Corrosion of Metals
Sections 16.2 – 16.6
Group 16
Aaron Salazar
Phillip Schneider
Kevin Stacy
Geoffrey Thiel
http://www.trekearth.com/gallery/Central_America/Cuba/East/Guantanamo/Baracoa/photo1001204.htm
Electrochemistry

For Metallic material the corrosion process is normally
electrochemical and is due to the transfer of electrons from
one species to another.

Oxidation - when Metal atoms characteristically lose or
give up electrons

Reduction - when a metal takes the electrons that are being
transferred, reducing its charge.
Electrochemical Cell
Figure 1
Iron corrodes while copper
electrodeposits, and the potential
(voltage) between the two cells is
the measure of the electron
transfer. For this system the
Potential is .78V
http://www.knovel.com/web/portal/basic_search/display?_EXT_KNOVEL_DISPLAY_bookid=656
Half Reactions and Overall Reaction
Figure 2

Overall Reaction (Figure2) is constructed of at least one
oxidation and one reduction reaction and will be the sum
of these two equations

The two separate reactions that are combined to make an
overall reaction are the Half Reactions.
Cu(s)  Cu2+(aq) + 2e2 Ag1+(aq) + 2e-  2 Ag(s)
2 Ag1+(aq) + Cu(s)  Cu2+(aq) + 2 Ag(s)
Figure 3

standard emf series (Figure 3)- a series generated by
coupling the standard hydrogen electrode to standard halfcells for various metals, and then ranking them according
to measured voltage. The difference between standard
voltages

galvanic couple- a pair of substances (ex. two different
metals) that when placed in a proper solution produces an
electromotive force by chemical action.
Summary: 1
http://www.matsceng.ohio-state.edu/mse205/lectures/chapter18/chap18_slide5.gif
Potential and Spontaneity
• Potential is a measure of electron transfer
• We use the potentials, V1 and V2, to calculate the change in
potential, ΔV, for a given system, an electrochemical cell
ΔV = V2 - V1
• The spontaneity of an electrochemical cell is determined by
the ΔV
ΔV < 0 not spontaneous
ΔV > 0 spontaneous
http://www.saskschools.ca/curr_content/chem30_05/6_redox/labs/e
lectrochem_cells.htm
Oxidation
Reduction
• Another way to calculate ΔV is using the Nernst Equation
RT [ M 1n  ]
V  (V  V ) 
ln
nF [ M 2n  ]
0
2
0
1
• Applying this equation to alloys we need the concentrations
of the metal in each alloy and the equation becomes
Shown are two examples of electrochemical cells displaying the
oxidation and reduction processes along with the direction of
electron flow, current flow, and the resulting voltage.
RT [ M 1n  ][ M 1 ]
V  (V  V ) 
ln
nF [ M 2n  ][ M 2 ]
0
2
0
1
• The temperature and concentration affect the cell potential
greatly
• The galvanic series is more practical than the emf series
because it is not an idealized situation
Summary: 2
http://np-apchemistry.wikispaces.com/chapter19
Corrosion Rates
Importance:
Corrosion rate, the rate of material removal as a consequence of
chemical actions, is an important design parameter for
engineers because corrosion can destroy process piping and
damage equipment if not accounted for properly.
We can calculate the corrosion rate as either:
Corrosion Penetration Rate (CPR)
- thickness loss over time
Or as Rate in (mol/m2-s)
KW
CPR 
At
i
r
nF
Based on current density
through material.
This figure shows the corrosion of steel over time.
Based on current density through material.
Summary: 3
Prediction of Corrosion Rates:
It is important to be able to
accurately predict corrosion rates in
order to adequately protect
equipment.
The most common prediction
method uses Polarization data. This
is based on the overvoltage or
displacement of an electrode’s
potential from its equilibrium value.
http://www.providenceri.com/publicworks/hurrica
ne/activation_polarization.gif
Polarization
Activation Polarization refers to when the reaction
rate is limited by the activation energy barrier is
associated with the slowest, rate limiting step in the
corrosion process.
1.
2.
3.
4.
Concentration Polarization is the polarization
component that is caused by concentration changes in the
environment adjacent to the surface.
Depletion Zone
Adsorption of H+ ions from the solution onto the metallic surface.
Electron transfer from metallic surface to form a hydrogen atom.
Combining of two hydrogen atoms to form hydrogen molecule
The coalescence of many hydrogen molecules to form a bubble.
H
+
Cathode
H
+
H
+
H
+
H
+
H
+
H
+
H
+
*Team generated graphic.
This figure illustrates the presence of a depletion zone
near the surface of a metal cathode. Depletion zones can
form when reaction rate is high or solution concentration
is low. This system is said to be concentration polarized.
http://www.providenceri.com/publicworks/hurricane
/activation_polarization.gif
Polarization data may be plotted according to:
 a    log
Summary: 4
i
i0
Intersection
The intersection of the extrapolation of the linear portions of the curves gives the
corrosion potential, Ecorr , and the corrosion current icorr.
The corrosion rate may be determined from the corrosion current.
http://www.energy-cie.ro/archives/2010/n1-1-14.pdf
Effects of Environment

Metal corrosion is greatly affected by it’s environment.

