Redox Reactions and Electrochemistry

1 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Electrochemical processes are oxidation-reduction reactions in which: • the energy released by a spontaneous reaction is converted to electricity or • electrical energy is used to cause a nonspontaneous reaction to occur

0 0 2+ 2-

2Mg (s) + O2 (g) 2MgO (s)

2Mg 2Mg2+ + 4e- Oxidation half-reaction (lose e-)

- 2- - O2 + 4e 2O Reduction half-reaction (gain e )

2 Oxidation number

The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred.

1. Free elements (uncombined state) have an oxidation number of zero.

Na, Be, K, Pb, H2, O2, P4 = 0 2. In monatomic ions, the oxidation number is equal to the charge on the ion.

Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2

3. The oxidation number of oxygen is usually –2. In H2O2 2- and O2 it is –1. 3 4. 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.

5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.

6. 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. − HCO3 Identify the oxidation numbers of O = −2 H = +1 − all the atoms in HCO3 ? 3x(−2) + 1 + ? = −1 C = +4 4 Balancing Redox Equations

2+ 3+ 2- The oxidation of Fe to Fe by Cr2O7 in acid solution?

1. Write the unbalanced equation for the reaction ion ionic form.

2+ 2- 3+ 3+ Fe + Cr2O7 Fe + Cr

2. Separate the equation into two half-reactions. +2 +3 Oxidation: Fe2+ Fe3+ +6 +3 2- 3+ Reduction: Cr2O7 Cr

3. Balance the atoms other than O and H in each half-reaction.

2- 3+ Cr2O7 2Cr 5 Balancing Redox Equations

4. Add electrons to one side of each half-reaction to balance the charges on the half-reaction. Fe2+ Fe3+ + 1e- - 2- 3+ 6e + Cr2O7 2Cr 5. For reactions in acid, add H+ to balance electronic charge and H2O to balance O atoms and H atoms - + 2- 3+ 6e +14H + Cr2O7 2Cr

- + 2- 3+ 6e + 14H + Cr2O7 2Cr + 7H2O 6. If necessary, equalize the number of electrons in the two half- reactions by multiplying the half-reactions by appropriate coefficients. 6Fe2+ 6Fe3+ + 6e-

- + 2- 3+ 6 6e + 14H + Cr2O7 2Cr + 7H2O Balancing Redox Equations

7. Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel.

Oxidation: 6Fe2+ 6Fe3+ + 6e-

- + 2- 3+ Reduction: 6e + 14H + Cr2O7 2Cr + 7H2O

+ 2- 2+ 3+ 3+ 14H + Cr2O7 + 6Fe 6Fe + 2Cr + 7H2O 8. Verify that the number of atoms and the charges are balanced. 14x1 – 2 + 6 x 2 = 24 = 6 x 3 + 2 x 3 9. For reactions in basic solutions, add OH- to instead of H+ to balance electronic charges. 10. Balance the reaction in the molecular form.

7 Galvanic Cells

anode cathode oxidation reduction

spontaneous redox reaction

8 Galvanic Cells The difference in electrical potential between the anode and cathode is called: • cell voltage • electromotive force (emf) • cell potential Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq) [Cu2+] = 1 M and [Zn2+] = 1 M Cell Diagram phase boundary Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)

anode salt bridge cathode 9 Standard Reduction Potentials

2+ + Zn (s) | Zn (1 M) || H (1 M) | H2 (1 atm) | Pt (s) Anode (oxidation): Zn (s) Zn2+ (1 M) + 2e-

- + Cathode (reduction): 2e + 2H (1 M) H2 (1 atm) + 2+ Zn (s) + 2H (1 M) Zn + H2 (1 atm) 10 Standard Reduction Potentials

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

Reduction Reaction

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

E° = 0 V

Standard hydrogen electrode (SHE) 11 Standard Reduction Potentials

0 Ecell = 0.76 V

° Standard emf (Ecell )

° ° ° Ecell = Ecathode - Eanode

2+ + Zn (s) | Zn (1 M) || H (1 M) | H2 (1 atm) | Pt (s) ° ° ° + 2+ Ecell = EH /H 2 - EZn /Zn ° 0.76 V = 0 - EZn 2+ /Zn ° EZn 2+ /Zn = -0.76 V Zn2+ (1 M) + 2e- Zn E° = -0.76 V 12 Standard Reduction Potentials

° Ecell = 0.34 V ° ° ° Ecell = Ecathode - Eanode

E° = E ° 2+ – E ° + cell Cu /Cu H /H 2 ° 0.34 = ECu2+ /Cu - 0 ° ECu 2+ /Cu = 0.34 V

