CORROSION BY The Science And Engineering Of Materials 6th Edition

Teks penuh

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and other power utilities, just to name a few. Corrosion can be prevented or contained using different tech-niques that include the use of coatings and anodic or cathodic protection.

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Have You Ever Wondered?

Why does iron rust?

What does the acronym “WD-40

TM

” stand for?

Is the process of corrosion ever useful?

Why do household water-heater tanks contain Mg rods?

Do metals like aluminum and titanium undergo corrosion?

How does wear affect the useful life of different components such

as crankshafts?

What process makes use of mechanical erosion and chemical corrosion in

the manufacture of semiconductor chips?

Corrosion and Wear

T

he composition and physical integrity of a solid material is altered in a corrosive

environment. In chemical corrosion, a corrosive liquid dissolves the material. In

electrochemical corrosion, metal atoms are removed from the solid material

as the result of an electric circuit that is produced. Metals and certain ceramics react

with a gaseous environment, usually at elevated temperatures, and the material may

be destroyed by the formation of oxides or other compounds. Polymers degrade when

exposed to oxygen at elevated temperatures. Materials may be altered when exposed

to radiation or even bacteria. Finally, a variety of wear and wear-corrosion mechanisms

alter the shape of materials. According to a study concluded in 2001, in the United

States, combating corrosion costs about 6% of the gross domestic product (GDP).

This amount, which includes direct and indirect costs, was approximately $550 billion

in 1998.

The corrosion process occurs in order to lower the free energy of a system.

The corrosion process occurs over a period of time and can occur either at high or low

temperatures. Chemical corrosion is an important consideration in many sectors

including transportation (bridges, pipelines, cars, airplanes, trains, and ships), utilities

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(electrical, water, telecommunications, and nuclear power plants), and production and

manufacturing (food industry, microelectronics, and petroleum refining).

In some applications, corrosion or oxidation is useful. The processes of

chem-ical corrosion and erosion are used to make ultra-flat surfaces of silicon wafers for

computer chips. Similarly, degradation and dissolution of certain biopolymers is

use-ful in some medical applications, such as dissolvable sutures.

The goal of this chapter is to introduce the principles and mechanisms by

which corrosion and wear occur under different conditions. This includes the aqueous

corrosion of metals, the oxidation of metals, the corrosion of ceramics, and the

degra-dation of polymers. We will offer a summary of different technologies that are used

to prevent or minimize corrosion and associated problems.

23-1

Chemical Corrosion

In

chemical corrosion

, or direct dissolution, a material dissolves in a corrosive liquid

medium. The material continues to dissolve until either it is consumed or the liquid is

sat-urated. An example is the development of a green patina on the surface of copper-based

alloys. This is due to the formation of copper carbonate and copper hydroxides and is

why, for example, the Statue of Liberty looks greenish. The chemical corrosion of copper,

tantalum, silicon, silicon dioxide, and other materials can be achieved under extremely

well-controlled conditions. In the processing of silicon wafers, for example, a process

known as chemical mechanical polishing uses a corrosive silica-based slurry to provide

mechanical erosion. This process creates extremely flat surfaces that are suitable for the

processing of silicon wafers. Chemical corrosion also occurs in nature. For example, the

chemical corrosion of rocks by carbonic acid (H

2

CO

3

) and the mechanical erosion of

wind and water play an important role in the formation of canyons and caverns.

Liquid Metal Attack

Liquid metals first attack a solid at high-energy

locations such as grain boundaries. If these regions continue to be attacked preferentially,

cracks eventually grow (Figure 23-1). Often this form of corrosion is complicated by the

presence of fluxes that accelerate the attack or by electrochemical corrosion. Aggressive

metals such as liquid lithium can also attack ceramics.

Figure 23-1

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2 3 - 1 Chemical Corrosion

853

Figure 23-2

Micrograph of a copper deposit in brass, showing the effect of dezincification (*50). (Reprinted courtesy of Don Askeland.)

Selective Leaching

One particular element in an alloy may be selectively

dissolved, or leached, from the solid.

Dezincification

occurs in brass containing more than

15% Zn. Both copper and zinc are dissolved by aqueous solutions at elevated

tempera-tures; the zinc ions remain in solution while the copper ions are replated onto the brass

(Figure 23-2). Eventually, the brass becomes porous and weak.

Graphitic corrosion

of gray cast iron occurs when iron is selectively dissolved in

water or soil, leaving behind interconnected graphite flakes and a corrosion product.

Localized graphitic corrosion often causes leakage or failure of buried gray iron gas lines,

sometimes leading to explosions.

Dissolution and Oxidation of Ceramics

Ceramic

refracto-ries used to contain molten metal during melting or refining may be dissolved by the slags

that are produced on the metal surface. For example, an acid (high SiO

2

) refractory is

rap-idly attacked by a basic (high CaO or MgO) slag. A glass produced from SiO

2

and Na

2

O

is rapidly attacked by water; CaO must be added to the glass to minimize this attack.

Nitric acid may selectively leach iron or silica from some ceramics, reducing their strength

and density. As noted in Chapters 7 and 15, the strength of silicate glasses depends on

flaws that are often created by corrosive interactions with water.

Chemical Attack on Polymers

Compared to metals and oxide

ceram-ics, plastics are considered corrosion resistant. Teflon

TM

and Viton

TM

are some of the most

corrosion-resistant materials and are used in many applications, including the chemical

pro-cessing industry. These and other polymeric materials can withstand the presence of many

acids, bases, and organic liquids. Aggressive solvents do, however, often diffuse into

low-molecular-weight thermoplastic polymers. As the solvent is incorporated into the polymer,

the smaller solvent molecules force apart the chains, causing swelling. The strength of the

bonds between the chains decreases. This leads to softer, lower-strength polymers with low

glass-transition temperatures. In extreme cases, the swelling leads to stress cracking.

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23-2

Electrochemical Corrosion

Electrochemical corrosion

, the most common form of attack of metals, occurs when metal

atoms lose electrons and become ions. As the metal is gradually consumed by this process,

a byproduct of the corrosion process is typically formed. Electrochemical corrosion occurs

most frequently in an aqueous medium, in which ions are present in water, soil, or moist

air. In this process, an electric circuit is created, and the system is called an

electrochemi-cal cell

. Corrosion of a steel pipe or a steel automobile panel, creating holes in the steel

and rust as the byproduct, are examples of this reaction.

Although responsible for corrosion, electrochemical cells may also be useful.

By deliberately creating an electric circuit, we can

electroplate

protective or decorative

coatings onto materials. In some cases, electrochemical corrosion is even desired. For

example, in etching a polished metal surface with an appropriate acid, various features

in the microstructure are selectively attacked, permitting them to be observed. In fact,

most of the photographs of metal and alloy microstructures in this text were obtained

in this way, thus enabling, for example, the observation of pearlite in steel or grain

boundaries in copper.