For instance, higher temperatures and higher velocity (or
motion) of a metal correlates with higher corrosion rates.

Velocity effects are generally known as erosion corrosion.
This is caused by the relative motion of the metal or it’s
environment.
Effects on Business
•
•
Erosion has many negative effects on the economy.
The main negative effect of corrosion on the economy
is that many companies are often having to replace parts
or machines due to corrosion which costs a large sum
of money.
Temperature and Corrosion
Cooler temperatures
cause less corrosion.
Higher temperatures
result in accelerated
corrosion.
http://4.bp.blogspot.com/_C_N3x2dSff0/SbJq8PC_dPI/AAAAAAA
AAGc/2JxWr93cozc/s320/erosion_corrosion.jpg
Summary: 5
4. An
1.
2.
3.
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actual tower
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corrosion
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support.
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theanchors
anchor
near the
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base.
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breaks.
phone towers.
http://www.corrosion-doctors.org/Corrosion-Factors-Cells/images/
Characteristics of Corrosion

Corrosion is defined as destructive and unintentional attack of metal

Corrosion is electrochemical and ordinarily begins at the surface.

Corrosion starts because of reactions with its surroundings.

Rust is a very common and well known form of corrosion

Some materials are intrinsically more resistant to corrosion (such as
pure gold and silver), due to the nature of the electrochemical process.
http://www.etftrends.com/wpcontent/uploads/2010/10/gold-bars.jpg
Pure gold and silver are intrinsically more resistant to
corrosion
http://www.ideaconnection.com/images/inventions/lg_coating-containing-friendly-bacteria-fight-corrosion.jpg
http://www.monex.com/images/photos/prodSilver01.jpg
Factors of Corrosion
http://www.corrosion-doctors.org/Corrosion-Factors-Cells/images/image011.jpg
Most of the time corrosion is
destructive and undesirable.
 Sometimes corrosion
processes are used to our
advantage. An example of this
is etching procedures.
 As shown in the table:
Environment, Stress,
Geometry, Temperature, and
Time are important factors
for different types of
corrosion.

http://nrqm.pbworks.com/f/1261611696/etching.jpg
The Copper plate is being placed in a Ferrous Chloride solution which is quickly corroding the
majority of the copper. A Protective ink layer has been placed over a specific area which allows
other Microcircuits to be attached carrying current between only the specified places. This is how
most large PCB computer components are made on a small scale.
Electrochemical Considerations
 For a metallic material the corrosion process is normally electrochemical in nature.
 This means that there is a transfer of electrons from one species to another.
 In the figure electrons are being transferred from the Zinc anode to the Copper cathode,
providing the current to light the bulb between the two electrodes.
 Zinc is then electrochemically corroded and releases zinc ions into the aqueous solution.
Electron
current
Release of
Zn+2 ions
(corrosion)
Electrodepositing
on surface of Cu(s)
http://t1.gstatic.com/images?q=tbn:WrINNqxuXP-mFM:http://upload.wikimedia.org/wikipedia/en/thumb/2/2f/Galvanic_cell_with_no_cation_flow.png/400px-Galvanic_cell_with_no_cation_flow.png&t=1
Oxidation
 Oxidation is when metal atoms
characteristically lose or give up electrons.
 General Equation for Oxidation is
M  Mn+ + ne The site at which oxidation takes place is
called the Anode
Figure: Oxidation of an iron bike over a large time. Inset
shows the oxidation reaction between iron and the
oxygen present in the atmosphere.
Figure above shows the oxidation half
reaction using a zinc electrode. The solid zinc
gives up electrons and ions are released into
solution.
http://nobelprize.org/nobel_prizes/chemistry/laureates/1992/illpres/oxidation.html
http://t1.gstatic.com/images?q=tbn:WrINNqxuXPmFM:http://upload.wikimedia.org/wikipedia/en/thumb/2/2f/Galvanic_cell_with_no_cation_flow.png/400pxGalvanic_cell_with_no_cation_flow.png&t=1
Example: Oxidation of Zinc

Oxidation of Zinc- The solid zinc (as indicated on the left side of the arrow) donates two electrons.

The two electrons along with the Aqueous Zn2+ are indicated on the right side of the arrow
Valence Charge
0
2+
Zn (s)  Zn2+(aq) + 2e-

The 2e- indicates that the Zinc is losing two electrons.

