Standard Voltages

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

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)

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

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

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)

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)

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

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

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

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.

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

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.

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

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)

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

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

Standard Voltages