Redox Reactions and Electrochemistry

I. Redox Reactionsa) Oxidation Numberb) Oxidizing and Reducing Reagents

II. Galavanic or Voltaic Cellsa) Anode/Cathode/Salt Bridgeb) Cell Notationsc) Determining Cell Potential/Cell Voltage/Electromotive

force (emf)

III. Relating Cell Potential to K and DG0

IV. Effect of Concentration on Cell Potential

Redox Reactions and Electrochemistry

V. Corrosion

VI. Batteries

VII. Fuel Cells

VIII.Electrolytic Cellsa) Calculating amounts of substances reduced or

oxidized

Electrochemistry: Interconversion of electrical and chemical energy using redox reactions

Redox (Oxidation-Reduction) Reaction: Type of electrontransfer reaction. One substance gives up electrons;

the other accepts electrons.

OIL RIG

• Oxidation Half-Reaction; Oxidation Involves Loss of electrons

• Reduction Half-Reaction; Reduction Involves Gain of electrons

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Net Redox Rxn; 2Mg + O2 -> 2 Mg+2 + 2 O-2

Oxidation numberThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferredto the more electronegative atom.

1. Oxidation number equals ionic charge for monoatomic ions in ionic compound

2. Metal ions in Family A have one, positive oxidation number; Group IA metals are +1, IIA metals are +2

Li+, Li = +1; Mg+2, Mg = +2

4.4

CaBr2; Ca = +2, Br = -1

Oxidation number,continuedThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferredto the more electronegative atom.

3. The oxidation number of a transition metal ion is positive, but can vary in magnitude.

4. Nonmetals can have a variety of oxidation numbers,both positive and negative numbers which can vary in magnitude.

4.4

5. Free elements (uncombined state) have an oxidation number of zero. Each atom in O2, F2, H2, Cl2, K, Be has the same oxidation number; zero.

7. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.

IF; F= -1; I = +1

8. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1 or when it’s in elemental form (H2; oxidation # =0).

HF; F= -1, H= +1

NaH; Na= +1, H = -1

6. The oxidation number of fluorine is always –1. (unless fluorine is in elemental form, F2)

H2O ; H=+1, O= -2SO3; O = -2; S = +6

9. The oxidation number of oxygen is usually –2. In H2O2 and O2

2- it is –1, in elemental form (O2 or O3) it is 0.

HCO3-

O = -2 H = +1

3x(-2) + 1 + ? = -1

C = +4

Oxidation numbers of all the atoms in HCO3

- ?

4.4

NaIO3

Na = +1 O = -2

3x(-2) + 1 + ? = 0

I = +5

IF7

F = -1

7x(-1) + ? = 0

I = +7

K2Cr2O7

O = -2 K = +1

7x(-2) + 2x(+1) + 2x(?) = 0

Cr = +6

Oxidation numbers of all the elements in the following ?

4.4

Determination of Oxidizing and Reducing Agents

I. Determine oxidation # for all atoms in both the reactants and products.

II. Look at same atom in reactants and products and see if oxidation # increased or decreased.• If oxidation # decreased; substance reduced• If oxidation # increased; substance oxidized

Determination of Oxidizing and Reducing Agents, continued

• Oxidizing Agent: Substance that oxidizes the other substance by accepting electrons. It is reduced in reaction.

• Reducing Agent: Substance that reduces the other substance by donating electrons. It is oxidized in reaction.

Spontaneous Redox ReactionZn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq)

Cu+2

Zn

time Zn+2Cu

Gets Smaller -> <- Gets Larger

Voltaic Cell Animation

Anode; Site of OxidationCathode; Site of Reduction

AnOx or both vowelsRed Cat or both consonants

Direction of electron flow; anode to cathode (alphabetical)

Salt Bridge; Maintains electrical neutrality+ ion migrates to cathode- ion migrates to anode

Cell Notation

1. Anode

2. Salt Bridge

3. Cathode

Anode | Salt Bridge | Cathode

| : symbol is used whenever there is a different phase

19.2

Cell Notation

Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)

[Cu2+] = 1 M & [Zn2+] = 1 M

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)anode cathode

Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s)

anode cathodeSalt bridge

More detail..

K(NO3)

Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)

Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt

Electrochemical Cells

19.2

The difference in electrical potential between the anode and cathode is called:

• cell voltage

• electromotive force (emf)

• cell potential

000reductionoxidationCell EEE

UNITS: Volts Volt (V) = Joule (J) Coulomb, C

Standard Electrode Potentials

19.3

Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.

E0 = 0 V

Standard hydrogen electrode (SHE)

2e- + 2H+ (1 M) H2 (1 atm)

Reduction Reaction

Determining if Redox Reaction is Spontaneous

• + E°CELL ; spontaneous reaction

• E°CELL = 0; equilibrium• - E°CELL; nonspontaneous

reaction

More positive E°CELL ; stronger oxidizing agent ormore likely to be reduced

19.3

• E0 is for the reaction as written

• The half-cell reactions are reversible

• The sign of E0 changes when the reaction is reversed

• Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0

• The more positive E0 the greater the tendency for the substance to be reduced

Relating E0Cell to DG0

ech

workECell arg

Unitswork, Joulecharge, CoulombEcell; Volts

charge = nFFaraday, F; charge on 1 mole e-F = 96485 C/mole

work = (charge)Ecell = -nFEcell

DG = work (maximum)

DG = -nFEcell

Relating EoCELL to the

Equilibrium Constant, KDG0 = -RT ln K

DG0 = -nFE0cell

-RT ln K = -nFE0cell

K

nF

RTECell ln0

0257.0

96485

29831.8

moleC

KmolK

J

F

RT

Kn

Kn

ECell log0592.0

ln0257.00

Effect of Concentration on Cell Potential

DG =DG0 + RTlnQ

DG0 = -nFE0cell

-nFEcell= -nFE0cell + RTln Q

Ecell= E0cell - RTln Q

nF

Ecell= E0cell - 0.0257ln Q

nEcell= E0

cell – 0.0592log Q n

Corrosion – Deterioration of Metals by Electrochemical Process

Corrosion – Deterioration of Metals by Electrochemical Process

Corrosion – Deterioration of Metals by Electrochemical Process

Cathodic Protection

Abbreviated Standard Reduction Potential Table

Batteries

19.6

Leclanché cell

Dry cell

Zn (s) Zn2+ (aq) + 2e-Anode:

Cathode: 2NH4+ (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)

Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

Batteries

Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-Anode:

Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)

Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

Mercury Battery

19.6

Batteries

19.6

Anode:

Cathode:

Lead storagebattery

PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e- PbSO4 (s) + 2H2O (l)

Pb (s) + SO42- (aq) PbSO4 (s) + 2e-

Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq) 2PbSO4 (s) + 2H2O (l)

Fuel Cell vs. Battery

• Battery; Energy storage device– Reactant chemicals already in device– Once Chemicals used up; discard (unless rechargeable)

• Fuel Cell; Energy conversion device– Won’t work unless reactants supplied– Reactants continuously supplied; products continuously

removed

Fuel Cell

A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning

Anode:

Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)

2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-

2H2 (g) + O2 (g) 2H2O (l)

Types of Electrochemical Cells

• Voltaic/Galvanic Cell; Energy released from spontaneous redox reaction can be transformed into electrical energy.

• Electrolytic Cell; Electrical energy is used to drive a nonspontaneous redox reaction.

Faraday’s Constant Redox Eqn

Molar Mass

Charge =(Current)(Time)