Chapt. 20 – Electrochemistry 19.1Oxidation and Reduction (review) 19.2Balancing Redox Equations...
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Transcript of Chapt. 20 – Electrochemistry 19.1Oxidation and Reduction (review) 19.2Balancing Redox Equations...
![Page 1: Chapt. 20 – Electrochemistry 19.1Oxidation and Reduction (review) 19.2Balancing Redox Equations (intro only on half reactions) 20.1Voltaic Cells 20.2.](https://reader030.fdocuments.in/reader030/viewer/2022032600/56649da05503460f94a8c0e1/html5/thumbnails/1.jpg)
Chapt. 20 – Electrochemistry
19.1 Oxidation and Reduction (review)19.2 Balancing Redox Equations (intro
only on half reactions)20.1 Voltaic Cells20.2 Batteries20.3 Electrolysis
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Section 19.1 Oxidation and Reduction
• Describe the processes of oxidation and reduction.
• Identify oxidizing and reducing agents.
• Interpret redox reactions in terms of change in oxidation state.
Oxidation and reduction are complementary—as an atom is oxidized, another atom is reduced.
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Key Concepts
• Oxidation-reduction reactions involve the transfer of electrons from one atom to another.
• When an atom or ion is reduced, its oxidation number is lowered. When an atom or ion is oxidized, its oxidation number is raised.
Section 19.1 Oxidation and Reduction
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Section 19.2 Balancing Redox Reactions
• Relate changes in oxidation number to the transfer of electrons.
Redox equations are balanced when the total increase in oxidation numbers equals the total decrease in oxidation numbers of the atoms involved in the reaction.
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Key Concepts
• The oxidation-number method is based on the number of electrons transferred from atoms equaling the number of electrons accepted by other atoms.
• A half-reaction is one of the two parts of a redox reaction.
Section 19.2 Balancing Redox Reactions
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Section 20.1 Voltaic Cells
• Describe a way to obtain electrical energy from a redox reaction.
• Identify the parts of a voltaic cell, and explain how each part operates.
• Calculate cell potentials, and determine the spontaneity of redox reactions.
In voltaic cells, oxidation takes place at the anode, yielding electrons that flow to the cathode, where reduction occurs.
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Key Concepts
• In a voltaic cell, oxidation and reduction take place at electrodes separated from each other.
• The standard potential of a half-cell reaction is its voltage when paired with a standard hydrogen electrode under standard conditions.
• The reduction potential of a half-cell is negative if it undergoes oxidation when connected to a standard hydrogen electrode. The reduction potential of a half-cell is positive if it undergoes reduction when connected to a standard hydrogen electrode.
• The standard potential of a voltaic cell is the difference between the standard reduction potentials of the half-cell reactions.
Section 20.1 Voltaic Cells
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Redox Reactions
Practical /everyday examples:Corrosion of iron (rust formation)
Forest fire
Charcoal grill
Natural gas burning
Batteries
Production of Al metal from Al2O3 (alumina)
Metabolic processes
combustion
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Redox Reactions
LEO
LEO says GER
LEO:• Lose Electrons
Oxidation
GER:• Gain Electrons
Reduction
GER
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Redox Reactions
Single replacement – zinc in acid
Zn(s) + 2 H+(aq) → Zn2+ (aq) + H2(g)
0 +2 0+1
Zn (s) oxidized
H+(aq) reduced
Zn = reducing agent
H+(aq) = oxidizing agent
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Summary of Terminology for
Reduction-Oxidation (Redox)
Reactions
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Half-ReactionsSpecies – any kind of chemical unit involved in a chemical reaction
NH3(g) + H2O(l) NH4+(aq) + OH-(aq)
4 species in above: 2 molecules & 2 ions
Redox reaction occurs when species that can give up electrons comes in contact with species that can accept them
Half-reaction: one of 2 parts of redox equation (oxidation or reduction half)
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Half-ReactionsRedox equation
2Fe(s) + 3Cl2(g) 2FeCl3(s)
Half-reactions:
Oxidation: Fe Fe3+ + 3e-
Reduction: Cl2 + 2e- 2Cl-
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Redox Reaction Example
On right—Zn metal is dipped in Cu2+ solution (blue)
After a bit, blue solution is lighter, and metal strip is covered with Cu metal
