Electron Transfer ReactionsElectron Transfer ReactionsElectron Transfer ReactionsElectron Transfer Reactions

• Electron transfer Electron transfer reactions are reactions are oxidation-oxidation-

reductionreduction or or redoxredox reactions. reactions.

• Redox reactions can result in the generation Redox reactions can result in the generation

of an electric current or be caused by of an electric current or be caused by

imposing an electric current. imposing an electric current.

• Therefore, this field of chemistry is often Therefore, this field of chemistry is often

called called ELECTROCHEMISTRYELECTROCHEMISTRY..

Dry Cell BatteryDry Cell Battery

Anode (-)Anode (-)

Zn Zn ZnZn2+2+ + + 2e-2e-

Cathode (+)Cathode (+)

2 NH2 NH44++ + 2e- + 2e-

2 NH2 NH33 + H + H22

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

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

Alkaline BatteryAlkaline Battery

Anode (-):Anode (-): Zn + 2 OHZn + 2 OH-- ZnO + HZnO + H22O + O +

2e-2e-Cathode (+): Cathode (+): 2 MnO2 MnO22 + H + H22O + 2e- O + 2e-

MnMn22OO33 + 2 OH + 2 OH--

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Lead Storage BatteryLead Storage Battery

• Secondary batterySecondary battery • Uses redox reactions Uses redox reactions

that can be reversed.that can be reversed.• Can be restored by Can be restored by

rechargingrecharging

Lead Storage BatteryLead Storage BatteryAnode (-)Anode (-) EEoo = +0.36 V = +0.36 V

Pb + HSOPb + HSO44-- PbSOPbSO44 + H + H++ + 2e- + 2e-

Cathode (+) Cathode (+) EEoo = +1.68 V = +1.68 V

PbOPbO22 + HSO + HSO44-- + 3 H + 3 H++ + 2e- + 2e- PbSO PbSO44 + 2 H + 2 H22OO

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Ni-Cad BatteryNi-Cad BatteryAnode (-)Anode (-)

Cd + 2 OHCd + 2 OH-- Cd(OH)Cd(OH)22 + 2e- + 2e-

Cathode (+) Cathode (+)

NiO(OH) + HNiO(OH) + H22O + e- O + e- Ni(OH)Ni(OH)22 + OH + OH--

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Review of Terminology Review of Terminology for Redox Reactionsfor Redox Reactions

Review of Terminology Review of Terminology for Redox Reactionsfor Redox Reactions

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

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

• OXIDIZING AGENTOXIDIZING AGENT—electron acceptor; species —electron acceptor; species is reduced.is reduced.

• REDUCING AGENTREDUCING AGENT—electron donor; species is —electron donor; species is oxidized.oxidized.

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

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

• OXIDIZING AGENTOXIDIZING AGENT—electron acceptor; species —electron acceptor; species is reduced.is reduced.

• REDUCING AGENTREDUCING AGENT—electron donor; species is —electron donor; species is oxidized.oxidized.

20.1

OXIDATION-REDUCTION REACTIONS

OXIDATION-REDUCTION REACTIONS

Understand a Understand a DirectDirect Redox Reaction Redox Reaction

Oxidizing and reducing agents in Oxidizing and reducing agents in directdirect contact. contact.Cu(s) + 2 AgCu(s) + 2 Ag++(aq) Cu(aq) Cu2+2+(aq) + 2 Ag(s)(aq) + 2 Ag(s)

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Know how to Balance Equations Know how to Balance Equations in a in a neutral neutral solutionsolutionCu + AgCu + Ag++ Cu Cu2+2+ + Ag + Ag

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

OxOx Cu Cu Cu Cu2+2+

RedRed Ag Ag++ Ag Ag

Step 2: Step 2: Balance each for mass. Already done in Balance each for mass. Already done in this case.this case.

Step 3: Step 3: Balance each half-reaction for charge by Balance each half-reaction for charge by adding electrons.adding electrons.

