electro chemical corrosion.pdf

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ElectrochemicalCorrosion

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Electrochemical corrosion - Corrosion produced by thedevelopment of a current in an electrochemical cell thatremoves 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 orplating.

Oxidation reaction - The anode reaction by whichelectrons are given up to the electrochemical cell.

Reduction reaction - The cathode reaction by whichelectrons are accepted from the electrochemical cell.


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Figure 22.3 The components in an electrochemical cell: (a) asimple electrochemical cell and (b) a corrosion cell between asteel water pipe and a copper fitting.

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Figure 22.4 The anode and cathode reactions in typicalelectrolytic corrosion cells: (a) the hydrogen electrode, (b) theoxygen electrode, and (c) the water electrode.


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Electrode potential - Related to the tendency of amaterial to corrode. The potential is the voltageproduced between the material and a standardelectrode.

emf series - The arrangement of elements according totheir electrode potential, or their tendency to corrode.

Nernst equation - The relationship that describes theeffect of electrolyte concentration on the electrodepotential in an electrochemical cell.

Faradays equation - The relationship that describes therate at which corrosion or plating occurs in anelectrochemical cell.

The Electrode Potential in

Electrochemical Cells

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Figure 22.5 The half-cellused to measured theelectrode potential ofcopper under standardconditions. The electrodepotential of copper is thepotential differencebetween it and thestandard hydrogenelectrode in an opencircuit. Since E0 is greatthan zero, copper iscathodic compared withthe hydrogen electrode.


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Suppose 1 g of copper as Cu2+ is dissolved in 1000 g of waterto produce an electrolyte. Calculate the electrode potential ofthe copper half-cell in this electrolyte.

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 theatomic mass of copper) to 1000 g of water. The atomic mass ofcopper is 63.54 g/mol. The concentration of the solution whenonly 1 g of copper is added must be:

Example 22.1Half-Cell Potential for Copper

From the Nernst equation, with n = 2 and E0 = +0.34 V:

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Design a process to electroplate a 0.1-cm-thick layer of copperonto a 1 cm 1 cm cathode surface.


In order for us to produce a 0.1-cm-thick layer on a 1 cm2

surface area, the weight of copper must be:

Example 22.2Design of a Copper Plating Process

From Faradays equation, where MCu = 63:54 g/mol and n = 2:


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Therefore, we might use several different combinations ofcurrent and time to produce the copper plate:

Our choice of the exact combination of current and timemight be made on the basis of the rate of production andquality of the copper plate.

A current of ~ 1 A and a time of ~ 45 minutes arenot uncommon in electroplating operations.

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An iron container 10 cm 10 cm at its base is filled to a heightof 20 cm with a corrosive liquid. A current is produced as aresult of an electrolytic cell, and after 4 weeks, the containerhas decreased in weight by 70 g. Calculate (1) the current and(2) the current density involved in the corrosion of the iron.

Example 22.3 SOLUTION1. The total exposure time is:

Example 22.3Corrosion of Iron

From Faradays equation, using n = 2 and M = 55.847 g/mol:


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2. The total surface area of iron in contact with thecorrosive liquid and the current density are:

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Suppose that in a corrosion cell composed of copper and zinc,the current density at the copper cathode is 0.05 A/cm2. Thearea of both the copper and zinc electrodes is 100 cm2.Calculate (1) the corrosion current, (2) the current density atthe zinc anode, and (3) the zinc loss per hour.


1. The corrosion current is:

Example 22.4Copper-Zinc Corrosion Cell

2. The current in the cell is the same everywhere. Thus:


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3. The atomic mass of zinc is 65.38 g/mol. From Faradaysequation:

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Consider a copper-zinc corrosion couple. If the currentdensity at the copper cathode is 0.05 A/cm2, calculate theweight loss of zinc per hour if (1) the copper cathode areais 100 cm2 and the zinc anode area is 1 cm2 and (2) the

copper cathode area is 1 cm2

and the zinc anode area is100 cm2.

Example 22.8Effect of Areas on Corrosion Rate for

Copper-Zinc Couple


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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 significantlywhen the zinc anode is much larger than the cathode.

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Polarization - Changing the voltage between the anodeand cathode to reduce the rate of corrosion.

Activation polarization is related to the energyrequired to cause the anode or cathode reaction

Concentration polarization is related to changes in thecomposition of the electrolyte

Resistance polarization is related to the electricalresistivity of the electrolyte.

The Corrosion Current andPolarization


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A brass fitting used in a marine application is joined bysoldering with lead-tin solder. Will the brass or the soldercorrode?


From the galvanic series, we find that all of the copper-basedalloys are more cathodic than a 50% Pb-50% Sn solder. Thus,the solder is the anode and corrodes. In a similar manner, thecorrosion of solder can contaminate water in freshwaterplumbing systems with lead.

