Chapter 16
Electrochemistry
Mr. von Werder
Supplemental Notes
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Review some ideas that help us understand electrochemical cells:
 oxidation and reduction is the result of electron transfer
 electrical energy is a flow of electrons (current)
 metals are good conductors (because of the loosely held valence electrons)
 electrolytes can also conduct a current (because of mobile ions)
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Electrochemical cell: redox chemical system that either produces or uses electrical energy
uses a spontaneous flow of electrons to our advantage
voltaic cells
 use spontaneous redox reactions
to generate electricity
 water downhill analogy
 electrodes:
- anode, oxidation occurs
+ cathode, reduction occurs
 example: batteries of various types
fuel cells
electrolytic cells
 use electricity (from external source) to
drive a nonspontaneous redox reaction
 water uphill analogy
 electrodes:
+ anode, oxidation occurs
- cathode, reduction occurs
 examples: electrolysis
electroplating
Electrochemical Cells
 Voltaic Cell Setup
Intro: Metals can replace metal ions in solution that are below them on the activity series.
 This happens by electron transfer - more active metal to the less active metal.
 This transfer will happen directly if the metal is put directly in the solution
Example:
Zn(s) + CuSO4(aq) ---> ZnSO4(aq) + Cu(s)
 However, if the oxidation and reduction processes are separated the flow of electrons
can be made to go through an external wire - then useful work can be made of the
electron flow (electricity)
 Voltaic cell
- spontaneous redox reaction that produces electricity
- harness the electricity by rerouting electrons through an external wire,
to light a bulb, play a radio, etc.
- this is the basis for what we commonly call “batteries”
Components:
OXIDATION HALF-CELL
ANODE
- the electrode where oxidation occurs
- has a negative sign (-) in a voltaic cell (here the electrons are
released and move through a wire to the cathode)
- often is the metal that will be oxidized (part of the reaction)
- sometimes is just a necessary conductor (graphite or platinum
electrode), but is not part of the reaction
SOLUTION - contains the cations of the element being oxidized (necessary?)
- also anions are present for neutrality
REDUCTION HALF-CELL
CATHODE + electrode where reduction occurs
+ has a positive sign (+) in a voltaic cell (here the electrons are
accepted through a wire from the anode)
+ often is the metal that will be reduced (part of the reaction)
+ sometimes is just a necessary conductor (graphite or platinum
electrode), but is not part of the reaction
SOLUTION - contains the cations of the element being reduced
- also anions are present for neutrality
SALT BRIDGE
- connects the two half-cells
- contains an electrolyte solution (ions that will not interfere)
- allows for ions to flow between cells
- (completes the circuit; maintains neutrality in each half-cell)
- porous barrier between half-cells can serve the same purpose
(porous disk(p.684), or a porous cup (demo))
WIRE
- to connect the anode to the cathode
- allows for the transfer of electrons to occur indirectly
 Cell Potential
Electrons will spontaneously flow from higher to lower potential energy position.
 The difference in potential energy is released as electrical energy as the electrons move
and can be made to do work (light a bulb, play a radio, run a motor).
 Cell potential - difference in electrical potential energy between the two electrodes
- measured in volts (named after who?) (sometimes called cell voltage)
- measured with a voltmeter (or DC setting on a multimeter)
- DEPENDS on
● substances reacting ● temperature
● concentrations
● pressures of gases (if involved)
 Standard Cell Potential
 In order to compare different cells, standard conditions have been established.
● temperature of 25°C
● concentrations of 1 M
● pressures of 1 atm
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 Standard Electrode Potentials
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Standard cell potential (Ecell) is the sum of the potentials of the two half-cells
Eox is the oxidation potential
Ecell = Eox + Ered
Ered is the reduction potential
Eox and Ered cannot be measured directly, so they are measured against a standard
and then tabulated for reference
Standard Hydrogen Electrode (SHE) is the reference
reduction of H+(aq) to H2(g) is assigned a potential of 0.0 V
Table of Standard Reduction Potentials – Supplemental HANDOUT
- more positive the Ered (from table), the more likely to be reduced
- more negative the Ered (from table), the more likely to be oxidized
- calculating an overall cell potential (Ecell)
look up the two half-reactions involved on the table
flip the one that is more negative and its E to represent oxidation
Common Batteries
 The Common Dry Cell
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“Dry” because the electrolyte is in the form of a paste (rather than aqueous solution)
See Figure 21-13, page 698.
Anode:
Zinc casing:
zinc is oxidized
Zn ---> Zn2+ + 2eCathode:
Graphite electrode; paste consisting of MnO2, and other things
Mn (from MnO2) is reduced
Pros: low cost
Cons: not rechargeable, high current draw discharges cell quickly, poor shelf life
 The Alkaline Dry Cell
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“Alkaline” because the electrolyte paste contains KOH (a base – alkaline)
See Figure 21-13, page 698.