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Redox Reaction Example
Viewed as overall single replacement reaction (net ionic):
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Viewed as separate half reactions:
Zn(s) → Zn2+(aq) + 2e- (oxidation)
Cu2+(aq) + 2e- → Cu(s) (reduction)
Moving electrons from one species to another
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Reductionist View of Redox Reactions
Movement of electrons from a source to a recipient
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Half Reactions – Oxidation of Iron
Overall Rxn (unbal.) Reduction Half RxnFe + O2 → Fe2O3 O2 + 4e– → 2O2 –
Fe + Cl2 → FeCl3 Cl2 + 2e– → 2Cl–
Fe + HBr → H2 + FeBr3 2H+ + 4e– → H2
Fe + AgNO3 → Ag + Fe(NO3)3
Ag + e– → Ag
Fe + CuSO4 → Cu + Fe2(SO4)3
Cu2+ + 2e– → Cu
Oxidation half reaction: Fe → Fe3+ + 3e–
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Simple electrochemical cells are referred to as galvanic cells (after Galvani) or voltaic cells (after Volta)
Our textbook uses term voltaic
Terminology
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Voltaic Cells
What happens if Zn metal immersed in CuSO4(aq) solution?
Spontaneous* redox reaction occurs:
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
* doesn’t need to be driven by outside source of energy (DG for process < 0)
Zn(s) oxidized and Cu2+(aq) reduced
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Half Reactions → Half Cells
For this example, redox half reactions occur in same place (species in direct contact)
Can separate them as half cells by constructing electrochemical cell (indirect)
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
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Electrochemical Cell
An apparatus that uses a redox reaction to produce electrical energy or uses electrical energy to cause a chemical reaction
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Voltaic (Galvanic) Cells
Device in which spontaneous redox reaction occurs as electrons are transferred from reductant to oxidant through an external circuit
Used to perform electrical work using energy released during spontaneous redox reaction
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Voltaic Cell: Half-Cells
Two half reactions occur in separate compartments called half-cells• 1 half-cell contains oxidation half reaction• 1 half-cell contains reduction half reaction
Each half cell contains:• metal electrode• electrolyte solution
of ion of electrode
1M Zn2+
1M Cu2+
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Voltaic Cell: Half-Cells
Connect half-cells
Redox happens - electrons can now transfer from 1 electrode to another
Zn Cu
Zn+2 Cu+2e-
e-
e-
e-
e-
e-
e- e-
e-
Process stops almost immediately – buildup of excess negative charge on Cu electrode prevents further transfer
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Electron flow
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Ion flow keeps the charge neutral
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Salt bridge contains strong electrolyte
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KCl Salt BridgeCl- K+
Voltaic Cell: Complete CellAdd salt bridge – allows ion transport
Bridge has soluble salt (e.g., KCl) in agar gel (like jello) with ion permeable plugs
Ion movement neutralizes charge created by electron movement
e-e-
e-
e-
e-
e-
K+Cl-
Cl- K+
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Porous Disk Also Allows Ion Flow
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Voltaic Cell Using Porous Barrier
Daniell cell: Cu & Zn electrodes dipping into solutions of copper(II) sulfate and zinc sulfate, respectively
Solutions make contact through porous pot, which allows ions to pass through to complete circuit
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Anodes and Cathodes
Anode
Electrode at which oxidation occurs• located in oxidation half-cell• the “negative” electrode • electrons are released here• anions move toward anode
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Anodes and Cathodes
Cathode
Electrode at which reduction occurs• located in reduction half-cell• the “positive” electrode • electrons move toward here• cations move toward anode
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General Voltaic Cell
Anode half-cell
Cathode half-cell
Oxidation occurs here
Reduction occurs here
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Voltaic Cell (with nitrates)Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Oxidation half cell Reduction
half cell
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Anodes and Cathodes
For following reaction:
Zn(s) + Ni2+(aq) Zn2+(aq) + Ni(s)• Which metal will be the anode?
