Materials 2H03 on-line Laboratory Manual
Updated: 24 May 2000
Experiment L
Corrosion
Objective
1.
• To examine and understand the basic electrode processes that control the
corrosion rate of a metal.
Introduction
2.
Corrosion can be defined as the deterioration of a material because of reaction with
its environment. Although the cost of such loss is extremely high each year, the
same electrochemical processes, if used under controlled conditions, can both refine
materials and provide corrosion protection. Corrosion resistance depends on many
factors. Thermodynamics indicates the direction of the reaction, electrochemical
kinetics control its rate, and microstructural features often determine its location.
This introduction describes the simple electrode processes that take place during
corrosion of metals in aqueous solutions.
2.a:
Single Electrode Processes
In general, when one material is placed in contact with another, the atoms will
diffuse into the contacted material until the equilibrium solubility is reached. Thus,
when a metal plate is put in water, some of the metal atoms will dissolve in the
water. These become ionized, i.e., lose one or more electrons, and these remain on
the metal. The positive M+ ions, will be attracted back to the negatively charged
plate to produce a charged ionic layer that is near the plate. The reaction is
M → M+ + e- ....(1)
which is called an "anodic" or "oxidation" reaction, since electrons are produced. If
nothing else occurred, a negative charge would build up on the plate that would
stop the reaction. However, as the concentration of M+ builds up, some of the ions
will move back to the metal, i.e.,
M+ + e- → M ....(2)
which is called a "cathodic" or "reduction" reaction, since electrons are consumed.
An equilibrium will be reached when the flow of M+ ions from reaction (2) arriving
at the metal equals the flow from reaction (1) that leave the metal. We can
represent this flow as an electrical current, io, which flows in both directions.
Because the ions must overcome a potential barrier in the jumping process (onto or
Single Electrode Processes
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off the metal), io is proportional to exp(-Ej/kT). Ej is the jump energy required, k is
Boltzmann's constant, and T is degrees Kelvin.
At equilibrium, an excess number of electrons has built up on the metal plate,
which can be measured with a voltmeter whose other terminal is connected to a
reference potential. "Voltage" or "electrostatic potential" is the potential energy per
unit positive charge. Thus, the more ions that dissolve; the more excess electrons on
the plate; and the more negative its voltage. The voltages produced by different
metals are listed as the electrochemical series.
All practical corrosion situations are not as ideal as described above. The water can
have three types of "impurities" that may replace part of reaction (2). They all
produce cathodic reactions.
Excess H+, as in an acid, will cause
H+ + e - → H ....(3)
followed by the formation of hydrogen molecules
H + H → H2....(4)
Dissolved O2 in acidic solutions will cause
O2 + 4H+ +4e- → 2H2O ....(5)
and in basic solutions the predominant cathodic reaction will be
O2 + 2H2O + 4e- → 4OH- ....(6)
These "impurities" will have two effects in our previous description. Reactions 3, 5,
and 6 can replace part of reaction 2. Thus, a steady state dissolution (equation 1)
will occur as the loss of metal atoms is not canceled by reaction 2. The number of
excess electrons on the metal plate will also decrease, which will raise the voltage
on the plate to a new volume determined by an equal rate of the anodic reaction 1,
with the sum of cathodic reactions 2, 3, 5, and 6.
Introduction
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2.b:
Galvanic Cells
An open-circuited galvanic cell is produced by putting two different metals, that are
not in electrical contact, into water as shown in the sketch. Metal M1 which
dissolves more readily than M2, will have more excess electrons and be at a lower
voltage. Thus, a positive deflection will be seen on a voltmeter if M1 is connected to
the negative terminal and M2 to the positive terminal. M2 is said to be more
"noble" than M1, because it does not dissolve as readily. By comparing different
pairs of metals, they can be arranged in a sequence of nobility, called the
Electromotive Force Series.
In the present experiments, NaCl is added to increase the conductivity of the
electrolyte solution.
In a short-circuited galvanic cell, the metals M1 and M2 are in electrical contact,
e.g. by a wire connecting them, as in the sketch of Figure 2. Before the wire is
connected (Figure 1), M1 had more excess electrons than M2. Thus, after
connection, electrons will flow from M1 to M2.
Reducing the negative charge on M1 will make it easier for its anodic reaction,
M1 → M1++ e-
to proceed. Similarly, the flow of electrons to M2 will reduce the rate of its anodic
reaction
M2 → M2++ e-.
