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Electrochemistry
Chapter 20
Oxidation-Reduction Review
oxidation: lose e-increase
oxidation number
reduction: gain e-reduces
oxidation number
LEO goes GER
Oxidation Numbers Review
Review rules p. 132
1. HNO3
6. Ag
10. CO2
1. CuCl2
7. PbSO4
11. (NH4)2Ce(SO4)3
1. O2
8. PbO2
12. Cr2O3
1. H2O2
9. Na2C2O4
1. MgSO4
Oxidation-Reduction Review
1. CH4(g) + H2O(g) → CO(g) + 3H2(g)
1. 2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag(s)
1. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
1. 2H+(aq) + 2CrO42-(aq) → Cr2O72-(aq) + H2O(l)
Balancing Redox Reactions
Half Reactions: show only the oxidation or only the
reduction part of the reaction
Sn2+(aq) + 2Fe3+(aq) → Sn4+(aq) + 2Fe2+(aq)
ox.: Sn2+(aq) → Sn4+(aq) + 2ered.: 2e- + 2Fe3+(aq) → 2Fe2+(aq)
Balancing Redox Reactions
1. split into the half-reactions
2. balance each half-reactions
a. first balance elements other than H and O
b. balance O atoms by adding H2O as needed
c. balance H atoms by adding H+ as needed
d. balance charge by adding e- as needed
3. multiply half-reactions by as needed to make the # of e- lost in ox.
equal to # of e- gained in red.
4. add half reactions, simplify by canceling
5. double check everything
Balancing Redox Reactions
What is reduced?
What is oxidized?
Reducing agent?
Oxidizing agent?
Balancing Redox Reactions
MnO4-(aq) + C2O42-(aq) → Mn2+(aq) + CO2(aq)
Balancing Redox Reactions
Do free electrons appear anywhere in the balanced
equation for a redox reaction? Explain
Balancing Redox Reactions
Cr2O72-(aq) + Cl-(aq) → Cr3+(aq) + Cl2(g) (acidic solution)
Balancing Redox Reactions
Cu(s) + NO3-(aq) → Cu2+(aq) + NO2(aq)
(acidic)
Balancing Redox Reactions
Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO4-(aq)
(acidic)
Balancing Redox Reactions
CN-(aq) + MnO4-(aq) → CNO-(aq) + MnO2(s)
(basic)
Balancing Redox Reactions
NO2-(aq) + Al(s) → NH3(aq) + Al(OH)4-(aq)
(basic)
Balancing Redox Reactions
Cr(OH)3(s) + ClO-(aq) → CrO42-(aq) + Cl2(g)
Electrochemistry
study of the interchange of chemical and
electrical energy
Galvanic or Voltaic Cells
-Device in which chemical energy is changed to electrical
energy
-Uses a spontaneous redox reaction to produce a current
that can be used to do work
-oxidizing agent is separated from the reducing
agent and the electrons are forced to transfer through a
wire
-the current produced through the wire can then be
used to do work
Galvanic or Voltaic Cells
Galvanic or Voltaic Cells
Salt Bridge: allows ions to flow so the net
charge is zero
-without the salt bridge, the current would
stop flowing due to a charge buildup
-negative at red. side and positive at
ox. side
Galvanic or Voltaic Cells
Anode: where oxidation occurs
Cathode: where reduction occurs
AN OX, Red Cat
Why does Na+ migrate into the cathode half-cell as the cell reaction
proceeds?
Cell Potential
Cell Potential: (E°cell)
-the oxidizing agent pulls electrons through the wire from the reducing
agent.
