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Basics of Electrochemistry: Redox Reactions, Current, and Potential

This chapter provides an introduction to electrochemistry, covering topics such as redox reactions, current flow, electrical potential, electron charge, Ohm's law, and Galvanic cells. It also explores the concept of electrical power, the use of voltmeters, and the importance of a salt bridge in electrochemical cells. Additionally, the Nernst equation and the role of standard potential in determining cell potential are discussed, along with the use of electrochemical cells as chemical probes.

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Basics of Electrochemistry: Redox Reactions, Current, and Potential

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  1. Chapter 14 Electrochemistry

  2. Basic Concepts • Chemical Reaction that involves the transfer of electrons. A Redox reaction. • Loss of electrons – oxidation • Gain of electrons – reduction • Oxidizing agent. A species that takes electrons. • Reducing agent. A species that gives electrons.

  3. Basics • Na(s) + H+ -> Na+ + H2(g) • Sodium is a reducing agent • Hydrogen ion is the oxidizing agent.

  4. Basics • We are donating and gaining electrons. If we could use these electrons perhaps we could do some useful work. • If we can make the electron travel in an electrical circuit then the amount of current can be measured. • Current is related to reaction rate or amount of reaction • Potential is related to free energy change of the reaction.

  5. Electron Charge • q used to denote. Unit is Coulombs (C) • Charge on a single electron is • 1.602x10-19 C which will allow us to determine the charge on a mole of electrons. • 1.602x10-19 C * 6.022x1023 mol-1) = 96490 C mol-1 • This is called the Faraday Constant • q = nF • n is the number of moles

  6. Current • Charge flowing through a circuit • One ampere, the charge of one coulomb per second flowing past a given point.

  7. Electrodes • The interface between a solution and an electrical circuit. Can be actively involved or just serve as a source or sink for electons.

  8. Electrical Potential • Work required when moving and electric charge from one point to another. • Electrical potential (E) is measured in Volts (V). • Work is a measure of energy, measured in joules (J). • Work = E * q • Joules volts coulombs

  9. Free Energy • Maximum amount of work that can be done on the surroundings is equal to the Gibbs free energy change. • then DG = -work = -Eq • Or DG = -nFE

  10. Ohm’s Law • Current is proportional to the potential and and inversely proportional to the resistance. • I = E/R

  11. Electric Circuit

  12. Power • Work done per unit time. Unit is the J/s which is know as the watt (W). • P = work/sec = Eq/sec = E(q/sec) = EI • P = EI = I2R = E2/R

  13. Galvanic Cells • Spontaneous chemical reaction used to generate electricity. • An example might be

  14. Voltmeter • A device to measure electrical potential. When electrons tend to flow into the negative terminal then a positive voltage is measured. • In this cell • 2 AgCl (s) + 2 e- = 2 Ag + 2 Cl- (aq) Red • Cd (s) + = Cd2+ + 2 e- Oxidation • Cd (s) + 2 AgCl (s) = Cd2+ + 2 Cl- Net • For this reaction we have a DG of -150 kJ/mole per mole of Cd oxidized.

  15. Potential of this System • DG = -150 kJ/mole then we have • E = - DG/nF = -150 x 103 J / (2 mol)(9.649x104 C/mol) • E = + 0.777 J/C = +0.777 V

  16. Cathode/anode • Cathode electrode where reduction occurs • Anode electrode where oxidation occur • Put both terms in alphabetical order to remember

  17. Salt Bridge • Any bridge in upstate New York in the winter. • Used to isolate the half cells so the work can be forced out into an external circuit. • The following cell has a problem.

  18. What is it? • The silver ions in solution can go directly to the cadmium electrode surface and be reduced there. • We need to put in a barrier to rapid ionic transfer.

  19. What about this cell

  20. Isn’t this cute • Chemistry paper dolls?

  21. Line Notation - Instead of Having to Draw the Cells • | phase boundary || salt bridge • For First Cell • Cd(s) | CdCl2(aq) | AgCl(s) | Ag(s) • For Second Cell • Cd(s) | Cd(NO3)2(aq) || AgNO3(aq) | Ag(s)

  22. A Word of Connectors (Two common in USA)

  23. Standard Potential EoThe energy to a half cell at standard conditions (1 M and 25 C) • Let us look at the reduction of silver ion. • Ag+ + e- = Ag(s) • We will compare this to a fixed reference. • That is the SHE or NHE Standard or Normal Hydrogen Electrode. • H+ (aq, A=1) + e- = ½ H2 (g, A = 1)

  24. SHE - All other redox couples are compared to this half cell. It is assigned a value of 0.000 V • In our cell the left side electrode (Pt) is attached to the negative terminal. (Reference) • Value of E are collected into Tables (Appendix H)

  25. Nernst Equation • For the half reaction • aA + ne- = bB Eo = is the standard Potential R = gas constant (8.314472 (V*C)/(k*mol) T = Temp (K) N = # of electrons in the half reaction F = Faraday A = Activity

  26. We will often lump the constants and assume 25 C • Nernst equation (25 C and converting to log10

  27. Complete Reaction • E = E+ - E- for full cell • Steps • Write both half cells as reductions, make electrons equal • Half cell connected to positive terminal is E+ • Other half cell is E- • Net voltage is from the above equation • Balance equation (reversing the left half reaction and adding to other half cell) • E > 0 spontaneous as written • E < 0 spontaneous in reverse

  28. Eo and K

  29. Cells as Chemical Probes • Equilibria between the half cells • Equilibria within each half cell

  30. A Probe Cell

  31. Probe Cell • Right side: • We have our Ksp equilibrium • The electrochemical reaction under this is • AgCl(s) + e- = Ag(s) + Cl- (aq, 0.10 M) • Eo = 0.222 v • Left side: • We have our Ka for the weak acid. • The electrochemical reaction • 2 H+(aq) + e- = H2 (g, 1.00 bar) • E = 0.00, but H+ is not fixed at 1 M so E varies with H+

  32. Eo’ • Formal Potential • Since so many redox couples exist in the body and many have H+ we modify the potential that we use to pH 7. (A little more reasonable than 1 M acid.

  33. Homework • 14- 4 • 13, 14, 15 and 27

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