chapter 17 electrochemistry redox review (4.9) 17.1-17.2 17.4-17.5 17.6-17.7
TRANSCRIPT
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Chapter 17Electrochemistry
Redox review (4.9) 17.1-17.2 17.4-17.5 17.6-17.7
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Review Oxidation-Reduction
Involves transfer of electrons from reducing agent to oxidizing agent
Oxidation= loss of e- (increase in oxid #)
Reduction= gain of e- (decrease in oxid#)
GER and LEO
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REVIEW
1. atom in element = 02. monatomic ion = charge3. fluorine = -14. oxygen = -25. hydrogen = +1
6. sum of oxid. # in compound = 0
7. sum of oxid. # in polyatomic ion = charge on ion
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4–4
The Half-Reaction Method (Acidic Solution)
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Review- Balancing Oxidation-Reduction Reactions
1. Separate in ½ reactions2. Intermediate steps
a. balance all elements other than H and Ob. balance O with H2O
c. balance H with H+
d. balance charge with (e-)3. Multiply ½ rxn. so that the number of electrons
is same4. Add ½ rxns.
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Capture the Energy
MnO4- + 5Fe2+ Mn2+ 5Fe3+
MnO4- and Fe2+ will react directly in solution.
Electrons will be transferred and energy will be released as heat.
No useful work will result.
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Capture the Energy! Zn + Cu2+ - Zn2+ + Cu
Separate ½ reactions Connect metals w/ wire (to transfer
electrons)
Connect soln w/ bridge (keeps solns separate but allows ions to move)
Converts Chemical Energy to Electrical Energy!!- A Battery!!
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Galvanic Cell
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Capture the energy
You have separated the oxidizing agent from the reducing agent
Requires electron transfer through wire
Attach a motor, light bulb, bell etc-the current produced in the wire by e- flow provides work!!
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17–10
Figure 17.6 A Galvanic Cell involving the Half-Reactions
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Cell potential is…..
The pressure of a Galvanic cell to “push” the e- “driving force”
Electromotive Force, emf Symbol E Units: Joule/coulomb (=1Volt, V) Coulomb = unit of charge
Specifies # of e-
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E cell = E anode +E cathode
(oxidation) (reduction) pushing e- pulling e-
(black wire) (red wire)
A spontaneous rxn in a Galvanic cell must be positive.
E > 0
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E 1/2 reactions
P. 796 table Standard Reduction potentials
1M solutions 1atm gases 25 C
Hydrogen ½ rxn = 0.00V
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Table 17.1 Standard Reduction Potentials at 25°C (298K) for Many Common Half-Reactions
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17–14
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Helpful Info
Need balanced oxidation-reduction rxns from the reduction potentials.
One reduction ½ rxn must be reversed.* The ½ rxn with largest positive potential will
run as written (reduction).The other ½ rxn will run in reverse (oxidation).
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Reversing Direction Changes Sign of E
Because:E oxidation = -E reduction
Then: E cell = E cathode – E anode
Examples:
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Standard Reduction Potential Math Rules
# of e- lost must equal # e- gained
½ rxns must be multiplied by integers to balance equations
Value of E is not changed when ½ rxn multiplied by an integer.
Potential is NOT multiplied by integer.Example….
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Line Notation
Anode listed on leftCathode listed on right
Mg(s) l Mg2+ ll Al3+l Al(s)
Anode Mg0(s) - Mg2+
Cathode Al3+ - Al0(s)
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Cell Potential & Free Energy
A galvanic cell will run in the direction that gives a positive value for E
+E corresponds to - G
+E and - G indicates a spontaneous reaction.
G = -n FE
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G = nFE
n = # of e- (exchanged in overall rxn)
F = 96,485(c/mol e-) (Faraday’s constant)
Examples:
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Effects of Concentration on E
So far the cells have been under standard conditions….
Le Chatelier’s principle applies if not std. conditions..
Determine if E cell > or < E cell ??
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To summarize:
If E cell not at standard conditions:
[Reactant] > 1mol/L E cell > E *cell
[Product] < 1mol/L E cell> E *cell
Reverse is also true
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Concentration cell
Same components in cells, but different concentrations.
Equilibrium wants these concentrations to be Equal.
Examples:
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Nernst Equation
Establishes relationship b/t cell potential and concentration of cell components.
For cells not at 1M Concentration:E = E * - RT/nF ln (Q)
E * is std cell potentialRT/nF ln (Q) is correction factor
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Common form:
E = E * - RT/nF ln (Q)
Commonly written :
E = E * - 0.0591/n log (Q)
Examples:
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A Battery @ Equilibrium
At Equilibrium:Ecell = 0 (completely discharged)
Q = K and delta G = 0
Using the Nernst Equation:@Equilibrium: 0 =E * - 0.0591/n log(K)Or log K = nE */0.0591
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Corrosion Process of returning metals to their natural state.
Metals oxidize readily resulting in corrosion.
Metal ½ rxn is reversed for oxidation. Combined with Oxygen ½ rxn. to give (+) Ecell
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Electrolysis
Involves forcing current through a cell to produce a chemical change resulting in (-) cell potential.
Example:
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17–29
Figure 17.19 a-b (a) A Standard Galvanic Cell Based on the Spontaneous Reaction Zn + Cu2+ - Zn2+ + Cu (b) A Standard Electrolytic Cell. A Power Source Forces the Opposite Reaction Cu + Zn2+ - Cu2+ + Zn.