1. Fundamentals of Electrochemistry It’s shocking!

2. Electroanalytical Chemistry: group of analytical methods based upon electrical properties of analytes when part of an electrochemical cell Potentiometry involves the measurement of potential for quantitative analysis. Electrolytic electrochemical phenomena involve the application of a potential or current to drive a chemical phenomenon, resulting in some measurable signal which may be used in an analytical determination

3. The majority of chemical reactions can be classified as one of two kind of major reaction types. Acid/Base Reactions: proton transfer Oxidation/Reduction (Redox) Reactions : electron transfer If you think about it!......

4. Basic Concepts Redox reactions involve a species which is oxidized and another that is reduced. In the above, Fe 3+ is reduced to Fe 2+. It is the oxidizing agent. Since  G<0 for this reaction we can say that V 3+ wants the extra electron less than Fe 3+.

5. Galvanic Cell

6. An Aside Why won’t this cell work? Ag + will go to left electrode and ask for e from Cd(s) directly.

7. Will this cell work?

8. How badly do the electrons want to flow? q = n x F I = current in amps R = resistance in ohms E = potential difference in Volts

9. Voltaic Cells Electrochemical cells that use an oxidation- reduction reaction to generate an electric current are known as galvanic or voltaic cells.

10. Voltaic Cells

11. The voltaic cell consist of the two reactions. reduction oxidation We can only measure E for the full reaction. We would like to calculate E for the half reactions. Before doing this, we must recognize the E depends on concentrations. + Or equivalently we can write the reactions as follows

12. Voltaic Cells

13. Since reactants and products are in their standard states, we call the E for this cell the standard reduction potential (E o ). Here E o =.76V. + We arbitrarily define the potential for, one half reaction, the second reaction above to be exactly 0V when reactants and products are in their standard states. Since E o for the cell is the sum of E o ’s for the two half reactions we see that E o for the first half reaction is.76V.

14. Oxidizing Power Increases

15. Voltaic Cells

16. This voltaic cell on the previous slide is fully described with the following notation +

17. Voltaic cells can be described by a line notation based on the following conventions.  Single vertical line indicates change in state or phase.  Within a half-cell, the reactants are listed before the products.  Activities of aqueous solns are written in parentheses after the symbol for the ion or molecule.  A double vertical line indicates a junction between half-cells.  The line notation for the anode (oxidation) is written before the line notation for the cathode (reduction). Line Notation For Voltaic Cells

18. Zn|Zn 2+ (1.0 M)||Cu 2+ (1.0 M)|Cu anode (oxidation) cathode (reduction) Electrons flow from the anode to the cathode in a voltaic cell. (They flow from the electrode at which they are given off to the electrode at which they are consumed.) Reading from left to right, this line notation therefore corresponds to the direction in which electrons flow.

20. The Nernst Equation The Nernst equation relates the potential of a cell in its standard state to that of a cell not in its standard state.. We know from Le Chatelier’s principle that increasing the concentration of Zn 2+ should drive the reaction to the right. In other words it should decrease the potential of the half cell. The Nernst equation allows us to calculate this increase for the above half reaction as

21. At 25 o C this equation simplifies to The Nernst Eq. for the reaction is

22. The Nernst Equation For Complete Cell Here E + and E - are the potentials of the half cells connected to the positive and negative terminals of potentiometer respectively. Let’s consider an example.

23. Taken from http://chemed.chem.purdue.edu/genchem Voltaic Cells -+

24. E + and E - are potentials of half cells connected to positive and negative terminals of potentiometer respectively

27. Calculating Equilibrium Constants E o =1.700V E o =0.767V made up of the following two half reactions Since E o is greater for cerium this reaction will be the reduction reaction. The standard potential for the galvanic cell would be

28. Calculating Equilibrium Constants Continued In a galvanic cell we would have At equilibrium E=0 and This connection to free energy is important

29. Calculating Equilibrium Constants for Nonredox Reactions This is a K sp problem. Not a redox problem. Nonetheless we can use electrochemistry to calculate K sp by considering (at 25 o C)

30. Electrochemistry Skills Understand how voltaic cells work. Be able to calculate standard reduction potentials for voltaic cells, given the chemical reactions. Be able to describe a voltaic cell using the line notation and visa versa. Know which way electrons flow and where the anode and cathode are. Know how to work with the Nernst Eq. to include concentration dependencies and calculate equilibirum constants