George Mason University General Chemistry 212 Chapter 21 Electrochemistry Acknowledgements

1. George Mason University General Chemistry 212 Chapter 21 Electrochemistry Acknowledgements Course Text: Chemistry: the Molecular Nature of Matter and Change, 6thed, 2011, Martin S. Silberberg, McGraw-Hill The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material.Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

2. Electrochemistry • Redox Reactions and Electrochemical Cells • Review of Oxidation Reduction Concepts • Half-Reaction Method for Balancing Redox Reactions • Electrochemical Cells • Voltaic Cell: Using Spontaneous Reactions to Generate Electrical Energy • Construction and Operation • Cell Notation • Why Does the Cell Work

3. Electrochemistry • Cell Potential: Output of a Voltaic Cell • Standard Cell Potentials • Strengths of Oxidizing and Reducing Agents • Free Energy and Electrical Work • Standard Cell Potential • Effect of Concentration of Ecell • Changes in Ecell During Cell Operation • Concentration Cells • Electrochemical Processes in Batteries • Primary (Nonrechargeable Batteries) • Fuel Cells

4. Electrochemistry • Corrosion: A case of Environmental Electrochemistry • Corrosion of Iron • Protecting Against Corrosion • Electrolytic Cells: Using electrical energy to drive Nonspontaneous Reactions • Construction and Operation • Predicting Electrolysis Products • Stoichiometry of Electrolytes

5. Electrochemistry • Electrochemistry • The study of the relationship between chemical change (reactions) and the flow of electrons (electrical work) • Electrochemical Systems • Electrolytic – Work done by absorbing free energy from a source (passage of an electrical current through a solution) to drive a nonspontaneous reaction • Voltaic/Galvanic – Release of free energy from a spontaneous reaction to produce electricity (Batteries)

6. Electrochemistry • Oxidation-Reduction Concepts Review • Oxidation – Loss of Electrons • Reduction – Gain of Electrons • Oxidizing Agent – Species that causes another species to be oxidized (lose electrons) Oxidizing agent is reduced (gains e-) • Reducing Agent – Species that cause another species to be reduced (gain electrons) Reducing agent is oxidized (loses e-) • Oxidation (e- loss) always accompaniesReduction (e- gain) • Total number of electrons gained by the atoms/ions of the oxidizing agent always equals the total number of electrons lost by the reducing agent

7. Electrochemistry

8. Electrochemistry • Oxidation Number • A number equal to the magnitude of the charge an atom would have if its shared electrons were held completely by the atom that attracts them more strongly • The oxidation number in a binary ionic compoundequals the ionic charge • The oxidation number for each element in a covalent compound (or polyatomic ion) are assigned according to the relative attraction of an atom for electrons • See next slide for a summary of the rules for assigning oxidation numbers

9. Electrochemistry

10. Electrochemistry • Balancing Redox Reactions • Oxidation Number Method • Half-Reaction Method • The balancing process must insure that: The number of electrons lost by the reducing agent equals the number of electrons gained by the oxidizing agent

11. Electrochemistry • Oxidation Number Method • Assign oxidation numbers to all elements in the reaction • From changes in oxidation number of given elements, identify oxidized and reduced species • For each element that undergoes a change of oxidation number, compute the number of electrons lost in the oxidation and gained in the reduction from the oxidation number change (Draw tie-lines between these atoms) • Multiply one or both these number by appropriate factors to make the electrons lost equal to the electrons gained • Use factors as coefficients in reaction equation

12. Practice Problem Balance equation with Oxidation Number method: +5 +5 -2 0 +1 -2 +2 -2 +4 +1 -2 Loses 2e- Accounts for the two electrons needed to balance the 2 electrons from the Copper oxidation Gains 1e-

13. Electrochemistry • Half-Reaction Method • Applicable to Acid or Base solutions • Does not usually require Oxidation Numbers (ON) • Procedure • Divide the overall reaction into: • Oxidation Half-Reaction • Reduction Half-Reaction • Balance each half-reaction for atoms & charge • Multiply one or both reactions by some integer to make electrons gained equal to electrons lost • Recombine to given balanced redox equation