The Valence charge is also a good indication of how many electrons are being transferred
The picture on the left
shows a zinc bracket that
has undergone slight
oxidation (as evidenced
by the spots of rust).
The picture on the right
shows a zinc tube that has
not undergone oxidation.
http://farm5.static.flickr.com/4058/4410488572_5d8d8e2c24.jpg
http://www.insteellimited.com/images/DSCF0023.jpg
Reduction
 Reduction is when a metal takes the
electrons that are being transferred, reducing
its charge.
 The General equation for Reduction is
Mn+ + ne-  M
 The Location at which the reduction occurs
is called the Cathode
Figure above shows the reduction half
reaction using a copper electrode. Copper
ions from solution are deposited on the
surface of the metal as it accepts electrons.
http://t1.gstatic.com/images?q=tbn:WrINNqxuXPmFM:http://upload.wikimedia.org/wikipedia/en/thumb/2/2f/Galvanic_cell_
with_no_cation_flow.png/400px-Galvanic_cell_with_no_cation_flow.png&t=1
http://cltad.arts.ac.uk/groups/camberwellmateriall
ibrary/wiki/53404/Electro-plating_.html
Plating is an example of industrial and
household application of the reduction concept
Example: Reduction of Copper

For the reduction of Copper, the process will add electrons to the copper as shown on the left hand
side of the arrow

The two electrons are being added to the copper on the left side of the equation
Valence Charge

2+
0
Cu2+(aq) + 2e-  Cu (s)
The copper solid being produced has a neutral Valence charge as indicated on the right hand side of
the arrow..
The picture on the left shows a piece of
copper metal that has undergone an
oxidation reaction with atmospheric
oxygen. The black surface of the metal is
copper oxide.
In the picture on the right, the metal is
placed in a hydrogen gas environment.
The hydrogen gas is a reducing agent. It
reduces the copper oxide on the surface
of the copper and causes pure copper to
be formed.
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/
STILLS/REDOXCU/REDOXCU/64JPG48/5.JPG
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/STILLS
/REDOXCU/REDOXCU/64JPG48/14.JPG
Overall and Half Reactions
Electrical Potential
Direction of Electron Flow
 An Overall Reaction is constructed of at
least one oxidation and one reduction
reaction and will be the sum of these two
equations
 The two separate reactions that are
combined to make an overall reaction are
the Half Reactions.
 The oxidation reaction donates electrons
and the reduction reaction accepts the
electrons.
 All electrons generated through oxidation
must be consumed by reduction, thus there
is no accumulation of electrons forms in
solution. – Conservation of Net Charge
Direction of Cation Flow
Direction of Anion Flow
Half Reactions
http://www.webassign.net/blb8/20-05.gif
Example: Overall and Half Reactions
 Reduction Half-Reaction
Valence Charge
0
2+
Zn (s)  Zn2+(aq) + 2e-
 Oxidation Half-Reaction
Valence Charge
2+
0
Cu2+(aq) + 2e-  Cu (s)
 Overall Reaction
.
Cu2+(aq) + 2e- + Zn(s)  Cu(s) + Zn2+(aq) + 2eCu2+(aq) + Zn(s)  Cu(s) +Zn2+(aq)
http://www.simarzincorame.com/imag
es/zinc-rolled-sections-3.jpg
This is an example of zinc that has been rolled into
sheets for easy use. Zinc’s largest use is in corrosion
protection. Galvanising is the main method of
protection against corrosion of steel, i.e. the steel is
coated with a layer of zinc in order to protect it from
decay.
This is an example of copper that is normally produced in industry for either
household or business use. Usually metals that are going to sit on a shelf for a
long period of time are coated with a corrosion resistant wax or petroleum jelly
depending on the metal.
http://www.etftrends.com/wp-content/uploads/2010/11/copper_1.jpg
Example 2: Overall and Half Reactions
 Reduction Half-Reaction
Cu(s)  Cu2+(aq) + 2e Oxidation Half-Reaction
2 Ag1+(aq) + 2e-  2 Ag(s)
 Overall Reaction
Cu(s)  Cu2+(aq) + 2e2 Ag1+(aq) + 2e-  2 Ag(s)
2 Ag1+(aq) + Cu(s)  Cu2+(aq) + 2 Ag(s)
http://www.plumbinghelp.ca/images/Copper-Pipe.jpg
http://static.seekingalpha.co
m/uploads/2009/7/13/saupl
oad_silver_bars.jpg
These silver bricks are used to keep track of large
amounts of silver that is very highly valued. Silver is a
natural anti-bacterial and is not very corrosive unless
introduced to other metals and thus in industry silver
rings are often used in processes where water is reused
to keep bacteria from growing.
This is another example of common copper tubes that
would be produced in industry. Copper pipe is
commonly used for plumbing/air conditioning/
refrigeration because of the amount of time that
copper tubing lasts and its versatility.
Electrode Potentials

Not all metallic substances oxidize to form ions with the same degree of ease which causes a movement
of electrons when connecting two half cells.

