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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