14. Electrochemistry • Redox Half-Reaction Method – Example • Divide steps into Half-Reactions

15. Electrochemistry • Balance Atoms & Charges for Cr2O72- / Cr3+ • Balance Atoms & Charges for I- / I2 Add 7 Water molecules to balance Oxygen Add 14 H+ ions on left tobalance 14 H on right Add 6 electrons (e-) on leftto balance reaction charges (6 electrons gained  this is the reduction reaction No need to add H2O or H+ Add 2 electrons (e-) on rightto balance reaction charges (2 electrons lost  this is the oxidation reaction

16. Electrochemistry • Redox Half-Reaction Method – Example (con’t) • Multiply each half-reaction, if necessary, by an integer to balance electrons lost/gained • 2 e- lost in oxidation reaction and 6 e- gained in reduction •  Multiply oxidation half-reaction by 3 • Add 2 half-reactions together

17. Electrochemistry • Half-Reaction Method in a “Basic” solution Sodium Permanganate & Sodium Oxalate NaMnO4 Na2C2O4 Half-Reactions Multiply each reaction by appropriate integer

18. Electrochemistry • Sodium Permanganate & Sodium Oxalate (con’t) • Add reactions • Add OH- to neutralize H+ , balance H2O, and form “basic” solution

19. Electrochemistry • Electrochemical Cells • Voltaic (Galvanic) Cells • Use spontaneous reaction (G < 0) to generate electrical energy • Difference in Chemical Potential energy between higherenergy reactants and lower energy products is converted to electrical energy to power electrical devices • Thermodynamically - Thesystem does work on the surroundings

20. Electrochemistry • Electrochemical Cells • Electrolytic Cells • Uses electrical energy to drive nonspontaneous reaction (G > 0) • Electrical energy from an external power supply converts lower energy reactants to higherenergy products • Thermodynamically – The surroundings do work on the system • Examples – Electroplating and recovering metals from ores

21. Electrochemistry

22. Electrochemistry • Electrochemical Cells • Cell notation is used to describe the structure of a voltaic (galvanic) cell • For the Zn/Cu cell, the cell notation is: Zn(s) Zn2+(aq) Cu2+(aq) Cu(s) = phase boundary (solid Zn vs. Aqueous Zn2+) = salt bridge • Anode reaction (oxidation) is left of the salt bridge • Cathode reaction (reduction) is right of the salt bridge • Half-cell components usually appear in the same order as in the half-reactions (Zn(s) + 2e- Zn2+). • Zinc solid loses 2 e- (oxidized) to produce zinc(II) at the negative ANODE • Copper(II) gains 2e- (reduced) to form copper metal at positiveCATHODE

23. Electrochemistry • Voltaic (Galvanic) Cells • Zinc metal (Zn) in solution of Cu++ ions • Construction of a Voltaic Cell • The oxidizing agent (Zn) and reducing agent (Cu2+) in the same beaker will not generate electrical energy • Separate the half-reactions by a barrier and connect them via an external circuit (wire) • Set up salt bridge between chambers to maintain neutral charge in electrolyte solutions

24. Electrochemistry • Oxidation Half-Cell • Anode Compartment – Oxidation of Zinc (An Ox) • Zinc metal in solution of Zn2+ electrolyte (ZnSO4) • Zn is reactant in oxidation half-reaction • Conducts released electrons (e-) out of its half-cell • Reduction Half-Cell • Cathode Compartment – Reduction of Copper (Red Cat) • Copper bar in solution of Cu2+ electrolyte (CuSO4) • Copper metal is product in reduction half-cell reaction • Conducts electrons into its half-cell

25. Electrochemistry Zinc-Copper Voltaic Cell

26. Electrochemistry • Relative Charges on the Anode/Cathode electrodes • Electrode charges are determined by the source of the electrons and the direction of electron flow • Zinc atoms are oxidized (lose 2 e-) to form Zn2+ at the anode • Anode – negative charge (e-rich) • Released electrons flow to right toward cathode to be accepted by Cu2+ to form Cu(s) • Cathode – positive charge (e-deficient)