Components of an Electrochemical Cell

There are four

components of an electrochemical cell (Figure 23-3):

1. The

anode

gives up electrons to the circuit and corrodes.

2. The

cathode

receives electrons from the circuit by means of a chemical, or cathode,

reaction. Ions that combine with the electrons produce a byproduct at the cathode.

3. The anode and cathode must be electrically connected, usually by physical contact,

to permit the electrons to flow from the anode to the cathode and continue the

reaction.

4. A liquid

electrolyte

must be in contact with both the anode and the cathode. The

electrolyte is conductive, thus completing the circuit. It provides the means by

which metallic ions leave the anode surface and move to the cathode to accept the

electrons.

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2 3 - 2

Electrochemical Corrosion

855

This description of an electrochemical cell defines electrochemical corrosion. If

metal ions are deposited onto the cathode,

electroplating

occurs.

Anode Reaction

The anode, which is a metal, undergoes an

oxidation

reac-tion

by which metal atoms are ionized. The metal ions enter the electrolytic solution, while

the electrons leave the anode through the electrical connection:

M

:

M

n+

+

ne

-

(23-1)

Because metal ions leave the anode, the anode corrodes, or oxidizes.

Cathode Reaction in Electroplating

In electroplating, a

cathodic

reduction reaction,

which is the reverse of the anode reaction, occurs at the

cathode:

M

n+

+

ne

-

:

M

(23-2)

The metal ions, either intentionally added to the electrolyte or formed by the anode

reac-tion, combine with electrons at the cathode. The metal then plates out and covers the

cath-ode surface.

Cathode Reactions in Corrosion

Except in unusual conditions,

plating of a metal does not occur during electrochemical corrosion. Instead, the reduction

reaction forms a gas, solid, or liquid byproduct at the cathode (Figure 23-4).

1.

The hydrogen electrode

: In oxygen-free liquids, such as hydrochloric acid (HCl) or

stagnant water, hydrogen gas may be evolved at the cathode:

2H

+

+

2

e

-

:

H

2

q

(gas)

(23-3)

If zinc were placed in such an environment, we would find that the overall reaction is

Zn

:

Zn

2+

+

2

e

-

(anode reaction)

2H

+

+

2

e

-

:

H

2

q

(cathode reaction)

(23-4)

Zn

+

2H

+

:

Zn

2+

+

H

2

q

(overall reaction)

The zinc anode gradually dissolves, and hydrogen bubbles continue to evolve at the

cathode.

2.

The oxygen electrode

: In aerated water, oxygen is available to the cathode, and

hydroxyl, or (OH)

-

, ions form:

(23-5)

The oxygen electrode enriches the electrolyte in (OH)

-

ions. These ions react with

posi-tively charged metallic ions and produce a solid product. In the case of rusting of iron:

(23-6)

Fe

+

1

2

O

2

+

H

2

O

:

Fe(OH)

2

(overall reaction)

Fe

2+

+

2(OH)

-

:

Fe(OH)

2

(cathode reactions)

1

2

O

2

+

H

2

O

+

2

e

-

:

2(OH)

-Fe

:

Fe

2+

+

2

e

-

(anode reaction)

1

2

O

2

+

H

2

O

+

2

e

-

:

2(OH)

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The reaction continues as the Fe(OH)

2

reacts with more oxygen and water:

(23-7)

Fe(OH)

3

is commonly known as

rust.

3.

The water electrode

: In oxidizing acids, the cathode reaction produces water as a

byproduct:

O

2

+

4H

+

+

4

e

-

:

2H

2

O

(23-8)

If a continuous supply of both oxygen and hydrogen is available, the water electrode

pro-duces neither a buildup of solid rust nor a high concentration or dilution of ions at the

cathode.

2Fe(OH)

2

+

1

2

O

2

+

H

2

O

:

2Fe(OH)

3

n

n

n

n

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2 3 - 3 The Electrode Potential in Electrochemical Cells

857

23-3

The Electrode Potential in Electrochemical

Cells

In corrosion, a potential naturally develops when a material is placed in a solution. Let’s

see how the potential required to drive the corrosion reaction develops.

Electrode Potential

When a pure metal is placed in an electrolyte, an

electrode potential

develops that is related to the tendency of the material to give up its

electrons; however, the driving force for the oxidation reaction is offset by an equal but

opposite driving force for the reduction reaction. No net corrosion occurs. Consequently,

we cannot measure the electrode potential for a single electrode material.

Electromotive Force Series

To determine the tendency of a metal

to give up its electrons, we measure the potential difference between the metal and a

standard electrode using a half-cell (Figure 23-5). The metal electrode to be tested is placed

in a 1 molar (M) solution of its ions. A reference electrode is also placed in a 1 M solution

of ions. The reference electrode is typically an inert metal that conducts electrons but

does not react with the electrolyte. The use of a 1 M solution of hydrogen (H

+

) ions with

the reference electrode is common. The reaction that occurs at this hydrogen electrode is

2H

+

+

2

e

-

=

H

2

. H

2

gas is supplied to the hydrogen electrode. An electrochemical cell such

as this is known as a standard cell when the measurements are made at 25°C and

atmos-pheric pressure with 1 M electrolyte concentrations. The two electrolytes are in electrical

contact but are not permitted to mix with one another. Each electrode establishes its own

electrode potential. By measuring the voltage between the two electrodes when the circuit

is open, we obtain the potential difference. The potential of the hydrogen electrode is

arbi-trarily set equal to zero volts. If the metal has a greater tendency to give up electrons than

the hydrogen electrode, then the potential of the metal is negative—the metal is anodic

with respect to the hydrogen electrode. If the potential of the metal is positive, the metal

is cathodic with respect to the hydrogen electrode.

The

electromotive force

(or

emf

)

series

shown in Table 23-1 compares the standard

electrode potential

E

0

for each metal with that of the hydrogen electrode under standard

Figure 23-5

The half-cell used to measure 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 E0is greater than zero, copper is cathodic

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conditions of 25°C and a 1 M solution of ions in the electrolyte. Note that the

measure-ment of the potential is made when the electric circuit is open. The voltage difference begins

to change as soon as the circuit is closed.

The more negative the value of potential for the oxidation of metal, the more

elec-tropositive is the metal; this means the metal will have a higher tendency to undergo an

oxi-dation reaction. For example, alkali and alkaline earth metals (e.g., Li, K, Ba, Sr, and Mg)

are so reactive that they have to be kept under conditions that prevent any contact with

oxy-gen. On the other hand, metals that are toward the bottom of the chart (e.g., Ag, Au, and

Pt) will not tend to react with oxygen. This is why we call them “noble metals.” Metals such

as Fe, Cu, and Ni have intermediate reactivities; however, this is not the only consideration.