A Standard Half-cell is a pure metal electrode immersed in a 1M solution of its ions at 25oC

When we connect the Standard Half-Cell to another Half-cell with a different metal, we can produce a
Potential based on the difference in ease of movement of electrons.
These are examples of multiple electrodes being used in a lab. These electrodes can be used to measure potential difference
between metals as seen in the right picture, or an electrical charge can be induced to force a reaction as seen on the left
http://www.discoverarmfield.co.uk/data/ceq/images/ceq2.jpg http://www.imagestate.com/Preview/PreviewPage.aspx?id=2218302
Electrode Potentials: Iron corrosion
 For the picture to the right Iron corrodes
while copper electrodeposits.
 The Potential (voltage) between the two cells
is the measure of the electron transfer, and for
this system the Potential is .78V
 As shown in the picture, both Electrodes are
in a 1M solution of the electrodes’ ions.
 A membrane is between the 2 solutions which
allows the charge in the solution to stay
constant
http://www.knovel.com/web/portal/basic_search/displa
y?_EXT_KNOVEL_DISPLAY_bookid=656
Electrode Potentials: Iron electrodepositing
 This is different than the last example
because now zinc is the electrode that is
corroding and the iron has
electrodepositing.
 Notice that the Voltage direction and
value has also changed. Electrons are
now transferring from the Zinc to the
Iron with a Potential of .323
http://www.knovel.com/web/portal/basic_search/display?_EXT_KNOVEL
_DISPLAY_bookid=656
Galvanic Couple
•
The electron flow indicates that this is a
spontaneous reaction which transfers
electron from the Zinc to the copper.
Anode
Zinc Sulfate Solution.
•
Electrolyte- any substance that dissociates
into its separate ions when dissolved in a
suitable medium or melted and thus
makes the medium conductive.
Cathode
Galvanic couple- a pair of
substances (ex. two different
metals) that when placed in a
proper solution produces an
electromotive force by
chemical action.
Copper Sulfate Solution.
The Diaphragm allows ions to be
transferred which equalizes charge and
allows the continuation of electrons to flow
http://www.pinkmonkey.com/studyguides/subjects/chem/chap9/c0909501.asp
Standard emf series

Standard emf Series- a series generated by coupling the standard hydrogen electrode to standard half-cells
for various metals, and then ranking them according to measured voltage.

The standard hydrogen electrode consists of an inert platinum electrode in a 1M solution of H+ ions with
Hydrogen gas bubbled through the solution.
Standard
Hydrogen
Electrode
http://www.ktfsplit.hr/glossary/image/standard_hydrogen_electrode.gif
http://www.matsceng.ohio-state.edu/mse205/lectures/chapter18/chap18_slide5.gif
Standard emf Series of Metals
Electrode potential vs.
Metal-metal ion
normal hydrogen
equilibrium
electrode at
(unit activity)
25°C,volts
Au-Au+3
+1.498
This is the Standard emf Series Table for
Pt-Pt+2
+1.2
common Metals. This table is used to
Noble or Pd-Pd+2
+0.987
approximate Potential differences in a
Cathodic Ag-Ag+
+0.799
spontaneous cell for standard pressure and
Hg-Hg+2
+0.788
temperature conditions. The metal listed on
Cu-Cu+2
+0.377
the left corresponds with the electrode vs
potential on the right. As you travel from
H2-H+
0.000
the top of the table to the bottom of the
table, the closer to the top that you are the
Pb-Pb+2
-0.126
more noble/Cathodic , the closer to the
Sn-Sn+2
-0.136
bottom the more active/anodic.
Ni-Ni+2
-0.250
Co-Co+2
Cd-Cd+2
Fe-Fe+2
Cr-Cr+3
Zn-Zn+2
Active or Al-Al+3
anodic Mg-Mg+2
Na-Na+
K-K+
-2.777
-0.403
-0.440
-0.744
-0.763
-1.662
-2.363
-2.714
-2.925
Source: A.J. de Bethune and N.A.S. Loud
“Standard Aqueous Electrode Potentials and
Temperature Coefficient at 25 Celsius” Clifford
A. Hampel Skokie, III. 1964 (Table 9-1)
Ni-Cu Cell
 The reaction is a reduction reaction because
one of the metals accepts electrons that flow
from the other metal because it has a lower
potential.
Example Problem:
Compute the voltage of an electrochemical cell at 25⁰C. Pure copper is
Example of a simple
2
immersed in a 1 M solution of Cu ⁺ ions and pure nickel is immersed in a 1 M Ni-Cu voltaic cell.
solution of Ni2⁺ ions.
http://www.tutorvista.com/topic/difference-betweengalvanic-cell-and-voltaic-cell
Sample Calculation:
Ni2⁺ + 2e⁻
Cu2⁺ + 2e⁻
Ni
Cu
V1⁰ = - 0.250 V
V2⁰ = 0.340 V
Ni
Ni2⁺ + 2e⁻
Cu2⁺ + 2e⁻
Cu
Ni + Cu2⁺
0.250 V
0.340 V
Ni2⁺ + Cu
ΔV⁰ = V2⁰ - V1⁰
0.340 V - (-0.250 V) = 0.590 V
0.590 V
Spontaneity
 Zn + Cu2+
Zn2+ + Cu
 ΔV is positive in one direction, the
direction in which the reaction is
spontaneous.