+ 2+ Pt (s) | H2 (1 atm) | H (1 M) || Cu (1 M) | Cu (s) + - Anode (oxidation): H2 (1 atm) 2H (1 M) + 2e Cathode (reduction): 2e- + Cu2+ (1 M) Cu (s)

2+ + H2 (1 atm) + Cu (1 M) Cu (s) + 2H (1 M) 13 • E° is for the reaction as written • The more positive E° the greater the tendency for the substance to be reduced • The half-cell reactions are reversible • The sign of E° changes when the reaction is reversed • Changing the stoichiometric coefficients of a half-cell reaction does not change ° the value of E 14 What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr electrode in a 1.0 M Cr(NO3)3 solution?

Cd2+ (aq) + 2e- Cd (s) E° = -0.40 V Cd is the stronger oxidizer

Cr3+ (aq) + 3e- Cr (s) E° = -0.74 V Cd will oxidize Cr Anode (oxidation): Cr (s) Cr3+ (1 M) + 3e- x 2 Cathode (reduction): 2e- + Cd2+ (1 M) Cd (s) x 3 2Cr (s) + 3Cd2+ (1 M) 3Cd (s) + 2Cr3+ (1 M)

° ° ° Ecell = Ecathode - Eanode ° Ecell = -0.40 – (-0.74) ° Ecell = 0.34 V 15 The electrochemical cell

The electrochemical cell

Redox reactions can be used to generate electric current Electrode processes Electrode processes The metallic electrode is dipped into a solution containing a salt of the metal. Some atoms of the metal can leave the electrode and form the cation in solution, leaving electrons in the metal. This form a double layer of opposite charges to the electrode surface. The electrochemical potential of the metal and its ion should be the same at the equilibrium.

n+ - M(s)  M (aq) + ne

There is the formation of an electric potential proportional at the ion concentration in solution

+ - - + + - - + + - - + + - - + + - - + + - - - + + + + Electrode potential Nernst law

+ - M (aq) + e  M(s)

GM = G°M + RTlnaM GM+ = G°M+ + RTlnaM+

ΔG = G°M + RTlnaM - G°M+ + RTlnaM+

ΔG = ΔG° + RTln(aM/aM+)

ΔG = -nFE

-nFE = ΔG° + RTln(aM/aM+)

E = -ΔG°/nF + RT/nF ln(aM+/aM) aM = 1

E = E° + RT/nF ln aM+

The Nernst equation

The potential of an electrode is expressed by the Nernst law:

2.3RT [Ox] E = E 0 + log nF [Re d] Where Ox and Red are oxidized and reduced forms of Red-Ox couple in equilibrium: ! Oxn+ + ne-  Red0

R is the universal gas constant, T is the absolute temperature in Kelvins, n is a number of electrons transferred in reaction, F is Faraday constant (~ 96500 C) The electromotive force

• It is useful to separate the overall redox reaction in two separated processes: the oxidation and reduction semi- reaction.

• In the electrochemical cell we have two electrodes and we indicate as Cathode the electrode where the reductions occur and Anode the electrode of oxidation processes.

• The electromotive force (EMF) of the cell is the electric potential difference among the cathode and anode.

The Standard Hydrogen Electrode

We cannot know the absolute potential of a single electrode (it is not possible to measure half reaction), so the E° = 0 V was assigned to the semi-reaction

+ - 2 H3O + 2e  H2 Reference electrodes: Ag/AgCl and SCE Metallic electrodes

3 main groups: • First kind - wire of active metal immersed in solution, contained the ions of this metal (Cu, Zn, Co, Fe, etc)

Al, Cu, Sn, inox, brass and Fe electrodes • Second kind – wire of metal covered by precipitate of hardly soluble salt or oxide: M/Mn+ (Ag/AgCl for instance)

• Third kind - inert metallic electrodes (Pt, Au, etc)

Pt electrode Electrochemical cell

Minimum 2 electrodes are required for electrochemical measurements. Dipped in electrolyte solution these electrodes constitute an electrochemical cell.