Anode:
Paste consisting of powdered Zn, KOH, water: zinc is oxidized
Cathode:
Paste consisting of MnO2, and other things: Mn (from MnO2) is reduced
Pros: longer shelf life, maintains a steady voltage, more energy than dry cell
Cons: more expensive than dry cell, also not rechargeable
 Lead Storage Battery
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commonly found in your car
Sulfuric acid is the electrolyte, so it is known as “battery acid”
Anode: Pb plates; Pb is oxidized (0 to +2)
Cathode: PbO2 plates; Pb is reduced (+4 to +2)
Pb + SO42- --> PbSO4 + 2ePbO2 + 4H+ + SO42- + 2e- --> PbSO4 + 2H2O
Pb + PbO2 + 4H+ + 2SO42- --> PbSO4 + 2H2O
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Eox = +0.35
Ered = +1.69
Ecell = +2.04
How do you get a 12-V car battery? Arrange six of these cells together --- a “battery”
Alternator in your car reverses the electron flow in the battery and recharges it.
Pros: rechargeable, relatively cheap
Cons: massive, disposal concerns will all of the lead (a heavy metal contaminant)
 Nickel-Cadmium battery
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Rechargeable batteries (nicad for nickel-cadmium) used in many ‘cordless’ devices.
Anode:
Cd metal is oxidized
Cathode:
Nickel (from NiO2) is reduced
Pros: lightweight, constant voltage, rechargeable
Cons: ‘discharge memory’, disposal concerns (cadmium is toxic)
 Fuel Cells
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Most common fuel cell is the hydrogen/oxygen fuel cell
Hydrogen and oxygen must be continually supplied to the cell
Anode:
Hydrogen is oxidized
2 H2 + 4 OH- --> 4 H2O + 4eCathode:
Oxygen is reduced
O2 + 2 H2O + 4e- --> 4 OH-
Eox = +0.83
Ered = +0.40
2 H2 + O2 --> 2 H2O
Ecell = +1.23
Biggest selling point is the efficiency – a reaction that may be combusted to produce heat
to heat water to spin a turbine to generate electricity … can directly create electricity.
Pros: don’t run down (as long as fuel is provided), very efficient, drinking water
Cons: expensive, not very portable
Electrolytic Cells
Electrolytic cells use electricity to cause a nonspontaneous redox reaction to occur
Components, See
source of direct current (for example, a battery - voltaic cell)
electrodes
anode;
where oxidation occurs, but is in the electrolytic cell
cathode;
where reduction occurs, but is
- often the electrodes are inert, and are not part of the reaction
- electrons are driven from anode to cathode by the battery (an electron
pump)
solution (or liquid) containing the reactants
 Electrolysis of Molten Sodium Chloride
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First, NaCl needs to be melted, so it be heated to 801C (Why so high a m.p.?)
Anode:
chlorine is oxidized
2 Cl-(l) --> Cl2(g) + 2eEox = -1.36
+
Cathode:
sodium is reduced
2 Na (l) + 2e --> 2 Na(l)
Ered = -2.71
2 NaCl(l) --> 2 Na(l) + Cl2(g)
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Ecell = -4.07
Notice the negative Ecell -- a nonspontaneous reaction, requiring electricity to occur
Figure 21-18, page 706, Down’s cell
Practical way of getting metallic sodium
 Electrolysis of Water
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First, an electrolyte needs to be added to the water -> to provide more ions -> to allow
more current to flow.
Anode:
O2 is formed
2 H2O --> 4 H+ + O2 + 4eCathode:
H2 is formed
4 H2O + 4e- --> 2 H2 + 4 OH2 H2O --> 2 H2 + O2
NOT a practical way to obtain quantities of oxygen (easier to liquefy and distill air)
- but, the reduction to form hydrogen in aqueous solutions is commonly used
 Electrolysis of Aqeous Sodium Chloride
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Anode:
Cathode:
Cl2 is formed
H2 is formed
2 Cl- --> Cl2 + 2e2 H2O + 2e- --> H2 + 2 OH2 Cl- + 2 H2O --> Cl2 + H2 + 2 OHCommercially this is an important reaction: it produces three valuable products:
Chlorine gas, hydrogen gas, and sodium hydroxide solution
 Electroplating
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Electroplating: uses electrolysis to deposit a thin coating of metal on an object
Uses: - to protect another metal (chrome plating, tin cans)
- decorative (silver and gold plating)
Anode: plating metal, to be oxidized
Cu --> Cu2+ + 2eCathode: object to be plated, reduction occurs here
Cu2+ + 2e- --> Cu(s)