Aluminum, for example, has a strongly negative standard electrode potential and does react

easily with oxygen to form aluminum oxide; it also reacts easily with fluoride to form

alu-minum fluoride. Both of these compounds form a tenacious and impervious layer that helps

stop further corrosion. Titanium also reacts readily with oxygen. The quickly formed

tita-nium oxide creates a barrier that prevents the further diffusion of species and, thus, avoids

further oxidation. This is why both aluminum and titanium are highly reactive, but can resist

corrosion exceptionally well. Note that the emf series tells us about the thermodynamic

fea-sibility and driving force, it does

not

tell us about the kinetics of the reaction.

Effect of Concentration on the Electrode Potential

The electrode potential depends on the concentration of the electrolyte. At 25°C, the

Nernst equation

gives the electrode potential in nonstandard solutions:

(23-9)

where

E

is the electrode potential in a solution containing a concentration

C

ion

of the

metal in molar units,

n

is the charge on the metallic ion, and

E

0

is the standard electrode

potential in a 1 M solution. Note that when

C

ion

=

1,

E

=

E

0

. The example that follows

illustrates the calculation of the electrode potential.

E

=

E

0

+

0.0592

n

log (

Cion

)

TABLE 23-1The standard reduction potentials for selected elements and reactions

Metal Electrode Potential E0(Volts)

Anodic Li++e-:Li -3.05

Mg2++2e-:Mg -2.37

Al3++3e-:Al -1.66

Ti2++2e-:Ti -1.63

Mn2++2e-:Mn -1.63

Zn2++2e-:Zn -0.76

Cr3++3e-:Cr -0.74

Fe2++2e-:Fe -0.44

Ni2++2e-:Ni -0.25

Sn2++2e-:Sn -0.14

Pb2++2e-:Pb -0.13

2H++2e-:H

2 0.00 — (defined)

Cu2++2e-:Cu +0.34

O2+2H2O +4e-:4OH- +0.40

Ag++e-:Ag +0.80

Pt4++4e-:Pt +1.20

O2+4H++4e-:2H2O +1.23

(11)

2 3 - 3 The Electrode Potential in Electrochemical Cells

859

Example 23-1

Half-Cell Potential for Copper

Suppose 1 g of copper as Cu

2+

is dissolved in 1000 g of water to produce an

elec-trolyte. Calculate the electrode potential of the copper half-cell in this elecelec-trolyte.

Assume

T

=

25°C.

SOLUTION

From chemistry, we know that a standard 1 M solution of Cu

2+

is obtained when we

add 1 mol of Cu

2+

(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

solu-tion when only 1 g of copper is added must be

From the Nernst equation, with

n

=

2 and

E

0

= +

0.34 V,

Rate of Corrosion or Plating

The amount of metal plated on the

cathode in electroplating, or removed from the anode by corrosion, can be determined

from

Faraday’s equation

,

(23-10)

where

w

is the weight plated or corroded (g),

I

is the current (A),

M

is the atomic mass of

the metal,

n

is the charge on the metal ion,

t

is the time (s), and F is Faraday’s constant

(96,500 C). This law basically states that one gram equivalent of a metal will be deposited

by 96,500 C of charge. Often the current is expressed in terms of current density,

i

=

I>A

,

so Equation 23-10 becomes

(23-11)

where the area

A

(cm

2

) is the surface area of the anode or cathode.

The following examples illustrate the use of Faraday’s equation to calculate the

current density.

w

=

iAtM

n

F

w

=

ItM

n

F

=

0.34

+

(0.0296)(

-

1.8)

=

0.29

V

E

=

E

0

+

0.0592

n

log(

C

ion

)

=

0.34

+

0.0592

2

log(0.0157)

Cion

=

1

63.54

=

0.0157

M

Example 23-2

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 cm

2

surface area, the mass of

copper must be

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Volume of copper

=

(1 cm

)(0.1 cm)

=

0.1 cm

Mass of copper

=

(8.93 g

>

cm

3

)(0.1 cm

3

)

=

0.893 g

From Faraday’s equation, where

M

Cu

=

63.54 g

>

mol and

n

=

2:

Therefore, we might use several different combinations of current and time to

pro-duce the copper plate:

Current Time

0.1 A 27,124 s =7.5 h 1.0 A 2,712 s =45.2 min 10.0 A 271.2 s =4.5 min 100.0 A 27.12 s =0.45 min

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 plated copper. Low currents

require very long plating times, perhaps making the process economically

unsound. High currents, however, may reduce plating efficiencies. The plating

effectiveness may depend on the composition of the electrolyte containing the

cop-per ions, as well as on any impurities or additives that are present. Currents such

as 10 A or 100 A are too high—they can initiate other side reactions that are not

desired. The deposit may also not be uniform and smooth. Additional background

or experimentation may be needed to obtain the most economical and efficient

plating process. A current of

!

1 A and a time of

!

45 minutes are not uncommon

in electroplating operations.

It

=

wn

F

M

=

(0.893)(2)(96,500)

63.54

=

2712

A

#

s

Example 23-3

Corrosion of Iron

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

four weeks, the container has decreased in weight by 70 g. Calculate (a) the current

and (b) the current density involved in the corrosion of the iron.

SOLUTION

(a) The total exposure time is

t

=

(4 wk)(7 d

>

wk)(24 h

>

d)(3600 s

>

h)

=

2.42

*

10

6

s

From Faraday’s equation, using

n

=

2 and

M

=

55.847 g

>

mol,

=

0.1

A

I

=

wn

F

tM

=

(13)

2 3 - 4 The Corrosion Current and Polarization

861

(b) The total surface area of iron in contact with the corrosive liquid and the

current density are

i

=

I

A

=

0.1

900

=

1.11

*

10

-4

A

>

cm

2

A

=

(4

sides)(10

*

20)

+

(1

bottom)(10

*

10)

=

900

cm

2

Example 23-4

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

>

cm

2

. The area of both the copper and zinc

elec-trodes is 100 cm

2

. Calculate (a) the corrosion current, (b) the current density at the

zinc anode, and (c) the zinc loss per hour.

SOLUTION

(a) The corrosion current is

I

=

i

Cu

A

Cu

=

(0.05 A

>

cm

2

)(100 cm

2

)

=

5 A

(b) The current in the cell is the same everywhere. Thus,

(c) The atomic mass of zinc is 65.38 g

>

mol. From Faraday’s equation:

=

6.1

g

>

h

w

zinc loss

=

ItM

n

F

=

a

5

A

cm

2

b

(3600

s

>

h)(65.38

g

>

mol)

(2)(96,500

C)

i

Zn

=

I

A

Zn

=

5

100

=

0.05

A

>

cm

2

23-4

The Corrosion Current and Polarization

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Activation Polarization

This kind of polarization is related to the

energy required to cause the anode or cathode reactions to occur. If we can increase the

degree of polarization, these reactions occur with greater difficulty, and the rate of

cor-rosion is reduced. Small differences in composition and structure in the anode and

cath-ode materials dramatically change the activation polarization. Segregation effects in the

electrodes cause the activation polarization to vary from one location to another. These

factors make it difficult to predict the corrosion current.