• Which metal will be the cathode?
Oxidation occurs at anode: Zn (s)
Reduction occurs at cathode: Ni (s)
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Spontaneity and Potential Energy
Redox reactions occurring in voltaic cell are spontaneous
Why do electrons flow spontaneously from one electrode to other?
• Flow spontaneously due to difference in potential energy between anode and cathode
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Current, Voltage (Cell Potential), and Electrical Potential Energy
Electric potential energy: measure of amount of current (flow of charge) that can be generated from voltaic cell to do work
Current can flow between two points only when a difference in electric potential energy exists between the two points
Volt: unit used to measure cell potential—driving force from difference in electric potential energy between two electrodes
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Electron Flow and Potential Energy
Anode • higher
potential energy
Cathode• lower
potential energy
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High Voltage
Low Voltage
Height is an analogy for voltage
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Current = charge passing/unit time
Low Amps High Amps
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Voltaic Cells – Cell Potential
Difference in electrical potential between anode and cathode called electromotive force (emf)
• Also known as cell potential or cell voltage
• measured in volts (V)• Indication of energy available to move
electrons from anode to cathode
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Voltaic Cells – Standard emf
Standard cell potential (emf) (Eocell)
Cell potential measured under standard conditions
• 25oC • 1M concentrations of reactants and
products in solution • 1 atm pressure for gases
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Voltaic Cells - Eocell
Eocell depends on half-cells or half-
reactions present
Standard potentials have been assigned to each individual half-cell – cannot be directly determined
By convention, standard reduction potential (Eo
red) for each half cell is used
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Voltaic Cells - Eocell
Must measure reduction potential of half cell against something – a reference
Use standard hydrogen electrode
Platinum electrode in 1 M HCl, 25oC, 1 atm H2(g)
Differs from previous half cells – metal not involved in redox process
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Standard Hydrogen Electrode
Redox process for SHE:
2H+ (aq, 1M) + 2 e- H2 (g, 1 atm)
Reduction potential (at 25oC) of this half-cell assigned a potential of exactly 0.000 V
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A Hydrogen Electrode= SHE if [H+] = 1 M (from HCl),
p(H2(g)) = 1 atm,
T = 25oC
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Cell Potentials – Cu vs SHE
e
e
1 M Cu2+ 1 M H+
Cu2+H+
e- flow
Cu = cathode (reduction)
Eocell =
+0.342 V
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Cell Potentials – Zn vs SHE
Zn = anode (oxidation)
Eocell =
0.762 V
Cu2+H+
e- flow e- flow
1 M Zn2+ 1 M H+
Zn2+
H+
e
e
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Eored Values
Eored tabulated table 25.1, p 712
Ordered from most negative (Li+ - strong reducing agent) to most positive (F2 – strong oxidizing agent)
Includes many reactions that don’t involve metals
O2 + 2H+ +2e- H2O2
For these half-cells, unreactive metal used as electrode (e.g., platinum)
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Eored Values (table 20.1)
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Table of Standard Reduction PotentialsTable of Standard Reduction Potentials
Eo for reaction as written
The more positive Eo, the greater tendency for substance to be reduced
Sign of Eo changes when reaction reversed
Changing stoichiometric coefficients of half-cell reaction does not change value of Eo
Eo (V)
Cu2+ + 2e- Cu +0.34
2 H+ + 2e- H2 0.00
Zn2+ + 2e- Zn -0.76
Greater reducing tendency
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Sign of Eored and Spontaneity
As Eored becomes increasingly positive,
driving force for reduction increases
F2(g) + 2e- 2 F-(aq) Eored = +2.87 V
Ag+(aq) + e- Ag(s) Eored = +0.80 V
Fluorine (most EN element) wants to be reduced more than silver ion
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Sign of Eored and Spontaneity
As Eored becomes increasingly negative,
driving force for oxidation increases
Li+ (aq) + e- Li (s) Eored = -3.05 V
Negative reduction potential indicates that oxidation half-reaction is spontaneous
Li (s) Li+ (aq) + e-
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Standard Reduction PotentialsGiven following potentials, which metals will be most easily oxidized?
