Fundamentals of Electrochemistry Introduction 1 Electrical Measurements of

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

Fundamentals of Electrochemistry Basic Concepts 1. ) A Redox titration is an analytical technique based on the transfer of electrons between analyte and titrant Ø Reduction-oxidation reaction Ø A substance is reduced when it gains electrons from another substance - Ø 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

Fundamentals of Electrochemistry Basic Concepts 2. ) The first two reactions are known as “ 1/2 cell reactions” Ø 3. ) Include electrons in their equation The net reaction is known as the total cell reaction Ø No free electrons in its equation ½ cell reactions: Net Reaction: 4. ) In order for a redox reaction to occur, both reduction of one compound and oxidation of another must take place simultaneously Ø Total number of electrons is constant

Fundamentals of Electrochemistry Basic Concepts 5. ) Electric Charge (q) Ø Measured in coulombs (C) Ø Charge of a single electron is 1. 602 x 10 -19 C Ø Faraday constant (F) – 9. 649 x 104 C is the charge of a mole of electrons Relation between charge and moles: Coulombs moles 6. ) Electric current Ø Quantity of charge flowing each second through a circuit Ampere: unit of current (C/sec)

Fundamentals of Electrochemistry Galvanic Cells 1. ) Galvanic or Voltaic cell Ø Spontaneous chemical reaction to generate electricity - Ø 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: 2+ Ag. Cl(s) toto. Ag(s) Cd(s)isisreduced oxidized Cd Electrons travel from Cd Ag deposited onto electrode and Cl. Cd 2+ goes into electrode Ag solution electrode goes into solution

Fundamentals of Electrochemistry Galvanic Cells 2. ) Cell Potentials Ø Reaction is spontaneous if it does not require external energy Reaction Type E Cell Type Spontaneous + Galvanic Nonspontaneous - Electrolytic Equilibrium 0 Dead battery Potential of overall cell = measure of the tendency of a reaction to proceed to equilibrium ˆ Larger the potential, the further the reaction is from equilibrium and the greater the driving force that exists

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

Fundamentals of Electrochemistry Galvanic Cells 4. ) Salt Bridge Ø Ø Connects & separates two half-cell reactions Prevents charge build-up and allows counter-ion migration Salt Bridge § Contains electrolytes not involved in redox reaction. Two. Cd half-cell reactions 2+) moves § K+ (and to cathode with e through salt bridge (counter balances –charge build-up § NO 3 - moves to anode (counter balances +charge build-up) § Completes circuit

Fundamentals of Electrochemistry Galvanic Cells 5. ) Short-Hand Notation Ø Representation of Cells: by convention start with anode on left Phase boundary Electrode/solution interface anode Zn|Zn. SO 4(a. ZN 2+ = 0. 0100)||Cu. SO 4(a. Cu 2+ = 0. 0100)|Cu Solution in contact with anode & its concentration 2 liquid junctions due to salt bridge cathode Solution in contact with cathode & its concentration

Fundamentals of Electrochemistry Standard Potentials 1. ) Predict voltage observed when two half-cells are connected Ø 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 Ag+ + e- » Ag(s) Ø Eo = +0. 799 V Potentials measured at standard conditions - All concentrations (or activities) = 1 M 25 o. C, 1 atm pressure Standard Hydrogen Electrode (S. H. E) Pt(s)|H 2(g)(a. H = 1)|H+(aq)(a. H+ = 1)|| 2 Hydrogen gas is bubbled over a Pt electrode

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

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) Electrons always flow towards more positive potential Place values on number line to determine the potential difference

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

Fundamentals of Electrochemistry Nernst Equation 1. ) Reduction Potential under Non-standard Conditions Ø Ø E determined using Nernst Equation Concentrations not-equal to 1 M For the given reaction: a. A + ne- » b. B Eo The ½ cell reduction potential is given by: Where: at 25 o. C 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. 314 J/(K-mol) F = Faraday’s constant (9. 649 x 104 C/mol) A = activity of A or B

Fundamentals of Electrochemistry Nernst Equation 2. ) Example: Ø Calculate the cell voltage if the concentration of Na. F and KCl were each 0. 10 M in the following cell: Pb(s) | Pb. F 2(s) | F- (aq) || Cl- (aq) | Ag. Cl(s) | Ag(s)

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

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

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

Fundamentals of Electrochemistry Cells as Chemical Probes 1. ) Two Types of Equilibrium in Galvanic Cells Ø Ø 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.

Fundamentals of Electrochemistry Cells as Chemical Probes 2. ) Example: Ø If the voltage for the following cell is 0. 512 V, find Ksp for Cu(IO 3)2: Ni(s)|Ni. SO 4(0. 0025 M)||KIO 3(0. 10 M)|Cu(IO 3)2(s)|Cu(s)

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

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