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 • Basic Concepts 7.) Electric Potential (E) • Measured in volts (V) • Work (energy) needed when moving an electric charge from one point to another • Measure of force pushing on electrons Higher potential difference Relation between free energy, work and voltage: Volts Coulombs Joules Higher potential difference requires more work to lift water (electrons) to higher trough

6. Fundamentals of Electrochemistry • Basic Concepts 7.) Electric Potential (E) • Combining definition of electrical charge and potential Relation between free energy difference and electric potential difference: Describes the voltage that can be generated by a chemical reaction

7. Fundamentals of Electrochemistry • Basic Concepts 8.) Ohm’s Law • Current (I) is directly proportional to the potential difference (voltage) across a circuit and inversely proportional to the resistance (R) • Ohms (W) - units of resistance 9.) Power (P) • Work done per unit time • Units: joules per second J/sec or watts (W)

8. 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

9. Fundamentals of Electrochemistry • Galvanic Cells 1.) Galvanic or Voltaic cell • Example: Calculate the voltage for the following chemical reaction DG = -150kJ/mol of Cd n – number of moles of electrons Solution:

10. Fundamentals of Electrochemistry • Galvanic Cells 2.) Cell Potentials vs. DG • Reaction is spontaneous if it does not require external energy

11. Fundamentals of Electrochemistry • Galvanic Cells 2.) Cell Potentials vs. DG • Reaction is spontaneous if it does not require external energy Potential of overall cell = measure of the tendency of this reaction to proceed to equilibrium ˆLarger the potential, the further the reaction is from equilibrium and the greater the driving force that exists Similar in concept to balls sitting at different heights along a hill

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

13. 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

14. 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

15. 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

16. 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

17. 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

18. Fundamentals of Electrochemistry • Standard Potentials 3.) Example: Calculate Eo, and DGo for the following reaction:

19. 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

20. 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)

21. 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:

22. 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

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

24. 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.

25. 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)

26. 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

27. 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!