Standard Voltages

1. Standard Voltages • Reading: Masterson 18.2 • Outline • What is a standard voltage (cell potential) • SHE, the electrochemical zero. • Using standard reduction potentials to calculate standard voltage of a voltaic cell

2. Standard Voltage/Cell Potentials • In a voltaic cell, a species is oxidized at the anode, a species is reduced at the cathode, and electrons flow from anode to cathode. • The force on the electrons causing them to flow is referred to as the electromotive force (EMF). The unit used to quantify this force is the volt (V)

3. Standard Voltage/Cell Potentials (cont.) • We can measure the magnitude of the EMF causing electron (i.e., current) flow by measuring the voltage. e- Anode Cathode

4. Standard Voltage/Cell Potentials (cont.) Eo = Eored + Eoox In the case below we have experimentally determined that 1.06 volts = Eored + Eoox e- Anode Cathode

5. 1/2 Cell Potentials • What we seek is a way to predict what the voltage will be between two 1/2 cells without having to measure every possible combination. • To accomplish this, what we need to is to know what the inherent potential for each 1/2 cell is. (i.e. knowing the Eored and Eoox values • The above statement requires that we have a reference to use in comparing 1/2 cells. That reference is the standard hydrogen electrode (SHE)

6. 1/2 Cell Potentials • Consider the following galvanic cell • Electrons are spontaneously flowing from the Zn/Zn+2 half cell (anode) to the H2/H+ half cell (cathode)

7. 1/2 Cell Potentials (cont.) • We define the 1/2 cell potential of the hydrogen 1/2 cell as zero. SHE P(H2) = 1 atm [H+] = 1 M 2H+ + 2e- H2 E°1/2(SHE) = 0 V

8. 1/2 Cell Potentials • With our “zero” we can then measure the voltages of other 1/2 cells. • In our example, Zn/Zn+2 is the anode: oxidation Zn Zn+2 + 2e- E°Zn/Zn+2 = 0.76 V 2H+ + 2e- H2 E° SHE = 0 V Zn + 2H+ Zn+2 + H2 E°cell = E°SHE + E°Zn/Zn+2 = 0.76 V 0

9. Standard Reduction Potentials • Standard Reduction Potentials: The 1/2 cell potentials that are determined by reference to the SHE. • These potentials are always defined with respect to reduction. Zn+2 + 2e- Zn E° = -0.76 V Cu+2 + 2e- Cu E° = +0.34 V Fe+3 + e- Fe+2 E° = 0.77 V

10. Standard Potentials (cont.) • If in constructing an electrochemical cell, you need to write the reaction as a oxidation instead of a reduction, the sign of the 1/2 cell potential changes. Zn+2 + 2e- Zn E° = -0.76 V Zn Zn+2 + 2e- E° = +0.76 V • 1/2 cell potentials are intensive variables. As such, you do NOT multiply them by any coefficients when balancing reactions.

12. Writing Galvanic Cells For galvanic cells, Ecell > 0 In this example: Zn/Zn+2 is the anode Zn Zn+2 + 2e- E° = +0.76 V Cu/Cu+2 is the cathode Cu+2 + 2e- Cu E° = 0.34 V

13. Writing Galvanic Cells (cont.) Zn Zn+2 + 2e- E° = +0.76 V Cu+2 + 2e- Cu E° = 0.34 V Cu+2 + Zn Cu + Zn+2 E°cell = 1.10 V Notice, we “reverse” the potential for the anode.

14. Writing Galvanic Cells (cont.) Shorthand Notation Zn|Zn+2||Cu+2|Cu Anode Cathode Salt bridge

15. Predicting Galvanic Cells • Given two 1/2 cell reactions, how can one construct a galvanic cell? • Need to compare the reduction potentials of the two half cells. • The stronger reducing agent will become the anode and get oxidized (flip this equation) while the stronger oxidizing agent will become the cathode and get reduced. ALL OLD CARS RUST= Anode (oxidation); Cathode (reduction)

16. What about our copperplating lab? Was it an example of galvanic/voltaic cell? The anode was the copper metal and the iron nail was the cathode ….. Yet copper is a weaker reducing agent than iron so you would expect their roles to be reversed? The answer is NO! It is not a galvanic cell. We accomplished this feat by using the battery as an external electron pump. In an electrolytic cell, a nonspontaneous redox reaction is made by pumping electrical energy into the system … electrolysis. (Section 18.5 of your textbook) reducing agent Weaker

17. Homework #116 DUE MONDAY #19, 21b, 23 reducing agent Weaker