Example 22.5

Corrosion of a Soldered Brass Fitting

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Intergranular corrosion - Corrosion at grain boundariesbecause grain boundary segregation or precipitationproduces local galvanic cells.

Stress corrosion - Deterioration of a material in which anapplied stress accelerates the rate of corrosion.

Oxygen starvation - In the concentration cell, low-oxygen regions of the electrolyte cause the underlyingmaterial to behave as the anode and to corrode.

Crevice corrosion - A special concentration cell in whichcorrosion occurs in crevices because of the lowconcentration of oxygen.

Types of Electrochemical Corrosion


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Figure 22.6 Example of microgalvanic cells in two-phase alloys:(a) In steel, ferrite is anodic to cementite. (b) In austeniticstainless steel, precipitation of chromium carbide makes the lowCr austenite in the grain boundaries anodic.

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Figure 22.7 Photomicrograph of intergranular corrosion in azinc die casting. Segregation of impurities to the grainboundaries produces microgalvanic corrosion cells (x50).


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Figure 22.8 Examples of stress cells. (a) Cold work required tobend a steel bar introduces high residual stresses at the bend,which then is anodic and corrodes. (b) Because grainboundaries have a high energy, they are anodic and corrode.

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A cold-drawn steel wire is formed into a nail by additionaldeformation, producing the point at one end and the head atthe other. Where will the most severe corrosion of the nailoccur?


Since the head and point have been cold-worked an additionalamount compared with the shank of the nail, the head andpoint serve as anodes and corrode most rapidly.

Example 22.6Corrosion of Cold-Drawn Steel


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Figure 22.9 Concentration cells: (a) Corrosion occurs beneath awater droplet on a steel plate due to low oxygen concentrationin the water. (b) Corrosion occurs at the tip of a crevicebecause of limited access to oxygen.

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Two pieces of steel are joined mechanically by crimpingthe edges. Why would this be a bad idea if the steel isthen exposed to water? If the water contains salt, wouldcorrosion be affected?

Example 22.7 SOLUTIONBy crimping the steel edges, we produce a crevice. Theregion in the crevice is exposed to less air and moisture,so it behaves as the anode in a concentration cell. Thesteel in the crevice corrodes.

Salt in the water increases the conductivity of thewater, permitting electrical charge to be transferred at amore rapid rate. This causes a higher current density and,thus, faster corrosion due to less resistance polarization.

Example 22.7Corrosion of Crimped Steel


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Figure 22.10 (a) Bacterialcells growing in a colony(x2700). (b) Formation ofa tubercule and a pitunder a biological colony.

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Inhibitors - Additions to the electrolyte that preferentiallymigrate to the anode or cathode, cause polarization, andreduce the rate of corrosion.

Sacrificial anode - Cathodic protection by which a moreanodic material is connected electrically to the materialto be protected. The anode corrodes to protect thedesired material.

Passivation - Producing strong anodic polarization by

causing a protective coating to form on the anodesurface and to thereby interrupt the electric circuit.

Protection Against

Electrochemical Corrosion

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Figure 22.11 Alternative methods for joining two pieces ofsteel: (a) Fasteners may produce a concentration cell, (b)brazing or soldering may produce a composition cell, and (c)welding with a filler metal that matches the base metal mayavoid the formation of galvanic cells (for Example 22.8)


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Figure 22.12 Zinc-plated steel and tin-plated steel areprotected differently. Zinc protects steel even when the

coating is scratched, since zinc is anodic to steel. Tin doesnot protect steel when the coating is disrupted, since steel isanodic with respect to tin.

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Figure 22.13 Cathodic protection of a buried steel pipeline:(a) A sacrificial magnesium anode assures that the galvaniccell makes the pipeline the cathode. (b) An impressed voltagebetween a scrap iron auxiliary anode and the pipeline assuresthat the pipeline is the cathode.


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Figure 22.14 (a) Intergranularcorrosion takes place in austeniticstainless steel. (b) Slow coolingpermits chromium carbides toprecipitate at grain boundaries.(c) A quench anneal to dissolvethe carbides may preventintergranular corrosion.

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Steel troughs are located in a field to provide drinking waterfor a herd of cattle. The troughs frequently rust through andmust be replaced. Design a system to prevent or delay thisproblem.


We might, for example, fabricate the trough using stainlesssteel or aluminum. Either would provide better corrosionresistance than the plain carbon steel, but both areconsiderably more expensive than the current material.

We might suggest using cathodic protection; a smallmagnesium anode could be attached to the inside of thetrough. The anode corrodes sacrificially and preventscorrosion of the steel.

Example 22.9Design of a Corrosion Protection System


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Example 22.9 SOLUTION (Continued)

Another approach would be to protect the steel troughusing a suitable coating. Painting the steel (that is, introducinga 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 galvanizedsteel, taking advantage of the protective coating and thesacrificial behavior of the zinc. Corrosion is very slow dueto 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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