27. Electrochemistry • Purpose of Salt Bridge • Electrons from oxidation of Zn leave neutral ZnSO4 solution producing net positive charge • Incoming electrons to CuSO4 solution would produce net negative charge in solution as copper ions are reduced to copper metal • Resulting charge imbalance would stop reaction • Salt bridge provides “liquid wire” allowing ions to flow through both compartments completing circuit • Salt bridge constructed of an inverted “U-tube” containing a solution of non-reactingNa+ & SO42- ions in a gel

28. Electrochemistry • Active vs Inactive Electrodes • Active Electrodes • Electrodes in Zn/Cu2+ cell are active • Zinc & Copper bars are components of the cell reactions • Mass of Zn bar decreases as Zn2+ ions in cell solution increase • Mass of Copper bar increases as Cu2+ ions accept electron to form more copper metal

29. Electrochemistry • Active vs Inactive Electrodes • Inactive Electrodes • In many Redox reactions, one or the other reactant/product is not capable of serving as an electrode • Inactive electrodes - Graphite or Platinum • Can conduct electrons into and out of half-cells • Cannot take part in the half-reactions

30. Electrochemistry Voltaic Cell with Inactive Graphite Electrodes

31. Practice Problem A mercury battery, used for hearing aids and electric watches, delivers a constant voltage (1.35 V) for long periods. The half reactions are given below. Which half reaction occurs at the Anode and which occurs at the Cathode? What is the overall cell reaction? HgO(s) + H2O(l) + 2e- Hg(l) + 2 OH-(aq) Zn(s) + 2 OH-(aq)  Zn(OH)2(s) + 2e- Ans: Reduction occurs at Cathode (Red Cat) Hg2+ gains 2 e- (reduced) to form Hg Oxidation occurs at the Anode (An Ox) Zn loses 2 e- (oxidized) to form Zn2+

32. Practice Problem Write the cell notation for a voltaic cell with the following cell reaction Ans: Anode (oxidation) is represented on left side of Cell notation Cathode (reduction) is represented on right side of cell notation

33. Practice Problem Write the cell reaction for the following voltaic cell Pt|H2(g) | H+(aq) ║ Br2(l) | Br-(aq)|Pt Note: Platinum (Pt) serves as a reaction site at the anode, but does not participate in the reaction Ans:

34. Electrochemistry • Cell Potential • The movement of electronsis analogous to the pumping of water from one point to another • Water moves from a point of high pressure to a point of lower pressure. Thus, a pressure differenceis required • The work expended in moving the water through a pipe depends on the volume of water and the pressure difference

35. Electrochemistry • Cell Potential • Movement of Electrons • An electric charge moves from a point of high electrical potential (high electrical pressure) to one of lower electrical potential • The work expended in moving the electrical charge through a conductor depends on the potential difference and the amount of charge

36. Electrochemistry • Cell Potential • Purpose of a voltaic cell is to convert the free energy of a spontaneous reaction into the kinetic energy of electrons moving through an external circuit (electrical energy) • Electrical energy is proportional to the difference in the electrical potential between the two cell electrodes Cell Potential

37. Electrochemistry • Cell Potential • Positive Cell Potential – Electrons flow spontaneously from the negative electrode (Anode) to the positive electrode (Cathode) • Negative cell potential is associated with a “nonspontaneous” cell reaction • Cell potential for a cell reaction at equilibrium would be “0” • As with Entropy, there is a clear relationship between Ecell , K, and G

38. Electrochemistry • Units of Cell Potential • The SI (metric) unit of electrical charge is the: Coulomb (C) • The SI (metric) unit of current is the: Ampere (A) • The SI (metric) unit of electrical potential is the: “Volt (V)” • By definition, the energy released by a potential difference of one volt moving between the anode and cathode of a voltaic cell releases 1 joule of work per coulomb of charge

39. Electrochemistry • The charge (F) that flows through a cell equals the number of moles of electrons (n) transferred times the charge of 1 mol of electrons

40. Electrochemistry • Standard Cell Potential • Eocell – The potential measured at a specific temperature (298 K) with no current flowing and all concentrations in their “Standard States” • 1 atm for gases • 1 M for solutions • Pure solids for electrodes

