Appears 5 out of 75 MC Questions, and Almost Every Year on FR

Electrochemistry Review

Electrochemistry Review

AP Chemistry 2011

Appears 5 out of 75 MC questions, and almost every year on FR

Main Topics

1.  Oxidation Numbers & Redox Reactions

2.  Balancing Redox Reactions (on attached handout)

3.  Galvanic Cells

4.  Reduction Potentials & Cell Potential

5.  Relationships between Voltage/ Spontaneity/ Equilibrium

6.  Electrolytic Cells

Tips for success!

1.  Memorize those acronyms! (FAT CAT, AN OX, RED CAT)

2.  Understand what is occurring at each of the major components in a galvanic cell

3.  Check this animation out (or Google “Galvanic Cell Animation”): http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

4.  Memorize the cell potential equation! E°cell = E°cathode - E°anode (others are given on AP sheet)

5.  Know how to use the Nernst equation both quantitatively and qualitatively

6.  In electrolytic cells, always consider what is actually present the solution

I.  Oxidation Numbers & Redox Reactions

A.  Oxidation Numbers indicate number of electrons gained/lost when forming a bond

B.  Rules for Assigning Oxidation Numbers

1.  Oxidation number of nonbonded atoms are zero (including diatomics!)

2.  Oxidation for atoms in molecules add up to their charge

3.  Oxidation numbers of the following will not change:

Element / Oxidation Number
Alkali Metals (Li, Na, etc) / +1
Alkaline Earth (Be, Mg, etc) / +2
Group 3A (B, Al, etc) / +3
Oxygen / -2 (except in peroxide: -1)
Halogens (F, Cl, etc) / -1

a)  Exception! More electronegative element will keep # above (ClO4-)

4.  Determine the oxidation numbers of the following underlined atoms:

a)  MnO41-

b)  N2

c)  H2SO3

C.  Redox Reactions

1.  Redox Reactions are reactions in which there is a transfer of electrons

a)  One atom gains electrons (reduced) and one atom loses electrons (oxidized) in redox reactions. They go hand in hand!

2.  LEO the lion says GER

a)  LEO: Lose Electrons, Oxidation

b)  GER: Gain Electrons, Reduced

3.  Oxidizing Agent: molecule containing atom that is being reduced and oxidizes another atom.

4.  Reducing Agent: molecule containing atom that is being oxidized and reduces another atom.

a)  Redox reactions can be written as two half reactions: one for reduction, one for oxidation.

5.  Determine the atom being oxidized, atom being reduced, reducing agent, oxidizing agent:

a)  Fe + HCl → FeCl2 + H2

b)  Cr2O72- + 6 I- + 14 H+ → 2 Cr3+ + 3 I2 + 7 H2O

c)  6 HOCl + 2 NO → 3 Cl2 + 2 NO31- + 2 H2O + 2H+

II.  Balancing Redox Reactions (see attached handout)

III.  Galvanic Cells

A.  Galvanic cells (aka voltaic cells) are a spontaneous redox reaction that generates electric current

B.  Know the parts of the Galvanic Cell!

1.  Anode is the place where oxidation occurs (AN OX!)

a)  Things to Note! Will be negative (in galvanic), will appear smaller

2.  Cathode is the place where reduction occurs (RED CAT!)

a)  Things to Note! Will be positive (in galvanic), will appear larger

3.  Electron flows from the anode to the cathode (FAT CAT!)

C.  Build Your Own! Sample Galvanic Cell Diagram:

Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)

D.  Salt Bridge allows passage of ions that maintains electrical neutrality between cells (must be a salt that does not react!)

E.  Electron flow is from Anode to Cathode (FAT CAT!)

IV.  Reduction Potentials & Cell Potentials

A.  Every half reaction has a potential, or voltage

1.  Standard reduction potentials (E°) are given, but the oxidation potentials are just the negative of these

B.  The larger the E° value, the greater its ability to reduce

1.  Vice versa: smaller the E°, greater ability to oxidize

C.  Cell Potential (aka Cell EMF) at standard conditions (E°cell)

1.  Memorize!! E°cell = E°cathode - E°anode

a)  Important tip: Ignore coefficients of half reaction!