INDICATOR (or WORKING) electrode is an electrode responding to a target analyte REFERENCE electrode has a stable well defined potential value, independent on analyzed solution composition Reference electrodes: Ag/AgCl and SCE Spontaneity of Redox Reactions

ΔG = -nFEcell n = number of moles of electrons in reaction J ΔG° = -nFE ° F = 96,500 = 96,500 C/mol cell V • mol

° ° ΔG = -RT ln K = -nFEcell RT (8.314 J/K•mol)(298 K) E ° = ln K = ln K cell nF n (96,500 J/V•mol) 0.0257 V E ° = ln K cell n 0.0592 V E ° = log K cell n

29 Spontaneity of Redox Reactions

° ° ΔG = -RT ln K = -nFEcell

30 What is the equilibrium constant for the following reaction at 25°C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq) 0.0257 V E ° = ln K cell n

Oxidation: 2Ag 2Ag+ + 2e- n = 2 Reduction: 2e- + Fe2+ Fe

° ° ° E = EFe 2+ /Fe – EAg + /Ag ° E = -0.44 – (0.80) ° Ecell x n -1.24 V x 2 E° = -1.24 V 0.0257 V 0.0257 V K = e = e

K = 1.23 x 10-42

31 The Effect of Concentration on Cell Emf

ΔG = ΔG° + RT ln Q ΔG = -nFE ΔG° = -nFE °

-nFE = -nFE° + RT ln Q

Nernst equation

RT E = E° - ln Q nF

At 298 K

0.0257 V 0.0592 V E = E ° - ln Q E = E ° - log Q n n

32 Will the following reaction occur spontaneously at 250C if [Fe2+] = 0.60 M and [Cd2+] = 0.010 M? Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)

Oxidation: Cd Cd2+ + 2e- n = 2 Reduction: 2e- + Fe2+ 2Fe ° ° ° E = EFe 2+ /Fe – ECd 2+ /Cd E° = -0.44 – (-0.40) 0.0257 V E = E ° - ln Q n E° = -0.04 V 0.0257 V 0.010 E = -0.04 V - ln 2 0.60 E = 0.013

E > 0 Spontaneous 33 Concentration Cells

Galvanic cell from two half-cells composed of the same material but differing in ion concentrations.

34 Electrolysis

• Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.

• Previously our lectures on electrochemistry were

involved with voltaic cells i.e. cells with Ecell > 0 and ΔG < 0 that were spontaneous reactions. • Today we discuss electrochemical cells where

Ecell < 0 and ΔG > 0 that are non-spontaneous reactions and require electricity for the reactions to take place. We can take a voltaic cell and reverse the electrodes to make an electrochemical cell.

Houghton Mifflin Company and G. 37 Hall. All rights reserved. Houghton Mifflin Company and G. 38 Hall. All rights reserved. Houghton Mifflin Company and G. 39 Hall. All rights reserved. Electrolytic conductors

Increase reducing power

AgNO3(aq)

• At the anode (oxidation): - - • 2H2O(l) + 2e  H2(g) + 2OH (aq) E= -0.42V

• At the cathode (reduction): + - • O2(g) + 4H (aq) + 4e  2H2O(l) E= 0.82V • Overall reaction after multiplying anode reaction by 2,

• 2H2O(l)  2H2(g) + O2(g) o • E cell = -0.42 -0.82 = -1.24 V

• Oxidation possibilities follow: – – • Cl2(g) + 2e  2Cl (aq) E° = +1.358 V 2– – 2– • S2O8 (aq) + 2e  2SO4 (aq) E° = +2.010 V + – • O2(g) + 4H (aq) + 4e  2H2O E° = +1.229 V

• Reduction possibilities follow: • Na+(aq) + e–  Na(s) E° = –2.713 V • Cu2+(aq) + 2e–  Cu(s) E° = +0.337 V – – • 2H2O + 2e  H2(g) + 2OH (aq) E° = -0.428 V

• We would choose the production of O2(g) and Cu(s). • But the voltage for producing O2(g) from solution is considerably higher than the standard potential, because of the high activation energy needed to form O2(g). • The voltage for this half cell seems to be closer to –1.5 V in reality.

• The result then is the production of Cl2(g) and Cu(s). – – anode, oxidation: Cl2(g) + 2e  2Cl (aq) E° = +1.358 V • cathode, reduction: Cu2+(aq) + 2e–  Cu(s) E° = +0.337 V

• overall: CuCl2(aq)  Cu(s) + Cl2(g) E = –1.021 V • We must apply a voltage of more than +1.021 V to cause this reaction to occur.

Houghton Mifflin Company and G. 53 Hall. All rights reserved. Houghton Mifflin Company and G. 54 Hall. All rights reserved. Stoichiometry of electrolysis: Relation between amounts of charge and product

• Faraday’s law of electrolysis relates to the amount of substance produced at each electrode is directly proportional to the quantity of charge flowing through the cell (half reaction). • Each balanced half-cell shows the relationship between moles of electrons and the product.

