Redox Reactions and Electrochemistry - uniroma2.it · Redox Reactions and Electrochemistry...

transcript

Redox Reactions and Electrochemistry

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

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Oxidation half-reaction (lose e-)

Reduction half-reaction (gain e-)

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-

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

and O22- it is –1.

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.

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.

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

HCO3−

O = −2 H = +1

3x(−2) + 1 + ? = −1

C = +4

Identify the oxidation numbers of all the atoms in HCO3

Balancing Redox Equations

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

The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?

Fe2+ + Cr2O72- Fe3+ + Cr3+

2.  Separate the equation into two half-reactions.

Oxidation:

Cr2O72- Cr3+

+6 +3 Reduction:

Fe2+ Fe3+ +2 +3

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

Cr2O72- 2Cr3+

4.  Add electrons to one side of each half-reaction to balance the charges on the half-reaction.

6e- +14H+ + Cr2O72- 2Cr3+

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

5.  For reactions in acid, add H+ to balance electronic charge and H2O to balance O atoms and H atoms

Fe2+ Fe3+ + 1e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 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-

6e- + Cr2O72- 2Cr3+

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

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

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

Reduction:

14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 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.

Galvanic Cells

spontaneous redox reaction

anode oxidation

cathode reduction

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

•  cell voltage

•  electromotive force (emf)

•  cell potential

Cell Diagram

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

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

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

anode cathode salt bridge

phase boundary

Standard Reduction Potentials

Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

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

Zn (s) Zn2+ (1 M) + 2e- Anode (oxidation):

Cathode (reduction):

Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

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.

E° = 0 V

Standard hydrogen electrode (SHE)

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

Reduction Reaction

E0 = 0.76 V cell

Standard emf (E° ) cell

0.76 V = 0 - EZn /Zn ° 2+

EZn /Zn = -0.76 V ° 2+

Zn2+ (1 M) + 2e- Zn E° = -0.76 V

E° = EH /H - EZn /Zn cell ° ° + 2+

E° = Ecathode - Eanode cell ° °

Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)

2e- + Cu2+ (1 M) Cu (s)

H2 (1 atm) 2H+ (1 M) + 2e- Anode (oxidation):

H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)

E° = 0.34 V cell

Ecell = ECu /Cu – EH /H 2+ + 2

° ° °

0.34 = ECu /Cu - 0 ° 2+

ECu /Cu = 0.34 V 2+ °

•  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°

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

Cr3+ (aq) + 3e- Cr (s) E° = -0.74 V

Cd is the stronger oxidizer

Cd will oxidize Cr

2e- + Cd2+ (1 M) Cd (s)

Cr (s) Cr3+ (1 M) + 3e- Anode (oxidation):

2Cr (s) + 3Cd2+ (1 M) 3Cd (s) + 2Cr3+ (1 M)

E° = -0.40 – (-0.74) cell

E° = 0.34 V 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.

M(s) Mn+(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:

Where Ox and Red are oxidized and reduced forms of Red-Ox couple in equilibrium: Oxn+ + ne- Red0

E = E 0 +2.3RTnF

log [Ox][Red]

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

•  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)

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

Pt electrode

Al, Cu, Sn, inox, brass and Fe electrodes

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

Minimum 2 electrodes are required for electrochemical measurements. Dipped in electrolyte solut ion these e lec t rodes const i tu te an electrochemical cell.

Reference electrodes: Ag/AgCl and SCE

Spontaneity of Redox Reactions ΔG = -nFEcell

ΔG° = -nFEcell °

n = number of moles of electrons in reaction

F = 96,500 J

V • mol = 96,500 C/mol

ΔG° = -RT ln K = -nFEcell °

Ecell ° = RT nF

ln K (8.314 J/K•mol)(298 K)

n (96,500 J/V•mol) ln K =

= 0.0257 V n ln K Ecell °

= 0.0592 V n log K Ecell °

Spontaneity of Redox Reactions

ΔG° = -RT ln K = -nFEcell °

2e- + Fe2+ Fe

2Ag 2Ag+ + 2e- Oxidation:

Reduction:

What is the equilibrium constant for the following reaction at 25°C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq)

= 0.0257 V n ln K Ecell °

E° = -0.44 – (0.80)

E° = -1.24 V e

0.0257 V x n E° cell

= e 0.0257 V

x 2 -1.24 V

K = 1.23 x 10-42

E° = EFe /Fe – EAg /Ag ° ° 2+ +

The Effect of Concentration on Cell Emf

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

-nFE = -nFE° + RT ln Q

E = E° - ln Q RT nF

Nernst equation

At 298 K

- 0.0257 V n ln Q E ° E = - 0.0592 V

n log Q E ° E =

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)

2e- + Fe2+ 2Fe

Cd Cd2+ + 2e- Oxidation:

Reduction: n = 2

E° = -0.44 – (-0.40)

E° = -0.04 V

E° = EFe /Fe – ECd /Cd ° ° 2+ 2+

- 0.0257 V n ln Q E ° E =

- 0.0257 V 2 ln -0.04 V E = 0.010

0.60 E = 0.013

E > 0 Spontaneous

Concentration Cells

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

Electrolysis

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

Electrolysis

•  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.

Voltaic Electrolytic

Electrolytic conductors

Fig. 21.18: Car battery, both voltaic and electrochemical cell.

Increase oxidizing power

Increase reducing power

A standard electrolytic cell. A power source forces the opposite reaction

Electrolysis

(a) A silver-plated teapot. (b) Schematic of the electroplating of a spoon.

Schematic of the electroplating of a spoon.

AgNO3(aq)

The electrolysis of water produces hydrogen gas at the cathode (on the right) and oxygen gas at the anode

(on the left).

Electrolysis of water

•  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) •  Eo

cell = -0.42 -0.82 = -1.24 V

Electrolysis: Consider the electrolysis of a solution that is 1.00 M in each of CuSO4(aq) and NaCl(aq)

•  Oxidation possibilities follow: •  Cl2(g) + 2e– 2Cl–(aq) E° = +1.358 V •  S2O8

2–(aq) + 2e– 2SO42–(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

Electrolysis •  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.

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.

Application of Faraday’s law

•  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

Stoichiometry of Electrolysis  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

Fig. 21.20

Doing work with electricity.

Electrolysis and Mass Changes

charge (C) = current (A) x time (s)

1 mol e- = 96,500 C

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:

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

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

Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g)

2 mole e- = 1 mole Ca

mol Ca = 0.452 C s x 1.5 hr x 3600

s hr 96,500 C

1 mol e- x 2 mol e- 1 mol Ca x

= 0.0126 mol Ca

= 0.50 g Ca

Batteries

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)

Alkaline ba*ery •  Electrolyte is a concentrated solu5on of KOH

•  The anode is inside the ba*ery as a powder paste

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

Alkaline ba*ery

•  Lower polariza5on •  Higher dura5on •  Lower self discharge

Rechargeable Alkaline Ba*ery

•  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

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 (Silver Oxide)

High energy and stable discharge, ideal for long time operation with low A

Batteries

Anode:

Cathode:

Lead storage battery

PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e- PbSO4 (s) + 2H2O (l) 4

Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4

Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l) 4

Acidic ba*eries

•  Pb ba*eries have been first reported in 1859 •  Electrode reac5ons: Pb + H2SO4 PbSO4 + 2H+ + 2e-‐ ( -‐0.356 V) PbO2 + H2SO4 + 2H+ + 2e-‐ PbSO4 + 2H2O (1.685 V) Total reac5on: Pb + PbO2 + H2SO4 2PbSO4 + 2H2O (2.041 V)

Pb ba*eries

•  Advantages: Low cost Well known Technology Good Ah

•  Disadvantages: Low energy density Deposi5on 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 •  Alterna5ve to Ni-‐Cd cells •  Developed ader Ni-‐H2 cells, for military applica5ons

•  Electrode reac5ons: 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 Ba*ery

•  Lightest metal; •  High nega5ve standard poten5al

•  Easy to oxidize •  Unstable and not compa5ble with water

Li Ba*ery

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

Li Ba*ery •  Cathode MnO2

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

Batteries

Solid State Lithium Battery

Li-‐ion Ba*eries

Batteries

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)

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) - 25-35% (stationary)

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

Polymer Electrolyte Fuel Cell (PEMFC)

Applications :

•  Backup power

•  Portable power

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)

•  Small distributed generation

•  Transportation

•  Operating Temperature: 90-100ºC

•  Anode – Zn H2(g) + 2OH-(aq) → 2H2O + 2e-

•  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%

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).

