corrosion of steel, which consists mainly of iron, is a major problem in our society. but steps can...

Corrosion of steel, which consists mainly of iron, is a major problem in our society. But steps can be taken to prevent it.

Corrosion

Prevention

The causes of corrosion or iron and factors that increase it were outlined in the previous video, Corrosion-Its causes. Knowing these will help us find ways to prevent corrosion.

Corrosion of Iron can be caused by:.

Exposure to oxygen and water

Exposure to stronger oxidizing agents than Fe2+

Being attached to a metal higher than Fe on the table in the presence of an electrolyte

And is enhanced by:.

The presence of acids The presence of an

electrolyte like salt water

Higher temperatures

Causes of Corrosion

Corrosion of iron, or rusting is primarily caused by exposure of an iron surface to oxygen and water.








Higher temperatures

We can deal with this by painting the surface or coating it with an anti-rust coating and keeping it dry. Painted steel objects like bridges and vehicles must be re-painted periodically to maintain a continuous coating. When cracks in paint occur, moisture can collect and rusting can take place quite rapidly








Higher temperatures

Corrosion of Iron can be prevented by:.

Painting or coating the surface and keeping it dry

Eliminating exposure to oxidizing agents

Being careful not to attach iron to less active metals, when an electrolyte is present.

Keeping acids away from iron or steel objects

Avoiding exposure to salt water

Keeping objects cool

Another cause of corrosion of iron or steel is exposure to any stronger oxidizing agent than Fe2+. This includes (click) all the species above Fe2+ on the left side of the reduction table.








Higher temperatures








So we must do what we can to eliminate exposure of iron or steel objects to these oxidizing agents.








Higher temperatures








Another potential problem is galvanic corrosion. This is where iron or steel is attached to a metal higher than iron on the reduction table, or a less active metal, in the presence of an electrolyte








Higher temperatures








So we must be careful not to attach iron or steel to a less active metal, such as an alloy containing copper, when an electrolyte if present. For example, we must pay attention to what type (click) of bolts or fasteners we use when building steel structures.








Higher temperatures




Being careful not to attach iron to a less active metal, when an electrolyte is present.




Corrosion of iron, or rusting is enhanced by the presence of acids,








Higher temperatures








So we can limit rusting by keeping acids away from iron or steel objects. This would include protection of the objects from acid rain.








Higher temperatures








When the surface of an iron or steel object is exposed to an electrolyte like salt water, rusting will be enhanced.








Higher temperatures








It can be slowed down or prevented by avoiding exposure to salt water and other electrolytes. Road salt should be washed off of vehicles when possible. Also protective coatings can be used on vehicles








Higher temperatures








Finally, when some factors are causing rusting to occur, having higher temperatures can make it happen faster.








Higher temperatures








So when it is possible to do so, we try to keep exposed iron and steel objects relatively cool








Higher temperatures








Another method that is widely used to prevent corrosion of iron is called cathodic protection. This first type of cathodic protection we’ll look at uses what is called a sacrificial anode.

Cathodic Protection

using a Sacrificial

Anode

This type of cathodic protection is achieved by attaching a metal that is (click) below Fe in the reduction table, to the iron or steel we’re trying to protect.

Cathodic Protection:Attaching a metal that is below Fe in the reduction table

Some metals we could use for this are: (click) chromium, (click) zinc, (click) manganese, (click) aluminum, or (click) magnesium

Cathodic Protection:Attaching a metal that is below Fe in the reduction table

It would not be practical to use these metals below magnesium on the table. They react rapidly with water to form hydrogen gas.

It would not be practical to use these metals because they

react rapidly with water to form hydrogen gas

Metals below iron on the right side of the table are (click) stronger reducing agents than iron

Stronger Reducing Agents than Fe

Which means they’re more easily oxidized than iron

Stronger Reducing Agents than Fe

More Easily

Oxidized than Fe

Remember, oxidation potentials of these are just these values with their signs switched, (click), so the oxidation potentials of these metals are all positive. Notice that as we move down, (click) their oxidation potentials increase from +0.45 volts for iron, to +2.37 Volts for magnesium.

