1 / 21

Lecture 11: Cell Potentials

Lecture 11: Cell Potentials. Reading: Zumdahl 11.2 Outline What is a cell potential? SHE, the electrochemical zero. Using standard reduction potentials. Reminder. “Redox” Chemistry: Reduction and Oxidation. • Oxidation: Loss of electrons (increase in oxidation number).

landen
Télécharger la présentation

Lecture 11: Cell Potentials

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Lecture 11: Cell Potentials • Reading: Zumdahl 11.2 • Outline • What is a cell potential? • SHE, the electrochemical zero. • Using standard reduction potentials.

  2. Reminder • “Redox” Chemistry: Reduction and Oxidation • Oxidation: Loss of electrons (increase in oxidation number) • Reduction: Gain of electrons (a reduction in oxidation number) • Electrons are transferred from the reducing agent to the oxidizing agent • Electrons are transferred from the species being oxidized to that being reduced.

  3. Galvanic Cells (cont.) 8H+ + MnO4- + 5e- Mn2+ + 4H2O Fe2+ Fe3+ + e- x 5

  4. Galvanic Cells (cont.) • Galvanic Cell: Electrochemical cell in which chemical reactions are used to create spontaneous current (electron) flow

  5. Galvanic Cells (cont.) Anode: Electrons are lost Oxidation Cathode: Electrons are gained Reduction

  6. Cell Potentials • In a galvanic cell, we had a species being oxidized at the anode, a species being reduced at the cathode, and electrons flowing from anode to cathode. • The force on the electrons causing them to full is referred to as the electromotive force (EMF). The unit used to quantify this force is the volt (V) • 1 volt = 1 Joule/Coulomb of charge V = J/C

  7. Cell potential and work • From the definition of electromotive force (emf): Volt = work (J)/charge (C) • 1 J of work is done when 1 C of charge is transferred between a potential difference of 1 V.

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

  9. 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. • The above statement requires that we have a reference to use in comparing 1/2 cells. That reference is the standard hydrogen electrode (SHE)

  10. 1/2 Cell Potentials (cont.) • Consider the following galvanic cell • Electrons are spontaneously flowing from the Zn/Zn+2 half cell (anode) to the H2/H+ half cell (cathode)

  11. 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

  12. 1/2 Cell Potentials (cont.) • 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

  13. 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

  14. 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.

  15. 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

  16. 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. E°cell = E°cathode - E°anode

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

  18. 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. • Turn the reaction for the weaker reduction (smaller E°1/2) and turn it into an oxidation. This reaction will be the anode, the other the cathode.

  19. Predicting Galvanic Cells (cont.) • Example. Describe a galvanic cell based on the following: Ag+ + e- Ag E°1/2 = 0.80 V Fe+3 + e- Fe+2 E°1/2 = 0.77 V Weaker reducing agent – turn it around Ag+ + Fe+2 Ag + Fe+3 E°cell = 0.03 V E°cell > 0….cell is galvanic

  20. Another Example • For the following reaction, identify the two half cells, and use these half cells to construct a galvanic cell 3Fe+2(aq) Fe(s) + 2Fe+3(aq) oxidation +2 0 +3 reduction Fe+2(aq) + 2e- Fe(s) E° = -0.44 V Fe+3(aq) + e- Fe+2(aq) E° = +0.77 V

  21. Another Example (cont.) weaker reduction – turn it around Fe+2(aq) + 2e- Fe(s) E° = -0.44 V Fe+3(aq) + e- Fe+2(aq) E° = +0.77 V 2 x 2Fe+3(aq) + Fe(s) 3Fe+2(aq) E°cell = 1.21 V E° = +0.44 V Fe(s) Fe+2(aq) + 2e-

More Related