Ag+(aq) + e- Ag(s) Eored = +0.80 V
Zn2+(aq) + 2 e- Zn(s) Eored = - 0.76 V
Na+(aq) + e- Na(s) Eored = - 2.71 V
Na(s) most easily oxidized
Potentials are quantitative version of “metal activity series” previously used
?
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Cell Potential – Anodes & Cathodes
Make voltaic cell by combining two half-cells and calculate standard cell potential, Eo
cell
Cell must have both anodic process (anode) and cathodic process (cathode)
Look at standard reduction potential, Eored
for each half-cell: cell with more positive value occurs as written (as a reduction = cathode)
Other half-cell will occur in reverse (as an oxidation = anode)
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Standard cell potential, Eocell
Once anode and cathode have been identified, calculate Eo
cell from
Eocell = Eo
red (cathode) – Eored (anode)
reduction oxidation
Book uses equation
Eocell = Eo
reduction – Eooxidation
Misleading - second term is still really a reduction potential (from table)
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Example: Eocell for Cu/Zn
Step 1 – Identify anode and cathode Cu2+(aq) + 2 e- Cu (s) Eo
red = +0.342 V
Zn2+(aq) + 2 e- Zn(s) Eored = - 0.762 V
Potential for Cu more positive
Cu will be cathode Zn will be anode
Step 2 – Calculate standard cell potential
Eocell = Eo
red (cathode) – Eored (anode)
Eocell = 0.342 V - (-0.762 V)
Eocell = 1.104 V
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Overall Cell Potential – Cu/Zn Cell
E0cell = +1.104 V
E0Cu = +0.342 V
E0SHE = 0.000 V
E0Zn = 0.762 V
anode cathode
Zn│Zn+2║Cu+2│Cu Cell notation
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Voltaic Cells – Line Notation
Anode components listed on left
Cathode components listed on right
Anode and cathode have reactants on left, products on right
Separate half cells with double vertical lines: ll
Indicate phase difference with single vertical line: l
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Line Notation & Overall Reaction
Al(s) | Al3+ (1.00 M) ║Cu2+ (1.00 M) | Cu(s)
Anode: Al(s) Al3+(aq) + 3e- Eored = -1.662 V
Cathode: Cu2+(aq) Cu(s) + 2e- Eored = 0.342 V
Note: Cell notation does not denote balance
To get balanced overall reaction, multiply anodic ×2 and cathodic ×3 to give
2Al(s) + 3Cu2+(aq) 3Cu(s) + 2Al3+(aq)
Eocell = Eo
red(cathode) – Eored(anode)
Eocell = 0.342 V – (-1.662 V) = 2.004 V
anode|anode solution║cathode solution|cathode
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Practice
Using Cell (Line) Notation
Problems 34, 39(a-d), 40(a-c) page 736
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Voltaic CellsGiven following reduction half-cells, identify anode, cell line notation, balanced reaction for cell, and Eo
cell
Al3+(aq) + 3 e- Al(s) Eored = -1.66 V
Fe2+(aq) + 2 e- Fe(s) Eored = -0.440 V
More positive half-cell = Fe cathode
So Al is anode: Al (s) Al3+ (aq) + 3 e-anode|anode solution║cathode solution|cathode
Al|Al+3║Fe2+|Fe
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Voltaic CellsBalanced equation:
Al(s) Al3+(aq) + 3 e-
Fe2+(aq) + 2 e- Fe(s)
× 2
× 3
2 Al(s) + 3 Fe2+(aq) + 6e- 2 Al3+(aq) + 3 Fe(s) + 6e-
2 Al(s) + 3 Fe2+(aq) 2 Al3+(aq) + 3 Fe(s)
Eocell = Eo
red (cathode) - Eored (anode)
Eocell = -0.440 V (-1.66 V) = 1.22 V
Reduction potentials not multiplied by anything
Anode:
Cathode:
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Practice
Calculating standard cell potentials
Problems 1 – 4, page 716
Problem 41(a-c) page 736
Problem 68 page 738
Problems 1(a-c) page 991
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Is Proposed Reaction Spontaneous?Used E0
red data to get E0cell; process
guaranteed to give spontaneous overall reaction (E0
cell always > 0)
Can reverse process to determine if a given redox reaction is spontaneous as written
Step 1 – Write reaction in form of half-cells
Step 2 – Find E0red for each
Step 3 – Calculate E0cell
