Corrosion is the unwanted oxidation of a metal.
Oxidation of All Metals
is called Corrosion
Oxidation of all Metals in general is called corrosion
Oxidation of
Oxidation of All Metals
Iron Metal is
is called Corrosion
called Rusting
Whereas, oxidation of iron metal specifically, is usually called rusting.
Looking at the standard reduction table, (click) near the bottom, we see that (click)
iron, zinc, and aluminum are all near the bottom right,
(click after fading)
Ox. Pot. = +0.45 V
Ox. Pot. = +0.76 V
Ox. Pot. = +1.66 V
They all have high oxidation potentials
Ox. Pot. = +0.45 V
Ox. Pot. = +0.76 V
Ox. Pot. = +1.66 V
Which means they are strong reducing agents,
Strong
Reducing
Agents
Ox. Pot. = +0.76 V
Strong
Reducing
Agents
Ox. Pot. = +1.66 V
Readily
Oxidized
Ox. Pot. = +0.45 V
And are readily oxidized.
The iron in steel show obvious corrosion,
but aluminum and zinc do not
But notice that aluminum and zinc don’t seem to corrode like the iron in steel does!
4Al(s) + 3O2(g)  2Al2O3(s)
2Zn(s) + O2(g)  2ZnO(s)
It’s because aluminum and zinc both readily oxidize (click),
4Al(s) + 3O2(g)  2Al2O3(s)
2Zn(s) + O2(g)  2ZnO(s)
Hard oxide coatings that
are difficult to penetrate
To form hard oxide coatings that are difficult for more oxygen to penetrate.
4Al(s) + 3O2(g)  2Al2O3(s)
2Zn(s) + O2(g)  2ZnO(s)
So these metals will not corrode much under normal atmospheric conditions, even
when they get rained on.
Iron (Fe(s)) is different than
these other metals
But iron, or Fe solid is different than other metals like aluminum and zinc,
Iron (Fe(s)) is different than
these other metals
It’s oxide does not form an
impenetrable coating. Instead it
can flake off, exposing a fresh
iron suface to oxidizing agents.
It’s oxide does not form an impenetrable coating.
Iron (Fe(s)) is different than
these other metals
It’s oxide does not form an
impenetrable coating. Instead it
can flake off, exposing a fresh
iron surface to oxidizing agents.
Instead it can flake off, exposing a fresh iron surface to oxidizing agents.
Steel is made up
primarily of iron.
Coating steel with tin
will protect its
surface and prevent
oxidation of the iron
in it. An example is a
tin can.
Steel is mainly iron.
Steel is made up
primarily of iron.
Coating steel with tin
will protect its
surface and prevent
oxidation of the iron
in it. An example is a
tin can.
Coating steel with tin protects it’s surface and prevents oxidation of the iron that’s in it. This is why
tin cans, which are tin coated steel, are normally good for keeping food without rusting.
However, if the cans are dented or pitted, the tin coating is broken and water can collect
in the indentations. Now rusting occurs relatively quickly
Water
droplet
iron
We’ll consider a piece of iron where a dent or a crack in a painted surface makes it easy
for (click) a water droplet to come in contact with the iron surface.
–
e
2+
Fe
e–
iron
Here an iron atom can oxidize to an iron 2+ cation, which will dissolve in the water.
Fe2+
Fe(s)  Fe2+ + 2e–
–
e
–
e
iron
The equation for this oxidation is Fe(s)  Fe2+ + 2e–. Further oxidation of iron here
causes pits in the surface to increase in size.
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
–
e
–
e
iron
What is formed here is an electrochemical cell. (click) Because oxidation takes place
here, (click) this is called the anode region.
O2
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
–
e
–
e
iron
Oxygen from the air can come into contact with the edge of the water droplet and the
iron metal surface.
If we could take this section of our standard reduction table.
And split it right here,
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
We could add in this half-reaction.
½O2(g) + H2O + 2e– ⇄ 2OH– … +0.40
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
½O2(g) + H2O + 2e– ⇄ 2OH– … +0.40
Reduction of oxygen to form
hydroxide ions
It is the reduction of oxygen in the presence of water, to produce hydroxide ions.