Oxidation

Potentials

Increase

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

Notice that these metals below iron, (click) all have higher oxidation potentials than iron

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

These all have higher

oxidation

potentials than

Fe

Which means (click)

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

These all have higher

oxidation

potentials than

Fe

They all oxidize more readily than iron

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

These all

oxidize more

readily than Fe

So if any of these are present with iron and an oxidizing agent like oxygen appears, these will oxidize instead of the iron, thus saving the iron from oxidation. Let’s look at an example.

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

If these are

present with Fe, and an

oxidizing agent

appears, they will oxidize

instead of the Fe,

thus saving the

Fe from being

oxidized

Recall from the video on the causes of corrosion, that when an iron surface is exposed to water and oxygen,

iron

Water droplet

O2

An iron atom oxidizes to form an iron 2+ cation, which dissolves in the water, and the electrons produced travel through the metal.

iron

FeFe2+e–

e–

O2

The equation for this oxidation is Fe(s) Fe2+ + 2e–.

iron

e–

e–

Fe2+

Fe(s) Fe2+ + 2e–

O2

This is the anode region

iron

e–

e–

Fe2+

Fe(s) Fe2+ + 2e–

Anode Region

O2

Oxygen from the air then comes into contact with the edge of the water droplet and the iron metal surface.

iron

e–

e–

Fe2+

Fe(s) Fe2+ + 2e–

Anode Region

O2

And it combines with water and the electrons formed by the oxidizing iron, to produce hydroxide ions.

iron

e–

e–

Fe2+

Fe(s) Fe2+ + 2e–

Anode Region

O2

H2O

½O2(g) + H2O + 2e– 2OH–

OH–OH–

Because reduction of oxygen takes place here, this is called the cathode region.

iron

Fe2+

Fe(s) Fe2+ + 2e–

Anode Region

½O2(g) + H2O + 2e– 2OH–

OH–

OH–

Cathode Region

So the anode region and the cathode region are both on the iron. The anode is the place where iron is being oxidized to Fe2+ and giving off electrons and the cathode is the area where oxygen is reduced, using up those electrons.

iron

Fe(s) Fe2+ + 2e–½O2(g) + H2O + 2e– 2OH–

Cathode Region

Anode Region

e–

We see that rust is formed at the cathode region and the iron is eaten up at the anode region. The anode and the cathode are in different places, but they are both on the iron

iron

Fe(s) Fe2+ + 2e–½O2(g) + H2O + 2e– 2OH–

Cathode Region

Anode Region

e–

Remember, we had shown that metals below iron on the right side of the reduction table, (click) will oxidize more readily than iron does

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

These all

oxidize more

readily than Fe

We’ll choose magnesium, with an oxidation potential of +2.37 volts

Ox. Pot. = +0.74 V

Ox. Pot. = +0.76 V

Ox. Pot. = +1.19 V

Ox. Pot. = +1.66 V

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

These all

oxidize more

readily than Fe

Oxidizes more readily than

iron

We’ll attach a piece of magnesium to the iron metal

iron

Water droplet

O2

magnesium

Because it’s lower on the right side of the table and therefore, has a higher oxidation potential, (click) magnesium will oxidize more readily than iron

Ox. Pot. = +2.37 V

Ox. Pot. = +0.45 V

Mg will oxidize more

readily than Fe

A magnesium atom will lose 2 electrons as it oxidizes to an Mg2+ ion. These electrons flow into the iron.

iron

Water droplet

O2

magnesium Mge–Mg2+e–H2O

The magnesium ion will leave the metal and dissolve in the water

iron

O2

magnesium

e–

e–

Mg2+

H2O

This oxidation half-reaction taking place on the magnesium is (click) Mg(s) Mg2+ + 2e–

iron

O2

magnesium

e–

e–

Mg2+

Mg(s) Mg2+ + 2e–

H2O

Because oxidation of magnesium occurs instead of iron, (click) magnesium is now the anode instead of the iron.

iron

O2

magnesium

e–

e–

Mg2+

Mg(s) Mg2+ + 2e–

Anode H2O

Oxygen from the air will move to the spot where the iron metal, the water, and air all meet.

iron

magnesium

e–

e–

Mg2+

Mg(s) Mg2+ + 2e–

Anode H2O

O2

Oxygen in the presence of water will gain the electrons lost by the magnesium and undergo reduction, to form hydroxide ions.