There are some important practical applications of these processes.
Cathodic protection of metal M2 has occurred since the rate of its dissolution
reaction is reduced. Zinc plates attached to steel will protect the steel. In this case
Galvanic Cells
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zinc, or M1 in our example, would function as a sacrificial anode. A similar result
occurs if a metal is made cathodic by attaching it to the negative terminal of a
battery, which is sometimes done for buried pipelines.
In a wet environment, you should not have two metals, which have a large
separation in the electromotive force series, in contact with each other, e.g.,
couplings between pipes in the ground, nails holding the metal roofing, etc.
A galvanic cell shunted by a resistor is similar to the short-circuited cell, but
the current flow is limited by the value of the resistance. Before contact, there are
more excess electrons on M1 than M2. After contact, the electrons will flow through
the resistor at a rate inversely proportional to the resistance. This transfer of
electrons, enhances the anodic reaction
M1 → M1+ + e-
on M1, and the cathodic reactions on M2. Its continuation as a current depends on
both the continual supply of electrons by the anodic reaction on M1, and removal of
electrons by the cathodic reaction at M2.
A galvanic cell driven by an external voltage At open circuit, M1 and M2 have
their own equilibrium anodic and cathodic currents, io, which are equal and
opposite, and proportional to exp(-Ej/kT) as discussed previously. In order to force
M1 to have a greater anodic current requires an added voltage η , called the over
voltage. The net current will then be
η
−η 

i = i 0 exp( ) − exp( )
β
β 

where b depends on kT and the charge on the M1 ion. Thus, for a given net current,
i, the actual voltage needed across the cell (using only the first exponential term) is
V = V 0 + η = V 0 + βln(
i
)
i0
where Vo is the open-circuit voltage between M1 and M2. Therefore, a plot of V
applied vs ln i should be linear, a so-called Tafel plot. This is called "activation
polarization" and is not dependent on time.
2.c:
Concentration Polarization
This type of polarization usually increases with time, because it depends on
diffusion of ions in the liquid. Dissolved ions in the liquid will be located in places
Introduction
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L-Corrosion
that minimize the free energy of the system. This produces an ionic concentration
that decreases exponentially with distance from the metal surface, as in the sketch.
If the arrival or removal rate of ions from the metal is slow, they have time to
redistribute by diffusion in the liquid. (D is commonly 10-5cm2 s-1 in a liquid). If the
current is too high, they either pile up near the metal surface for a net anodic
current, or are depleted near the surface for a net cathodic current. This causes the
current to drop with time if the voltage is held constant, or causes the over potential
to increase with time, if the current is held constant.
2.d:
Inhibitors and Passivation
All the above arguments assumed a relatively easy transfer of ions across the
interface between the metal and liquid. If this transfer is made difficult, then the
corrosion rate is lower. One way to inhibit this transfer is to add complex molecules
that chemisorb on the metal surface. For example, an inhibitor is added to the
coolant for cars which have an engine block made from aluminum and radiator or
couplings made from other metals. Another way to reduce this transfer of ions is to
develop a chemisorbed or oxide layer during the corrosion process. This method,
called "Passivation", occurs naturally beyond a critical current density for some
metals, if there is a good supply of dissolved oxygen in the water.
A typical curve of applied voltage vs. ln i is shown here. Low voltages give the usual
straight line, which indicates a logarithmic increase in corrosion rate with applied
potential, but the current drops suddenly once a critical current is reached.
Current and hence corrosion almost cease for applied voltages within this "passive"
region. At much higher potentials oxygen evolution gives a linear rise again in the
"transpassive" region. If the very minimal current in the passive region is not
maintained, the chemisorbed oxygen or oxide on the metal slowly dissolves.
Ni, Ti, Cr, and many stainless steels containing Cr, have very low corrosion rates
due to this Passivation process. If the water or atmosphere is sufficiently oxidizing,
they will be in the passive region without an applied voltage.
Inhibitors and Passivation
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2.e:
Non-uniform "pitting" and "crevice" corrosion
If there are non-uniformities in the chemical composition of either the material or
the environment at its surface, then both local anodes and cathodes can exist on the
same piece of metal. The region that is locally anodic will corrode forming a pit. The
inhomogeneities are frequently caused by depletion of one of the reactants. For
instance, the oxygen supply under the centre of a water droplet, or at the bottom of
the crevices is quickly used up. This slows the local cathodic reaction, which makes
it locally anodic with respect to other areas.