Concentration Polarization

After corrosion begins, the

concentra-tion of ions at the anode or cathode surface may change. For example, a higher

concen-tration of metal ions may exist at the anode if the ions are unable to diffuse rapidly into

the electrolyte. Hydrogen ions may be depleted at the cathode in a hydrogen electrode, or

a high (OH)

-

concentration may develop at the cathode in an oxygen electrode. When this

situation occurs, either the anode or cathode reaction is stifled because fewer electrons

are released at the anode or accepted at the cathode.

In any of these examples, the current density, and thus the rate of corrosion,

decreases because of concentration polarization. Normally, the polarization is less

pro-nounced when the electrolyte is highly concentrated, the temperature is increased, or the

electrolyte is vigorously agitated. Each of these factors increases the current density and

encourages electrochemical corrosion.

Resistance Polarization

This type of polarization is caused by

the electrical resistivity of the electrolyte. If a greater resistance to the flow of the

cur-rent is offered, the rate of corrosion is reduced. Again, the degree of resistance

polar-ization may change as the composition of the electrolyte changes during the corrosion

process.

23-5

Types of Electrochemical Corrosion

In this section, we will look at some of the more common forms taken by

electrochemi-cal corrosion. First, there is

uniform attack.

When a metal is placed in an electrolyte, some

regions are anodic to other regions; however, the location of these regions moves and even

reverses from time to time. Since the anode and cathode regions continually shift, the

metal corrodes uniformly.

Galvanic attack

occurs when certain areas always act as anodes, whereas other

areas always act as cathodes. These electrochemical cells are called galvanic cells and can

be separated into three types:

composition cells, stress cells,

and

concentration cells

.

(15)

2 3 - 5

Types of Electrochemical Corrosion

863

TABLE 23-2The galvanic series in seawater

Active, Anodic End Magnesium and Mg alloys Zinc

Galvanized steel 5052 aluminum 3003 aluminum 1100 aluminum Alclad

Cadmium 2024 aluminum Low-carbon steel Cast iron

50% Pb-50% Sn solder 316 stainless steel (active) Lead

Tin

Cu-40% Zn brass

Nickel-based alloys (active) Copper

Cu-30% Ni alloy

Nickel-based alloys (passive) Stainless steels (passive) Silver

Titanium Graphite Gold Platinum

Noble, Cathodic End

(After ASM Metals Handbook, Vol. 10, 8th Ed., Copyright 1975 ASM International.)

Example 23-5

Corrosion of a Soldered Brass Fitting

A brass fitting used in a marine application is joined by soldering with lead-tin

sol-der. 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.

Composition cells also develop in two-phase alloys, where one phase is more

anodic than the other. Since ferrite is anodic to cementite in steel, small microcells cause

steel to galvanically corrode (Figure 23-6). Almost always, a two-phase alloy has less

resist-ance to corrosion than a single-phase alloy of a similar composition.

(16)

Intergranular corrosion

occurs when the precipitation of a second phase or

seg-regation at grain boundaries produces a galvanic cell. In zinc alloys, for example,

impuri-ties such as cadmium, tin, and lead segregate at the grain boundaries during solidification.

The grain boundaries are anodic compared to the remainder of the grains, and corrosion

of the grain boundary metal occurs (Figure 23-7). In austenitic stainless steels, chromium

carbides can precipitate at grain boundaries [Figure 23-6(b)]. The formation of the

car-bides removes chromium from the austenite adjacent to the boundaries. The

low-chromium (

"

12% Cr) austenite at the grain boundaries is anodic to the remainder of the

grain and corrosion occurs at the grain boundaries. In certain cold-worked aluminum

alloys, grain boundaries corrode rapidly due to the presence of detrimental precipitates.

This causes the grains of aluminum to peel back like the pages of a book or leaves. This

is known as exfoliation.

Stress Cells

Stress cells develop when a metal contains regions with

differ-ent local stresses. The most highly stressed or high-energy regions act as anodes to the

less-stressed cathodic areas (Figure 23-8). Regions with a finer grain size, or a higher

Figure 23-7

Micrograph of intergranular corrosion in a zinc die casting. Segregation of impurities to the grain boundaries produces

microgalvanic corrosion cells (* 50). (Reprinted courtesy of Don Askeland.)

(17)

2 3 - 5

Types of Electrochemical Corrosion

865

density of grain boundaries, are anodic to coarse-grained regions of the same material.

Highly cold-worked areas are anodic to less cold-worked areas.

Stress corrosion

occurs by galvanic action, but other mechanisms, such as the

adsorption of impurities at the tip of an existing crack, may also occur. Failure occurs as

a result of corrosion and an applied stress. Higher applied stresses reduce the time required

for failure.

Fatigue failures are also initiated or accelerated when corrosion occurs.

Corrosion

fatigue

can reduce fatigue properties by initiating cracks (perhaps by producing pits or

crevices) and by increasing the rate at which the cracks propagate.

Figure 23-8 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 high energy, they are anodic and corrode.

Example 23-6

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.

(18)

contact with the most concentrated solution is the cathode; the metal in contact with the

dilute solution is the anode.

The oxygen concentration cell (often referred to as

oxygen starvation

) occurs when

the cathode reaction is the oxygen electrode, H

2

O

+

O

2

+

2

e

-

:

2(OH)

-

. Electrons flow

from the low oxygen region, which serves as the anode, to the high oxygen region, which

serves as the cathode.

Deposits, such as rust or water droplets, shield the underlying metal from oxygen.

Consequently, the metal under the deposit is the anode and corrodes. This causes one

form of pitting corrosion. Waterline corrosion is similar. Metal above the waterline is

exposed to oxygen, while metal beneath the waterline is deprived of oxygen; hence, the

metal underwater corrodes. Normally, the metal far below the surface corrodes more

slowly than metal just below the waterline due to differences in the distance that electrons

must travel. Because cracks and crevices have a lower oxygen concentration than the

sur-rounding base metal, the tip of a crack or crevice is the anode, causing

crevice corrosion.

Pipe buried in soil may corrode because of differences in the composition of the

soil. Velocity differences may cause concentration differences. Stagnant water contains

low oxygen concentrations, whereas fast-moving, aerated water contains higher oxygen

concentrations. Metal near stagnant water is anodic and corrodes. The following is an

example of corrosion due to water.

1 2

Figure 23-9 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 23-7

Corrosion of Crimped Steel

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.

