Chapter 19Chapter 19

Electrochemistry

Mr. Watson HST

Mr. Watson HST

Redox ReactionsRedox ReactionsOxidationloss of electrons

Reductiongain of electrons

oxidizing agentsubstance that cause oxidation by being reduced

reducing agentsubstance that cause oxidation by being oxidized

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ElectrochemistryElectrochemistry

In the broadest sense, electrochemistry is the study of chemical reactions that produce electrical effects and of the chemical phenomena that are caused by the action of currents or voltages.

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Oxidation-Reduction Oxidation-Reduction ReactionsReactions

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Voltaic CellsVoltaic Cells

harnessed chemical reaction which produces an electric current

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Voltaic CellsVoltaic Cells

Cells and Cell Reactions

Daniel's Cell

Zn(s) + Cu+2(aq) ---> Zn+2

(aq) + Cu(s)

oxidation half reaction

anode Zn(s) ---> Zn+2(aq) + 2 e-

reduction half reaction

cathode Cu+2(aq) + 2 e- ---> Cu(s)

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Voltaic CellsVoltaic Cells

copper electrode dipped into a solution of copper(II) sulfate

zinc electrode dipped into a solution of zinc sulfate

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Voltaic CellsVoltaic Cells

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Hydrogen ElectrodeHydrogen Electrode consists of a platinum

electrode covered with a fine powder of platinum around which H2(g) is bubbled. Its potential is defined as zero volts.

Hydrogen Half-Cell

H2(g) = 2 H+(aq) + 2 e-

reversible reaction

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Standard Reduction Standard Reduction PotentialsPotentials

the potential under standard conditions (25oC with all ions at 1 M concentrations and all gases at 1 atm pressure) of a half-reaction in which reduction is occurring

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Some Standard Reduction Some Standard Reduction Potentials Table 18-1, pg 837Potentials Table 18-1, pg 837

Li+ + e- ---> Li -3.045 v

Zn+2 + 2 e- ---> Zn -0.763v

Fe+2 + 2 e- ---> Fe -0.44v

2 H+(aq) + 2 e- ---> H2(g) 0.00v

Cu+2 + 2 e- ---> Cu +0.337v

O2(g) + 4 H+(aq) + 4 e- ---> 2 H2O(l) +1.229v

F2 + 2e- ---> 2 F- +2.87v

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If the reduction of mercury (I) in a voltaic cell is desired, the half reaction is:

Which of the following reactions could be used as the anode (oxidation)?

A, B

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Cell PotentialCell Potential

the potential difference, in volts, between the electrodes of an electrochemical cell

Direction of Oxidation-Reduction Reactionspositive value indicates a spontaneous

reaction

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Standard Cell PotentialStandard Cell Potential

the potential difference, in volts, between the electrodes of an electrochemical cell when the all concentrations of all solutes is 1 molar, all the partial pressures of any gases are 1 atm, and the temperature at 25oC

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Cell DiagramCell Diagram

the shorthand representation of an electrochemical cell showing the two half-cells connected by a salt bridge or porous barrier, such as:

Zn(s)/ZnSO4(aq)//CuSO4(aq)/Cu(s)

anode cathode

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Metal Displacement Metal Displacement ReactionsReactions

solid of more reactive metals will displace ions of a less reactive metal from solution

relative reactivity based on potentials of half reactions

metals with very different potentials react most vigorously

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Ag+ + e- --->Ag E°= 0.80 V

Cu2+ + 2e- ---> Cu E°= 0.34 V

Will Ag react with Cu2+?

yes, no

Will Cu react with Ag+?

yes, no

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Gibbs Free EnergyGibbs Free Energyand Cell Potentialand Cell Potential

G = - nFE

where n => number of electrons changed

F => Faraday’s constant

E => cell potential

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Applications of Applications of Electrochemical CellsElectrochemical Cells

Batteries– device that converts chemical energy into

electricity

Primary Cells– non-reversible electrochemical cell– non-rechargeable cell

Secondary Cells– reversible electrochemical cell– rechargeable cell

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Applications of Applications of Electrochemical CellsElectrochemical Cells

BatteriesPrimary Cells

"dry" cell & alkaline cell 1.5 v/cellmercury cell 1.34 v/cellfuel cell 1.23v/cell

Secondary Cellslead-acid (automobile battery) 2

v/cellNiCad 1.25 v/cell

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““Dry” CellDry” Cell

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““Dry” CellDry” Cell

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““Flash Light” BatteriesFlash Light” Batteries

"Dry" CellZn(s) + 2 MnO2(s) + 2 NH4

+ ----->

Zn+2(aq) + 2 MnO(OH)(s) + 2 NH3

Alkaline CellZn(s) + 2 MnO2(s) ---> ZnO(s) + Mn2O3(s)

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““New” Super Iron BatteryNew” Super Iron Battery

Mfe(VI)O4 + 3/2 Zn 1/2 Fe(III)2O3 + 1/2 ZnO + MZnO2

(M = K2 or Ba)Environmentally friendlier than MnO2 containing batteries.

