Electrochemistry
Electrochemistry = the interchange of chemical and electrical energy
= used constantly in batteries, chemical instruments, etc…
Galvanic Cells
A.
Definitions
1) Redox Reaction = oxidation/reduction reaction = chemical reaction in which
electrons are transferred from a reducing agent (which gets oxidized) to an
oxidizing agent (which gets reduced)
2) Oxidation = loss of electron(s) to become more positively charged
3) Reduction = gain of electron(s) to become more negatively charged
B.
Using Redox Reactions to generate electric current (moving electrons)
1) Zno + Cu2+
Cuo + Zn2+
a) Zno is oxidized and Cu2+ is reduced
b) Half Reaction = oxidation or reduction process only
Reduction: Cu2+ + 2eCuo
Sum = Redox Rxn
o
2+
Oxidation: Zn
Zn + 2e-)
2)
In solution:
4) Zno and Cu2+ collide and electrons are transferred
a) No work can be obtained; only heat is generated
3)
In separate compartments, electrons must go through a wire = Galvanic Cell
a) Generates a current = moving electrons from Zno side to Cu2+ side
b) Current can produce work in a motor or light up a light bulb
c) Salt Bridge = allows ion flow without mixing solutions (Jello-like matrix)
d) Chemical reactions occur at Electrodes = conducting solid dipped into solution
i) Anode = electrode where oxidation occurs (production of e-)
ii) Cathode = electrode where reduction occurs (using up e-)
C.
Cell Potential
1) Think of the Galvanic Cell as an oxidizing agent “pulling” electrons off of the
reducing agent. The “pull” = Cell Potential
a) ecell = Cell Potential = Electromotive Force = emf
b) Units for ecell = Volt = V
1 V = 1 Joule/1 Coulomb
2) Voltmeter = instrument drawing current through a known resistance to find V
Potentiometer = voltmeter that doesn’t effect V by measuring it
3)
Standard Hydrogen Electrode: must have a standard to compare emf to
Cathode = Pt electrode in 1 M H+ and 1 atm of H2(g)
Half Reaction: 2H+ + 2eH2(g)
e1/2 = 0
4)
Standard Reduction Potentials can be found in your text appendices
a) Always given as a reduction process
b) All solutes are 1M, gases = 1 atm
5)
Combining Half Reactions to find Cell Potentials
a) Reverse one of the half reactions to an oxidation; this reverses the sign of e1/2
b) Don’t need to multiply for coefficients = Intensive Property (color, flavor)
c) Example: 2Fe3+(aq) + Cuo
2Fe2+(aq) + Cu2+(aq)
i. Fe3+ + eFe2+ e1/2 = +0.77 V
ii. Cu2+ + 2eCuo e1/2 = +0.34 V
iii. Reverse of (ii) added to (i) = -0.34 V + +0.77 V = +0.43 V = e1/2
D.
Direction of electron flow in a cell
1) Cell always runs in a direction to produce a positive ecell
2) Fe2+ + 2eFeo
e1/2 = -0.44 V
MnO4- + 5e- + 8H+
Mn2+ + 4H2O
e1/2 = +1.51 V
3)
We put the cell together to get a positive potential: DG = -nFecell
a) 5(Feo
Fe2+ + 2e-)
e1/2 = +0.44 V
b) 2(MnO4- + 5e- + 8H+
Mn2+ + 4H2O)
e1/2 = +1.51 V
16H+(aq) + 2MnO4-(aq) + 5Feo(s)
2Mn2+(aq) + 5Fe2+(aq) + 8H2O(l)
ecell = 1.95V
E.
Notes on the Experimental Procedure
1. Do all of Part I: Direct Redox Reactions
2. Part II Indirect Spontaneous Redox Reactions: Only do procedures 1-16
3. Skip 17-20 of Part II and all of Part III: Indirect Non-Spontaneous Reactions
4. Make sure to clean metal electrodes with sandpaper to get best results
5. Potentials may not be identical to predicted, but the relative sizes will be
6. Use Nickel in place of Tin (Sn); it works better
7. Don’t throw away any metal pieces; clean them, dry them, put them back
8. Use the same piece of filter paper (NaNO3 soaked) for all galvanic cells
9. Place solutions in wells on plate, so that salt bridge can reach all needed
V