Reference: Jastrzebski, "The Nature and Properties of Engineering Materials", J.
Wiley, 1987.
Procedure for the First Afternoon
3.
Read parts 1-5, and 7 of the introduction before the first lab period. If uncertain
about anything ask the demonstrator before you do the experiments.
3.a:
Galvanic couples
Sheets of Cu, Zn, Sn and mild steel are provided for electrodes. Clean them with
abrasive paper. Place each possible combination (6 pairs) of the four electrodes into
a 3% NaCl solution that fills one third of a 600 ml beaker. Measure the peak voltage
generated and the sign of each electrode. If the voltage is still increasing after 2
minutes, take the value at 2 minutes. (A voltmeter draws very little current
because it has an extremely high resistance). Knowing the polarity, deduce the
direction of the electron flow, and decide which metal is the anode and cathode in
each pair. Which metal in each pair is corroding? Using your results, list these four
metals in a galvanic series from least noble to most noble. Note: the "electric
current" is defined as the flow of positive charge from a positive to negative voltage.
The meter reads this electric current. Electrons will be flowing in the opposite
direction, from the negative terminal to the positive terminal. For a positive reading,
Procedure for the First Afternoon
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L-Corrosion
the positive terminal on the voltmeter or ammeter is connected to the positive
electrode (i.e., the cathode).
3.b:
Concentration polarization
Using the Cu and Zn electrodes, set up the circuit shown. Measure the open-circuit
voltage (i.e., switch not closed). Now set the resistance to 1000 ohms, close the
switch and record the current and voltage at 15 second intervals for 5 minutes.
Remeasure the open circuit voltage. If this does not agree with earlier readings, stir
the electrolyte with the electrodes and try again.
Repeat this for resistance of 500 and 200 ohms, including the remeasurement of
open circuit voltage. Add a few ml of peroxide (H2O2), which forms H2O+ ½ O2
when put in water, and repeat the experiment using 200 ohms resistance. Clean
and dry all electrodes. Throw out the electrolyte.
In your writeup, compare the results by plotting current and voltage vs. time.
Explain the results in terms of the flow of electrons, and of specific ions at the
cathode surface. Also explain why stirring the electrolyte restores the voltage.
3.c:
Anodes and cathodes
At the beginning of the session, prepare the mixed indicator gel containing 7.5 g
NaCl dissolved in 250 ml distilled water. Boil it and add 5 g agar. Boil and stir until
agar dissolves. Add 5 ml potassium ferricyanide and 5 ml of 1% phenolphthalein
solution.
Pour some of the fluid gel into a petri dish and drop into it, without disturbance,
Samples
• A short piece of the steel wire.
Anodes and cathodes
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• A similar piece that has been bent back and forth many times to form a
sharp U shape, and
• A steel/zinc combination
• A steel/tin combination.
Clean an area 75 mm square in the centre of a piece of steel sheet with abrasive
paper. Pour into the middle of this a small pool of the indicator gel and immediately
press the convex face of the watch onto the steel surface to make contact..
Place a few drops of indicator on the cleaned surface of another steel sheet.
Observe what you see during the afternoon, and save the samples for another look
at the second session.
Note: In the presence of Fe++, the ferricyanide will turn blue. The
phenolphthalein turns pink with excess OH-.
With this knowledge, plus the observations from part (b), and considering the
crevice between the glass and steel, explain the pink, white, and blue colour
changes observed.
When your examination is complete, remember to remove the specimens,
wash and dry them, put the gel in the container provided for disposal,
wash the Petri dish !
In the next experiments, the change in potential (voltage), of the steel is measured by
comparison with a reference potential, that is produced by a Cu plate immersed in a
saturated Cu2SO4 solution, i.e., a simple reference electrode.
Procedure for the First Afternoon
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L-Corrosion
3.d:
Sacrificial Anode
1.
Measure the weight of the ZINC electrode.
2. Place the steel electrode into a 3% NaCl solution and set up
the Cu/Cu2SO4 reference cell and circuit as shown by solid
lines in the accompanying figure. Connect the two beakers
with the salt bridge. Measure current and voltage. Record
polarity.
3. Connect an immersed piece of Cu wire to "above" the
ammeter, as shown dotted in the sketch. Remeasure current,
voltage and record polarity.
4. Connect an immersed piece of Zn sheet to "above" the
ammeter, as shown dotted in the sketch. Remeasure current,
voltage and record polarity.