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Mercury CellMercury Cell

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Lead-AcidLead-Acid(Automobile Battery)(Automobile Battery)

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Lead-AcidLead-Acid(Automobile Battery)(Automobile Battery)

Pb(s) + PbO2(s) + 2 H2SO4 = 2 PbSO4(s) + 2 H2O

2 v/cell

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Nickel-Cadmium (Ni-Cad)Nickel-Cadmium (Ni-Cad)

Cd(s) + 2 Ni(OH)3(s) = Cd(OH)2(s) + 2 Ni(OH)2(s)

NiCad 1.25 v/cell

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Hydrogen-Oxygen Fuel CellHydrogen-Oxygen Fuel Cell

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Automobile Oxygen SensorAutomobile Oxygen Sensor

Air, constant [O ]2

porous Pt electrodes

migrating O ions2-

measured potential difference

exhaust gas, unknown [O ]2

ZrO / CaO2

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Automobile Oxygen SensorAutomobile Oxygen Sensor

see Oxygen Sensor Movie from Solid-State Resources CD-ROM

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pH = (Eglass electrode - constant)/0.0592

pH MeterpH Meter

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Effect of Concentration on Cell Effect of Concentration on Cell Voltage: The Nernst EquationVoltage: The Nernst EquationEcell = Eo

cell - (RT/nF)ln Q

Ecell = Eocell - (0.0592/n)log Q

where Q => reaction quotient

Q = [products]/[reactants]

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EXAMPLE:EXAMPLE: What is the cell potential for What is the cell potential for the Daniel's cell when the the Daniel's cell when the [Zn[Zn+2+2] = 10 [Cu] = 10 [Cu+2+2]] ??

Q = ([ZnQ = ([Zn+2+2]/[Cu]/[Cu+2+2] = (] = (10 [Cu10 [Cu+2+2])/[Cu])/[Cu+2+2]] = 10 = 10

EEoo = (0.34 V) = (0.34 V)Cu coupleCu couple + (-(-0.76 V) + (-(-0.76 V)Zn coupleZn couple

n = 2, 2 electron changen = 2, 2 electron change EEcellcell = E = Eoocellcell - (0.0257/n)ln Q - (0.0257/n)ln Q

thus Ecell = (1.10 - (0.0257/2)ln 10) V

Ecell = (1.10 - (0.0257/2)2.303) V

Ecell = (1.10 - 0.0296) V = 1.07 V

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Nernst EquationNernst Equation

+0.83 V

salt bridge

H in2

1 atm

voltmeter

Pt electrode

Pt electrodeNaOH

1 M

H in2

1 atm

e– e–

anode (–)

HCl 1 M

cathode (+)

pn+

+

+

+

–

–

–

–

[H+]acid side [H+]base side

E = Eo – RTnF ln Q =

– 2.3 RTF log

[H+]base side[H+]acid side

[h+]p-type side [h+]n-type side

E (in volts) = – 2.3 RT

F log [h+]n-type side [h+]p-type side

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ElectrolysisElectrolysis

non-spontaneous reaction is caused by the passage of an electric current through a solution

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Electrolysis of KIElectrolysis of KI(aq)(aq)

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ElectrolysisElectrolysis

Electrolysis of Sodium chloride

(chlor-alkali process)

molten reactants => liquid sodium and chlorine gas

aqueous reactants => caustic soda (sodium hydroxide) and chlorine gas

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Aqueous Reactants => Aqueous Reactants => caustic soda (sodium hydroxide) and chlorine caustic soda (sodium hydroxide) and chlorine

gasgas

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Molten Reactants =>Molten Reactants =>liquid sodium and chlorine gasliquid sodium and chlorine gas

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ElectrolysisElectrolysis

Preparation of Aluminum (Hall process)

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Electrolytic Refining of Electrolytic Refining of CopperCopper

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

(aq) + Cu(s)

impure pure anode cathode

impurities: anode mud; Ag, Au, Pb

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Copper PurificationCopper Purification

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Copper PurificationCopper Purification

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Hall Process for AluminumHall Process for Aluminum

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Quantitative Aspects of Quantitative Aspects of ElectrolysisElectrolysis

1 coulomb = 1 amp sec1 mole e- = 96,500 coulombs

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

(45 min)#g Cr = ------------

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

definition of minute

(45 min)(60 sec)#g Cr = ---------------------

(1 min)

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

(45) (60 sec) (25 amp)#g Cr = ---------------------------

(1)

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

definition of a coulomb

(45)(60 sec)(25 amp)(1 C)#g Cr = -----------------------------

(1) (1 amp sec)

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

Faraday’s constant

(45)(25)(60)(1 C)(1 mol e-)#g Cr = ----------------------------------

(1)(1)(96,500 C)

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

atomic weight

(45)(60)(25)(1)(1 mol e-)(52 g Cr)#g Cr = -------------------------------------------

(1)(1)(96,500) (6 mol e-)

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ElectroplatingElectroplating

EXAMPLE: How many grams of chromium can be plated from a Cr+6 solution in 45 minutes at a 25 amp current?

(45)(60)(25)(1)(1 mol e-)(52 g Cr)#g Cr = -------------------------------------------

(1)(1)(96,500)(6 mol e-)

= 58 g Cr

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Rusting of IronRusting of Iron

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CorrosionCorrosion

O2(g) + 4 H+(aq) + 4 e- -----> 2 H2O(l)

Eo = 1.23 V

RustingFe(s) -----> Fe+2

(aq) + 2 e- Eo = 0.44 V

O2(g) + 4 H+(aq) + 4 e- -----> 2 H2O(l) Eo = 1.23 V

------------------------------------------- --------------2 Fe(s) + O2(g) + 4 H+

(aq) ----->

2 H2O(l) + Fe+2(aq) Eo = 1.67 V

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Preventing CorrosionPreventing Corrosion

painting

galvanizing

sacrificial anode