If ΔV is negative then the reaction is
spontaneous in the opposite direction.
 The metal that has the higher potential
will be reduced more. The reaction will
proceed in the direction where that metal
will act as a cathode.
http://www.chem.tamu.edu/class/majors/tutorialnotefiles/electrochem.htm
Sample Calculation:
Zn2⁺ + 2e⁻
Cu2⁺ + 2e⁻
ΔV⁰ = V2⁰ - V1⁰ spontaneous if ΔV⁰ > 0
non-spontaneous if ΔV⁰ < 0
The result of this calculation shows that electrons
will spontaneously flow from Zn to Cu.
Zn
Cu
V1 = - 0.763 V
V2⁰ = - 0.440 V
ΔV⁰ = V2⁰ - V1⁰
(-0.440 V) - (-0.763 V) = 0.323 V
Nernst Equation
 The Nernst Equation is used to find the cell potential of two half cells that are
electrically coupled and for which solution ion concentrations are other than 1 M
RT [ M 1n  ]
V  (V  V ) 
ln
nF [ M 2n  ]
0
2
0
1
 V2 and V1- are the standard potentials as taken





from the standard emf series
R- gas constant
T- temperature (K)
n- # of electrons participating in either of the
half-cell reactions
F- faraday constant (96,500 C/mol)
[M1n+] and [M2n+]- molar ion concentrations
of the two pure metals
The solid line is a plot of the
Nernst equation for potassium.
http://www.unmc.edu/physiology/Mann/pix_3/f3-13.gif
Nernst Equation at Room Temperature
n
.
0592
[
M
0
0
1 ]
V  (V2  V1 ) 
ln
n
[ M 2n  ]
Ni2⁺ + 2e⁻
Cu2⁺ + 2e⁻
Ni
Cu
V1⁰ = - 0.250 V Conc. 2 M
V2⁰ = 0.340 V Conc. 1 M
 ΔV= (0.340 V – (-0.250 V)) 

= 0.590 V - 0.00891 V
= 0.58109 V (volts)
0.0592 2
ln
2
1
An example illustrating the Nernst equation
and ion flow through a membrane.
This is a special case of the Nernst equation
which is only applicable at Room
Temperature.
http://science.widener.edu/~svanbram/chem366/echem.pdf
Nernst Equation Applied to Alloys
These values, [M1] and [M2], will be
= 1M for pure metals
< 1M for alloys
n
1
n
2
RT [ M ][ M 1 ]
V  (V  V ) 
ln
nF [ M ][ M 2 ]
0
2
0
1
•V2 and V1- are the standard potentials
•R- gas constant
•T- temperature (K)
•n- # of electrons participating
•F- faraday constant
•[M1n+] and [M2n+]- molar ion concentrations
•[M1] and [M2] - molar concentrations of metals in the two alloys
For alloys composition must be considered, so the Nernst Equation is
modified to include [M1] and [M2]. This is because the molar
concentrations of the metal ions in solution will not have
concentrations of 1. Therefore it is necessary to include the initial
concentrations of the metals in the alloys.
http://www.doitpoms.ac.uk/tlplib/pourbaix/nersnt_detailed.php
Several Metal Alloys
http://t2.gstatic.com/images?q=tbn:Efj-LqaM8KC4M:http://en.fukesi.com/img/product_pic/50156036_Alumi
Influence of concentration and
temperature on cell potential
 The cell potential relies heavily on the temperature and molar ion concentrations
 Altering the temperature or solution concentrations will change the cell potential and in
some cases reverse a spontaneous reaction.
 According to the Nernst equation, as you increase temperature the potential decreases
and thus becomes less and less spontaneous.
Concentration of a solution
refers to the weight or
volume of solute in a
specified amount of solution
or solvent. According to
chem.purdue.edu, it is a
macroscopic property,
meaning it describes the
behaviors or characteristic
of a sample which is large
enough to see, weigh,
manipulate, handle, etc.
Solutions of varying concentrations.
http://4.bp.blogspot.com/_e096J02yVY/TAtuocstqgI/AAAAAAAAAkU/iBpiQp4s5fI/s400/concentration.jpg
http://www.thedail
ygreen.com/media/
cm/thedailygreen/i
mages/refridgeratorthermometer-l.jpg
Galvanic Series

