Fundamentals of Electrochemistry

1. Fundamentals of Electrochemistry • Introduction • 1.)Electrical Measurements of Chemical Processes • Redox Reaction involves transfer of electrons from one species to another. • Chemicals are separated • Can monitor redox reaction when electrons flow through an electric current • Electric current is proportional to rate of reaction • Cell voltage is proportional to free-energy change • Batteriesproduce a direct current by converting chemical energy to electrical energy. • Common applications run the gamut from cars to ipods to laptops

2. Fundamentals of Electrochemistry • Basic Concepts • 1.)A Redox titration is an analytical technique based on the transfer of electrons between analyte and titrant • Reduction-oxidation reaction • A substance is reduced when it gains electrons from another substance • gain of e- net decrease in charge of species • Oxidizing agent (oxidant) • A substance is oxidized when it loses electrons to another substance • loss of e- net increase in charge of species • Reducing agent (reductant) (Reduction) (Oxidation) Oxidizing Agent Reducing Agent

3. Fundamentals of Electrochemistry • Basic Concepts • 2.)The first two reactions are known as “1/2 cell reactions” • Include electrons in their equation 3.) The net reaction is known as the total cell reaction • Nofreeelectrons in its equation 4.) In order for a redox reaction to occur, both reduction of one compound and oxidation of another must take place simultaneously • Total number of electrons is constant ½ cell reactions: Net Reaction:

4. Fundamentals of Electrochemistry • Basic Concepts 5.) Electric Charge (q) • Measured in coulombs (C) • Charge of a single electron is 1.602x10-19C • Faraday constant (F) – 9.649x104C is the charge of a mole of electrons 6.) Electric current • Quantity of charge flowing each second through a circuit • Ampere: unit of current (C/sec) Relation between charge and moles: Coulombs moles

5. Fundamentals of Electrochemistry • Galvanic Cells 1.) Galvanic or Voltaic cell • Spontaneous chemical reaction to generate electricity • One reagent oxidized the other reduced • two reagents cannot be in contact • Electrons flow from reducing agent to oxidizing agent • Flow through external circuit to go from one reagent to the other Reduction: Oxidation: Net Reaction: AgCl(s) is reduced to Ag(s) Ag deposited on electrode and Cl- goes into solution Cd(s) is oxidized to Cd2+ Cd2+ goes into solution Electrons travel from Cd electrode to Ag electrode

6. Fundamentals of Electrochemistry • Galvanic Cells 2.) Cell Potentials • Reaction is spontaneous if it does not require external energy Potential of overall cell = measure of the tendency of a reaction to proceed to equilibrium ˆLarger the potential, the further the reaction is from equilibrium and the greater the driving force that exists

7. Fundamentals of Electrochemistry • Galvanic Cells 3.) Electrodes Cathode: electrode where reduction takes place Anode: electrode where oxidation takes place

8. Fundamentals of Electrochemistry • Galvanic Cells 4.) Salt Bridge • Connects & separates two half-cell reactions • Prevents charge build-up and allows counter-ion migration Salt Bridge • Contains electrolytes not involved in redox reaction. • K+ (and Cd2+) moves to cathode with e- through salt bridge (counter balances –charge build-up • NO3- moves to anode (counter balances +charge build-up) • Completes circuit Two half-cell reactions

9. Fundamentals of Electrochemistry • Galvanic Cells 5.) Short-Hand Notation • Representation of Cells: by convention start with anode on left Zn|ZnSO4(aZN2+ = 0.0100)||CuSO4(aCu2+ = 0.0100)|Cu Phase boundary Electrode/solution interface anode cathode 2 liquid junctions due to salt bridge Solution in contact with anode & its concentration Solution in contact with cathode & its concentration

10. Fundamentals of Electrochemistry • Standard Potentials 1.) Predict voltage observed when two half-cells are connected • Standard reduction potential (Eo) the measured potential of a half-cell reduction reaction relative to a standard oxidation reaction • Potential arbitrary set to 0 for standard electrode • Potential of cell = Potential of ½ reaction • Potentials measured at standard conditions • All concentrations (or activities) = 1M • 25oC, 1 atm pressure Ag+ + e-»Ag(s) Eo = +0.799V Standard Hydrogen Electrode (S.H.E) Pt(s)|H2(g)(aH2 = 1)|H+(aq)(aH+ = 1)|| Hydrogen gas is bubbled over a Pt electrode

