Electrochemical CellElectrochemical Cell  Electrochemical device with 2 half-cells connecting electrodes and solutions  Electrode —metal strip in electrochemical cell

2 types of electrochemical cells2 types of electrochemical cells 1)Voltaic Cells/Galvanic Cell 2)Electrolytic Cells  Still dealing with oxidation-reduction reactions  Physical separation of oxidation and reduction processes

1) Voltaic Cells/Galvanic Cell  “simple battery”  Electric current generated from a redox reaction Pathway of electron transfer  Redox reactions in this cell are always SPONTANEOUS ( Δ G < 0 )  Physically separates oxidation process from reduction process

Voltaic Cell—Oxidation Process  Anode  Electrode where oxidation occurs  Negative charge, source of electrons  Metal electrode dissolves and metallic ions form in solution  Electrons released into solution and buildup NEGATIVE charge at electrode---electrons migrate out through connecting wire  GIVING UP ELECTRONS !!

Voltaic Cell—Reduction Process  Cathode  Electrode where reduction occurs  Positive charge, electron receiver, ion source  Metallic ions (cations) attracted to electrons on electrode surface and accept electrons coming from the anode through the connecting wire.  Metallic ions converted to solid metal on the electrode

Salt BridgeSalt Bridge  U-shaped tube containing a soluble salt in a saturated solution (ex. KNO 3 )  Salt solution MUST be soluble, not form precipitate  Maintains electrical neutrality within cell solutions  Electrons do NOT go through bridge, only through wire  Salt dissociates into ions, ions move to balance charges  Negative ions move to ANODE, minimize POSITIVE charge  Positive ions move to CATHODE, minimize NEGATIVE charge  No acting role in redox reaction

General Points for Voltaic Cells  Electrons move from ANODE (-) to CATHODE (+)  Electrons have HIGH potential energy at anode, not cathode  Naturally favor a state of low potential energy  Electron movement through the connecting wire generates an electric current that can be utilized.

Cell DiagramsCell Diagrams  Representation of an electrochemical cell, short-hand method to actual drawing of cell  Anode---  Left portion  Cathode---  Right portion  Single line—  Boundary between electrode and solution  Double line—  Represents salt bridge

Example 1:Example 1:  Oxidation: Zn (s)  Zn e -  Reduction: Cu e -  Cu (s) Remember to combine reactions by balancing elements and electrons in half reactions.

Example 2:Example 2:  Write the equation for the redox reaction occurring in this voltaic cell. Al (s) Al +3 (aq) H + (aq) H 2(g) Pt (s)

Cell Potential (E cell )Cell Potential (E cell )  Also called “ cell voltage ”  “Driving force” transferring electrons from anode to cathode  Difference between the electric potential between the electrodes in an electrochemical cell  Magnitude indicates amount of current generated through redox reaction  Measured by voltmeter, units = volts (V)

How do we cell potential/voltage? 1)Voltmeter 2)Calculation of cell voltages  Find cell potentials for each half-cell reaction and combine these potentials  Need to set a baseline or zero point for measuring electrode potentials for half-cell reactions

Standard Hydrogen Electrode (SHE)  Assigned zero point/baseline for electrode potentials  All electrode potentials based on this point  H 2 gas passed over Pt electrode at standard conditions  1 atm, 1M, 25°C°  2H + + 2e -  H 2(g) E° = 0V

Standard Electrode Potential (E°)  Also known as “reduction potentials”  Tendency for reduction to happen at an electrode  Measured with solutions at 1M and gases at 1atm, 25°C  Used to determine standard cell potential (E°cell) for an overall reaction

Standard Cell Potential (E°cell)  Difference between the standard potential of the cathode and the standard potential of the anode.  Measured with a voltmeter  E°cell = E°cathode –E°anode OR E° cell = E° ox + E° red  Enables us to indirectly calculate the standard electrode potentials for chemical compounds with unknown potentials.

Standard Electrode Potentials  Located in reference tables for common reduction half- reactions (Table 18.1 p. 762, Appendix C)  Arranged from increasing to decreasing E° values  Compounds favoring reduction, high on table  Compounds favoring oxidation, low on table

Example 1:Example 1:  Find E° Cu+2/Cu based on the following reaction.  Pt H 2(g) H + (aq) Cu +2 (aq) Cu (s) E°cell = 0.340V

Example 2:Example 2:  Find E° Zn+2/Zn based on the following reaction.  Pt H 2(g) H + (aq) Zn +2 (aq) Zn (s) E°cell = V

Example 3:Example 3:  Calculate the standard cell potential (E°cell) for the following voltaic cell:  Zn (s) Zn +2 (aq) Cu +2 (aq) Cu (s)

Example 4:Example 4:  Determine the standard electrode potential (E° Sm+2/Sm )  Sm (s) Sm +2 I - I 2(s) Pt (s) E°cell = 3.21V

Example 5:Example 5:  Balance the following redox reaction and determine the E°cell.  O 2(g) + H + (aq) + I - (aq)  H 2 O (l) + I 2(s)

Example 6:Example 6:  Zn e -  Zn (s) E° = V  Zn (s)  Zn e - E° = V