potential for the following electrochemical cell? What type of cell is it?Ni(s) | Ni2+ (aq, 0.1 M) || Co2+ (aq, 2.5 M) | Co(s)

  41. Nernst Equation

  42. Nernst Equation • The Nernst equation underlies the operating principle of potentiometric sensing electrodes and reference electrodes • Electrolysis vs. battery is determined by Eo sign

  43. Electrolytic cell (example): Au cathode (inert surface for e.g. Ni deposition) Graphite anode (not attacked by Cl2) Two electrode cells (anode, cathode, working and reference or counter electrode) e.g. for potentiometric measurements (voltage measurements) (A) Three electrode cells (working, reference and counter electrode) e.g. for amperometric measurements (current measurements)(B) Two-electrode and three-eletrode cells, potentiostat, galvanostat

  44. At equilibrium the exchange current density is given by: The reaction polarization is then given by: The measurable current density is then given by: For large enough overpotential: Cyclic voltammetry: activation control (Butler-Volmer) (Tafel law)

  45. From activation control to diffusion control: Concentration difference leads to another overpotential i.e. concentration polarization: Using Faraday’s law we may write also: At a certain potential C x=0=0 and then: Cyclic voltammetry: diffusion control • Since we get :

  46. Cyclic voltammetry and potentiometric and amperometric sensors • Scan the voltage at a given speed (e.g. from + 1 V vs SCE to -0.1 V vs SCE and back at 100 mV/s) and register the current • Potentiometric: the voltage between the sensing electrode and a reference electrode is registered • Amperometric: the current at a fixed voltage in the diffusion plateau is registered Ferricyanide

  47. Cyclic voltammetry (also polarography) and potentiometric and amperometric sensors

  48. Calculate the potential of a battery with a Zn bar in a 0.5 M Zn 2+ solution and Cu bar in a 2 M Cu 2+ solution. Show in a cyclic voltammogram the transition from kinetic control to diffusion control and why does it really happen ? Derive how the capacitive charging of a metal electrode depends on potential sweep rate. What do you expect will be the influence of miniaturization on a potentiometric sensor and on an amperometric sensor? Homework

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