• 1. First balance the half-reactions to find number of moles of electrons needed per mole of product. • 2. Use Faraday constant (F = 9.65E4 C/mol e-) to find corresponding charge. • 3. Use the molar mass of substance to find the charge needed for a given mass of product. – 1 ampere = 1 coulomb/second or 1 A = 1 C/s – A x s = C

 How much chemical change occurs with the flow of a given current for a specified time?

• current and time → quantity of charge → • moles of electrons → moles of analyte → • grams of analyte

1 mol e- = 96,500 C

60 How much Ca will be produced in an electrolytic cell of molten CaCl2 if a current of 0.452 A is passed through the cell for 1.5 hours?

- - Anode: 2Cl (l) Cl2 (g) + 2e

Cathode: Ca2+ (l) + 2e- Ca (s)

2+ - Ca (l) + 2Cl (l) Ca (s) + Cl2 (g) 2 mole e- = 1 mole Ca

C s 1 mol e- 1 mol Ca mol Ca = 0.452 x 1.5 hr x 3600 x x s hr 96,500 C 2 mol e- = 0.0126 mol Ca

Dry cell

Leclanché cell

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

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

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

63 Alkaline baery

• Electrolyte is a concentrated soluon of KOH • The anode is inside the baery as a powder paste

• MnO2 is pasted with graphite around the Zn anode and contacted with the external steel electrode Alkaline baery

• Lower polarizaon • Higher duraon • Lower self discharge Rechargeable Alkaline Baery

• Rechargeable Alkaline Manganese (RAM) cell • The interest is due to the higher A: Alkaline manganese 2 – 3 Ah Nickel / cadmium 0.5 – 1.0 Ah Nickel / metal hydride 1 – 1.5 Ah The number of charge-discharge is lower than usual Ni/Cd or Ni/MH cells Button Batteries High energy and stable discharge, ideal for long time operation with low A

Mercury Battery

(Silver Oxide)

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

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

67 Batteries

Lead storage battery

2- - Anode: Pb (s) + SO 4 (aq) PbSO4 (s) + 2e

+ 2- - Cathode: PbO2 (s) + 4H (aq) + SO4 (aq) + 2e PbSO4 (s) + 2H2O (l)

+ 2- Pb (s) + PbO2 (s) + 4H (aq) + 2SO4 (aq) 2PbSO4 (s) + 2H2O (l)

68 Acidic baeries

• Pb baeries have been ﬁrst reported in 1859 • Electrode reacons: + - Pb + H2SO4 PbSO4 + 2H + 2e ( -0.356 V) + - PbO2 + H2SO4 + 2H + 2e PbSO4 + 2H2O (1.685 V) Total reacon:

Pb + PbO2 + H2SO4 2PbSO4 + 2H2O (2.041 V) Pb baeries

• Advantages: Low cost Well known Technology Good Ah • Disadvantages: Low energy density

Deposion of low soluble PbSO4 Ni-Cd

. Cd + 2NiOOH + 4H2O  Cd(OH)2 + 2Ni(OH)2 H2O e.f.m. = 1.20 V

High number of cycles, reliable, low maintenance Energy density not high Cd is toxic and costly Ni-MH • Alternave to Ni-Cd cells

• Developed aer Ni-H2 cells, for military applicaons • Electrode reacons: - H2 + 2OH 2H2O + 2e - 2NiOOH + 2H2O + 2e 2Ni(OH)2 + 2OH e.f.m. = 1.2 – 1.3 V Metallic hydride is used as hydrogen source Li Baery

• Lightest metal;

• High negave standard potenal

but:

• Easy to oxidize

• Unstable and not compable with water Li Baery

• Need non aqueous solvents • Advantages: • High voltage ( >4V) • Uniform T discharge • Long shelf-life • Loss of capacity < 10% • Wide range of working T Li Baery • Cathode MnO2

• Long self discharge (up to 10 years) • Working T around -40 °C and 60 °C Batteries

Solid State Lithium Battery 76 Li-ion Baeries Batteries

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

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

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

2H2 (g) + O2 (g) 2H2O (l) 78

History

• The fuel cells had been conceived in 1839 by the British scientist Mr. William Grove.

• Developed practical applications during years 60 and 70, for NASA.

• The American astronauts consumed the water produced for the electric generators of its ships.