aka. electromotive force
-unit is the volt: 1V = 1J of work per coulomb of charge (J/coulomb)
-measured with a voltmeter:
-drawing of current through a known resistance
-or a potentiometer: measures opposition to current
(compares to a known emf)
Standard Reduction Potentials
All galvanic cell reactions are redox reactions that can be broken down
into 2 half reactions
-can assign a potential to each half reaction and calculate the
cell potential by adding the potentials of the two half reactions
The standard:
-platinum electrode in contact with 1M H+ and bathed by H2(g)
at 1atm has a reduction potential of 0V → called a standard hydrogen
electrode
2H+ + 2e- → H2 E° = 0v
Standard Hydrogen Electrode
Standard Reduction Potentials
-Compare all half reactions to the standard hydrogen electrode to
obtain the reduction potential of the half-reaction
E° means standard state: 1M, 1atm
-Potentials of half reactions are given as reduction potentials - pink
sheet
-oxidation potentials are the reverse
E° is an intensive property: does not change when multiplied
E°cell = E°cathode - E°anode
SRP Calculation
SRP Calculation
SRP Calculation
Calculating Cell Potential
Voltaic cell based on the reaction:
Fe3+(aq) + Cu(s) → Cu2+ + Fe2+
Calculating Cell Potential
Mg anode and Sn cathode
Calculating Cell Potential
Cu cathode and Zn anode
Calculating Cell Potential
Cr anode and Ag cathode
Strength of Oxidizing & Reducing Agents
more negative E = most likely to be reversed and run as an
oxidation reaction
-Li+: E = -3.05: poor oxidizing agent
-difficult to reduce (gain e-)
- it is a very good reducing agent: will easily
lose e-
Would the halogens be good reducing agents or good
oxidizing agents? Why
SRP Table
In the reduction half reaction:
the reactants are oxidizing agents
the products are reducing agents
SRP Table
Relative Strengths of Oxidizing Agents
Rank the following ions in order of increasing strength as
oxidizing agents: NO3-, Ag+, Cr2O72Rank the following ions in order of decreasing strength as
reducing agents: I-(aq), Fe(s), Al(s)
Determining Spontaneity
+E = spontaneous
-E = nonspontaneous
Are the following reactions spontaneous?
1.Cu(s) + 2H+(aq) → Cu2+(aq) + H2(g)
2. Cl2(g) + 2I-(aq) → 2Cl-(aq) + I2(s)
Free Energy and Redox
Gibbs Free Energy: measure of the spontaneity of a process at
standard conditions
-EMF (electromotive force) also indicates spontaneity
ΔG = -nFE
n = moles of electrons transferred in balanced equation
F = Faraday’s constant = 96485C/mol or 96485 J/Vmol
Units of ΔG: J/mol
-means per mole of reaction
Free Energy and Redox
does it make sense?
ΔG = -nFE
What sign of ΔG indicates a spontaneous reaction?
What about E?
Free Energy and Redox
Calculate ΔG and K at 298K for:
4Ag(s) + O2(g) + 4H+(aq) → 4Ag+(aq) + 2H2O
Free Energy and Redox
Would the answer be different if the reaction was written:
2Ag(s) + ½ O2(g) + 2H+(aq) → 2Ag+(aq) + H2O(l)
Nonstandard Conditions
Use LeChatelier’s Principle to predict the direction the reaction
will shift
(there is something called the Nernst equation, but that has
been removed from the AP curriculum - you will probably come
across it at some point)
Cell Potential and Concentration
Standard conditions: 1M concentration
Given: 2Al(s) + 3Mn2+(aq)
2Al3+(aq) + 3Mn(s); E°cell = 0.48V
Change concentrations:
1.[Al3+] = 2.0M and [Mn2+] = 1.0M
2. [Al3+] = 1.0M and [Mn2+] = 3.0M
Electrolysis
non spontaneous redox reaction
-requires an outside source of
electrical energy
electrolytic cell - 2 electrodes in
molten salt or a solution
-the electrodes are inert (don’t
react, just allow the reaction to
occur
Electrolysis
Electroplating
-the cathode is active
-metal deposits on the cathode
Ni2+ + 2e- → Ni(s)
-2 moles of e- needed to plate 1
mole of Ni from Ni2+
Electrolysis
1 mole of electrons has a charge of 96,485C (C for coulombs)
coulomb is the amount of charge
Current is the charge per second = ampere
I = q/t
I = current (ampere)
q = charge (coulombs)
t = time (seconds)
Electrolysis
Calculate the number of grams of aluminum produced in 1.00h
by the electrolysis of molten AlCl3 if the electrical current is
10.0A
Electrolysis
The half-reaction for formation of magnesium metal upon electrolysis of molten
MgCl2 is Mg2+ + 2e- → Mg.
1.Calculate the mass of magnesium formed upon passage of a current of 60.0A for a
period of 4.00x103s.
b. How many seconds would be required to produce 50.0g of Mg from MgCl2 if the
current is 100.0A?