5. Remove the Zn sheet. Measure the current and voltage. Add a
commercial inhibitor and remeasure the current and voltage.
6. Remove all the steel and Cu electrodes. Wash them and
measure the approximate total immersed surface area of the
steel and zinc sheet.
7. Re-Measure the weight of the ZINC electrode.
In your writeup, explain the voltage, current, and polarity changes in parts (iii)
and (iv). Which metal is corroding in part (iii) and in part (iv)? If the steel sample
represents a ship hull to be protected, what would be the weight of zinc anode per
m2 of structure consumed in a year to protect it?
According to Faraday's Law:
Sacrificial Anode
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L-Corrosion
(weightofzincdissolved )
it
=
(GAW Zn )
nF
where GAW = gramatomic weight, i = current in amps, t = seconds, n = charge on Zn
ion, and F = Faraday's constant =96,500 coulomb per mole.
If the steel is protected now by a paint layer, whose pores cover 0.1% of the area,
what would be the new consumption of zinc?
Procedure for the Second Afternoon
4.
4.a:
Polarization Curve
The Measurement of the Polarization Curve of 430 Stainless Steel with a
Potentiostat (HA-151)
1. Read part 6 of the introduction and the description of the
potentiostat used in this experiment (available in the lab).
2. Before connecting or disconnecting the power cable, be sure to
turn off the power switch of the potentiostat.
3. There are two pieces of 430 stainless steel to be used as
electrodes. Set them up with the straight piece connected to
the WE1 abd WE2 of the TO-CELL-OUT cable of the
potentiostat (one connection is used to measure voltage while
the other is used to measure current), and the curved piece
connected to the CE terminal.
Procedure for the Second Afternoon
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L-Corrosion
4. The reference electrode to be used is Cu/Cu2SO4. Connect the
piece of copper to the RE terminal of the TO-CELL-OUT cable
of the potentiostat.
5. Pour the 6% H2SO4 solution into the beaker with the two
pieces of stainless steel, and pour the Cu2SO4 solution into
the beaker with the piece of copper. Connect the two beakers
with a salt bridge.
6. Before turning on the potentiostat, make sure the
FUNCTION selector is set to ZERO ADJ. Turn on the
potentiostat.
7. Set the P-STAT/G-STAT selector to P-STAT (potentiostat),
and set the EXT.SET selector to the OFF position. The
potentiostat is now ready to be used.
The starting value of the potential will be the equilibrium potential,
defined as the potential of the working electrode with respect to that
of the reference electrode at zero current.
8. To measure the equilibrium potential, the FUNCTION
selector should be set to REST POT, and the
POTENTIAL/CURRENT selector should be set at 2V. The
green voltage light should be on.
9. The polarity can be set at either + or - using the POLARITY
selector. A positive potential reading with the POLARITY set
to + , means that the potential of the working electrode is
higher than that of the reference electrode. The potential
value is displayed on the digital meter on the front panel of
the potentiostat.
10. Once the equilibrium potential has been measured, set the
FUNCTION selector to OPERATION. You can now maintain
the potential at any value you choose.
11. Increase the potential by 50 mV. Take readings of both
potential and current from the digital meter. To obtain the
current values, switch the POTENTIAL/CURRENT selector
to CURRENT with a suitable current range. If the current
reading is unstable, take the current reading after 30 seconds.
12. Repeat step (k) until you have obtained enough data to plot
the Passivation curve.
13. Dissolve 10 g of NaCl in the H2SO4 solution. Repeat steps (h)
through (l). The presence of Cl- is expected to diminish the
passivity range and increase the passive current density.
Once you have completed the above steps, you must close down the
system and clean the equipment:
Polarization Curve
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14. Set the FUNCTION selector to ZERO ADJ. Turn off the
power to the potentiostat.
15. Return the Cu2SO4 solution to the bottle. Remove the salt
bridge, and store it in a beaker filled with distilled water.
Clean the metal sheets and the beakers used.
16. Measure the approximate immersed surface area of the
straight piece of stainless steel used as the working electrode
(don’t forget to account for both sides of the steel).
5.
Writeup
Plot the data obtained for both cases (i.e., with and without NaCl) as E vs log i ,
where E and i are potential and current density (A/cm2) respectively.
Answer the following questions for each curve:
1. What is the potential range corresponding to passivity?
2. What is the minimum potential necessary to preserve
passivity?
3. Discuss the differences observed in the two curves.
Writeup
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