ElectroChemistry PPT

7/28/2019 ElectroChemistry PPT

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CHEM 160 General Chemistry II

Lecture Presentation

Electrochemistry

Chapter 20


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Electrochemistry

Electrochemistry

deals with interconversion between chemical and

electrical energy


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Electrochemistry

Electrochemistry

deals with the interconversion between chemical and

electrical energy involves redox reactions


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Electrochemistry

Electrochemistry

deals with interconversion between chemical and

electrical energy involves redox reactions

electron transfer reactions

Oh No! Theyre back!


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Redox reactions (quick review)

Oxidation

Reduction

Reducing agent

Oxidizing agent


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Oxidation

loss of electrons

Reduction

Reducing agent

Oxidizing agent


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Oxidation

loss of electrons

Reduction gain of electrons

Reducing agent

Oxidizing agent


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Oxidation

loss of electrons


Reducing agent

donates the electrons and is oxidized

Oxidizing agent


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Oxidation

loss of electrons


Reducing agent

donates the electrons and is oxidized

Oxidizing agent

accepts electrons and is reduced


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Redox Reactions

Direct redox reaction


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Redox Reactions


Oxidizing and reducing agents are mixed together


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CuSO4(aq)

(Cu2+)

Zn rod

Direct Redox Reaction


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CuSO4(aq)

(Cu2+)

Zn rod

Deposit of

Cu metal

forms

Direct Redox Reaction


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Redox Reactions


Oxidizing and reducing agents are mixed together

Indirect redox reaction Oxidizing and reducing agents are separated but

connected electrically

Example

Zn and Cu2+ can be reacted indirectly

Basis for electrochemistry Electrochemical cell


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Electrochemical Cells


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Voltaic Cell

cell in which a spontaneous redox reaction generates

electricity

chemical energy electrical energy


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Voltaic Cell



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Electrolytic Cell

electrochemical cell in which an electric current

drives a nonspontaneous redox reaction

electrical energy chemical energy


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Cell Potential


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Cell Potential

Cell Potential (electromotive force), Ecell (V) electrical potential difference between the two

electrodes or half-cells

Depends on specific half-reactions, concentrations, and

temperature

Under standard state conditions ([solutes] = 1 M, Psolutes =

1 atm), emf = standard cell potential, Ecell

1 V = 1 J/C

driving force of the redox reaction


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high electrical

potential

low electrical

potential

Cell Potential


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Cell Potential

Ecell = Ecathode - Eanode = Eredn - Eox


(Ecathode and Eanode are reduction potentials by definition.)


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Cell Potential


Ecell can be measured

Absolute Ecathode and Eanode values cannot

Reference electrode

has arbitrarily assigned E

used to measure relative Ecathode and Eanode for half-

cell reactions Standard hydrogen electrode (S.H.E.)

conventional reference electrode


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Standard Hydrogen Electrode

E = 0 V (by

definition; arbitrarily

selected)

2H+ + 2e- H2


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Example 1

A voltaic cell is made by connecting a standard

Cu/Cu2+ electrode to a S.H.E. The cell potential

is 0.34 V. The Cu electrode is the cathode.

What is the standard reduction potential of the

Cu/Cu2+ electrode?


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Example 2

A voltaic cell is made by connecting a standard

Zn/Zn2+ electrode to a S.H.E. The cell potential

is 0.76 V. The Zn electrode is the anode of the

cell. What is the standard reduction potential of

the Zn/Zn2+ electrode?


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Standard Electrode Potentials

Standard Reduction Potentials, E

Ecell measured relative to S.H.E. (0 V)

electrode of interest = cathode

If E < 0 V:

Oxidizing agent is harder to reduce than H+

If E > 0 V:

Oxidizing agent is easier to reduce than H+


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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e

- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e

- 2Cl-(aq) 1.36Cr2O7

2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H

+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04


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Uses of Standard Reduction

Potentials

Compare strengths of reducing/oxidizing agents.