OxOx Cu Cu Cu Cu2+2+ + + 2e-2e-

RedRed Ag Ag++ + + e-e- Ag Ag

Know how to Balance Equations Know how to Balance Equations in a in a neutralneutral solution solution

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

Reducing agentReducing agent Cu Cu Cu Cu2+2+ + 2e- + 2e-Oxidizing agentOxidizing agent 22 Ag Ag++ + + 22 e- e- 22 Ag AgStep 5:Step 5: Add half-reactions to give the overall Add half-reactions to give the overall

equation.equation.Cu + 2 AgCu + 2 Ag++ Cu Cu2+2+ + 2Ag + 2AgThe equation is now balanced for both charge and The equation is now balanced for both charge and

mass.mass.

Know how to Balance Equations Know how to Balance Equations in a in a neutralneutral solution solution

Balance the following in Balance the following in acidacid solution—solution—

VOVO22++ + Zn VO + Zn VO2+ 2+ + Zn+ Zn2+2+

Step 1:Step 1: Write the half-reactionsWrite the half-reactions

OxOx Zn ZnZn Zn2+2+

RedRed VOVO22++ VO VO2+2+

Step 2:Step 2: Balance each half-reaction for mass.Balance each half-reaction for mass.

OxOx Zn ZnZn Zn2+2+

RedRedVOVO22

++ VO VO2+2+ + + HH22OO2 H2 H++ ++

Add HAdd H22O on O-deficient side and add HO on O-deficient side and add H++

on other side for H-balance.on other side for H-balance.

Know how to Balance Equations Know how to Balance Equations in an in an acidicacidic solution solution

Step 3:Step 3: Balance half-reactions for charge.Balance half-reactions for charge.OxOx Zn ZnZn Zn2+2+ + + 2e2e--RedRed e-e- + 2 H+ 2 H++ + VO + VO22

++ VO VO2+2+ + H + H22OOStep 4:Step 4: Multiply by an appropriate factor.Multiply by an appropriate factor.OxOx Zn ZnZn Zn2+2+ + + 2e-2e-RedRed 22e-e- + + 44 H H++ + + 22 VO VO22

++ 22 VO VO2+2+ + + 22 HH22OO

Step 5:Step 5: Add Add balancedbalanced half-reactions half-reactionsZn + 4 HZn + 4 H++ + 2 VO + 2 VO22

++ Zn Zn2+2+ + 2 VO + 2 VO2+2+ + 2 + 2 HH22OO

Know how to Balance Equations Know how to Balance Equations in an in an acidicacidic solution solution

Two approaches:Two approaches:

1)1)Follow the same steps as for an acid solution Follow the same steps as for an acid solution but add OHbut add OH-- rather than H rather than H++ in step 2 in step 2

2)2)Balance as if it were an acidic solution then add Balance as if it were an acidic solution then add enough OHenough OH-- to both sides to neutralize the H to both sides to neutralize the H++

Know how to Balance Equations Know how to Balance Equations in a in a basicbasic solution solution

Tips on Balancing Tips on Balancing EquationsEquations

• Never add ONever add O22, O atoms, or , O atoms, or

OO2-2- to balance oxygen. to balance oxygen.

• Never add HNever add H22 or H atoms to or H atoms to

balance hydrogen.balance hydrogen.• Be sure to write the correct Be sure to write the correct

charges on all the ions.charges on all the ions.• Check your work at the end to Check your work at the end to

make sure mass and charge make sure mass and charge are balanced.are balanced.

• PRACTICEPRACTICE

20.2 Know the following:

What a voltaic cell is The features of a voltaic

o Salt Bridge o Anode (oxidation) o Cathode (reduction) o Direction of the flow of electrons

How to draw a voltaic cell including the features

The purpose of a salt bridge

Terms Used for Voltaic CellsTerms Used for Voltaic Cells

See Figure 20.6See Figure 20.6

The Cu|CuThe Cu|Cu2+2+ and Ag|Ag and Ag|Ag++ CellCell

ElectrochemicalElectrochemical CellCell

ElectrochemicalElectrochemical CellCell

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

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20.4 Cells at Standard Conditions Understand the following:

Electromotive force (emf) Voltage potential Standard Conditions & o Standard Reduction Potentials The use of Eo

cell = Eo

cathode - Eoanode

Predict the relative strengths of oxidizing & reducing agents

CELL POTENTIAL, ECELL POTENTIAL, E

• Electrons are “driven” from anode to cathode by Electrons are “driven” from anode to cathode by

an an electromotive forceelectromotive force or or emfemf..• For Zn/Cu cell, this is indicated by a voltage of For Zn/Cu cell, this is indicated by a voltage of