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
½O2(g) + H2O + 2e– ⇄ 2OH– … +0.40
Reduction of oxygen to form
hydroxide ions
And its reduction potential is + 0.40 volts
H2O
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
–
e
–
e
–
OH
O
OH2–
½O2(g) + H2O + 2e–  2OH–
Here, oxygen combines with water and the electrons formed by the oxidizing iron, to
form hydroxide ions.
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
Which then dissolve in the water droplet.
OH–
OH–
½O2(g) + H2O + 2e–  2OH–
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
Because reduction has taken place near the point where the water, iron, and air meet,
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
This is called the cathode region.
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
Including Fe2+ & Fe3+
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
Looking at the solubility table, we see that Fe2+ and OH– form a low solubility
compound.
Including Fe2+ & Fe3+
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
The Fe2+ ion formed at the anode can combine with the OH– ions formed at the
cathode,
Including Fe2+ & Fe3+
– Fe2+ OH–
Fe(OH)
OH
2(s)
Anode
Region
Fe(s)  Fe2+ + 2e–
iron
To produce the precipitate Fe(OH)2.
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
Including Fe2+ & Fe3+
Fe(OH)2(s)
Anode
Region
Fe(s)  Fe2+ + 2e–
iron
Which can build up on the iron surface.
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
H2O
Fe(OH)2(s)
Anode
Region
Fe(s)  Fe2+ + 2e–
iron
O2
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
What can happen is Oxygen from the air and water in the droplet can act to (click)
further oxidize the iron in iron(II) hydroxide,
Fe2O
•xH
Fe(OH)
O
2O
3H
(s)
22O
2(s)
Anode
Region
Fe(s)  Fe2+ + 2e–
iron
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
And through a fairly complex series of reactions, form what is called hydrated iron(III)
oxide, which has the formula Fe2O3 dot x H2O.
Variable numbers of
water molecules can
occur in this formula
Fe2O3•xH2O(s)
Anode
Region
Fe(s)  Fe2+ + 2e–
iron
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
The “x” here means that variable numbers of water molecules can appear in this formula, depending
on conditions. Fe2O3 dot x H2O, or hydrated iron(III) oxide is one of the major components of rust.
Neutral or
Basic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Anode: Fe(s)Fe2+ + 2e–
The process that caused the oxidation of iron that we just discussed takes place when neutral or basic water
contacts iron. The anode half-rx (click) is the reverse of this one found lower on the reduction table.
Neutral or
Basic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Cathode:
½O2(g) + H2O + 2e–  2OH–
Anode: Fe(s)Fe2+ + 2e–
And the cathode half-reaction is the one we added to the table proceeding in a forward direction
Neutral or
Basic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Cathode:
½O2(g) + H2O + 2e–  2OH–
2e–
2OH–
½O2(g) + H2O +

Fe(s)  Fe2+ + 2e–
E°
+0.40 V
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
Adding the cathode half-reaction, with it’s E naught value of +0.40 Volts
Neutral or
Basic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Cathode:
½O2(g) + H2O + 2e–  2OH–
2e–
2OH–
½O2(g) + H2O +

Fe(s)  Fe2+ + 2e–
E°
+0.40 V
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
To the anode half-reaction, which because it is the reverse of the one on the table, has
an E naught value of (click) + 0.45 volts.
Neutral or
Basic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Cathode:
½O2(g) + H2O + 2e–  2OH–
2e–
2OH–
½O2(g) + H2O +

Fe(s)  Fe2+ + 2e–
E°
+0.40 V
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
Gives the overall redox equation: ½O2(g) + H2O + Fe(s)  Fe2+ + 2OH–, which has an E naught value of 0.40 V plus 0.45 V,
which is (click) positive 0.85 volts. The + value for the E° of the overall reaction, means this process is spontaneous.
Acidic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Cathode:
½O2(g) + 2H+ + 2e–  H2O
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
Now if the water contacting the iron is acidic, the cathode becomes this half-reaction up
here.
