1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver Principles of Reactivity:

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© 2006 Brooks/Cole - Thomson

Chemistry and Chemical Reactivity 6th Edition

John C. Kotz Paul M. Treichel

Gabriela C. Weaver

Principles of Reactivity: Electron Transfer Reactions

© 2006 Brooks/Cole Thomson

Lectures written by John Kotz

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ELECTROCHEMISTRYChapter 19

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TRANSFER REACTIONS

Atom/Group transfer

HCl + H2O ---> Cl- + H3O+

Electron transfer

Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)

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Electron Transfer Reactions

• Electron transfer reactions are oxidation-

reduction or redox reactions.

• Redox reactions can result in the

generation of an electric current or be

caused by imposing an electric current.

• Therefore, this field of chemistry is often

called ELECTROCHEMISTRY.

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Review of Terminology for Redox Reactions

• OXIDATION—loss of electron(s) by a species; increase in oxidation number.

• REDUCTION—gain of electron(s); decrease in oxidation number.

• OXIDIZING AGENT—electron acceptor; species is reduced.

• REDUCING AGENT—electron donor; species is oxidized.

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OXIDATION-REDUCTION REACTIONS

Direct Redox ReactionOxidizing and reducing agents in direct

contact.Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)

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Balancing Equations

Cu + Ag+ --give--> Cu2+ + Ag

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Balancing Equations

Step 1:Divide the reaction into half-reactions, one for oxidation and the other for reduction.

Ox Cu ---> Cu2+

Red Ag+ ---> AgStep 2:Balance each for mass. Already done in

this case.Step 3:Balance each half-reaction for charge

by adding electrons.Ox Cu ---> Cu2+ + 2e-Red Ag+ + e- ---> Ag

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Balancing Equations

Step 4:Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires.

Reducing agent Cu ---> Cu2+ + 2e-Oxidizing agent 2 Ag+ + 2 e- ---> 2 AgStep 5:Add half-reactions to give the overall

equation.Cu + 2 Ag+ ---> Cu2+ + 2Ag

The equation is now balanced for both charge and mass.

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OXIDATION-REDUCTION REACTIONS

Indirect Redox Reaction

A battery functions by transferring electrons through an external wire from the reducing

agent to the oxidizing agent.



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ElectrochemistryAlessandro Volta, 1745-1827, Italian scientist and inventor.

Luigi Galvani, 1737-1798, Italian scientist and inventor.

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Oxidation: Zn(s) ---> Zn2+(aq) + 2e-Reduction: Cu2+(aq) + 2e- ---> Cu(s)--------------------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)

Electrons are transferred from Zn to Cu2+, but there is no useful electric current.

CHEMICAL CHANGE --->ELECTRIC CURRENT

With time, Cu plates out onto Zn metal strip, and Zn strip

“disappears.”

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•To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs through an external wire.

CHEMICAL CHANGE --->ELECTRIC CURRENT

This is accomplished in a GALVANIC or VOLTAIC cell.

A group of such cells is called a battery.

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•Electrons travel through external wire.•Salt bridge allows anions and cations to move between electrode compartments.

Fe --> Fe2+ + 2e- Cu2+ + 2e- --> Cu

<--AnionsCations-->

OxidationAnodeNegative

ReductionCathodePositive

Fe

Fe

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The Cu|Cu2+ and Ag|Ag+ Cell

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Electrochemical Cell

Electrons move from anode to cathode in the wire.Anions & cations move thru the salt bridge.



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Terms Used for Voltaic Cells

Figure 20.6

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CELL POTENTIAL, E

• Electrons are “driven” from anode to cathode by an electromotive force or emf.

• For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M.

• Standard reduction potentials are measured at standard conditions (1 M, 25oC)

Zn and Zn2+,anode

Cu and Cu2+,cathode

1.10 V

1.0 M 1.0 M

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CELL POTENTIAL, E

• For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M.

• This is the STANDARD CELL POTENTIAL, Eo

• —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.

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Calculating Cell Voltage

• Balanced half-reactions can be added together to get overall, balanced equation.

Zn(s) ---> Zn2+(aq) + 2e-Cu2+(aq) + 2e- ---> Cu(s)--------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)

If we know Eo for each half-reaction, we could get Eo for net reaction.

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CELL POTENTIALS, Eo

Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL

2 H+(aq, 1 M) + 2e- <----> H2(g, 1 atm)

Eo = 0.0 V

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Zn/Zn2+ half-cell hooked to a SHE.Eo for the cell = +0.76 V

Negative electrode

Supplier of

electrons

Acceptor of

electrons

Positive electrode

2 H+ + 2e- --> H2

ReductionCathode

Zn --> Zn2+ + 2e- Oxidation

Anode

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Reduction of H+ by Zn

Active Figure 20.13

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Overall reaction is reduction of H+ by Zn metal.

Zn(s) + 2 H+ (aq) --> Zn2+ + H2(g) Eo = +0.76 V

Therefore, Eo for Zn ---> Zn2+ (aq) + 2e- is +0.76 V

Zn is a better reducing agent than H2.

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Zn/Cu Electrochemical Cell

Zn(s) ---> Zn2+(aq) + 2e- Eo = +0.76 VCu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 V---------------------------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)

Eo (calc’d) = +1.10 V

Cathode, positive, sink for electrons

Anode, negative, source of electrons

+

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Uses of Eo Values

Organize half-reactions by relative ability to act as oxidizing agents

• Use this to predict direction of redox reactions and cell potentials.