The galvanic series (or electropotential series) determines the nobility of metals and semi-metals. When two metals are submerged in an
electrolyte, while electrically connected, the less noble (base) will experience galvanic corrosion. The rate of corrosion is determined by the
electrolyte and the difference in nobility. The difference can be measured as a difference in voltage potential. Galvanic reaction is the
principle upon which batteries are based.

The Galvanic series is a much more practical and realistic ranking system than the standard emf series, because the standard emf series is
generated under highly idealized conditions.

In nature nearly all metals occur as compounds. This is because there is a net decrease in free energy in going from metallic to oxidized states,
which essentially means that the reaction is spontaneous. Common compounds are oxides, hydroxides, carbonates, silicates, sulfides, and
sulfates.

The exception to this rule is the noble metals gold and platinum. These metals are in their metallic state in nature.
Case One
This figure demonstrates different scenarios in which potential will or will not be generated. The galvanic series deals with creating potential
through connecting dissimilar metals in a conductive solution, as shown in Case One.
http://www.enviroscan.com/assets/images/galvanic_corrosion_cells.jpg
Galvanic Series of Metals
 More Cathodic Metals are Less
active (more inert) and have a
higher potential. Platinum is
listed as the most Cathodic metal
and is thus very unreactive
 More Anodic Metals are more
active (less inert) and have lower
potentials. The most Anodic matal
is listed as Magnesium and is thus
is very active and has the highest
potential
http://www.amacgroup.com.au/library/images/corrosion/The-Galvanic-Series-of-Meta.gif
Corrosion Rates

Corrosion rates are very important design parameter for engineers because corrosion can destroy process
piping and damage equipment if not accounted for properly.

The corrosion rate is calculated as the rate of material removal as a consequence of chemical actions.

We can calculate the corrosion penetration rate (CPR) which is the actual rate of removal of material.

If corrosion is not calculated correctly major processes can be completely shut down to fix a corroded
section of the process, which will cause a company to incur significant costs.

According to the US Department of Energy, Corrosion does more economic damage than the effect of all
natural disasters combined. Corrosion costs industry approximately 5% of the Gross National Product.
Examples of
Industrial
Corrosion
and
Corrosion
Prevention
http://www.bushman.cc/photos/Commercial_Hot_Water_Heat_Tube_Corrosion.j
pg
http://www.nansulate.sk/magyar/images/industrial_1.jpg
Corrosion Penetration Rate
 The rate of material removed as a consequence of chemical reaction is





given by:
KW
CPR 
At
W- is the weight loss after exposure time
p- Density
A- exposed surface area
t- exposure time
K- is a constant
 Defined as 87.6 for mm/year (mm/yr)
 Defined as 534 for mils/year (mpy)
 (where 1mil = .001 in)
For most applications, an acceptable Corrosion
Penetration Rate (CPR) is 20 mpy.
The figure above shows “Pitting
Corrosion.” Pitting is a particularly
dangerous form of corrosion. Rate of
penetration can be 10 to 100 times that by
general corrosion.
http://backup.exprobase.com/docs/Illustrations/Databases/Failure%20mechanisms/Pitting%20corrosion.png
Alternative Corrosion Rate Expression
i
r
nF
 r- rate in mol/m^2*s
 i- current density( current per unit surface area of material corroding)
 n- number of electrons with ionization of each metal atom
 F- 96,500 C/mol
This figure shows the corrosion of steel over time.
http://people.csail.mit.edu/wojciech/TSVBRDF/TSVBRDF.pdf
Prediction of corrosion rates
 When an electrochemical cell which has been
short circuited such that oxidation occurs at the
respective electrode surfaces, we can not use
standard emf values. This is because the system is
no longer at an equilibrium state. The difference
between the standard emf value and the actual
Potential is termed Polarization. The Magnitude
of Polarization is termed overvoltage.
These are pictures of corrosion in the form of
rust forming on the surface of metals.
http://wpcontent.answcdn.com/wikipedia/commons/thumb/5/55/Rust03102006.JPG/22
5px-Rust03102006.JPG
http://www.munters.com/AvanMediaBank/Image/%7B2C83A853-5A03-49AE-9E78-77A92967E3C5%7D/width_439/height_124/Bar_Corrosion02.jpg
Activation Polarization

An electrochemical reaction consists of a
sequence of steps at the interface between the
metal electrode and the electrolyte solution.