(19)

2 3 - 5

Types of Electrochemical Corrosion

867

Microbial Corrosion

Various microbes, such as fungi and bacteria,

cre-ate conditions that encourage electrochemical corrosion. Particularly in aqueous

environ-ments, these organisms grow on metallic surfaces. The organisms typically form colonies

that are not continuous. The presence of the colonies and the byproducts of the growth

of the organisms produce changes in the environment and, hence, the rate at which

cor-rosion occurs.

Some bacteria reduce sulfates in the environment, producing sulfuric acid, which

in turn attacks metal. The bacteria may be either aerobic (which thrive when oxygen is

available) or anaerobic (which do not need oxygen to grow). Such bacteria cause attacks

on a variety of metals, including steels, stainless steels, aluminum, and copper, as well as

some ceramics and concrete. A common example occurs in aluminum fuel tanks for

air-craft. When the fuel, typically kerosene, is contaminated with moisture, bacteria grow and

excrete acids. The acids attack the aluminum, eventually causing the fuel tank to leak.

The growth of colonies of organisms on a metal surface leads to the

develop-ment of oxygen concentration cells (Figure 23-10). Areas beneath the colonies are anodic,

whereas unaffected areas are cathodic. In addition, the colonies of organisms reduce the

rate of diffusion of oxygen to the metal, and—even if the oxygen does diffuse into the

colony—the organisms tend to consume the oxygen. The concentration cell produces

pit-ting beneath the regions covered with the organisms. Growth of the organisms, which may

include products of the corrosion of the metal, produces accumulations (called

tubercules

)

that may plug pipes or clog water-cooling systems in nuclear reactors, submarines, or

chemical reactors.

Figure 23-10

(20)

23-6

Protection Against Electrochemical

Corrosion

A number of techniques are used to combat corrosion, including proper design and

mate-rials selection and the use of coatings, inhibitors, cathodic protection, and passivation.

Design

Proper design of metal structures can slow or even avoid corrosion. Some

of the steps that should be taken to combat corrosion are as follows:

1. Prevent the formation of galvanic cells. This can be achieved by using similar

met-als or alloys. For example, steel pipe is frequently connected to brass plumbing

fixtures, producing a galvanic cell that causes the steel to corrode. By using

inter-mediate plastic fittings to electrically insulate the steel and brass, this problem can

be minimized.

2. Make the anode area much larger than the cathode area. For example, copper

ets can be used to fasten steel sheet. Because of the small area of the copper

riv-ets, a limited cathode reaction occurs. The copper accepts few electrons, and the

steel anode reaction proceeds slowly. If, on the other hand, steel rivets are used for

joining copper sheet, the small steel anode area gives up many electrons, which are

accepted by the large copper cathode area; corrosion of the steel rivets is then very

rapid. This is illustrated in the following example.

Example 23-8

Effect of Areas on Corrosion Rate for Copper-Zinc Couple

Consider a copper-zinc corrosion couple. If the current density at the copper

cath-ode is 0.05 A

>

cm

2

, calculate the weight loss of zinc per hour if (1) the copper

cath-ode area is 100 cm

2

and the zinc anode area is 1 cm

2

and (2) the copper cathode area

is 1 cm

2

and the zinc anode area is 100 cm

2

.

SOLUTION

For the small zinc anode area,

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.

wZn

=

ItM

n

F

=

(0.05

A

>

cm

2

)(3600

s)

Q

65.38

g

mol

R

(2)(96,500

C)

=

0.061g

>

h

I

=

i

Cu

A

Cu

=

(0.05

A

>

cm

2

)(1

cm

2

)

=

0.05

A

w

Zn

=

ItM

n

F

=

(5)(3600)(65.38)

(21)

2 3 - 6 Protection Against Electrochemical Corrosion

869

Figure 23-11 Alternative methods for joining two pieces of steel: (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 may avoid the formation of galvanic cells.

3. Design components so that fluid systems are closed, rather than open, and so that

stagnant pools of liquid do not collect. Partly filled tanks undergo waterline

corrosion. Open systems continuously dissolve gas, providing ions that participate

in the cathode reaction, and encourage concentration cells.

4. Avoid crevices between assembled or joined materials (Figure 23-11). Welding may

be a better joining technique than brazing, soldering, or mechanical fastening.

Galvanic cells develop in brazing or soldering since the filler metals have a

differ-ent composition from the metal being joined. Mechanical fasteners produce

crevices that lead to concentration cells. If the filler metal is closely matched to

the base metal, welding may prevent these cells from developing.

5. In some cases, the rate of corrosion cannot be reduced to a level that will not

inter-fere with the expected lifetime of the component. In such cases, the assembly

should be designed in such a manner that the corroded part can be easily and

eco-nomically replaced.

Coatings

Coatings are used to isolate the anode and cathode regions. Coatings

also prevent diffusion of oxygen or water vapor that initiates corrosion or oxidation.

Temporary coatings, such as grease or oil, provide some protection but are easily

dis-rupted. Organic coatings, such as paint, or ceramic coatings, such as enamel or glass,

pro-vide better protection; however, if the coating is disrupted, a small anodic site is exposed

that undergoes rapid, localized corrosion.

Metallic coatings include tin-plated and hot-dip galvanized (zinc-plated) steel

(Figure 23-12). This was discussed in earlier chapters on ferrous materials. A continuous

coating of either metal isolates the steel from the electrolyte; however, when the coating

is scratched, exposing the underlying steel, the zinc continues to be effective, because zinc

is anodic to steel. Since the area of the exposed steel cathode is small, the zinc coating

corrodes at a very slow rate and the steel remains protected. In contrast, steel is anodic to

tin, so a small steel anode is created when the tin is scratched, and rapid corrosion of the

steel subsequently occurs.

(22)

oxide layers form on the surface of aluminum, chromium, and stainless steel. These

oxides exclude the electrolyte and prevent the formation of galvanic cells. Components

such as reaction vessels can also be lined with corrosion-resistant Teflon

TM

or other

plastics.

Inhibitors

When added to the electrolyte, some chemicals migrate preferentially

to the anode or cathode surface and produce concentration or resistance polarization—

that is, they are

inhibitors

. Chromate salts perform this function in automobile radiators.

A variety of chromates, phosphates, molybdates, and nitrites produce protective films on

anodes or cathodes in power plants and heat exchangers, thus stifling the electrochemical

cell. Although the exact contents and the mechanism by which it functions are not

clear, the popular lubricant WD-40

TM

works from the action of inhibitors. The name

“WD-40

TM

” was designated water displacement (WD) that worked on the fortieth

exper-imental attempt!

Cathodic Protection

We can protect against corrosion by supplying the

metal with electrons and forcing the metal to be a cathode (Figure 23-13). Cathodic

pro-tection can use a sacrificial anode or an impressed voltage.

