Lecture 11: Electrochemistry Introduction Reading: Zumdahl 4.10, 4.11, 11.1 Outline –General...
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Transcript of Lecture 11: Electrochemistry Introduction Reading: Zumdahl 4.10, 4.11, 11.1 Outline –General...
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Lecture 11: Electrochemistry
Introduction• Reading: Zumdahl 4.10, 4.11, 11.1
• Outline– General Nomenclature– Balancing Redox Reactions (1/2 cell method)
– Electrochemical Cells
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Nomenclature• “Redox” Chemistry: Reduction and Oxidation
• Oxidation: Loss of electrons
• Reduction: Gain of electrons (a reduction in oxidation number)
LEO goes GER
Loss of Electrons is Oxidation
Gain of Electrons is Reduction
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Redox Reaction Example
• A redox reaction:
Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)
Ox. #: 0 +1 +2 0
oxidation reduction
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Redox Reaction Example (cont.)
• We can envision breaking up the full redox reaction into two 1/2 reactions:
Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)
oxidation reduction
Cu(s) Cu2+(aq) + 2e-
Ag+(aq) + e- Ag(s)
“half-reactions”
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Redox Reaction Example (cont.)
• Note that the 1/2 reactions are combined to make a full reaction:
Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)
oxidation reduction
• The important thing to remember: electrons are neither created or destroyed during a redox reaction. They are transferred from the species being oxidized to that being reduced.
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Example• Identify the species being oxidized and reduced in the following (unbalanced) reactions:
ClO3- + I- I2 + Cl-
+5 -1 0 -1
oxidation
reduction
NO3- + Sb Sb4O6 + NO
+5 0 +3 +2oxidationreduction
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Balancing Redox Reactions
• One method of half reactions
• Key idea….make sure e- are neither created or destroyed in final reaction.
• Let’s balance the following reaction:
CuS(s) + NO3-(aq) Cu2+(aq) + SO4
2-(aq) + NO(g)
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Balancing (cont.)
CuS(s) + NO3-(aq) Cu2+(aq) + SO4
2-(aq) + NO(g)
• Step 1: Identify and write down the unbalanced 1/2 reactions.
-2 +5 +6 +2
oxidation reduction
CuS(s) Cu2+(aq) + SO42-(aq)oxidation
NO3-(aq) NO(g)reduction
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Balancing (cont.)• Step 2: Balance atoms and charges in each 1/2
reactions. Use H2O to balance O, and H+
to balance H (assume acidic media). Use e- to balance charge.
CuS(s) Cu2+(aq) + SO42-(aq)
NO3-(aq) NO(g)
4H2O + + 8H+ + 8e-
+ 2H2O4H+ +3e- +
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Balancing (cont.)• Step 3: Multiply each 1/2 reaction by an integer
such that the number of electronic cancels.
CuS(s) Cu2+(aq) + SO42-(aq)
NO3-(aq) NO(g)
4H2O + + 8H+ + 8e-
+ 2H2O4H+ +3e- +
x 3
x 8
• Step 4: Add and cancel.
3CuS + 8 NO3- + 8H+ 8NO + 3Cu2+ + 3SO4
2- + 4H2O
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Balancing (end)• Step 5. For reactions that occur in basic solution, proceed as above. At the end, add OH- to both sides for every H+ present, combining to yield water on the H+ side.
3CuS + 8 NO3- + 8H+ 8NO + 3Cu2+ + 3SO4
2- + 4H2O
+ 8OH- + 8OH-
3CuS + 8 NO3- + 8H2O 8NO + 3Cu2+ + 3SO4
2- + 4H2O + 8OH-4
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ExampleBalance the following redox reaction occurring in basic media:
BH4- + ClO3
- H2BO3- + Cl-
BH4- H2BO3
-
-5 +3
oxidation
ClO3- Cl-
+5 -1reduction
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Example (cont)
BH4- H2BO3
-
ClO3- Cl-
3H2O + + 8H++ 8e-
+ 3H2O6H+ +6e- +
x 3
x 4
3
3 BH4- + 4 ClO3
- 4 Cl- + 3H2BO3
- + 3H2O • Done!
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Galvanic Cells• In redox reaction, electrons are transferred from the oxidized species to the reduced species.
• Imagine separating the two 1/2 cells physically, then providing a conduit through which the electrons travel from one cell to the other.
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Galvanic Cells (cont.)
8H+ + MnO4- + 5e- Mn2+ + 4H2O
Fe2+ Fe3+ + e- x 5
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Galvanic Cells (cont.)• In turns out that we still will not get electron flow in the example cell. This is because charge build-up results in truncation of the electron flow.
• We need to “complete the circuit” by allowing positive ions to flow as well.
• We do this using a “salt bridge” which will allow charge neutrality in each cell to be maintained.
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Galvanic Cells (cont.)
Salt bridge/porous disk: allows for ion migration such that the solutions will remain neutral.
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Galvanic Cells (cont.)
• Galvanic Cell: Electrochemical cell in which chemical reactions are used to create spontaneous current (electron) flow
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Salt bridge
Zn2+ Cu2+
Na+Zn CuSO4
2–
Voltmeter
(–) (+)
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Oxidation half-reactionZn(s)
Salt bridge
Zn2+ Cu2+
Na+Zn CuSO4
2–
Zn2+(aq) + 2e–
Voltmetere–
Anode(–) (+)
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Zn2+Zn
Oxidation half-reactionZn(s)
Salt bridge
Zn2+ Cu2+
Na+Zn CuSO4
2–
Zn2+(aq) + 2e–
Voltmetere–
2e– lostper Zn atomoxidized
Anode(–) (+)
e–
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Zn2+Zn
Oxidation half-reaction
Reduction half-reaction Cu2+(aq) + 2e–
Zn(s)
Salt bridge
Zn2+ Cu2+
Na+Zn CuSO4
2–
Zn2+(aq) + 2e–
Cu(s)
Voltmetere– e–
2e– lostper Zn atomoxidized
Anode(–)
Cathode (+)
e–
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Cu2+ e–Cu
2e– gainedper Cu2+ ionreduced
Zn2+Zn
Oxidation half-reaction
Reduction half-reaction Cu2+(aq) + 2e–
Zn(s)
Salt bridge Anode(–)
Cathode (+)
Zn2+ Cu2+
Na+Zn CuSO4
2–
Zn2+(aq) + 2e–
Cu(s)
Voltmetere– e–
2e– lostper Zn atomoxidized
e–
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Cu2+ e–Cu
2e– gainedper Cu2+ ionreduced
Zn2+Zn
Oxidation half-reaction
Reduction half-reaction
Overall (cell) reactionZn(s) + Cu2+(aq)
Cu2+(aq) + 2e–
Zn(s)
Salt bridge
Zn2+ Cu2+
Na+Zn CuSO4
2–
Zn2+(aq) + 2e–
Cu(s)
Zn2+(aq) + Cu(s)
Voltmetere– e–
Anode(–)
Cathode (+)
2e– lostper Zn atomoxidized
e–