41. Electrochemistry • Standard Electrode Half-Cell Potentials • Eohalf-cell – Potential associated with a given half-cell reaction (electrode compartment) when all components are in “Standard States” • Standard Electrode Potential for a half-cell reaction, whether anode (oxidation) or cathode (reduction) is written and presented in Appendix D as a “reduction” • Ex. would be written in the table as:

42. Electrochemistry • Standard Electrode Half-Cell Potentials • Electrons flow spontaneously from Anode (negative) to Cathode (positive) • Cathode must have a more “Positive” Eohalf-cell than the Anode • For a “positive” Eocell ; i.e., a spontaneous reaction • The standard cell potential is the difference between the standard electrode potential of the “Cathode” (reduction) half-cell and the standard electrode potential of the “Anode” (oxidation) half-cell • Standard half-cell potentials are “intensive” properties, thus their values do NOT have to be adjusted for stoichiometry (# of moles)

43. Practice Problem Write out the overall equation for the cell reaction and determine the standard cell potential for the following galvanic cell [Eo (Ag+/Ag) = 0.80 V; Eo (Ni2+/Ni) = - 0.26 V]

44. Electrochemistry • The Standard Hydrogen Electrode • Half-cell potentials are not absolute quantities • The values found in tables are determined relative to a “Standard” • The Standard Electrode potential is defined as zero (Eoreference) = 0.00 • The “standard reference half-cell” is a standard “Hydrogen” electrode • Specially prepared Platinum electrode immersed in a1 M aqueous solution of a strong acid through which H2 gas at 1 atm is bubbled

45. Electrochemistry • Reference Half-Cell and Unknown Half-Cell • The “Standard” electrode can act as either the “Anode” or the “Cathode” • Oxidation of H2 (lose e-) at anode half-cell and reduction of unknown at cathode half-cell • Reduction of H+ (gain e-) at cathode half-cell and oxidation of unknown at anode half-cell

46. Practice Problem Determine the standard electrode potential, Eozinc, using a voltaic cell consisting of the Zn/Zn2+ half-reaction and the H+/H2 half-reaction. Eocell = + 0.76 V Ans: Zinc is being oxidized (loses 2e-) producing electrons at the negative anode; H+ gains e- at positive cathode

47. Electrochemistry • Relative Strength of Oxidizing and Reducing Agents • The more positive the Eo value, the more readily the reaction occurs • Strength of Oxidizing Agents – Cu2+ > H+ > Zn2+ • Strength of Reducing Agents – Zn > H2 > Cu • Oxidizing agents decrease in strength as the value of Eodecreases, while the strength of the reducing agents increases as the value of Eodecreases • Cu2+ is the stronger Oxidizing agent • Zn metal (not the ion) is the stronger Reducing agent

48. Electrochemistry • Table of Standard Electrode Potentials (The emf Series) • All Values are relative to the “standard hydrogen (reference) electrode • All reactions are written as “reductions”

49. Electrochemistry • EMF Series • All Values are relative to the “standard hydrogen (reference) electrode • All reactions are written as “reductions” • F2 is the strongest oxidizing agent (high, positive Eo) • Fluorine is very electronegative with a high ionization potential (easily reduced by gaining electrons) • By gaining e- it forms the weakest reducing agent, F- , which is very reluctant to lose electrons) • Li metal is strongest reducing agent (low, more negative Eo) • Lithium has a Low ionization energy and is easily oxidized by losing electrons • By losing e- , Lithium forms the weakest oxidizing agent, Li+

50. Electrochemistry • Similarities – Acid/Base vs Redox • Acid Strength vs Base Strength using Ka & Kb values • Redox (Oxidizing agent vs Reducing agent) using Standard Electrode Potential (Eo) values • Appendix D (Table of Standard Electrode Potentials) • The stronger oxidizing agent (species on left side of table) has a half-reaction with a larger more positive (less negative) Eo than a species lower in the list • The stronger reducing agent (species on the right side of table) has a half-reaction with a smaller (less positive) Eo value than a species higher in the list