Cathode:
½O2(g) + 2H+ + 2e–  H2O
Acidic Water
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Which is ½O2(g) + 2H+ + 2e–  H2O
.
Anode: Fe(s)Fe2+ + 2e–
Cathode:
½O2(g) + 2H+ + 2e–  H2O
Acidic Water
Acid is
Present
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
The H+ here means acid is present
Anode: Fe(s)Fe2+ + 2e–
Cathode:
½O2(g) + 2H+ + 2e–  H2O
½O2(g)
.
In acidic water
+ 2H+ + 2e–  H2O
Fe(s)  Fe2+ + 2e–
½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O
Acidic Water
E°
+1.23 V
+0.45 V
+1.68 V
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
To determine the equation for the overall reaction and find it’s cell potential, we add (click) the halfreaction at the cathode, ½O2(g) + 2H+ + 2e–  H2O, with its E naught value of +1.23 volts,
Cathode:
½O2(g) + 2H+ + 2e–  H2O
½O2(g)
.
In acidic water
+ 2H+ + 2e–  H2O
Fe(s)  Fe2+ + 2e–
½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O
Acidic Water
E°
+1.23 V
+0.45 V
+1.68 V
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
To the half-reaction at the anode, which is Fe(s)Fe2+ + 2e–, with its E naught value of +0.45 Volts
Cathode:
½O2(g) + 2H+ + 2e–  H2O
½O2(g)
.
In acidic water
+ 2H+ + 2e–  H2O
Fe(s)  Fe2+ + 2e–
½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O
Acidic Water
E°
+1.23 V
+0.45 V
+1.68 V
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
And we get the overall reaction, ½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O, and its E naught value is (click) 1.23
V + 0.45 V, which is equal to (click) positive 1.68 volts.
Cathode:
½O2(g) + 2H+ + 2e–  H2O
½O2(g)
.
In acidic water
+ 2H+ + 2e–  H2O
Fe(s)  Fe2+ + 2e–
½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O
Acidic Water
E°
+1.23 V
+0.45 V
+1.68 V
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
We see that the overall cell potential for the reaction in acidic water, 1.68 volts.
Cathode:
½O2(g) + 2H+ + 2e–  H2O
½O2(g)
.
In acidic water
+ 2H+ + 2e–  H2O
Fe(s)  Fe2+ + 2e–
½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O
Acidic Water
E°
+1.23 V
+0.45 V
+1.68 V
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
Anode: Fe(s)Fe2+ + 2e–
Is significantly higher than the overall cell potential in neutral or basic water, which is
only 0.85 Volts.
Cathode:
½O2(g) + 2H+ + 2e–  H2O
½O2(g)
.
In acidic water
+ 2H+ + 2e–  H2O
Fe(s)  Fe2+ + 2e–
½O2(g) + 2H+ + Fe(s)  Fe2+ + H2O
Acidic Water
E°
+1.23 V
+0.45 V
+1.68 V
There is a greater
tendency for rusting
to occur when an acid
is present.
In neutral or basic water E°
½O2(g) + H2O + 2e–  2OH–
+0.40 V
Fe(s)  Fe2+ + 2e–
+0.45 V
.
½O2(g) + H2O + Fe(s)  Fe2+ + 2OH– +0.85 V
½O2(g) + H2O + 2e– ⇄ 2OH– ........……..+0.40
Anode: Fe(s)Fe2+ + 2e–
This means there is a greater tendency for rusting of iron to occur when an acid is
present, than when it is not. In acidic conditions, iron tends to rust more readily.
Will
Oxidize
Fe(s)
Recall that any oxidizing agent above Fe(s) on the left, will oxidize solid iron. So if iron
is in contact with any of these, this will increase iron’s tendency to corrode.
Acidified hydrogen peroxide
Chlorine
Acidified oxygen
Bromine
Nitric acid
Sulphuric acid
All acids
This includes some common things like (clk) all acids, which contain H+ , or hydrogen ions, (clk) sulphuric acid, with
hydrogen and sulphate ions, (clk) nitric acid with hydrogen and nitrate ions, (clk) bromine and chlorine sometimes
Salt or other electrolytes
dissolved in the water
increase conductivity
and speed up corrosion
of iron
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
Rusting occurs more quickly when salt or other electrolytes are present. (click) When
salt or other electrolytes are added to water,
Salt or other electrolytes
dissolved in the water
increases conductivity
and speed up corrosion
of iron
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
The conductivity is increased, so charged particles can move faster.
