Redox Reactions and ElectrochemistryRedox Reactions and ElectrochemistryI. Redox Reactions

a) 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 PotentialV. CorrosionVI. BatteriesVII. Fuel CellsVIII.Electrolytic Cells

a) Calculating amounts of substances reduced or oxidized

REDOX REACTIONREDOX REACTION

DEFINITIONS Oxidation: Loss of electrons. Reduction: Gain of electrons. LEO says GER

Oxidation: Gain of oxygenReduction: Lost of oxygen

Oxidation: Increasing of oxidation numberReduction: Reducing of oxidation number

CO CO2 (CO oxidized) why?CH3COOH CH3CHO (CH3COOH reduced) why?H2SO3 H2SO4 ??HNO3 HNO2 ??

Ca2+ Ca (Ca2+ reduced) why? Na Na+ (Na oxidized) why?Fe3+ Fe2+ Mn2+ MnO4

– (Mn2+ oxidized) why?Cl2 2 Cl– (Cl2 reduced) why? H2O2 H2O ??NaH H2 ??

LEO says GER :LEO says GER :

Lose Electrons = Oxidation

Sodium is oxidized

Gain Electrons = Reduction

Chlorine is reduced

Na0 Na+1 + 1e-

Cl0 + 1e– Cl–

Rules for Assignment of Oxidation Number (ON)

1) The ON of all pure elements is zero.2) The ON of H is +1, except in hydrides, where it is -1.3) The ON of O is -2, except in peroxides, where it is -1.4) The algebraic sum of ON must equal zero for a

neutral molecule or the charge on an ion.

Variable Oxidation Number of Elements Sulfur: SO4

2-(+6), SO32-(+4), S(0), FeS2(-1), H2S(-2)

Carbon: CO2(+4), C(0), CH4(-4) Nitrogen: NO3

-(+5), NO2-(+3), NO(+2), N2O(+1), N2(0), NH3(-3)

Iron: Fe2O3(+3), FeO(+2), Fe(0) Manganese: MnO4

-(+7), MnO2(+4), Mn2O3(+3), MnO(+2), Mn(0) Copper: CuO(+2), Cu2O(+1), Cu(0) Tin: SnO2(+4), Sn2+(+2), Sn(0) Uranium: UO2

2+(+6), UO2(+4), U(0) Arsenic: H3AsO4

0(+5), H3AsO30(+3), As(0), AsH3(-1)

Chromium: CrO42-(+6), Cr2O3(+3), Cr(0)

Gold: AuCl4-(+3), Au(CN)2-(+1), Au(0)

BALANCING OVERALL REDOX REACTIONSBALANCING OVERALL REDOX REACTIONSExample balance the redox reaction below:

Fe + Cl2 Fe3+ + Cl-

Step 1: Assign oxidation number,Fe0 + Cl20 Fe3+ + Cl-

Step 2: Determine number of electrons lost or gained by reactants.

Fe0 + Cl20 Fe3+ + Cl- 3e- 2e-

Step 3: Cross multiply.2Fe + 3Cl20 2Fe3+ + 6Cl-

may be written as the sum of two half-cell reactions:

2Fe 2Fe3+ + 6e- (oxidation)3Cl20 + 6e- 6Cl- (reduction)

All overall redox reactions can be expressed as the sum of two half-cell reactions, one a reduction and one an oxidation.

The overall reaction:

2Fe + 3Cl20 2Fe3+ + 6Cl-

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Another example - balance the redox reaction:

FeS2 + O2 Fe(OH)3 + SO42-

Fe+2S20 + O2

0 Fe+3(OH)3 + S+6O42-

15e- 4e-

4FeS2 + 15O2 Fe(OH)3 + SO42-

4FeS2 + 15O2 4Fe(OH)3 + 8SO42-

4FeS2 + 15O2 +14H2O 4Fe(OH)3 + 8SO42- + 16H+

This reaction is the main cause of acid generation in drainage from sulfide ore deposits. Note that we get 4 moles of H+ for every mole of pyrite oxidized!

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Electrochemistry: Interconversion of electrical and chemical energy using redox reactions

Oxidation Half-Reaction: Oxidation Involves Loss of electrons

Reduction Half-Reaction: Reduction Involves Gain of electrons

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Net Redox Reaction: 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 formonoatomic 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

CaBr2; Ca = +2, Br = -1

The 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.

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 = +18. 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 = -2H = +1

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

C = +4

IF7

F = -17x(-1) + ? = 0I = +7

NaIO3

Na = +1O = -23x(-2) + 1 + ? = 0I = +5

Determination of Oxidizing and Reducing Agents

Determine oxidation number for all atoms in both the reactants and products.Look at same atom in reactants and products and see if oxidation number increased or decreased.

If oxidation number decreased:substance reduced Oxidizing AgentIf oxidation number increased; substance oxidized Reducing Agent

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

Cu+2

Zn

time Zn+2Cu

Spontaneous Redox Reaction

Gets Smaller Gets Larger

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.

Voltaic CellAnode: 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 NotationCell Notation1. Anode2. Salt Bridge3. Cathode

Anode | Salt Bridge | Cathode

| : symbol is used whenever there is a different phase

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

• cell voltage

• electromotive force (emf)

• cell potential000reductionoxidationCell EEE

E Units: Volts Volt (V) = Joule (J) Coulomb ( C )

Standard Electrode PotentialsStandard Electrode PotentialsStandard 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 reaction0 E°CELL equilibrium- E°CELL nonspontaneous reaction

More positive E°CELL stronger oxidizing agent or more likely to be reduced

• 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 G0

echworkECell arg

UnitsWork: JouleCharge (Q): CoulombEcell : Volts

Charge = nF Faraday (F): charge on 1 mole e-F = 96485 C/mole

Work = (charge)Ecell = -nFEcell

G = work (maximum)

G = -nFEcell

Relating ERelating EooCELLCELL to the Equilibrium Constant, K to the Equilibrium Constant, K

G0 = -RT ln K

G0 = -nFE0cell

-RT ln K = -nFE0cell

KnFRTECell ln0

0257.0

96485

29831.8

moleC

KmolKJ

FRT

Kn

Kn

ECell log0592.0ln0257.00

Effect of Concentration on Cell Potential

G = G0 + RTlnQ

G0 = -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

Cathodic Protection

Batteries

Dry cell

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

C: 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

Batteries

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. BatteryBattery: 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 CellA 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)

Faraday’s Constant Redox Eqn

Molar Mass

Charge =(Current)(Time)