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Electrochemistry Introduction
Voltaic Cells
Electrochemical Cell
Electrochemical device with 2 half-cells connecting electrodes and solutions
Electrode—metal strip in electrochemical cell
2 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 (ΔG < 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 Bridge
U-shaped tube containing a soluble salt in a saturated solution (ex. KNO3) 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 Potential (Ecell)
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)
Cell 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:
Oxidation: Zn(s) Zn+2 + 2e-
Reduction: Cu+2 + 2e- Cu(s)
Remember to combine reactions by balancing elements and electrons in half reactions.
Example 2:
Write the equation for the redox reaction occurring in this voltaic cell.
Al(s) Al+3(aq) H+
(aq) H2(g) Pt(s)
Standard Electrode Potentials
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
H2 gas passed over Pt electrode at standard conditions 1 atm, 1M, 25°C° 2H+ + 2e- H2(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:
Find E°Cu+2/Cu based on the following reaction.
Pt H2(g) H+(aq) Cu+2
(aq) Cu(s) E°cell = 0.340V
Example 2:
Find E°Zn+2/Zn based on the following reaction.
Pt H2(g) H+(aq) Zn+2
(aq) Zn(s) E°cell = -0.763V
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:
Determine the standard electrode potential (E°Sm+2/Sm)
Sm(s) Sm+2 I- I2(s) Pt(s) E°cell = 3.21V
Example 5:
Balance the following redox reaction and determine the E°cell. O2(g) + H+
(aq) + I-(aq) H2O(l) + I2(s)
Example 6:
Zn+2 + 2e- Zn(s) E° = -0.7628 V
Zn(s) Zn+2 + 2e- E° = +0.7628V
Homework
p.759 #18.3B
p. 792-793 #33, 34, 35, 37-40