the more - E, stronger the red. agent

the more + E, stronger the ox. agent


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- 2Cl-(aq) 1.36Cr2O7



Ox.ag

entstrengthincreases

R

ed.agentstre

ngthincrease

s


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Potentials

Determine if oxidizing and reducing agent react

spontaneously

diagonal ruleox. agent

red. agent


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Potentials

Determine if oxidizing and reducing agent react

spontaneously

Cathode

(reduction)

more +

Anode

(oxidation)

more -


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- 2Cl-(aq) 1.36Cr2O7




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Potentials

Calculate Ecell Ecell = Ecathode - Eanode

Greater Ecell

, greater the driving force

Ecell > 0 : spontaneous redox reactions

Ecell < 0 : nonspontaeous redox reactions


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Example 3

A voltaic cell consists of a Ag electrode in 1.0 M

AgNO3 and a Cu electrode in 1 M Cu(NO3)2.

Calculate Ecell

for the spontaneous cell reaction

at 25C.

St d d R d ti P t ti l


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- 2Cl-(aq) 1.36Cr2O7




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Example 4

A voltaic cell consists of a Ni electrode in 1.0 M

Ni(NO3)2 and an Fe electrode in 1 M Fe(NO3)2.

Calculate Ecell

for the spontaneous cell reaction

at 25C.

St d d R d ti P t ti l


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- 2Cl-(aq) 1.36Cr2O7




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Cell Potential

Is there a relationship between Ecell and DG for a

redox reaction?


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Cell Potential

Relationship between Ecell and DG: DG = -nFEcell

F = Faraday constant = 96500 C/mol e-s, n = # e-s

transferred redox rxn.


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Cell Potential



transferred redox rxn.

1 J = CV

DG < 0, Ecell > 0 = spontaneous


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Equilibrium Constants from Ecell



transferred redox rxn

1 J = CV


Under standard state conditions:

DG = -nFEcell


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transferred redox rxn

1 J = CV



DG = -nFEcell


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Relationship between Ecell andD

G: DG = -nFEcell

F = Faraday constant = 96500 C/mol e-s, n = # e-s transferred redoxrxn

1 J = CV



DG = -nFEcell

and

DG = -RTlnK

so

-nFEcell = -RTlnK


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DH DSCalorimetric Data

DGElectrochemical

DataComposition

DataEcell

Equilibrium

constants

K


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Example 5

Calculate Ecell, DG, and K for the voltaic cell

that uses the reaction between Ag and Cl2 under

standard state conditions at 25C.


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The Nernst Equation

DG depends on concentrations DG = DG + RTlnQ

and

DG = -nFEcell and DG = -nFEcellthus

-nFEcell = -nFEcell + RTlnQ

or

Ecell = Ecell - (RT/nF)lnQ (Nernst eqn.)


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The Nernst Equation

Ecell = Ecell - (RT/nF)lnQ (Nernst eqn.)

At 298 K (25C), RT/F = 0.0257 V

so Ecell = Ecell - (0.0257/n)lnQ

or

Ecell = Ecell - (0.0592/n)logQ


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Example 7

Calculate the voltage produced by the galvanic

cell which uses the reaction below if [Ag+] =

0.001 M and [Cu2+] = 1.3 M.

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

Standard Reduction Potentials


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- 2Cl-(aq) 1.36Cr2O7



Ox.agentstrength

increases

R

ed.agentstre

ngthincrease

s


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Commercial Voltaic Cells

Battery

commercial voltaic cell used as portable source of

electrical energy

types

primary cell

Nonrechargeable

Example: Alkaline battery

secondary cell

Rechargeable

Example: Lead storage battery


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How Does a Battery Work

cathode (+)

anode (-)

Electrolyte

Paste

Seal/cap

Assume a generalized battery


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Battery

cathode (+):

Reduction occurs

here

anode (-):

oxidation

occurs here

e- flow

Electrolyte paste:ion migration occurs

here

Placing the battery into a flashlight,

etc., and turning the power oncompletes the circuit and allows

electron flow to occur


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How Does a Battery Work

Battery reaction when producing electricity

(spontaneous):Cathode: O1 + e

- R1

Anode: R2 O2 + e-

Overall: O1 + R2 R1 + O2 Recharging a secondary cell

Redox reaction must be reversed, i.e., current isreversed (nonspontaneous)

Recharge: O2 + R1 R2 + O1

Performed using electrical energy from an externalpower source


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Batteries

Read the textbook to fill in the details on

specific batteries.