Activation Polarization refers to when the
reaction rate is controlled by the slowest step in
the series.

The term “activation” is applied to this type of
polarization because an activation energy barrier
is associated with this slowest, rate limiting step.

There are 4 proposed steps for activation
polarization, which will be explained in the
following slides
1.
Adsorption of H+ ions from the solution
onto the metallic surface.
2.
Electron transfer from metallic surface to
form a hydrogen atom.
3.
Combining of two hydrogen atoms to
form hydrogen molecule
4.
The coalescence of many hydrogen
molecules to form a bubble.
This figure shows the activation
polarization process.
http://www.providenceri.com/publicworks/hurrica
ne/activation_polarization.gif
Activation Polarization
Step 1: Adsorption of H+ ions onto the zinc surface.
 The reaction takes place in a solution
containing H+ ions.
 These ions will adsorb to the surface of
the submerged metal.
 This step will be the rate limiting step if
there are a limited number of H+ ions in
solution. In an acidic solution, where
there are an abundance of H+ ions, this
will most likely not be the slow step.
http://www.providenceri.com/publicworks/hurrica
ne/activation_polarization.gif
Activation Polarization
Step 2: Electron transfer from the metal to form a hydrogen atom
 This step involves the rate at which
electrons travel through the metal.
 There is a wide range of transfer rates of
electrons by various metals and, as a result,
the rate of hydrogen evolution from
different metal surfaces can vary greatly.
 In most reaction sequences this step will be
the most important in determining overall
reaction rate.
http://www.providenceri.com/publicworks/hurrica
ne/activation_polarization.gif
Activation Polarization
Step 3: Combining of two hydrogen atoms to form a molecule of hydrogen
 This step depends on the location of hydrogen
ions in relation to one another on the surface
of the metal.
 If the concentration on of ions on the surface
is low, this step will be limiting as the ions
will not come into contact with other ions
with which they can form a molecule as
frequently as if the surface concentration is
high.
http://www.providenceri.com/publicworks/hurrica
ne/activation_polarization.gif
Activation Polarization
Step 4: Bubble formation
•
The final step is the combining of many
hydrogen molecules to form a
hydrogen gas bubble.
•
This bubble is then diffused through the
solution and evolved at the liquid
surface. This evolution of hydrogen gas
can be used by the observer to confirm
the reaction is proceeding.
•
This will very rarely be the rate
determining step.
http://www.providenceri.com/publicworks/hurrica
ne/activation_polarization.gif
Activation Polarization
 - overvoltage
a


i - current density
io and  -constants for a
𝑖
𝑛𝑎 = ±𝛽 ∙ 𝑙𝑜𝑔
𝑖0
particular half-cell
 A plot of overvoltage vs. the
logarithm of current density
exhibits a linear relationship with
a slope of β.
 This relationship holds for both
the anodic and cathodic reactions.
http://electrochem.cwru.edu/encycl/fig/c02/c02-f03b.gif
Activation Polarization
 Equilibrium is actually a dynamic state on the atomic level.
 Equilibrium exists for a half-cell when the rate of reduction is equal to the rate
of oxidation, so there is no net reaction.
 i0 is the “exchange current density” and is measured experimentally for each
system
 Equation at equilibrium:
rred = roxid
io
=
nF
This figures shows area in curve where corrosion
takes place and hydrogen is liberated as in the previous
example.
http://nautarch.tamu.edu/crl/images/fig9-1.jpg
Concentration Polarization
 Concentration polarization is
the polarization component
that is caused by concentration
changes in the environment
adjacent to the surface.
 When a chemical species
participating in a corrosion
process is in short supply, the
mass transport of that species
to the corroding surface can
become rate controlling.
http://corrosion-doctors.org/Corrosion-Kinetics/Overpotential-concentration.htm
Concentration Polarization
At slow reaction rates and/or high concentrations there is always a sufficient amount of ions
available near the electrode surface, allowing the reaction to proceed unhindered.
H
+
Cathode
H
+
H
+
H
+
H
+
H
+
H
+
H
+
H
+
H
+
H
+
H
+
*Team generated graphic.
NOT concentration limited.
This figure shows a situation
where hydrogen ions in
solution are present in ample
supply to feed the reduction
reaction. The ions will be
available at the surface of the
metal to accept electrons that
are produced during the
reaction.
Concentration
Polarization
The reaction rate in concentration polarization is limited by the diffusion of ions in solution.
At high reaction rates or low concentrations a depletion zone will form near the electrode
surface, in which the ions are not replaced fast enough to keep up with the rate of reaction.
Depletion Zone
H
+
Cathode
H
+
H
+
H
+
H
+
H
+
H
+
The figure to the left illustrates the presence of
a depletion zone near the surface of a metal
cathode. Depletion zones can form when
reaction rate is high or solution concentration is
low. This system is said to be concentration
polarized.
c 
H
+
A reaction that is concentration polarized is
described by the following equation:
c 
2.303RT
i
log( 1  )
nF
iL
When overvoltage is plotted vs.
the logarithm of current density,
the result is a nonlinear graph that
approaches a constant value, iL- the
limiting diffusion current density.
iL
0
-
Concentration limited
Overvoltage, nc
*Team generated graphic.
+
2.303RT
i
log( 1  )
nF
iL
Log current density,
i
*Team generated graphic.
Combination Polarization
 As shown in the figure below, this reaction is limited by both activation and
concentration polarization.
 This graph is representative of combination Polarization and is Overvoltage vs Log
current Density
The first part of the curve is linear with a slope of –β.
+
0
i
0
Overvoltage, nc
-
Activation Polarization
The slope of the second part of the curve
decreases rapidly as the current density
approaches iL.
Concentration Polarization
Log current density,
i
i
L
*Team generated graphic.
Corrosion rates from Polarization Data