Cu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 VZn2+(aq) + 2e- ---> Zn(s) Eo = –0.76 V

Note that when a reaction is reversed the sign of E˚ is reversed!

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© 2006 Brooks/Cole - ThomsonPotential Ladder for Reduction Half-Reactions

Best oxidizing agents

Best reducing agents

Figure 20.14

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TABLE OF STANDARD REDUCTION POTENTIALS

2

Eo (V)

Cu2+ + 2e- Cu +0.34

2 H+ + 2e- H 0.00

Zn2+ + 2e- Zn -0.76

oxidizingability of ion

reducing abilityof element

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Using Standard Potentials, Eo

Table 20.1

• Which is the best oxidizing agent:

O2, H2O2, or Cl2? _________________

• Which is the best reducing agent:

Hg, Al, or Sn? ____________________

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Standard Redox Potentials, Eo

Any substance on the right will reduce any substance higher than it on the left.

• Zn can reduce H+ and Cu2+.

• H2 can reduce Cu2+ but

not Zn2+

• Cu cannot reduce H+ or Zn2+.

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Standard Redox Potentials, Eo

Cu2+ + 2e- --> Cu +0.34

+2 H + 2e- --> H2 0.00

Zn2+ + 2e- --> Zn -0.76

Northwest-southeast rule: product-favored reactions occur between • reducing agent at southeast corner • oxidizing agent at northwest corner

Any substance on the right will reduce any substance higher than it on the left.

Ox. agent

Red. agent

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Cu(s) | Cu2+(aq) || H+(aq) | H2(g)

Cu2+ + 2e- --> CuOr

Cu --> Cu2+ + 2 e-

H2 --> 2 H+ + 2 e-or

2 H+ + 2e- --> H2

CathodePositive

AnodeNegativeElectrons

<----------

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Cu(s) | Cu2+(aq) || H+(aq) | H2(g)

Cu2+ + 2e- --> Cu H2 --> 2 H+ + 2 e-

CathodePositive

AnodeNegativeElectrons

<----------

The sign of the electrode in Table 20.1 is the polarity when hooked to the H+/H2 half-cell.

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Using Standard Potentials, Eo

• In which direction do the following reactions

go?

• Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)

–Goes right as written

• 2 Fe2+(aq) + Sn2+(aq) ---> 2 Fe3+(aq) + Sn(s)

–Goes LEFT opposite to direction written

• What is Eonet for the overall reaction?

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Cd --> Cd2+ + 2e-or

Cd2+ + 2e- --> Cd

Fe --> Fe2+ + 2e-or

Fe2+ + 2e- --> Fe

Eo for a Voltaic Cell

All ingredients are present. Which way does reaction proceed? Calculate Eo for this cell.

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E at Nonstandard Conditions

• The NERNST EQUATION• E = potential under nonstandard conditions

• n = no. of electrons exchanged

• F = Faraday’s constant

• R = gas constant

• T = temp in Kelvins

• ln = “natural log”

• Q = reaction quotient

QnF

RTcello

cell EE ln

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Eo and Thermodynamics

• Eo is related to ∆Go, the free energy change for the reaction.

• ∆G˚ is proportional to –nE˚

∆Go = -nFEo where F = Faraday constant

= 9.6485 x 104 J/V•mol of e-

(or 9.6485 x 104 coulombs/mol)and n is the number of moles of electrons

transferred

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Eo and ∆Go

∆Go = - n F Eo For a product-favored reaction Reactants ----> Products

∆Go < 0 and so Eo > 0Eo is positive

For a reactant-favored reaction Reactants <---- Products

∆Go > 0 and so Eo < 0Eo is negative

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Eo and Equilibrium Constant

DGo = -RT ln K

DGo = -nFEo

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Dry Cell Battery

Anode (-)

Zn ---> Zn2+ + 2e-

Cathode (+)

2 NH4+ + 2e- --->

2 NH3 + H2

Primary battery — uses redox reactions that cannot be restored by recharge.

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Nearly same reactions as in common dry cell, but under basic conditions.

Alkaline Battery

Anode (-): Zn + 2 OH- ---> ZnO + H2O + 2e-

Cathode (+): 2 MnO2 + H2O + 2e- --->

Mn2O3 + 2 OH-



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Lead Storage Battery

• Secondary battery • Uses redox

reactions that can be reversed.

• Can be restored by recharging

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Ni-Cad Battery

Anode (-)

Cd + 2 OH- ---> Cd(OH)2 + 2e-

Cathode (+)

NiO(OH) + H2O + e- ---> Ni(OH)2 + OH-



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Fuel Cells: H2 as a Fuel

• Fuel cell - reactants are

supplied continuously

from an external

source.• Cars can use

electricity generated

by H2/O2 fuel cells.

• H2 carried in tanks or

generated from

hydrocarbons.

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Hydrogen—Air Fuel Cell

Figure 20.12

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H2 as a Fuel

Comparison of the volumes of substances required to store 4 kg of hydrogen relative to car size. (Energy, p. 290)

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Storing H2 as a Fuel

One way to store H2 is to adsorb the gas onto a metal or metal alloy. (Energy, p. 290)

Documents

1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver Principles of Reactivity: