Transcript Chemical Corrosion

Chemical vs. Electrochemical
Reactions
 Chemical reactions are those in which elements are
added or removed from a chemical species.
 Electrochemical reactions are chemical reactions in
which not only may elements may be added or removed
from a chemical species but at least one of the species
undergoes a change in the number of valance electronS.
 Corrosion processes are electrochemical in nature.
Chemical Corrosion
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 Chemical corrosion: Removal of atoms from a material by virtue of
the solubility or chemical reaction between the material and the
surrounding liquid.
EXAMPLES:
 Dezincification: A special chemical corrosion process by which both
zinc and copper atoms are removed from brass, but the copper is
replated back onto the metal.
 Graphitic corrosion: A special chemical corrosion process by which
iron is leached from cast iron, leaving behind a weak, spongy mass of
graphite.
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Photomicrograph of a copper deposit in brass, showing the effect of
dezincification (x50).
Electrochemical Corrosion
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 Electrochemical corrosion - Corrosion produced by the
development of a current in an electrochemical cell that removes
ions from the material.
 Electrochemical cell - A cell in which electrons and ions can flow by
separate paths between two materials, producing a current which,
in turn, leads to corrosion or plating.
 Oxidation reaction - The anode reaction by which electrons are
given up to the electrochemical cell.
 Reduction reaction - The cathode reaction by which electrons are
accepted from the electrochemical cell.
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The components in an electrochemical cell: (a) a simple electrochemical cell and
(b) a corrosion cell between a steel water pipe and a copper fitting.
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The anode and cathode reactions in typical electrolytic corrosion cells:
(a) the hydrogen electrode, (b) the oxygen electrode, and (c) the water electrode.
Electrochemistry
Thermodynamics at the electrode
Redox (Review)
Oxidation is...
 Loss of electrons
Reduction is...
 Gain of electrons
Oxidizing agents oxidize and are reduced
Reducing agents reduce and are oxidized
Redox Review (Cu-zn)
 Zn displaces Cu from CuSO4(aq)
 In direct contact the enthalpy of reaction is dispersed as
heat, and no useful work is done
 Redox process:
 Zn is the reducing agent
 Cu2+ is the oxidizing agent
2
Zn( s)  Zn (aq)  2e
Cu 2 (aq)  2e  Cu(s)
Separating the combatants
 Each metal in touch with a solution of its own ions
 External circuit carries electrons transferred during the redox process
 A “salt bridge” containing neutral ions completes the internal circuit.
 The energy released by the reaction in the cell can perform useful work –
like lighting a bulb
Labelling the parts
Cell notation
 Anode on left, cathode on right
 Electrons flow from left to right
 Oxidation on left, reduction on right
 Single vertical = electrode/electrolyte boundary
 Double vertical = salt bridge
Anode:
Zn →Zn2+ + 2e
Cathode:
Cu2+ + 2e →Cu
Odes to a galvanic cell
 Cathode
 Where reduction occurs
 Where electrons are
consumed
 Where positive ions migrate
to
 Has positive sign
 Anode
 Where oxidation occurs
 Where electrons are
generated
 Where negative ions migrate
to
 Has negative sign
The role of inert electrodes
Fe(s)  2Fe3 (aq)  3Fe2 (aq)
 Not all cells start with elements as the redox agents
 Consider the cell
 Fe can be the anode but Fe3+ cannot be the cathode.
 Use the Fe3+ ions in solution as the “cathode” with an inert metal such as Pt
Anode
Catho
de
Oxidati
on
Reduct
ion
Connections: cell potential and free
energy
 The cell in open circuit generates an electromotive force (emf) or
potential or voltage. This is the potential to perform work
 Energy is charge moving under applied voltage
1J  1C 1V
Relating free energy and cell
potential
The Faraday :
F = 96 485 C/mole
Standard conditions (1 M, 1 atm, 25°C)
G  nFE
G  nFE
Standard Reduction Potentials
 The total cell potential is the sum of the potentials for the
two half reactions at each electrode
 Ecell = Ecath + Ean
 From the cell voltage we cannot determine the values of
either – we must know one to get the other
 Enter the standard hydrogen electrode (SHE)
 All potentials are referenced to the SHE (EH=0 V)
Unpacking the SHE
 The SHE consists of a Pt electrode in contact with
H2(g) at 1 atm in a solution of 1 M H+(aq).
 The voltage of this half-cell is defined to be 0 V.
 An experimental cell containing the SHE half-cell
with other half-cell gives voltages which are the
standard potentials for those half-cells
Ecell = 0 + Ehalf-cell
Zinc half-cell with SHE
Cell measures 0.76 V
Standard potential for
Zn(s) = Zn2+(aq) + 2e : 0.76 V
Where there is no SHE
 In this cell there is no SHE and the measured voltage is 1.10 V
2
2
Zn Zn (aq) C u (aq) Cu
2
2
Zn( s )  Cu (aq)  Zn (aq)  Cu ( s )
Zn(s)  Zn 2 (aq)  2e, E o  0.76V
2
Cu (aq)  2e  Cu(s), E  0.34V
o
Standard reduction potentials
 Any half reaction can be written in two ways:
 Oxidation:
M = M+ + e (+V)
 Reduction:
M+ + e = M (-V)
 Listed potentials are standard reduction potentials
Applying standard reduction
potentials
 Consider the reaction
Zn(s)  2 Ag  (aq)  Zn 2 (aq)  2 Ag (s)
 What is the cell potential?
 The half reactions are:
Ag  (aq)  e  Ag ( s)
Zn( s )  Zn 2 (aq)  2e
 E° = 0.80 V – (-0.76 V) = 1.56 V
 NOTE: Although there are 2 moles of Ag reduced for each
mole of Zn oxidized, we do not multiply the potential by 2.
Extensive VS intensive
 Free energy is extensive property so need to multiply by no of moles
involved
G  nFE
 But to convert to E we need to divide by no of electrons involved


G
E 

 E is an intensive property
nF
The Nernst equation
 Working in nonstandard conditions

G  G  RT ln Q

 nFE  nFE  RT ln Q

E  E  RT

nF
ln Q
E  E  0.0592 log Q
n
Electrode potentials and pH
 For the cell reaction
 The Nernst equation

H 2 ( g )  2H (aq)  2e
EH
2 2 H

E

H 2 2 H 
0.06V

n


 2

H
 log

pH 2

 Half-cell potential is proportional to pH
EH
2 2 H


  
0.06V
log H 
n
2




The pH meter is an electrochemical cell
Overall cell potential is proportional to pH
In practice, a hydrogen electrode is impractical
Ecell  0.06V  pH   Eref
pH 
Ecell  Eref
0.06V
Cell potentials and equilibrium
G  nFE
So then

G   RT ln K

nFE   RT ln K
and E   RT
2.303RT
ln K 
log 10 K
nF
nF
Cell potential a convenient way to
measure K
Many pathways to one ending
Measurement of K from different experiments
C  D
a
b
A B
c
 Concentration data
 Thermochemical data
 Electrochemical data
d

G   RT ln K

nFE   RT ln K
Summary
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 Electrode potential - Related to the tendency of a material to
corrode. The potential is the voltage produced between the material
and a standard electrode.
 emf series - The arrangement of elements according to their
electrode potential, or their tendency to corrode.
 Nernst equation - The relationship that describes the effect of
electrolyte concentration on the electrode potential in an
electrochemical cell.
 Faraday’s equation - The relationship that describes the rate at
which corrosion or plating occurs in an electrochemical cell.
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The half-cell used to measured the
electrode potential of copper under
standard conditions.
The electrode
potential of copper is the potential
difference between it and the standard
hydrogen electrode in an open circuit.
Since E0 is great than zero, copper is
cathodic compared with the hydrogen
electrode.
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Example
HalfCell Potential for Copper
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Suppose 1 g of copper as Cu2+ is dissolved in 1000 g of water to produce an
electrolyte. Calculate the electrode potential of the copper half-cell in this
electrolyte.
The atomic mass of copper is 63.54 g/mol.
Example 22.1 SOLUTION
From chemistry, we know that a standard 1-M solution of Cu2+ is obtained when we
add 1 mol of Cu2+ (an amount equal to the atomic mass of copper) to 1000 g of
water. The atomic mass of copper is 63.54 g/mol. The concentration of the solution
when only 1 g of copper is added must be:
From the Nernst equation, with n = 2 and E0 = +0.34 V:
Example
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Design of a Copper Plating Process
Design a process to electroplate a 0.1-cm-thick layer of copper onto a
1 cm  1 cm cathode surface.
SOLUTION
In order for us to produce a 0.1-cm-thick layer on a 1 cm2 surface
area, the weight of copper must be:
From Faraday’s equation, where MCu = 63:54 g/mol and n = 2:
SOLUTION
Therefore, we might use several different combinations of current and
time to produce the copper plate:
Our choice of the exact combination of current and time might be
made on the basis of the rate of production and quality of the
copper plate.
A current of ~ 1 A and a time of ~ 45 minutes are not uncommon in
electroplating operations.
Example
Corrosion of Iron
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An iron container 10 cm  10 cm at its base is filled to a height of 20 cm
with a corrosive liquid. A current is produced as a result of an electrolytic
cell, and after 4 weeks, the container has decreased in weight by 70 g.
Calculate (1) the current and (2) the current density involved in the
corrosion of the iron.
SOLUTION
1. The total exposure time is:
From Faraday’s equation, using n = 2 and M = 55.847 g/mol:
SOLUTION
2. The total surface area of iron in contact with the corrosive liquid
and the current density are:
Example
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Copper-Zinc Corrosion Cell
Suppose that in a corrosion cell composed of copper and zinc, the current
density at the copper cathode is 0.05 A/cm2. The area of both the copper
and zinc electrodes is 100 cm2. Calculate (1) the corrosion current, (2) the
current density at the zinc anode, and (3) the zinc loss per hour.
Example 22.4 SOLUTION
1. The corrosion current is:
2. The current in the cell is the same everywhere. Thus:
SOLUTION
3. The atomic mass of zinc is 65.38 g/mol. From Faraday’s equation:
Polarization
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 Polarization - Changing the voltage between the anode and cathode
to reduce the rate of corrosion.
–
–
–
Activation polarization is related to the energy required to cause
the anode or cathode reaction
Concentration polarization is related to changes in the
composition of the electrolyte
Resistance polarization is related to the electrical resistivity of
the electrolyte.
Types of Electrochemical Corrosion
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 Intergranular corrosion - Corrosion at grain boundaries because grain
boundary segregation or precipitation produces local galvanic cells.
 Stress corrosion - Deterioration of a material in which an applied
stress accelerates the rate of corrosion.
 Oxygen starvation - In the concentration cell, low-oxygen regions of
the electrolyte cause the underlying material to behave as the anode
and to corrode.
 Crevice corrosion - A special concentration cell in which corrosion
occurs in crevices because of the low concentration of oxygen.
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Example of microgalvanic cells in two-phase alloys: (a) In steel, ferrite is anodic to
cementite. (b) In austenitic stainless steel, precipitation of chromium carbide makes the
low Cr austenite in the grain boundaries anodic.
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Photomicrograph of intergranular corrosion in a zinc die casting. Segregation of
impurities to the grain boundaries produces microgalvanic corrosion cells (x50).
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Examples of stress cells. (a) Cold work required to bend a steel bar introduces high residual
stresses at the bend, which then is anodic and corrodes. (b) Because grain boundaries have a
high energy, they are anodic and corrode.
Corrosion of a Soldered Brass Fitting
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A brass fitting used in a marine application is joined by soldering with leadtin solder. Will the brass or the solder corrode?
SOLUTION
From the galvanic series, we find that all of the copper-based alloys are
more cathodic than a 50% Pb-50% Sn solder. Thus, the solder is the anode
and corrodes. In a similar manner, the corrosion of solder can contaminate
water in freshwater plumbing systems with lead.
Example
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Corrosion of Cold-Drawn Steel
A cold-drawn steel wire is formed into a nail by additional deformation,
producing the point at one end and the head at the other. Where will the
most severe corrosion of the nail occur?
SOLUTION
Since the head and point have been cold-worked an additional amount
compared with the shank of the nail, the head and point serve as anodes
and corrode most rapidly.
Concentration cells: (a) Corrosion occurs beneath a water droplet on a steel plate due to low
oxygen concentration in the water. (b) Corrosion occurs at the tip of a crevice because of
limited access to oxygen.
Example Corrosion of Crimped Steel
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Two pieces of steel are joined mechanically by crimping the edges.
Why would this be a bad idea if the steel is then exposed to water? If
the water contains salt, would corrosion be affected?
SOLUTION
By crimping the steel edges, we produce a crevice. The region in the
crevice is exposed to less air and moisture, so it behaves as the anode
in a concentration cell. The steel in the crevice corrodes.
Salt in the water increases the conductivity of the water, permitting
electrical charge to be transferred at a more rapid rate. This causes a
higher current density and, thus, faster corrosion due to less
resistance polarization.
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(a) Bacterial cells growing in a
colony (x2700). (b) Formation of a
tubercule and a pit under a
biological colony.
Protection Against Electrochemical
Corrosion
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 Inhibitors - Additions to the electrolyte that preferentially migrate
to the anode or cathode, cause polarization, and reduce the rate of
corrosion.
 Sacrificial anode - Cathodic protection by which a more anodic
material is connected electrically to the material to be protected.
The anode corrodes to protect the desired material.
 Passivation - Producing strong anodic polarization by causing a
protective coating to form on the anode surface and to thereby
interrupt the electric circuit.
Example
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Effect of Areas on Corrosion Rate for Copper-Zinc
Couple
Consider a copper-zinc corrosion couple. If the current density at the
copper cathode is 0.05 A/cm2, calculate the weight loss of zinc per
hour if (1) the copper cathode area is 100 cm2 and the zinc anode
area is 1 cm2 and (2) the copper cathode area is 1 cm2 and the zinc
anode area is 100 cm2.
SOLUTION
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1. For the small zinc anode area:
2. For the large zinc anode area:
The rate of corrosion of the zinc is reduced significantly when the zinc
anode is much larger than the cathode.
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Zinc-plated steel and tin-plated steel are protected differently. Zinc protects steel
even when the coating is scratched, since zinc is anodic to steel. Tin does not
protect steel when the coating is disrupted, since steel is anodic with respect to tin.
Example
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Design of a Corrosion Protection System
Steel troughs are located in a field to provide drinking water for a herd
of cattle. The troughs frequently rust through and must be replaced.
Design a system to prevent or delay this problem.
SOLUTION
We might, for example, fabricate the trough using stainless steel or
aluminum. Either would provide better corrosion resistance than the
plain carbon steel, but both are considerably more expensive than the
current material.
We might suggest using cathodic protection; a small magnesium anode
could be attached to the inside of the trough. The anode corrodes
sacrificially and prevents corrosion of the steel.
SOLUTION (Continued)
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Another approach would be to protect the steel trough using a suitable
coating. Painting the steel (that is, introducing a protective polymer
coating) and, using a tin-plated steel, provides protection as long as the
coating is not disrupted.
The most likely approach is to use a galvanized steel, taking advantage
of the protective coating and the sacrificial behavior of the zinc.
Corrosion is very slow due to the large anode area, even if the coating
is disrupted. Furthermore, the galvanized steel is relatively
inexpensive, readily available, and does not require frequent
inspection.
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The standard free energy of formation of selected oxides as a function of
temperature. A large negative free energy indicates a more stable oxide.
Example
59
Chromium-Based Steel Alloys
Explain why we should not add alloying elements such as chromium
to pig iron before the pig iron is converted to steel in a basic oxygen
furnace at 1700oC.
SOLUTION
In a basic oxygen furnace, we lower the carbon content of the metal
from about 4% to much less than 1% by blowing pure oxygen
through the molten metal. If chromium were already present before
the steel making began, chromium would oxidize before the carbon,
since chromium oxide has a lower free energy of formation (or is
more stable) than carbon dioxide (CO2). Thus, any expensive
chromium added would be lost before the carbon was removed from
the pig iron.
Three types of oxides may form,
depending on the volume ratio
between the metal and the oxide:
(a) magnesium produces a
porous oxide film
(b) aluminum forms a protective,
adherent, nonporous oxide
film,
(c) iron forms an oxide film that
spills off the surface and
provides poor protection.
Corrosion Cells
 Galvanic cell (Dissimilar electrode cell) – dissimilar metals
 Salt concentration cell – difference in composition of aqueous
environment
 Differential aeration cell – difference in oxygen concentration
 Differential temperature cell – difference in temperature
distribution over the body of the metallic material
Dissimilar Electrode Cell
 When a cell is produced due to two
dissimilar metals it is called dissimilar
electrode cell
 Dry cell
Zn anode
 Local action cell
 A brass fitting connected to a steel pipe
 A bronze propeller in contact with the
steel hull of a ship
Cu cathode
HCl Solution
Differential Temperature Cell
 This is the type of cell when two identical electrodes are
immersed in same electrolyte, but the electrodes are
immersed into solution of two different temperatures
 This type of cell formation takes place in the heat
exchanger equipment where temperature difference
exists at the same metal component exposed to same
environment
 For example for CuSO4 electrolyte & Cu electrode the
electrode in contact with hot solution acts as cathode.
Salt Concentration Cell
Differential Aeration Cell
Corrosion at the bottom of the
electrical poles
Local Action Cell