A

sacrificial anode

is attached to the material to be protected, forming an

electro-chemical circuit. The sacrificial anode corrodes, supplies electrons to the metal, and thereby

prevents an anode reaction at the metal. The sacrificial anode, typically zinc or

magne-sium, is consumed and must eventually be replaced. Applications include preventing the

Figure 23-12 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.

(23)

2 3 - 6 Protection Against Electrochemical Corrosion

871

corrosion of buried pipelines, ships, off-shore drilling platforms, and water heaters. A

mag-nesium rod is used in many water heaters. The Mg serves as an anode and undergoes

dis-solution, thus protecting the steel from corroding.

An

impressed voltage

is obtained from a direct current source connected between

an auxiliary anode and the metal to be protected. Essentially, we have connected a battery

so that electrons flow to the metal, causing the metal to be the cathode. The auxiliary

anode, such as scrap iron, corrodes.

Passivation or Anodic Protection

Metals near the anodic end

of the galvanic series are active and serve as anodes in most electrolytic cells; however, if

these metals are made passive or more cathodic, they corrode at slower rates than normal.

Passivation

is accomplished by producing strong anodic polarization, preventing the

nor-mal anode reaction; thus the term anodic protection.

We cause passivation by exposing the metal to highly concentrated oxidizing

solu-tions. If iron is dipped in very concentrated nitric acid, the iron rapidly and uniformly

corrodes to form a thin, protective iron hydroxide coating. The coating protects the iron

from subsequent corrosion in nitric acid.

We can also cause passivation by increasing the potential on the anode above

a critical level. A passive film forms on the metal surface, causing strong anodic

polar-ization, and the current decreases to a very low level. Passivation of aluminum is called

anodizing

, and a thick oxide coating is produced. This oxide layer can be dyed to

pro-duce attractive colors. The Ta

2

O

5

oxide layer formed on tantalum wires is used to make

capacitors.

Materials Selection and Treatment

Corrosion can be

pre-vented or minimized by selecting appropriate materials and heat treatments. In castings,

for example, segregation causes tiny, localized galvanic cells that accelerate corrosion. We

can improve corrosion resistance with a homogenization heat treatment. When metals are

formed into finished shapes by bending, differences in the amount of cold work and

resid-ual stresses cause local stress cells. These may be minimized by a stress-relief anneal or a

full recrystallization anneal.

The heat treatment is particularly important in austenitic stainless steels

(Figure 23-14). When the steel cools slowly from 870 to 425°C, chromium carbides

pre-cipitate at the grain boundaries. Consequently the austenite at the grain boundaries

may contain less than 12% chromium, which is the minimum required to produce a

passive oxide layer. The steel is

sensitized

. Because the grain boundary regions are small

and highly anodic, rapid corrosion of the austenite at the grain boundaries occurs. We

can minimize the problem by several techniques.

1. If the steel contains less than 0.03% C, chromium carbides do not form.

2. If the percent chromium is very high, the austenite may not be depleted to below

12% Cr, even if chromium carbides form.

3. Addition of titanium or niobium ties up the carbon as TiC or NbC, preventing the

formation of chromium carbide. The steel is said to be

stabilized

.

4. The sensitization temperature range—425 to 870°C—should be avoided during

manufacture and service.

(24)

The following examples illustrate the design of a corrosion protection system.

Figure 23-14 (a) Austenitic stainless steels become sensitized when cooled slowly in the temperature range between 870 and 425°C. (b) Slow cooling permits chromium carbides to precipitate at grain boundaries. The local chromium content is depleted. (c) A quench anneal to dissolve the carbides may prevent intergranular corrosion.

Example 23-9

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

pre-vent or delay this problem.

SOLUTION

The troughs are likely a low-carbon, unalloyed steel containing ferrite and

cemen-tite, producing a composition cell. The waterline in the trough, which is partially

filled with water, provides a concentration cell. The trough is also exposed to the

environment, and the water is contaminated with impurities. Consequently,

corro-sion of the unprotected steel tank is to be expected.

(25)

2 3 - 6 Protection Against Electrochemical Corrosion

873

provide better corrosion resistance than the plain carbon steel, but both are

consid-erably 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. This would require that the farm operator regularly check the

tank to be sure that the anode is not completely consumed. We also want to be sure

that magnesium ions introduced into the water are not a health hazard.

Another approach would be to protect the steel trough using a suitable coating.

Painting the steel (that is, introducing a protective polymer coating) or 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

gal-vanized steel is relatively inexpensive, readily available, and does not require frequent

inspection.

Example 23-10

Design of a Stainless Steel Weldment

A piping system used to transport a corrosive liquid is fabricated from 304 stainless

steel. Welding of the pipes is required to assemble the system. Unfortunately,

cor-rosion occurs, and the corrosive liquid leaks from the pipes near the weld. Identify

the problem and design a system to prevent corrosion in the future.

SOLUTION

Table 13-4 shows that 304 stainless steel contains 0.08% C, causing the steel to be

sensitized if it is improperly heated or cooled during welding. Figure 23-15 shows

the maximum temperatures reached in the fusion and heat-affected zones during

welding. A portion of the pipe in the HAZ heats into the sensitization temperature

Figure 23-15

The peak temperature surrounding a stainless steel weld and the

sensitized structure produced when the weld slowly cools (for

(26)

range, permitting chromium carbides to precipitate. If the cooling rate of the weld

is very slow, the fusion zone and other areas of the heat-affected zone may also be

affected. Sensitization of the weld area, therefore, is the likely reason for corrosion

of the pipe.

Several solutions to the problem may be considered. We might use a welding

process that provides very rapid rates of heat input, causing the weld to heat and cool

very quickly. If the steel is exposed to the sensitization temperature range for only a

brief time, chromium carbides may not precipitate. Joining processes such as laser

welding or electron-beam welding are high-rate-of-heat-input processes, but they

are expensive. In addition, electron beam welding requires the use of a vacuum, and

it may not be feasible to assemble the piping system in a vacuum chamber.

We might heat treat the assembly after the weld is made. By performing a

quench anneal, any precipitated carbides are re-dissolved during the anneal and do

not reform during quenching; however, it may be impossible to perform this

treat-ment on a large assembly.

We might check the original welding procedure to determine if the pipe was

preheated before joining in order to minimize the development of stresses due to

the welding process. If the pipe were preheated, sensitization would be more likely

to occur. We would recommend that any preheat procedure be suspended.

Perhaps our best design is to use a stainless steel that is not subject to

sensitiza-tion. For example, carbides do not precipitate in a 304L stainless steel, which

con-tains less than 0.03% C. The low-carbon stainless steels are more expensive than the

normal 304 steel; however, the extra cost does prevent corrosion and still permits us

to use conventional joining techniques.