1.10 V at 25 ˚C and when [Zn1.10 V at 25 ˚C and when [Zn2+2+] and [Cu] and [Cu2+2+] = ] = 1.0 M.1.0 M.

Zn and ZnZn and Zn2+2+,,anodeanode

Cu and CuCu and Cu2+2+,,cathodecathode

Zn

Zn2+ ions

Cu

Cu2+ ions

wire

saltbridge

electrons

Zn

Zn2+ ions

Cu

Cu2+ ions

wire

saltbridge

electrons

1.10 V1.10 V

1.0 M1.0 M 1.0 M1.0 M

TABLE OF STANDARD TABLE OF STANDARD REDUCTION POTENTIALSREDUCTION POTENTIALS

TABLE OF STANDARD TABLE OF STANDARD REDUCTION POTENTIALSREDUCTION POTENTIALS

2

Eo (V)

Cu2+ + 2e- Cu +0.34

2 H+ + 2e- H 0.00

Zn2+ + 2e- Zn-0.76oxidizing

ability of ion reducing abilityof element

Using Standard Potentials, EUsing Standard Potentials, Eoo

Table 20.1Table 20.1

• Which is the best oxidizing agent:

O2, H2O2, or Cl2? _________________

• Which is the best reducing agent:

Hg, Al, or Sn? ____________________

Standard Redox Potentials, EStandard Redox Potentials, Eoo

• Zn can reduce HZn can reduce H++ and Cu and Cu2+2+..

• HH22 can reduce Cu can reduce Cu2+2+ but not but not

ZnZn2+2+

• Cu cannot reduce HCu cannot reduce H++ or or ZnZn2+2+..

Eo (V)

Cu2+ + 2e- Cu +0.34

2 H+ + 2e- H2 0.00

Zn2+ + 2e- Zn -0.76

oxidizingability of ion

reducing abilityof element

Eo (V)

Cu2+ + 2e- Cu +0.34

2 H+ + 2e- H2 0.00

Zn2+ + 2e- Zn -0.76

oxidizingability of ion

reducing abilityof element

20.5 Cells at Nonstandard Conditions Know how to use the Nernst Equation

E = Eo – (RT)lnQ nF

Eocell at standard conditions Nonequilibrium concentrations

Number of moles of electrons Faraday’s constant

E = Eo – 0.0257lnQ n E = Eo – 0.0591logQ n

The Nernst Equation at 25oC

20.6 Know how to relate Keq & ΔGo to Eo

cell

ΔrxnGo = -nFEocell

lnK = nEocell

0.0257

20.7

Know that Electrolysis is the use of electrical current to carry out nonspontaneous redox reactions (Eo

cell < 0).

ElectrolysisElectrolysisUsing electrical energy to produce chemical change.

Sn2+(aq) + 2 Cl-(aq) Sn(s) + Cl2(g)

Electrolysis of Aqueous Electrolysis of Aqueous CuClCuCl22Anode (+) Anode (+)

2 Cl2 Cl-- ClCl22(g) + (g) +

2e-2e-

Cathode (-) Cathode (-)

CuCu2+2+ + 2e- Cu + 2e- Cu

EEoo for cell = -1.02 V for cell = -1.02 V

BATTERY

+

Cu2+Cl-

Anode Cathode

electrons

H2O

BATTERY

+

Cu2+Cl-

Anode Cathode

electrons

H2O

Electrolytic Refining of Electrolytic Refining of CopperCopper

See Figure 22.11See Figure 22.11

Impure copper is oxidized to CuImpure copper is oxidized to Cu2+2+ at the anode. at the anode. The aqueous CuThe aqueous Cu2+2+ ions are reduced to Cu metal ions are reduced to Cu metal at the cathode.at the cathode.

Producing AluminumProducing Aluminum2 Al2 Al22OO33 + 3 C + 3 C ff 4 Al + 3 CO4 Al + 3 CO22

Charles Hall (1863-1914) developed electrolysis Charles Hall (1863-1914) developed electrolysis process. Founded Alcoa.process. Founded Alcoa.

For an electrolysis reaction, Know how to determine the amount of current (amps), or the amount of ions needed or the amount metal plated

20.8

Electrical current is amps (I): coulombs/sec