Step 4 – If E0cell > 0, reaction spontaneous
if < 0 not spont., but reverse rxn is
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Practice
Determining spontaneity of redox reaction
Problems 5 – 9, page 716
Problems 13(a-c) page 717
Problems 67(a-d), page 738
Problems 2(a-c) page 991
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Chapt. 20 – Electrochemistry
19.1 Oxidation and Reduction (review)19.2 Balancing Redox Equations (intro
only on half reactions)20.1 Voltaic Cells20.2 Batteries20.3 Electrolysis
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Section 20.2 Batteries
• Describe the structure, composition, and operation of the typical carbon-zinc dry-cell battery.
• Distinguish between primary and secondary batteries, and give two examples of each type.
• Explain the structure and operation of the hydrogen-oxygen fuel cell.
• Describe the process of corrosion of iron and methods to prevent corrosion.
Batteries are voltaic cells that use spontaneous reactions to provide energy for a variety of purposes.
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Key Concepts
• Primary batteries can be used only once; secondary batteries can be recharged.
• When a battery is recharged, electric energy supplied to the battery reverses the direction of the battery’s spontaneous reaction.
• Fuel cells are batteries in which the substance oxidized is a fuel from an external source.
• Methods of preventing corrosion are painting, coating (plating) with another metal, or using a sacrificial anode.
Section 20.2 Batteries
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Battery
A battery is a:• voltaic (galvanic) cell• or group of voltaic cells connected in
series cell potentials of individual cells add up to
give total battery cell potential• source of direct current (DC)
http://www.powerstream.com/BatteryFAQ.html
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Batteries – Primary vs Secondary
Primary – cannot be recharged because one or both half-reactions not reversible
• can explode if recharge attempted
Secondary – can be recharged by reversing flow of current, regenerating reactants
• AKA storage batteries
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Batteries – Wet Cell vs Dry Cell
Wet cell – conventional electrolyte
“Dry” Cell – typically a moist paste
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Showed that If put electrolyte–soaked cloth in between 2 different metals, (e.g., Cu & Zn), current would flow
Alessandro Volta
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Types of Primary Cell Batteries
Alkaline
Aluminium (aluminum-air)
Lithium (not Lithium-ion)
Oxyride (Oxy Nickel Hydroxide (NiOOH)
Silver-oxide
Zinc-air
Zinc-carbon
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Zinc-Carbon Dry CellWet version 1866, “dry” version 1887Anode: Zn Cathode : carbon rod (inactive) in contact with moist, acidic paste of solid MnO2, solid NH4Cl, and powdered graphiteProduces 1.5 VInexpensive, safeShort shelf life due to reaction of Zn with acidic electrolyte
Can generate NH3(g) at high current drains
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Standard Zinc-Carbon
Dry CellBattery
2NH4+(aq) + 2 MnO2(s) + Zn(s)
Zn+2(aq) + Mn2O3(s) +2 NH3(aq) + H2O(l)
Simplified overall reaction
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Zinc-Carbon: Half & Other Reactions
Zn(s) Zn2+(aq) = 2e- (anode)
2MnO2 (s) + 2NH4+(aq) + 2e- Mn2O3(s) +
2NH3(aq) + H2O (l) (cathode)
Zn2+(aq) + 2NH3 (aq) + 2Cl-(aq) Zn(NH3)2Cl2(s) (complexation – no redox)
Overall:
2MnO2 (s) + 2NH4Cl(aq) + Zn(s) Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s)
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Alkaline Battery
Anode: Zn powder (not solid)
Cathode: moist paste of solid MnO2, KOH or NaOH, and carbon (no NH4Cl as in normal carbon-zinc) with current collector (various)
Lasts longer than standard carbon-zinc: zinc doesn’t corrode as fast under basic conditions
Has less voltage drop than carbon-zinc
More expensive; prone to leaking KOH
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Alkaline Battery
Steel Case (current
collector)MnO2, cathode
mixture
Zn-KOH anode paste
Brass current collector
KOH electrolyte
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Alkaline BatteryE°Cell = 1.5 V
Published half reactions vary [book uses first anode and first cathode reactions]:
Anode: Zn(s) + 2OH-(aq) ZnO(s) + H2O(l) + 2e-
Anode: Zn(s) + 2OH-(aq) Zn(OH)2(s) + 2e-
Cathode (+4 +2): MnO2 (s) + 2H2O(l) + 2e- Mn(OH)2 (s) + 2OH-
(aq)
Cathode(+4 +3):
2MnO2 (s) + H2O(l) + 2e- Mn2O3 (s) + 2OH-(aq)
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Lithium Batteries Wide variety of types of cathodes and electrolytesAll types have lithium metal as anode [book does not properly distinguish between lithium & lithium-ion (no solid Li) batteries!]Low metal density makes lightweight batteryLow standard reduction potential gives higher output voltage (2.3 V higher than with cell utilizing Zn anode) – typical 3 V outLong lasting
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Lithium Batteries Most common: MnO2 cathode, salt of lithium (LiClO4) dissolved in organic solvent (propylene carbonate and dimethoxy ethane) as electrolyte
For medical use: Li-I2, Iodine cathode, solid electrolyte (charge transfer complex, e.g., poly-2-vinylpyridine)
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Secondary Batteries
Secondary – can be recharged by reversing flow of current, regenerating reactants
• AKA storage batteries
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Types of Secondary Cell Batteries
Lead-acid (automobile)
Lithium
Lithium-ion
Lithium-ion polymer
Nickel-cadmium (NiCad)
Nickel metal hydride (NiMH)
Molten salt
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Lead-Acid (Storage) BatteryMost common secondary battery
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Lead Acid Battery
Can function for several years over temperature range from -30oF to 120oF
12 V battery (six cells)
anode = Pb
cathode = Pb coated with PbO2
electrolyte solution = H2SO4 solution
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Lead storage battery
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Lead Acid Battery (+2.041 V)Pb(s) + SO4
2-(aq) → PbSO4(s) + 2 e- +1.685 V
PbO2(s) + SO42-(aq) + 4H+(aq) + 2 e- → PbSO4(s)
+ 2 H2O(l) +0.356 V
Pb + PbO2 + 2SO42- + 4H+ → 2 PbSO4 + 2 H2O
PbSO4 (from both electrodes) adheres
During discharge, H2SO4 consumed at both electrodes and H2O produced
Battery condition can be determined by measuring electrolyte density – drops with drop in acid concentration
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Lead Acid Battery
Can be recharged because PbSO4(s)
products adhere to electrodes - alternator can force current through battery in opposite direction and reverse reactions
Even though battery can be recharged, physical damage from road shock and chemical side reactions (e.g. electrolysis of water) eventually cause battery failure
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NiCad Battery
1.4 V
Rechargeable (like lead acid battery, products adhere to electrodes)
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Anode (Cd to Cd2+)
NiCad Cell
2 NiOOH(s) + Cd(s) 2 Ni(OH)2(s) + Cd(OH)2(s)
Overall
Cathode (Ni3+ to Ni2+)
Ni(OH)2(s) + OH–(aq) NiOOH(s) + H2O(l) + e–
Cd(s) + 2 OH–(aq) Cd(OH)2(s) + 2 e–
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Lithium-Ion Cell3.6 V
One of most popular types for portable electronics (Sony commercialized in 1991)
High energy to weight ratio (2x NiCad)
No memory effect
Low self-discharge (5% month) compared to NiCad (10%) and NiMH (30%) but subject to aging starting from time of manufacture
Solid, polymeric electrolyte (salt bridge)solid electrolyte interphase (SEI).
solid lithium-salt electrolytes (LiPF6, LiBF4, or LiClO4) and organic solvent
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Lithium-Ion CellAnode: carbon (graphite)
Cathode: metal oxide of Li with Co (most), Mn, or Ni/Co/Mn (NCM) – no metallic Li
Electrolyte: lithium salt in organic solvent
Reaction:
Li+ transported to and from cathode/anode
Transition metal, Co, in LixCoO2 oxidized from Co3+ to Co4+ during charging, and reduced from Co4+ to Co3+ during discharge
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Fuel CellVoltaic cell for which reactants are continuously supplied
Used in U.S. space program
Based on reaction of hydrogen (and other fuels such as methane) with oxygen to form water
2 H2(g) + O2(g) → 2H2O(l)
Same as combustion, but done in way that electrical energy can be extracted
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Fuel CellElectrodes:Hollow chamber of porous carbon walls
Walls of chamber contain catalysts (Pt, Pd)
Electrolyte: KOH (alkaline cell)
Or
Proton-exchange membrane (PEM) – allow H+ ions to pass through
Safer & lighter than using liquid electrolyte
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Anode: Cathode: O2 (g) + 2H2O (l) + 4e- → 4OH- (aq)
2H2 (g) + 4OH- (aq) → 4H2O (l) + 4e-
2H2 (g) + O2 (g) → 2H2O (l)
Fuel Cell with KOH Electrolyte
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Hydrogen Fuel Cell with PEM
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Corrosion
Loss of metal resulting from redox reaction of metal with substances in environment
Ordinary rusting requires presence of both oxygen and water (electrolyte)
Iron surface naturally becomes inhomogeneous and develops anodic and cathodic regions
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Corrosion Cell
and oftenFe2+ Fe3+ + e-
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Corrosion
Oxidation: Fe(s) Fe2+(aq) + 2 e–
Reduction: O2(g) + 4 H+(aq) + 4 e– 2 H2O(l)
Overall: 2 Fe(s) + O2(g) + 4 H+(aq)
2Fe2+(aq) + 2 H2O(l)
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Corrosion
Can be slow in pure water
Faster in water with dissolved salts, especially with chloride or sulfate ions
Seawater especially corrosive
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Corrosion Prevention
Organic coatings (paint) can be effective
Problems occur at breaks, pinholes
Alternative: sacrificial anodes
Metal with more negative reduction potential than iron – Mg, Zn
This material corrodes and iron is cathode only (stays protected)
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Use of Zinc Sacrificial Anode
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Sacrificial Mg Anode to Protect Steel Pipeline
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Cathodic Protection of Iron Storage Tank Using Mg Sacrificial Anode
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Corrosion Prevention - Galvanizing
Combines coating idea (barrier) with sacrificial anode idea
Inexpensive and widely used
Especially important for parts of car bodies
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Intact Zinc Coating (Barrier)
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Breached Zinc Coating (Sacrificial Anode)
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Chapt. 20 – Electrochemistry
19.1 Oxidation and Reduction (review)19.2 Balancing Redox Equations (intro
only on half reactions)20.1 Voltaic Cells20.2 Batteries20.3 Electrolysis
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Section 20.3 Electrolysis
• Describe how it is possible to reverse a spontaneous redox reaction in an electrochemical cell.