• These generators had constituted the first operational use of fuel cells. What they are… • Electrochemical cell that converts chemical energy into electric energy;

• It can have taxes of conversion in the order of 90%;

• Cathode + anode + electrolyte + catalyst;

• Ex.: Combustible H2 and oxidant O2 + - Anode – H2(g) → 2H (aq) + 2e + - Cathode – 1/2O2(g) + 2H (aq) + 2e → H2O(g)

• It is important the selection of the electrolyte, and the dimensions of this and the electrodes. and its operating… Types of Fuel Cells • Polymer Electrolyte Fuel Cell (PEMFC) • Alkaline Fuel Cell (AFC) • Phosphoric Acid Fuel Cell (PAFC) • Molten Carbonate Fuel Cell (MCFC) • Intermediate Temperature Solid Oxide Fuel Cell (ITSOFC) • Solid Oxide Fuel Cell (SOFC) PEMFC

• Operating Temperature: 50-100ºC

• Appropriate for electric vehicles (Automobile Industry)

• Anode – Platinum (0.4mg/Pt cm2)

+ - H2(g) → 2H + 2e • Cathode – Platinum (0.4 mg/Pt cm2)

+ - 1/2O2(g) + 2H + 2e → H2O(aq) • Common electrolyte: - Solid organic polymer poly- perfluorosulfonic acid; - Membrane of Nafion.

• System Output: < 1kW - 250kW • Efficiency Electrical:

- 53-58% (transportation) [1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm - 25-35% (stationary) Polymer Electrolyte Fuel Cell (PEMFC)

Applications :

• Backup power • Small distributed generation • Portable power • Transportation

Advantages : • Solid electrolyte reduces corrosion & electrolyte management problems • Low temperature

• Quick start-up

Disadvantages : • Requires expensive catalysts

• High sensitivity to fuel impurities

• Low temperature waste heat

• Waste heat temperature not suitable for combined heat and power (CHP) AFC

• Operating Temperature: 90-100ºC

• Anode – Zn H (g) + 2OH-(aq) → 2H O + 2e- 2 2

• Cathode – MnO2 - - 1/2O2(g) + H2O + 2e → 2OH (aq)

• Common electrolyte: - Aqueous solution of potassium hydroxide soaked in a matrix

• System Output: 10kW - 100kW

• Efficiency Electrical: 60%

[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm

Alkaline Fuel Cell (AFC)

Applications :

• Military • Space

Advantages :

• Cathode reaction faster in alkaline electrolyte, higher performance.

Disadvantages :

• Expensive removal of CO2 from fuel and air streams required (CO2 degrades the electrolyte). PAFC

• Operating Temperature: 150-200ºC

• Anode – Platinum (0.1 mg/Pt cm2) H (g) → 2H+ + 2e- 2 • Cathode – Platinum (0.5 mg/Pt cm2) + - 1/2O2(g) + 2H + 2e → H2O(aq)

• Common electrolyte: - Liquid phosphoric acid soaked in a matrix

• System Output: 50kW – 1MW (250kW module typical)

• Efficiency Electrical: 32-38%

[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm

Phosphoric Acid Fuel Cell (PAFC)

Applications :

• Distributed generation

Advantages :

• Higher overall efficiency with CHP

• Increased tolerance to impurities in hydrogen

Disadvantages :

• Requires expensive platinum catalysts

• Low current and power • Large size/weight MCFC

• Operating Temperature: 600-700ºC

• Anode: Nickel 2- - H2(g) + CO3 → H2O(g) + CO2(g) + 2e • Cathode: Nickel

- 2- • 1/2O2(g)+CO2(g)+2e → CO3 • Common electrolyte: - Carbonate salt

• System Output: < 1kW – 1MW (250kW module typical)

• Efficiency Electrical: 45-47%

[1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm

Molten Carbonate Fuel Cell (MCFC)

Applications :

• Electric utility • Large distributed generation Advantages :

• High efficiency • Fuel flexibility • Can use a variety of catalysts • Suitable for CHP

Disadvantages :

• High temperature speeds corrosion and breakdown of cell components

• Complex electrolyte management

• Slow start-up TSOFC

• Operating Temperature: 800 -1000ºC

• Anode: Co-ZrO2 or Ni-ZrO2 cermet 2- - • H2(g) + O → H2O(l) + 2e

• Cathode: Sr-doped LaMnO3 - 2- • 1/2O2(g) + 2e → O

• Common electrolyte: - Solid zirconium oxide to which a small amount of Yttria is added