23-7

Microbial Degradation and Biodegradable

Polymers

Attack by a variety of insects and microbes is one form of “corrosion” of polymers. Relatively

simple polymers (such as polyethylene, polypropylene, and polystyrene),

high-molecular-weight polymers, crystalline polymers, and thermosets are relatively immune to attack.

Some polymers—including polyesters, polyurethanes, cellulosics, and plasticized

polyvinyl chloride (which contains additives that reduce the degree of polymerization)—

are particularly vulnerable to microbial degradation. These polymers can be broken into

low-molecular-weight molecules by radiation or chemical attack until they are small

enough to be ingested by the microbes.

(27)

2 3 - 8

Oxidation and Other Gas Reactions

875

23-8

Oxidation and Other Gas Reactions

Materials of all types may react with oxygen and other gases. These reactions can, like

cor-rosion, alter the composition, properties, or integrity of a material. As mentioned before,

metals such as Al and Ti react with oxygen very readily.

Oxidation of Metals

Metals may react with oxygen to produce an oxide at

the surface. We are interested in three aspects of this reaction: the ease with which the metal

oxidizes, the nature of the oxide film that forms, and the rate at which

oxidation

occurs.

The ease with which oxidation occurs is given by the standard free energy of

for-mation for the oxide. The standard free energy of forfor-mation for an oxide can be

deter-mined from an Ellingham diagram, such as the one shown in Figure 23-16. Figure 23-16

shows that there is a large driving force for the oxidation of magnesium and aluminum

compared to the oxidation of nickel or copper. This is illustrated in the following example.

Standard free ener

gy of formation (G

°

)

(kcal/mol O

2

)

Figure 23-16 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 23-11

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 1700°C.

SOLUTION

(28)

were already present before the steel making began, chromium would oxidize before

the carbon (Figure 23-16), since chromium oxide has a lower free energy of

forma-tion (or is more stable) than carbon dioxide (CO

2

). Thus, any expensive chromium

added would be lost before the carbon was removed from the pig iron.

The type of oxide film influences the rate at which oxidation occurs (Figure 23-17).

For the

oxidation reaction

n

M

+

m

O

2

:

M

n

O

2m

(23-12)

the

Pilling-Bedworth

(P-B)

ratio

is

(23-13)

where

M

is the atomic or molecular mass,

!

is the density, and

n

is the number of metal

atoms in the oxide, as defined in Equation 23-12.

If the Pilling-Bedworth ratio is less than one, the oxide occupies a smaller volume

than the metal from which it formed; the coating is therefore porous and oxidation

con-tinues rapidly—typical of metals such as magnesium. If the ratio is one to two, the

vol-umes of the oxide and metal are similar and an adherent, non-porous, protective film

P-B ratio

=

oxide volume per metal atom

metal volume per metal atom

=

(

M

oxide

)(

r

metal

)

n

(

Mmetal

)(

r

oxide

)

Figure 23-17

(29)

2 3 - 8

Oxidation and Other Gas Reactions

877

forms—typical of aluminum and titanium. If the ratio exceeds two, the oxide occupies a

large volume and may flake from the surface, exposing fresh metal that continues to

oxi-dize—typical of iron. Although the Pilling-Bedworth equation historically has been used

to characterize oxide behavior, many exceptions to this behavior are observed. The use of

the P-B ratio equation is illustrated in the following example.

Example 23-12

Pilling-Bedworth Ratio

The density of aluminum is 2.7 g

>

cm

3

and that of Al

2

O

3

is about 4 g

>

cm

3

. Describe the

characteristics of the aluminum-oxide film. Compare these with the oxide film that forms

on tungsten. The density of tungsten is 19.254 g

>

cm

3

and that of WO

3

is 7.3 g

>

cm

3

.

SOLUTION

For 2Al

+

3

>

2O

2

:

Al

2

O

3

, the molecular weight of Al

2

O

3

is 101.96 g

>

mol and that

of aluminum is 26.981 g

>

mol.

For tungsten, W

+

3

>

2O

2

:

WO

3

. The molecular weight of WO

3

is 231.85 g

>

mol and

that of tungsten is 183.85 g

>

mol.

Since P-B

M

1 for aluminum, the Al

2

O

3

film is nonporous and adherent, providing

protection to the underlying aluminum. Since P-B

#

2 for tungsten, the WO

3

should

be nonadherent and nonprotective.

P-B

=

M

WO3

r

W

nM

W

r

WO3

=

a

231.85

g

mol

b

(19.254

g

>

cm

3

)

(1)

a

183.85

g

mol

b

(7.3

g

>

cm

3

)

=

3.33

P-B

=

MAl

2O3

r

Al

nMAl

r

A12O3

=

a

101.96

g

mol

b1

2.7

g

>

cm

3

2

(2)

a

26.981

g

mol

b

(4

g

>

cm

3

)

=

1.28

The rate at which oxidation occurs depends on the access of oxygen to the metal

atoms. A linear rate of oxidation occurs when the oxide is porous (as in magnesium) and

oxygen has continued access to the metal surface:

y

=

k

t

(23-14)

where

y

is the thickness of the oxide,

t

is the time, and k is a constant that depends on the

metal and temperature.

A parabolic relationship is observed when diffusion of ions or electrons through

a nonporous oxide layer is the controlling factor. This relationship is observed in iron,

copper, and nickel:

(30)

Finally, a logarithmic relationship is observed for the growth of thin-oxide films that are

particularly protective, as for aluminum and possibly chromium:

y

=

k ln(c

t

+

1)

(23-16)

where k and c are constants for a particular temperature, environment, and composition.

The example that follows shows the calculation for the time required for a nickel sheet to

oxidize completely.

Example 23-13

Parabolic Oxidation Curve for Nickel

At 1000°C, pure nickel follows a parabolic oxidation curve given by the constant

k

=

3.9

*

10

-12

cm

2

>

s in an oxygen atmosphere. If this relationship is not affected by

the thickness of the oxide film, calculate the time required for a 0.1-cm nickel sheet

to oxidize completely.

SOLUTION

Assuming that the sheet oxidizes from both sides:

Temperature also affects the rate of oxidation. In many metals, the rate of

oxida-tion is controlled by the rate of diffusion of oxygen or metal ions through the oxide. If

oxy-gen diffusion is more rapid, oxidation occurs between the oxide and the metal; if the metal

ion diffusion is more rapid, oxidation occurs at the oxide-atmosphere interface.