•  Anode – Platinum (0.1 mg/Pt cm2) H2(g) → 2H+ + 2e-

•  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%

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

•  Anode: Nickel

H2(g) + CO32- → H2O(g) + CO2(g) + 2e-

•  Cathode: Nickel

•  1/2O2(g)+CO2(g)+2e-→ CO32-

•  Common electrolyte: - Carbonate salt

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

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

•  Operating Temperature: 800 -1000ºC

•  Anode: Co-ZrO2 or Ni-ZrO2 cermet •  H2(g) + O2- → H2O(l) + 2e-

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

•  Common electrolyte: - Solid zirconium oxide to which a small

amount of Yttria is added

•  System Output: 5kW – 3MW

[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

•  Fuel flexibility

•  Can use a variety of catalysts

Disadvantages :

•  High temperature enhances corrosion and breakdown of cell components

•  Slow start-up

•  Brittleness of ceramic electrolyte with thermal cycling

•  Suitable for CHP

•  Hybrid/GT cycle

•  Solid electrolyte

reduces electrolyte

management problems

ITSOFC •  Operating Temperature: 600-800ºC

•  Anode: Co-ZrO2 or Ni-ZrO2 cermet •  H2(g) + O2- → H2O(l) + 2e-

•  Cathode: Sr-doped LaMnO3 •  1/2O2(g) + 2e- → O2- •  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 is a spontaneous and irreversible electrochemical process, which results in the degrada5on of a metallic material, upon interac5on 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 kine5c (rate of the process) factors

Corrosion

The interac5on with the environment could lead to: 1. The corrosion of the metal (ac5ve condi5on): the process is both thermodynamic and kine5c favored. ΔE > 0

2. The forma5on of a protec5ve film (passive condi5on): the process is favored by thermodynamic but kine5cally inhibited

3. No modifica5on 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: O2 + 2H2O + 4e-‐ 4OH-‐

or O2 + 4H+ + 4e-‐ 2H2O

The molecular oxygen is more concentrated at the surface than in the bulk of the droplet, leading to a concentra5on cell. The oxygen reduc5on produces the hydroxide ions that lead to the rust forma5on

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 reduc5on 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 a*ack from an external inspec5on.

Temporal evolu5on of the corrosion Constant process: ex. Fe in HCl Self-‐cataly5c process: the hydrolysis of iron in the presence of Cl-‐ ion produces protons in the anodic zone that increases the corrosion rate Self-‐inhibi5ng process: the forma5on of carbonate salts in the alkaline region can produce low soluble salts that par5ally protect the metal surface from the oxygen reduc5on Passiva5ng process: the forma5on in the anodic zone of a compact oxide film that protect the metal: ex. Al

Galvanic corrosion The corrosion is produced by a junc5on of two metals having different E: the metal with lower E is oxidized. Lower the ra5o of the zone anode/cathode, higher and more penetra5ng the dissolu5on 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 characteris5c colored powder

Protec5on methods Cathodic protec5on: cathodic current or sacrificial anode Appica5on of films resistant to corrosion: metallic, non-‐metallic, polymers

Cathodic Protection of an Iron Storage Tank

Protec5on with metallic films Hot deposi5on: immersion or spray coa5ng Galvanic deposi5on: electrochemical deposi5on (problems: not homogeneous thickness) Chemical deposi5on: deposi5on of the film by redox reac5on

Protec5on with non-‐metallic films Conver5on layers: the film is formed in situ by forma5on of chemical bond with the metal surface Ex.: chromature

Protec5on with organic layers Thick films: gums or polymers Thin films: paints

Redox Reactions and Electrochemistry - uniroma2.it · Redox Reactions and Electrochemistry...

Documents