• System Output: 5kW – 3MW

• Efficiency Electrical: 35-43%

[2] http://www.treehugger.com/files/2007/06/biogas-powered_fuel_system.php Solid Oxide Fuel Cell (SOFC)

Applications :

• Auxiliary power • Electric utility • Large distributed generation Advantages :

• High efficiency • Solid electrolyte • Suitable for CHP • Fuel flexibility reduces electrolyte • Hybrid/GT cycle • Can use a variety of catalysts management problems

Disadvantages :

• High temperature enhances corrosion and breakdown of cell components

• Slow start-up • Brittleness of ceramic electrolyte with thermal cycling ITSOFC

• Operating Temperature: 600-800ºC

• Anode: Co-ZrO2 or Ni-ZrO2 cermet 2- - • H2(g) + O → H2O(l) + 2e

• Cathode: Sr-doped LaMnO3 - 2- • 1/2O2(g) + 2e → O

• Lower temperatures ⇒ increase the internal resistance of the cell

• Common electrolyte: - Solid zirconium oxide to which a small amount of Yttria is added

• System Output: 5kW – 3MW

• Efficiency Electrical: 35-43% [1] http://www.cogeneration.net/molten_carbonate_fuel_cells.htm Applications Corrosion

Corrosion is a spontaneous and irreversible electrochemical process, which results in the degradaon of a metallic material, upon interacon with the environment.

The corrosion could occur in the presence or in the absence of water: The ﬁrst one is called wet corrosion, the second dry corrosion

As for all the chemical processes, the corrosion depends on both thermodynamic (spontaneous or not process) and kinec (rate of the process) factors

The interacon with the environment could lead to:

1.The corrosion of the metal (acve condion): the process is both thermodynamic and kinec favored. ΔE > 0

2.The formaon of a protecve ﬁlm (passive condion): the process is favored by thermodynamic but kinecally inhibited

3.No modiﬁcaon of the metal: : the process is not thermodynamic favored. ΔE < 0 The corrosion is an electrochemical process, where a cathode and an anode are formed The metal is oxidized in the anodic region and leaves the electrons that migrate to the cathodic region, the corrosive region, where molecular oxygen is reduced

Anodic process:

Me  Men+ + ne-

Cathodic process:

O + 2H O + 4e-  4OH- 2 2 or

+ - O2 + 4H + 4e  2H2O

The molecular oxygen is more concentrated at the surface than in the bulk of the droplet, leading to a concentraon cell. The oxygen reducon produces the hydroxide ions that lead to the rust formaon

This eﬀect produces the ring morphology for the metal corrosion The corrosion can be:

Generalized: the anodic zone is big, while the cathodic zone is small

Localized: is the reverse case of the generalized corrosion. It is the most dangerous Generalized corrosion

This corrosion interests all the metallic surface and leads to a reducon of the metal thickness

Uniform

Not uniform Localized corrosion

This corrosion interests only small parts of the metal surface and it is the most dangerous because it is impossible to evaluate the gravity of the corrosive aack from an external inspecon. Temporal evoluon of the corrosion

Constant process: ex. Fe in HCl

Self-catalyc process: the hydrolysis of iron in the presence of Cl- ion produces protons in the anodic zone that increases the corrosion rate

Self-inhibing process: the formaon of carbonate salts in the alkaline region can produce low soluble salts that parally protect the metal surface from the oxygen reducon

Passivang process: the formaon in the anodic zone of a compact oxide ﬁlm that protect the metal: ex. Al Galvanic corrosion

The corrosion is produced by a juncon of two metals having diﬀerent E: the metal with lower E is oxidized. Lower the rao of the zone anode/cathode, higher and more penetrang the dissoluon of the less noble metal. Example of corrosion: the brass

The brass is an alloy of copper and zinc: the zinc is oxidized and copper forms the characterisc colored powder Protecon methods

Cathodic protecon: cathodic current or sacriﬁcial anode

Appicaon of ﬁlms resistant to corrosion: metallic, non-metallic, polymers Cathodic Protection of an Iron Storage Tank

107 Protecon with metallic ﬁlms

Hot deposion: immersion or spray coang

Galvanic deposion: electrochemical deposion (problems: not homogeneous thickness)

Chemical deposion: deposion of the ﬁlm by redox reacon Protecon with non-metallic ﬁlms

Converon layers: the ﬁlm is formed in situ by formaon of chemical bond with the metal surface

Ex.: chromature

Protecon with organic layers

Thick ﬁlms: gums or polymers

Thin ﬁlms: paints