iron

magnesium

Mg2+

Mg(s) Mg2+ + 2e–

Anode

e–

e–

O2

H2O

½O2(g) + H2O + 2e– 2OH–

OH–

OH–

The hydroxide ions formed will dissolve in the water

iron

magnesium

Mg2+

Mg(s) Mg2+ + 2e–

Anode

½O2(g) + H2O + 2e– 2OH–

OH–OH–

Because reduction occurs here on the iron, this is the location of the cathode

iron

magnesium

Mg2+

Mg(s) Mg2+ + 2e–

Anode

½O2(g) + H2O + 2e– 2OH–

OH–

OH–

Cathode

Because iron did not oxidize (click), there are no Fe2+ ions in solution, so no rust can form.

iron

magnesium

Mg2+

Mg(s) Mg2+ + 2e–

Anode

½O2(g) + H2O + 2e– 2OH–

OH–

OH–

Cathode

No Fe2+ ions

As magnesium is oxidized, the electrons it loses travel through the iron toward the cathode, where they are used up as oxygen is reduced.

iron

magnesium

Mg(s) Mg2+ + 2e–

Anode

½O2(g) + H2O + 2e– 2OH–

Cathode

e–

Looking at the magnesium block, we see that as it is oxidized, it is gradually eaten up. When the magnesium block has been consumed, it is simply replaced by a new one.

iron

magnesium

Mg(s) Mg2+ + 2e–

Anode

½O2(g) + H2O + 2e– 2OH–

Cathode

e–

Because magnesium is consumed in order to save the iron from rusting, it is sometimes called a sacrificial anode.

iron

magnesium

Mg(s) Mg2+ + 2e– ½O2(g) + H2O + 2e–

2OH–

Cathode

Called a “Sacrificial”

Anode

When magnesium is present it is the anode, so the (click) iron is NOT the anode. Instead, the surface of the iron acts (click) as the CATHODE, where the oxygen is reduced.

iron

magnesium

Mg(s) Mg2+ + 2e– ½O2(g) + H2O + 2e–

2OH–

Cathode

Anode

That’s why the process is called Cathodic Protection.

iron

magnesium

Mg(s) Mg2+ + 2e– ½O2(g) + H2O + 2e–

2OH–

Anode

The process is called Cathodic

Protection

Cathode

Here is a photograph of the propeller of a large ship. (click). The white blocks are zinc anodes attached to the ship. Their purpose is to cathodically protect the metal the propeller is made of.

By Knotnic at en.wikipedia [Public domain],from Wikimedia Commons

The white blocks are zinc

anodes attached to

this ship. They are there to cathodically protect the propeller

Here is a simplified diagram of an electric hot water tank. Inside most hot water tanks, we find (click) a sacrificial anode rod. This can be made of either aluminum, magnesium, or zinc. The anode will corrode rather than exposed steel inside the tank. It is recommended that these be replaced periodically.

Upper Heating Element

Lower Heating Element

Temperatureand Pressure Relief Valve

An Electric Hot Water

Tank

Drain Value

Hot Water Outlet

Cold Water Supply

Discharge Pipe

Sacrificial Anode

Rod

Galvanized steel is steel coated with zinc. It is made either by dipping steel in molten zinc or electroplating steel with a thin coating of zinc.

This work has been released into the public domain byits author, Splarka at the English Wikipedia project.

Galvanized Steel is steel coated with zinc. The object

may be dipped in molten zinc or the

zinc could be electroplated on

the object.

The zinc protects the steel in two ways: (click) first it coats the steel to prevent oxygen and water from reaching its surface.





the object.

Zinc coats the steel to protect it from oxygen and water

Zinc is more easily oxidized than iron, so zinc cathodically protects the iron in steel

Secondly, zinc is lower than iron on the reduction table. So it is a stronger reducing agent, and thus, is more easily oxidized than iron. So zinc cathodically protects iron in the steel


Zinc coats the steel to protect it from oxygen and water

Zinc is more easily oxidized than iron, so zinc cathodically protects the iron in steel




the object.

We see that chromium is below iron on the right side of the reduction table. (click) its oxidation potential is +0.74 V, which is higher than that of iron, at +0.45 volts. (click) so chromium oxidized more readily than iron

Ox. Pot. = +0.74 V

Ox. Pot. = +0.45 V

Oxidizes more readily than

iron

Stainless steel is an alloy containing iron, carbon, and other metals, including chromium. It is at least 11% chromium and can be anywhere up to 26% chromium. When chromium oxidizes, it forms a thin layer of chromium(III) oxide, which oxygen and water can’t penetrate

License: CC0 Public Domain / FAQFree for commercial use / No attribution requiredThank you to Taken for sharing this image on Pixabay.