When both oxidation and reduction reactions
are rate limited by activation polarization:

The intersection of the extrapolation of the
linear portions of the curves gives the corrosion
potential, Ecorr , and the corrosion current icorr.

Intersection
The corrosion rate may be determined from the
corrosion current.
The figures demonstrate the use of graphs to predict
corrosion rates from activation polarization data.
http://www.gamry.com/App_Notes/DC_Corrosion/GettingStartedWithEch
emCorrMeasurements.htm#Quantitative Corrosion Theory
http://www.energy-cie.ro/archives/2010/n1-1-14.pdf
Corrosion Rates from Polarization Data
 When both concentration and activation polarization control the reduction reaction:
i
(H+/H2)
Potential, V
0
The figure to the left shows a reduction
reaction under combined activationconcentration polarization control. The
hydrogen is concentration limited while the
metal corrosion behavior is described by
activation polarization.
Vcorr , icorr
(M/M2+)
Log current density,
i
*The total overvoltage is the
sum of both overvoltage contributions.
i
L
*Team generated graphic.
The intersection of the graphs of each
materials’ polarization data gives the
corrosion potential and corrosion current.
Corrosion rate is then calculated using the
equation:
rcorr 
icorr
nF
Passivity
 Passivity is the characteristic of a metal
exhibited when that metal does not
become active in the corrosion
reaction. Passivity is caused by the
buildup of a stable, tenacious layer of
metal oxide on the surface of the metal.
 Stainless steels and aluminum are highly
resistant to corrosion as a result of
passivation. Stainless steels contain
chromium allowing the formation of a
protective surface film in an oxidizing
atmosphere which reduces rusting.
Aluminum also forms a protective film
and its film is able to reform very
rapidly if damaged to prevent corrosion.
http://www.tpub.com/content/doe/h1015v1/css/h1015v1_108.htm
http://www.concretecorrosion.net/imgen/suite/corrosion/schema1.gif
The figure above illustrates the corrosion
process for a material exhibiting passivity
characteristics.
Passivity
As shown in the figure:
Transpassive Region: At very high
potentials, the current density again
begins to increase with increasing
potential.
Passive Region: With increasing potential,
current density suddenly drops to a very
low value that is independent of
potential.
Active Region: At low potentials, behavior
is linear for normal metals.
http://www.industrialheating.com/IH/Home/Images/ih0
308-vst-fig.1-lg.jpg
Effects of Environment on Corrosion
 Metal corrosion is greatly affected by it’s environment.
 For instance, higher temperatures and higher velocity (or motion) of a metal
correlates with higher corrosion rates.
 Velocity effects are generally known as erosion corrosion. This is caused by the
relative motion of the metal or it’s environment.
Temperature and Corrosion
Progression of corrosion.
EROSION
Cooler
temp.
Higher
temp.
http://4.bp.blogspot.com/_C_N3x2dSff0/SbJq8PC_dPI/AAAAAAAAAGc/2JxWr93cozc/s320/erosion_corrosion.jpg
http://www.thewatertreatments.com/wp-content/uploads/2009/12/corrosion-stages.JPG
Business Effects
 Corrosion has many negative effects on the economy.
 The main negative effect of corrosion on the economy is that many companies often have
to replace parts or machines due to corrosion. This costs a large sum of money.
 Oil refineries in the U.S. lose billions of dollars every year due to corrosion.
Chart shows a comparison of corrosion costs in the
U.S. in 1975 with 1995.
http://asmcommunity.asminternational.org/content/ASM/StoreFiles/06691G_Chapter_1.pdf
An oil refinery that has to protect
against corrosion to maintain optimal
production and profit.
http://granitegrok.com/pix/oil%20refinery.jpg