Consequently, we expect oxidation rates to follow an Arrhenius relationship, increasing

exponentially as the temperature increases.

t

=

(0.05

cm)

2

3.9

*

10

-12

cm

2

>

s

=

6.4

*

10

8

s

=

20.3

years

y

=

1

k

t

=

C

a

3.9

*

10

-12

cm

2

s

b

(

t

)

=

0.1

cm

2

sides

=

0.05

cm

(31)

2 3 - 9 Wear and Erosion

879

Polymers also degrade by the loss of side groups on the chain. Chloride ions [in

polyvinyl chloride (PVC)] and benzene rings (in polystyrene) are lost from the chain,

form-ing byproducts. For example, as polyvinyl chloride is degraded, hydrochloric acid (HCl)

is produced. Hydrogen atoms are bonded more strongly to the chains; thus, polyethylene

does not degrade as easily as PVC or polystyrene. Fluoride ions (in Teflon

TM

) are more

difficult to remove than hydrogen atoms, providing Teflon

TM

with its high temperature

resistance. As mentioned before, degradation of polymers, such as PCL, PLGA, and PLA,

actually can be useful for certain biomedical applications (e.g., sutures that dissolve) and

also in the development of environmentally friendly products.

23-9

Wear and Erosion

Wear and erosion remove material from a component by mechanical attack of solids or

liquids. Corrosion and mechanical failure also contribute to this type of attack.

Adhesive Wear

Adhesive wear

—also known as scoring, galling, or seizing—

occurs when two solid surfaces slide over one another under pressure. Surface projections,

or asperities, are plastically deformed and eventually welded together by the high local

pressures (Figure 23-18). As sliding continues, these bonds are broken, producing cavities

on one surface, projections on the second surface, and frequently tiny, abrasive particles—

all of which contribute to further wear of the surfaces.

Many factors may be considered in trying to improve the wear resistance of

mate-rials. Designing components so that loads are small, surfaces are smooth, and continual

lubrication is possible helps prevent adhesions that cause the loss of material.

The properties and microstructure of the material are also important. Normally,

if both surfaces have high hardnesses, the wear rate is low. High strength, to help resist the

applied loads, and good toughness and ductility, which prevent the tearing of material

from the surface, may be beneficial. Ceramic materials, with their exceptional hardness, are

expected to provide good adhesive wear resistance.

Figure 23-18

Figur

Figure 23-1

Figure 23-1

p.4
Figure 23-2

Figure 23-2

p.5
Figure 23-3 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.

Figure 23-3

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. p.6
Figure 23-4 The anode and cathode reactions in typical electrolytic corrosion cells: (a) the hydrogen electrode,(b) the oxygen electrode, and (c) the water electrode.

Figure 23-4

The anode and cathode reactions in typical electrolytic corrosion cells: (a) the hydrogen electrode,(b) the oxygen electrode, and (c) the water electrode. p.8
Figure 23-5

Figure 23-5

p.9
TABLE 23-1 I The standard reduction potentials for selected elements and reactions

TABLE 23-1

I The standard reduction potentials for selected elements and reactions p.10
TABLE 23-2 I The galvanic series in seawater

TABLE 23-2

I The galvanic series in seawater p.15
Figure 23-6Example of microgalvanic cells in two-phase alloys: (a) In steel, ferrite is anodicto cementite

Figure 23-6Example

of microgalvanic cells in two-phase alloys: (a) In steel, ferrite is anodicto cementite p.16
Figure 23-7

Figure 23-7

p.16
Figure 23-8Examples of stress cells. (a) Cold work required to bend a steel bar introduceshigh residual stresses at the bend, which then is anodic and corrodes

Figure 23-8Examples

of stress cells. (a) Cold work required to bend a steel bar introduceshigh residual stresses at the bend, which then is anodic and corrodes p.17
Figure 23-9Concentration cells: (a) Corrosion occurs beneath a water droplet on a steel platedue to low oxygen concentration in the water

Figure 23-9Concentration

cells: (a) Corrosion occurs beneath a water droplet on a steel platedue to low oxygen concentration in the water p.18
Figure 23-10

Figure 23-10

p.19
Figure 23-11Alternative methods for joining two pieces of steel: (a) Fasteners may producea concentration cell, (b) brazing or soldering may produce a composition cell, and (c) weldingwith a filler metal that matches the base metal may avoid the formation of galvanic cells.

Figure 23-11Alternative

methods for joining two pieces of steel: (a) Fasteners may producea concentration cell, (b) brazing or soldering may produce a composition cell, and (c) weldingwith a filler metal that matches the base metal may avoid the formation of galvanic cells. p.21
Figure 23-12Zinc-plated steel and tin-plated steel are protected differently. Zinc protectssteel even when the coating is scratched since zinc is anodic to steel

Figure 23-12Zinc-plated

steel and tin-plated steel are protected differently. Zinc protectssteel even when the coating is scratched since zinc is anodic to steel p.22
Figure 23-13Cathodic protection of a buried steel pipeline: (a) A sacrificial magnesiumanode ensures that the galvanic cell makes the pipeline the cathode

Figure 23-13Cathodic

protection of a buried steel pipeline: (a) A sacrificial magnesiumanode ensures that the galvanic cell makes the pipeline the cathode p.22
Figure 23-14(a) Austenitic stainless steels become sensitized when cooled slowly in the temperature rangebetween 870 and 425°C

Figure 23-14(a)

Austenitic stainless steels become sensitized when cooled slowly in the temperature rangebetween 870 and 425°C p.24
Figure 23-15The peak temperature surrounding astainless steel weld and thesensitized structure produced whenthe weld slowly cools (forExample 23-10).

Figure 23-15The

peak temperature surrounding astainless steel weld and thesensitized structure produced whenthe weld slowly cools (forExample 23-10). p.25
Figure 23-16The standard free energy of formation of selected oxides as a function oftemperature

Figure 23-16The

standard free energy of formation of selected oxides as a function oftemperature p.27
Figure 23-17Three types of oxides may form,depending on the volume ratiobetween the metal and the oxide:(a) magnesium produces a porousoxide film, (b) aluminum forms aprotective, adherent, nonporousoxide film, and (c) iron forms anoxide film that spalls off the surfaceand provides poor protection.

Figure 23-17Three

types of oxides may form,depending on the volume ratiobetween the metal and the oxide:(a) magnesium produces a porousoxide film, (b) aluminum forms aprotective, adherent, nonporousoxide film, and (c) iron forms anoxide film that spalls off the surfaceand provides poor protection. p.28
Figure 23-18

Figure 23-18

p.31
Figure 23-19 Abrasive wear, caused by either trapped or free-flying abrasives, produces troughsin the material, piling up asperities that may fracture into debris.

Figure 23-19

Abrasive wear, caused by either trapped or free-flying abrasives, produces troughsin the material, piling up asperities that may fracture into debris. p.32
Figure 23-20 Two steel sheets joined by an aluminumrivet (for Problem 23-30).

Figure 23-20

Two steel sheets joined by an aluminumrivet (for Problem 23-30). p.37
Figure 23-21 Cross-section through an integratedcircuit showing the external lead connection to thechip (for Problem 23-31).

Figure 23-21

Cross-section through an integratedcircuit showing the external lead connection to thechip (for Problem 23-31). p.38

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