Alkaline battery

Lead storage battery

Nicad battery

Fuel cell


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Corrosion

Corrosion deterioration of metals by a spontaneous redox

reaction

Attacked by species in environment

Metal becomes a voltaic cell

Metal is often lost to a solution as an ion

Rusting of Iron

Corrosion of Iron


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Corrosion of Iron


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Corrosion of Iron

Half-reactions

anode: Fe(s) Fe2+(aq) + 2e-

cathode: O2(g) + 4H+

(aq) + 4e-

2H2O(l)overall: 2Fe(s) + O2(g) + 4H

+(aq)

2Fe2+(aq) + 2H2O(l)

Ecell > 0 (Ecell = 0.8 to 1.2 V), so process isspontaneous!


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Corrosion of Iron

Rust formation:4Fe2+(aq) + O2(g) + 4H

+(aq) 4Fe3+(aq) + 2H2O(l)

2Fe3+(aq) + 4H2O(l) Fe2O3H2O(s) + 6H+(aq)


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Prevention of Corrosion

Cover the Fe surface with a protective coating Paint

Passivation

surface atoms made inactive via oxidation2Fe(s) + 2Na2CrO4(aq) + 2H2O(l) -->

Fe2O3(s) + Cr2O3(s) + 4NaOH(aq)

Other metal

Tin

Zn

Galvanized iron


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Prevention of Corrosion

Cathodic Protection

metal to be protected is brought into contact with a

more easily oxidized metal

sacrificial metal becomes the anode

Corrodes preferentially over the iron

Iron serves only as the cathode



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Half-reaction EF

2(g) + 2e- -> 2F-(aq) +2.87 V

Ag+(aq) + e- -> Ag(s) +0.80 V

Cu2+(aq) + 2e- -> Cu(s) +0.34 V

2H+(aq) + 2e- -> H2(g) 0 V

Ni2+(aq) + 2e- -> Ni(s) -0.25 V

Fe2+(aq) + 2e- -> Fe(s) -0.44 V

Zn2+(aq) + 2e- -> Zn(s) -0.76 V

Al3+(aq) + 3e- -> Al(s) -1.66 V

Mg2+(aq) + 2e- ->Mg(s) -2.38 V

Metals more

easily oxidized

than Fe havemore negative

Es

Cathodic Protection


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Cathodic Protection

galvanized steel (Fe)


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Cathodic Protection

(cathode)

(electrolyte)

(anode)


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Electrolysis

Electrolysis

process in which electrical energy drives a

nonspontaneous redox reaction

electrical energy is converted into chemical energy

Electrolytic cell

electrochemical cell in which an electric current

drives a nonspontaneous redox reaction


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Electrolysis

Same principles apply to both electrolytic and

voltaic cells

oxidation occurs at the anode

reduction occurs at the cathode

electrons flow from anode to cathode in the external

circuit

In an electrolytic cell, an external power source pumps theelectrons through the external circuit


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Electrolysis of Molten NaCl

Q tit ti A t f El t h i l C ll


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Quantitative Aspects of Electrochemical Cells

For any half-reaction, the amount of a substanceoxidized or reduced at an electrode is proportional to

the number of electrons passed through the cell

Faradays law of electrolysis

Examples Na+ + 1e- Na

Al3+ + 3e- Al

Number of electrons passing through cell is measured by

determining the quantity of charge (coulombs) that haspassed

1 C = 1 A x 1 s

1 F = 1 mole e- = 96500 C

Steps for Quantitative Electrolysis


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Steps for Quantitative Electrolysis

Calculations

current (A) and time

(s), A x s

charge in

coulombs

(C)

Number of

moles of e-moles of substance

oxidized or reduced

mass of substance

oxidized or reduced


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Example 8

What mass of copper metal can be produced by

a 3.00 A current flowing through a copper(II)

sulfate (CuSO4) solution for 5.00 hours?

E l 9


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Example 9

An aqueous solution of an iron salt is

electrolyzed by passing a current of 2.50 A for

3.50 hours. As a result, 6.1 g of iron metal are

formed at the cathode. Calculate the charge on

the iron ions in the solution.

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ElectroChemistry PPT