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

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.

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

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.

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

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.

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

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

Redox Reactions

LEO

LEO says GER

LEO:• Lose Electrons

Oxidation

GER:• Gain Electrons

Reduction

GER

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

Summary of Terminology for

Reduction-Oxidation (Redox)

Reactions

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)

Half-ReactionsRedox equation

2Fe(s) + 3Cl2(g) 2FeCl3(s)

Half-reactions:

Oxidation: Fe Fe3+ + 3e-

Reduction: Cl2 + 2e- 2Cl-

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

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

Reductionist View of Redox Reactions

Movement of electrons from a source to a recipient

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–

Simple electrochemical cells are referred to as galvanic cells (after Galvani) or voltaic cells (after Volta)

Our textbook uses term voltaic

Terminology

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

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)

Electrochemical Cell

An apparatus that uses a redox reaction to produce electrical energy or uses electrical energy to cause a chemical reaction

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

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+

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

Electron flow

Ion flow keeps the charge neutral

Salt bridge contains strong electrolyte

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+

Porous Disk Also Allows Ion Flow

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

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

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

General Voltaic Cell

Anode half-cell

Cathode half-cell

Oxidation occurs here

Reduction occurs here

Voltaic Cell (with nitrates)Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Oxidation half cell Reduction

half cell

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)

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

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

Electron Flow and Potential Energy

Anode • higher

potential energy

Cathode• lower

potential energy

High Voltage

Low Voltage

Height is an analogy for voltage

Current = charge passing/unit time

Low Amps High Amps

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

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

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

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

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

A Hydrogen Electrode= SHE if [H+] = 1 M (from HCl),

p(H2(g)) = 1 atm,

T = 25oC

Cell Potentials – Cu vs SHE

e

e

1 M Cu2+ 1 M H+

Cu2+H+

e- flow

Cu = cathode (reduction)

Eocell =

+0.342 V

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

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)

Eored Values (table 20.1)

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

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

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-

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

?

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)

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)

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

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

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

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

Practice

Using Cell (Line) Notation

Problems 34, 39(a-d), 40(a-c) page 736

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

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:

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

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

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

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

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.

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

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

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

Batteries – Wet Cell vs Dry Cell

Wet cell – conventional electrolyte

“Dry” Cell – typically a moist paste

Showed that If put electrolyte–soaked cloth in between 2 different metals, (e.g., Cu & Zn), current would flow

Alessandro Volta

Types of Primary Cell Batteries

Alkaline

Aluminium (aluminum-air)

Lithium (not Lithium-ion)

Oxyride (Oxy Nickel Hydroxide (NiOOH)

Silver-oxide

Zinc-air

Zinc-carbon

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

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

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)

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

Alkaline Battery

Steel Case (current

collector)MnO2, cathode

mixture

Zn-KOH anode paste

Brass current collector

KOH electrolyte

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)

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

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)

Secondary Batteries

Secondary – can be recharged by reversing flow of current, regenerating reactants

• AKA storage batteries

Types of Secondary Cell Batteries

Lead-acid (automobile)

Lithium

Lithium-ion

Lithium-ion polymer

Nickel-cadmium (NiCad)

Nickel metal hydride (NiMH)

Molten salt

Lead-Acid (Storage) BatteryMost common secondary battery

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

Lead storage battery

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

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

NiCad Battery

1.4 V

Rechargeable (like lead acid battery, products adhere to electrodes)

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–

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

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

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

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

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

Hydrogen Fuel Cell with PEM

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

Corrosion Cell

and oftenFe2+ Fe3+ + e-

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)

Corrosion

Can be slow in pure water

Faster in water with dissolved salts, especially with chloride or sulfate ions

Seawater especially corrosive

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)

Use of Zinc Sacrificial Anode

Sacrificial Mg Anode to Protect Steel Pipeline

Cathodic Protection of Iron Storage Tank Using Mg Sacrificial Anode

Corrosion Prevention - Galvanizing

Combines coating idea (barrier) with sacrificial anode idea

Inexpensive and widely used

Especially important for parts of car bodies

Intact Zinc Coating (Barrier)

Breached Zinc Coating (Sacrificial Anode)

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

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.

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

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)

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

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

Cell for electrolysis of waterGenerates H2(g) & O2(g) in 2:1 ratio

Electrolysis Applications

Down’s cell – molten NaCl to produce Na(l), Cl2(g)

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)

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

Brine Electrolysis

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.

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

Refining of Aluminum

Refining Al – Hall-Héroult Process

Power Source

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

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

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

Electroplating

Electroplating of silver

End of Chapter