Salt or other electrolytes
dissolved in the water
increases conductivity
and speeds up corrosion
of iron
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
This speeds up the corrosion of iron.
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
Chloride (Cl–) ions
present in salt solutions
tend to “eat” through any
protective films on the
iron surface
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
It is also known that chloride ions present in salt solutions tend to eat through any
protective films that may exist on the surface of iron, thus speeding up corrosion.
Thank you to Jacques
from Beloeil, Canada for
sharing this photo on Pixabay
In wintertime, salt is spread on roads to melt ice and snow. Salty water that is splashed
on vehicles can greatly increase the rate of rusting if they are unprotected.
Rate of Corrosion
Increase temperature
OH–
Anode
Region
Fe2+
Fe(s)  Fe2+ + 2e–
iron
OH–
Cathode
Region
½O2(g) + H2O + 2e–  2OH–
We should also point out that as (click) temperature increases, the rate of corrosion also
increases. All reactions increase in rate as the temperature increases.
Galvanic Corrosion
A type of corrosion we should be aware of is something called galvanic corrosion.
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Galvanic corrosion results when two dissimilar metals are attached in the presence of
an electrolyte
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Cu
Cu
Steel (iron)
For example, let’s say we have two pieces of steel (which is mainly iron) (click) bolted
together with bolts made of copper, or an alloy of copper.
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Cu
Cu
Steel (iron)
Saltwater
We’ll add some saltwater, which is an electrolyte.
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Cu
Cu
Steel (iron)
Saltwater
This establishes a type of electrochemical or galvanic cell, in which the metal lower on
the reduction table, (click) the iron in this case,
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Cu
Cu
ANODE
Steel (iron)
Fe(s)  Fe2+ + 2e–
Saltwater
Becomes the anode, and undergoes (click) oxidation, which slowly causes the steel to
corrode where the bolts contact it.
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Zn Cu Cu Cu Cu
Cu Cu Zn Cu Cu
Cu Cu Cu Cu Zn
Cu Zn Cu Cu Cu
Brass Propeller
Saltwater
Galvanic corrosion can also happen when a brass boat propeller is used in saltwater. If
we take a closer look at brass (click) we see it is an alloy of mainly copper and zinc.
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Zn Cu Cu Cu Cu
Cu Cu Zn Cu Cu
Cu Cu Cu Cu Zn
Cu Zn Cu Cu Cu
Brass Propeller
Saltwater
Because zinc is lower on the right side of the reduction table, (click) it is a stronger reducing agent and is more readily
oxidized than copper. You may recall that zinc normally forms a protective oxide layer. However, saltwater tends to
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Zn Cu Cu Cu Cu
Zn2+
Cu Cu Zn Cu Cu
Zn  Zn2+ + 2e–
Cu Cu Cu Cu Zn
Zn
2+
Zn
Cu Zn
Cu Cu Cu
2+
Brass Propeller
Zinc atoms on the surface (click) oxidize to Zn2+ ions
Saltwater
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Zn2+
Cu Cu Cu Cu
Cu Cu Zn Cu Cu
Cu Cu Cu Cu
Cu
Brass Propeller
Zn2+
Zn  Zn2+ + 2e–
Zn2+
Cu Cu Cu
Saltwater
These ions will leave the zinc and dissolve in the surrounding water,
Galvanic Corrosion
Galvanic corrosion results when
two dissimilar metals are attached
in the presence of an electrolyte
Zn2+
Cu Cu Cu Cu
Cu Cu Zn Cu Cu
Zn  Zn2+ + 2e–
Cu Cu Cu Cu
Cu
Brass Propeller
Gradually corroding the brass propeller.
Cu Cu Cu
Zn2+
Zn2+
Saltwater
.
To summarize,
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.