Stainless steel is an alloy with

anywhere from 11% to 26%

chromium. When chromium

oxidizes, it forms a thin layer of Cr2O3, which oxygen and

water cannot penetrate.

In addition to using sacrificial anodes, another method of cathodic protection is the use of what is called an impressed current.

Impressed Current Cathodic

Protection

Let’s say we have an underground steel pipeline we want to cathodically protect.

We bury an anode made of an inert metal near the pipe

An

od

e

Now we install a device called a rectifier. A rectifier is connected to an alternating current (or AC) supply line. What it does is convert AC to direct current (DC). So it has a (click) + and – terminal.

Rectifier+–

An

od

e

The positive terminal is connected by a wire to the anode and the negative terminal is connected by a wire to the steel pipe.

Rectifier+–

An

od

e

The rectifier pulls electrons from the anode, and pushes them onto the pipe,

Rectifier+–

e–

(click)

Rectifier+–

e–e–

(click)

Rectifier+–

e–e–

e–

(click).

Rectifier+–

e–e–

e–

e–

(click)

Rectifier+–

e–e–

e–

e–

e–

(click)

Rectifier+–

e–e–

e–

e–

e–

e–

Because the anode has lost electrons, it has (click) acquired a positive charge, and because the pipe has gained electrons, it has (click) acquired a negative charge.

Rectifier+–

e–

e–

e–

e–

e–

e–+

–

If oxygen makes its way through the ground, it will combine with water and excess electrons from the pipe to form hydroxide ions. As long as the pipe has extra electrons supplied by the rectifier, iron atoms in the steel pipe will not be oxidized, and in this way the pipe is cathodically protected.

Rectifier+–

e–

e–

e–

e–

e–

e–+

–

O2

H2OOH–

OH–

½O2(g) + H2O + 2e– 2OH–

The rectifier will keep supplying the pipe with electrons to replace the ones used for the reduction of oxygen. Because the pipe always has an excess of electrons, iron atoms will not need to lose any electrons, and thus will be spared from oxidation. So the pipe is cathodically protected.

Rectifier+–

e–

e–

e–

e–+

–

OH–

OH–

e–e–

½O2(g) + H2O + 2e– 2OH–

So we can summarize by saying that there are two main types of cathodic protection.

Two Types of Cathodic Protection:.

1. Using a Sacrificial Anode (attaching a metal that is more easily oxidized than Fe—on lower on the right side of the table).

2. Using a rectifier or DC power supply to create an impressed current (keeping an excess on electrons on the object being protected)

The first method is accomplished using a Sacrificial Anode (that is, attaching a metal that is more easily oxidized than Fe—or lower on the right side of the table)


1. Using a Sacrificial Anode (attaching a metal that is more easily oxidized than Fe—or lower on the right side of the table).


For example, if magnesium is attached to iron, the magnesium gets oxidized rather than the iron, thus protecting the iron from oxidation.

iron

magnesium

Mg(s) Mg2+ + 2e–

Anode

½O2(g) + H2O + 2e– 2OH–

Cathode

e–

The second method is accomplished using a rectifier or DC power supply to create an impressed current. Excess electrons pumped onto the object being protected will be used for the reduction of oxidizing agents, and will eliminate the need for any iron atoms to be oxidized.


1. Using a Sacrificial Anode (attaching a metal that is more easily oxidized than Fe—or lower on the right side of the table).


For example, a rectifier (click) takes electrons from an anode and pushes them onto a steel pipe making it negative.

Rectifier+–

e–e–

e–

e–

e–

e– +–

So when oxygen comes along to be reduced, it uses these excess electrons from the pipe, thus sparing iron atoms in the pipe from being oxidized and protecting it.

Rectifier+–

e–

e–

e–

e–

e–

e–+

–

O2

H2OOH–

OH–

½O2(g) + H2O + 2e– 2OH–

corrosion of steel, which consists mainly of iron, is a major problem in our society. but steps can...

Documents

exposure of iron

water exposure

iron surface

corrosion of steel

corrosion slide

presence of acids

causes of corrosion

stronger oxidizing agents