11. Fundamentals of Electrochemistry • Standard Potentials 1.) Predict voltage observed when two half-cells are connected As Eo increases, the more favorable the reaction and the more easily the compound is reduced (better oxidizing agent). Reactions always written as reduction Appendix H contains a more extensive list

12. Fundamentals of Electrochemistry • Standard Potentials 2.) When combining two ½ cell reaction together to get a complete net reaction, the total cell potential (Ecell) is given by: Where:E+ = the reduction potential for the ½ cell reaction at the positive electrode E+ = electrode where reduction occurs (cathode) E- = the reduction potential for the ½ cell reaction at the negative electrode E- = electrode where oxidation occurs (anode) Place values on number line to determine the potential difference Electrons always flow towards more positive potential

13. Fundamentals of Electrochemistry • Standard Potentials 3.) Example: Calculate Eo for the following reaction:

14. Fundamentals of Electrochemistry • Nernst Equation 1.) Reduction Potential under Non-standard Conditions • E determined using Nernst Equation • Concentrations not-equal to 1M For the given reaction: aA + ne-»bB Eo The ½ cell reduction potential is given by: Where: E = actual ½ cell reduction potential Eo = standard ½ cell reduction potential n = number of electrons in reaction T = temperature (K) R = ideal gas law constant (8.314J/(K-mol) F = Faraday’s constant (9.649x104 C/mol) A = activity of A or B at 25oC

15. Fundamentals of Electrochemistry • Nernst Equation 2.) Example: • Calculate the cell voltage if the concentration of NaF and KCl were each 0.10 M in the following cell: Pb(s) | PbF2(s) | F- (aq) || Cl- (aq) | AgCl(s) | Ag(s)

16. Fundamentals of Electrochemistry • Eo and the Equilibrium Constant 1.) A Galvanic Cell Produces Electricity because the Cell Reaction is NOT at Equilibrium • Concentration in two cells change with current • Concentration will continue to change until Equilibrium is reached • E = 0V at equilibrium • Battery is “dead” Consider the following ½ cell reactions: aA + ne-»cC E+o dD + ne-»bB E-o Cell potential in terms of Nernst Equation is: Simplify:

17. Fundamentals of Electrochemistry • Eo and the Equilibrium Constant 1.) A Galvanic Cell Produces Electricity because the Cell Reaction is NOT at Equilibrium SinceEo=E+o- E-o: At equilibrium Ecell =0: Definition of equilibrium constant at 25oC at 25oC

18. Fundamentals of Electrochemistry • Eo and the Equilibrium Constant 2.) Example: • Calculate the equilibrium constant (K) for the following reaction:

19. Fundamentals of Electrochemistry • Cells as Chemical Probes 1.) Two Types of Equilibrium in Galvanic Cells • Equilibrium between the two half-cells • Equilibrium within each half-cell If a Galvanic Cell has a nonzero voltage then the net cell reaction is not at equilibrium Conversely, a chemical reaction within a ½ cell will reach and remain at equilibrium. For a potential to exist, electrons must flow from one cell to the other which requires the reaction to proceed  not at equilibrium.

20. Fundamentals of Electrochemistry • Cells as Chemical Probes 2.) Example: • If the voltage for the following cell is 0.512V, find Kspfor Cu(IO3)2: Ni(s)|NiSO4(0.0025M)||KIO3(0.10 M)|Cu(IO3)2(s)|Cu(s)

21. Fundamentals of Electrochemistry • Biochemists Use Eo´ 1.) Redox Potentials Containing Acids or Bases are pH Dependent • Standard potential  all concentrations = 1 M • pH=0 for [H+] = 1M 2.) pH Inside of a Plant or Animal Cell is ~ 7 • Standard potentials at pH =0 not appropriate for biological systems • Reduction or oxidation strength may be reversed at pH 0 compared to pH 7 Metabolic Pathways

22. Fundamentals of Electrochemistry • Biochemists Use Eo´ 3.) Formal Potential • Reduction potential that applies under a specified set of conditions • Formal potential at pH 7 is Eo´ Eo´ (V) Need to express concentrations as function of Ka and [H+]. Cannot use formal concentrations!