• Compare the reactions involved in the electrolysis of molten sodium chloride with those in the electrolysis of brine.
• Discuss the importance of electrolysis in the smelting and purification of metals.
In electrolysis, a power source causes nonspontaneous reactions to occur in electrochemical cells.
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Key Concepts
• In an electrolytic cell, an outside source of power causes a nonspontaneous redox reaction to occur.
• The electrolysis of molten sodium chloride yields sodium metal and chlorine gas. The electrolysis of brine yields hydrogen gas, sodium hydroxide, and chlorine gas.
• Metals such as copper are purified in an electrolytic cell.
• Electrolysis is used to electroplate objects and to produce pure aluminum from its ore.
Section 20.3 Electrolysis
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ElectrolysisUse of electrical energy to bring about a chemical reaction
Electrolysis cell: electrochemical cell in which electrolysis occurs
Reverse of normal voltaic (galvanic) cell process - apply voltage to drive reactions in reverse (same as charging secondary battery)
Not possible for some systems (primary batteries)
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Electrolytic Cell Discharge (a) and Recharge (b)
e flow e flow
Note: this particular cell is not really rechargeable due to migration of Zn2+(aq) away from electrode
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Electrolysis Applications
Water – produces H2(g), O2(g)
Down’s cell – molten NaCl to produce Na(l), Cl2(g)
Chloralkali process – electrolyze brine to generate NaOH and Cl2(g)
Aluminum production (Hall process)
Electrorefining (copper)
Electroplating
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Cell for electrolysis of waterGenerates H2(g) & O2(g) in 2:1 ratio
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Electrolysis Applications
Down’s cell – molten NaCl to produce Na(l), Cl2(g)
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Electrolysis of Molten NaCl –
Down’s Cell
Only practical way to get Na
Na+(l) Na(l) + e-
2Cl-(l) Cl2(g) + 2e-
2Na+(l) + 2Cl-(l) 2Na(l) + Cl2(g)
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Brine Electrolysis Chloralkali Production
Brine = concentrated aqueous NaCl
Overall reaction:
2H2O(l) + 2NaCl(aq) H2(g) + Cl2(g) + 2NaOH(aq)
All 3 products commercially important
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Brine Electrolysis
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Producing Aluminum2 Al2O3 + 3 C 4 Al + 3 CO2
Charles Hall (1863-1914) developed processFounded Alcoa Price of Al dropped from $100,000/lb in 1855 to $2/lb in 1890.
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Producing AluminumL.T.Héroult discovered same process at same time; called Hall-Héroult processElectrolyze Al2O3 (from bauxite) - dissolved at 1000C in molten cryolite (Na3AlF6)
Cell lined with graphite = cathode: Al3+ +3e- Al(l)Graphite rods are also anode:
2O2- O2(g) + 4e- & 2C(s) + O2(g) CO(g)
Requires huge amount of energy – Al production uses 4% of US electrical power
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Refining of Aluminum
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Refining Al – Hall-Héroult Process
Power Source
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Anodes: slabs of impure Cu
Cathodes: thin sheets of pure Cu
Electrolyte: acidic copper sulfate
Voltage across electrodes designed to produce only Cu at cathode
Metallic impurities that can oxidize to ions (e.g., Zn) do not plate out on cathode
Insoluble metal impurities (don’t oxidize) collected in sludge at bottom of cell
Copper Electrorefining
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Copper Electrorefining
Cu(s) Cu2+
(aq) + 2e-
and other metals more
easily oxidized than Cu
Cu2+(aq) + 2e- Cu(s)
Voltage in range where this only can
happen for Cu
Impure Cu Anode
Pure Cu Cathode
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Used to produce decorative and/or protective layer of metal on top of a second (usually cheaper) metal
Metal being plated is cathode
Source of metal to plate may be anode or metal ion in solution (Au,Cr)
Electroplating
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Electroplating
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Electroplating of silver
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End of Chapter