D.  Determine the standard cell potentials for the following:

1.  Fe + Cu2+ → Fe2+ + Cu

2.  2 AgNO3 + Zn → Zn(NO3)2 + 2 Ag

3.  Zn + CuSO4 → ZnSO4 + Cu

E.  Cell Potential at non-standard conditions (Ecell)

1.  When not at 1M (or 25 °C) must use Nernst Equation!

The Nernst Equation

Ecell = cell potential at nonstandard conditions (V)
E°cell = cell potential at standard conditions (V)
R = “energy” gas constant, 8.31 (volt-coulomb/mol K)
T = absolute temperature (K)
n = number of moles of electrons transferred (mol)
F = Faraday’s constant, 96,500 coulombs/mole
Q = reaction quotient (same as equilib. constant but with initial concentrations)

F.  However, if at 25 °C, you get to use Nernst Lite!

G.  Key Point: as concentration of products increases, voltage decreases. As reactants increases, voltage increases (direct result of Le Chatelier’s Principle)

H.  What is the voltage of a Zn/Al cell when [Al3+] = 1.60 M and [Zn2+] = 0.10 M?

V.  Relationships between Voltage/Spontaneity/Equilibrium

A.  When cell potential is positive the reaction is spontaneous

B.  Relationship between free energy and cell potential:

∆G° = -nFE°
∆G° = Standard Gibbs free energy change (kJ/mol)
n = number of moles of electrons exchanged in the reaction (mol)
F = Faraday’s constant, 96,500 coulombs / mole
E° = standard cell potential

1.  Note how when ∆G is negative (implying spontaneous), E is positive

C.  Relationship between equilibrium and cell potential (at 25 °C):

E° = standard cell potential
n = number of moles of electrons exchanged in the reaction (mol)
K = equilibrium constant

VI.  Electrolytic Cells

A.  Electrolytic Cells are nonspontaneous redox reactions that require an outside source of voltage

B.  RED CAT/AN OX still apply, however the polarity of the electrodes switches

1.  Remember “E.P.A.” – Electrolytic Positive Anode

C.  Electrolysis is used for electroplating of metals. Given the electrical current (in amps) you can determine the amount of metal that “plates out”

Current

I = current (amperes, A)
q = charge (coulombs, C)
t = time (sec)

1.  Once the charge is determined, the moles of electrons can be found using Faraday’s constant

2.  From moles of electrons, stoichiometry can be used to find moles of the metal (and then grams if needed)

D.  An 800 mL solution containing 0.080 molar Ag+ is electrolyzed, resulting in the formation of silver and oxygen gas. The solution was subjected to a current of 2.00 amperes for 10.0 minutes.

1.  Write the two half-reactions that occur. Determine which occur at the cathode and which at the anode.

2.  How many grams of silver are produced in this reaction?

3.  If the oxygen gas produced was collected at STP, what was its volume?

4.  What is the pH change after the reaction has occurred?

Answers to Practice:
Section 1: 4.a) +7 b) 0 c) +4 5. a) Fe is oxid., H is red., HCl is OA, Fe is RA b) I is oxid., Cr is red., Cr2O72- is OA, I- is RA c) N is oxid., Cl is red., HOCl is OA, NO is RA
Section 3: Diagram: Anode – Zn / Species: Zn2+ and SO42- Cathode – Cu / Species: Cu2+ and SO42-
Salt Bridge: Nonreactive salt (aka K2SO4), Half Reaction Anode: Zn -> Zn2+ + 2e- Cathode: Cu2+ + 2e- -> Cu
Section 4: D 1) +0.7 V, 2) 1.56 V, 3) 1.10 V H 0.87 V
Section 5: D 1) Cathode: Ag+ + e- Ag, Anode: 2 H2O O2 + 4 H+ + 4 e- 2) 1.34 g 3) 0.0694 L 4) 1.81

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