Understanding Electrochemistry and Redox Reactions in Chemistry Lectures
Dive into the world of electrochemistry and redox reactions through this insightful lecture series. Explore fundamental equations, galvanic cells, and practical examples to enhance your understanding. Prepare for Exam 1 with in-depth explanations on topics like charge transfer, electrical work, and battery efficiency.
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Presentation Transcript
Announcements Exam 1 coming up (Mar. 7th) Friday s Seminar Gallo Winery (Internship opportunity) Today s Lecture - Electrochemistry Redox Reactions Fundamental Equations (relating moles of electrons to charge and G to electrical energy Galvanic Cells Standard Potential Nernst Equation (if time)
Electrochemistry Redox Reactions Reduction = loss of charge e.g. Fe3++ e- Fe2+ Oxidation = gain in charge e.g. Pb2++ 2H2O PbO2(s) + 4H++ 2e-(Pb goes from +2 to +4) Balancing reactions review steps in general chemistry book example: Zn(s) + Cr2O72- Zn2++ Cr3+ note: based on methods used for problems in this book, full cell balancing may not be needed
Electrochemistry Fundamental Equations Relationship between charge, energy and current redox reactions involve the exchange of electrons when the exchange occurs on an electrode surface, current can be measured Total charge transfer = q = nF, where n = moles of electrons in reaction and F = Faraday s constant = 96500 C/moles e- F = NAvogadro e (e = elementary charge = 1.6 x 10-19C) Current Produced = I = dq/dt or q = I dt (or = I t under constant current conditions) can be used to determine battery lifetime
Electrochemistry Fundamental Equations Relationship between charge, energy and current (continued) Electrical work (units = J) = E q (E = potential in volts and q in C) and G = -E q = -nFE under standard conditions (1 M reactant/product conc., 298K, etc.), G = -nFE G are given in Tables and allows calculation of K values E , standard reduction potential, also given in Tables (see Appendix H), but for half-reactions
Electrochemistry Fundamental Equations Example problem: A NiCad battery contains 12.0 g of Cd that is oxidized to Cd(OH)2. How long should the battery last if a motor is drawing 421 mA? Assume 100% efficiency.
Electrochemistry Galvanic Cells What are galvanic cells? Cells that use chemical reactions to generate electrical energy Batteries are examples of useful galvanic cells Example reaction GALVANIC CELL voltmeter Zn(s) Ag(s) Zn(s) + 2Ag+ Zn2+ + 2Ag(s) If reactants are placed in a beaker, only products + heat are produced When half reactions are isolated on electrodes, electrical work can be produced AgNO3(aq) ZnSO4(aq) Salt Bridge
Electrochemistry Galvanic Cells Description of how example cell works Reaction on anode = oxidation Anode = Zn electrode (as the E for Zn2+ is less than for that for Ag+) So, reaction on cathode must be reduction and involve Ag Oxidation produces e-, so anode has ( ) charge (galvanic cells only); current runs from cathode to anode Salt bridge allows replenishment of ions as cations migrate to cathode and anions toward anodes GALVANIC CELL voltmeter Ag+ + e- Ag(s) Zn(s) Ag(s) + AgNO3(aq) ZnSO4(aq) Zn(s) Zn2+ + 2e- Salt Bridge
Electrochemistry Galvanic Cells Cell notation Example Cell: Zn(s)|ZnSO4(aq)||AgNO3(aq)|Ag(s) GALVANIC CELL voltmeter Zn(s) Ag(s) | means phase boundary left side for anode (right side for cathode) || means salt bridge AgNO3(aq) ZnSO4(aq) Salt Bridge
Electrochemistry Galvanic Cells Example Questions Given the following cell, answer the following question: MnO2(s)|Mn2+(aq)||Cr3+(aq)|Cr(s) What compound is used for the anode? What compound is used for the cathode? Write out both half-cell reactions and a net reaction
Electrochemistry Standard Reduction Potential A half cell or electrode, is half of a galvanic cell A standard electrode is one under standard conditions (e.g. 1 M AgNO3(aq)) Standard reduction potential (E ) is cell potential when reducing electrode is coupled to standard hydrogen electrode (oxidation electrode) Large + E means easily reduced compounds on electrode Large E means easily oxidized compounds on anode Pt(s) Ag(s) H2(g) AgNO3(aq) H+(aq)
Electrochemistry Electrolytic Cells Used in more advanced electrochemical analysis (not covered in detail) Uses voltage to drive (unfavorable) chemical reactions Example: use of voltage to oxidize phenol in an HPLC electrochemical detector (E of 0 to 0.5 V needed) anode (note: oxidation driven by voltage, but now + charge) cathode (reduction, - charge)
Electrochemistry The Nernst Equation The Nernst Equation relates thermodynamic quantities to electrical quantities for a cell reaction Thermodynamics: G = G + RTlnQ G = free energy, Q = reaction quotient so, -nFE = -nFE + RTlnQ, or E = E (RT/nF)lnQ more often seen as: E = E (0.05916/n)logQ (although only valid at T = 298K) Note: in calculations, E is for reductions (even if oxidation actually occurs at that electrode) Equation for electrodes or full cells, although text uses Ecell = E+ E- where + and refer to voltmeter leads Best to use activities in Q (even though we will just use concentrations)
Electrochemistry The Nernst Equation Example: Determine the voltage for a Ag/AgCl electrode when [Cl-] = 0.010 M if E = 0.222 V (at T = 25 C)?
Electrochemistry Applications of The Nernst Equation Examples: The following electrode, Cd(s)|CdC2O4(s)|C2O42- is used to determine [C2O42-]. It is paired with a reference electrode that has an E value of 0.197 V (vs. the S.H.E.) with the reference electrode connected to the + end of the voltmeter. If E for the above reduction reaction is -0.522 V, and the measured voltage is 0.647 V, what is [C2O42-]?
