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

Causes of 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 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 Corrosion of iron, or rusting is primarily caused by exposure of an iron surface 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 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 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

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 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.

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

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 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 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

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool 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.

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool Corrosion of iron, or rusting is enhanced by the presence of acids,

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool So we can limit rusting by keeping acids away from iron or steel objects. This would include protection of the objects from acid rain.

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool When the surface of an iron or steel object is exposed to an electrolyte like salt water, rusting will be enhanced.

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool 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

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool Finally, when some factors are causing rusting to occur, having higher temperatures can make it happen faster.

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 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 a less active metal, when an electrolyte is present. Keeping acids away from iron or steel objects Avoiding exposure to salt water Keeping objects cool So when it is possible to do so, we try to keep exposed iron and steel objects relatively cool

Cathodic Protection using a Sacrificial Anode
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.

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

Attaching a metal that is below Fe in the reduction table
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

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

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

Stronger Reducing Agents than Fe More Easily Oxidized than Fe
Which means they’re more easily oxidized than iron

Oxidation Potentials Increase
Ox. Pot. = V Oxidation Potentials Increase Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V 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 volts for iron, to Volts for magnesium.

These all have higher oxidation potentials than Fe
Ox. Pot. = V These all have higher oxidation potentials than Fe Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Notice that these metals below iron, (click) all have higher oxidation potentials than iron

These all have higher oxidation potentials than Fe
Ox. Pot. = V These all have higher oxidation potentials than Fe Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Which means (click)

These all oxidize more readily than Fe
Ox. Pot. = V These all oxidize more readily than Fe Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V They all oxidize more readily than iron

Ox. Pot. = 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 Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V 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.

O2 Water droplet iron Recall from the video on the causes of corrosion, that when an iron surface is exposed to water and oxygen,

O2 e– Fe Fe2+ e– iron An iron atom oxidizes to form an iron 2+ cation, which dissolves in the water, and the electrons produced travel through the metal.

iron O2 e– e– Fe(s)  Fe2+ + 2e– Fe2+
The equation for this oxidation is Fe(s)  Fe2+ + 2e–.

iron O2 e– e– Fe(s)  Fe2+ + 2e– Anode Region Fe2+
This is the anode region

iron O2 e– e– Fe(s)  Fe2+ + 2e– Anode Region Fe2+
Oxygen from the air then comes into contact with the edge of the water droplet and the iron metal surface.

iron O2 e– e– Fe(s)  Fe2+ + 2e– H2O Anode Region Fe2+ OH– OH–
½O2(g) + H2O + 2e–  2OH– iron And it combines with water and the electrons formed by the oxidizing iron, to produce hydroxide ions.

iron Fe(s)  Fe2+ + 2e– Cathode Region OH– Anode Region Fe2+ OH–
½O2(g) + H2O + 2e–  2OH– iron Because reduction of oxygen takes place here, this is called the cathode region.

iron Fe(s)  Fe2+ + 2e– e– Cathode Region Anode Region
½O2(g) + H2O + 2e–  2OH– iron 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– e– Cathode Region Anode Region
½O2(g) + H2O + 2e–  2OH– iron 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

These all oxidize more readily than Fe
Ox. Pot. = V These all oxidize more readily than Fe Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Remember, we had shown that metals below iron on the right side of the reduction table, (click) will oxidize more readily than iron does

These all oxidize more readily than Fe Oxidizes more readily than iron
Ox. Pot. = V These all oxidize more readily than Fe Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Ox. Pot. = V Oxidizes more readily than iron Ox. Pot. = V We’ll choose magnesium, with an oxidation potential of volts

iron O2 Water droplet magnesium
We’ll attach a piece of magnesium to the iron metal

Mg will oxidize more readily than Fe
Ox. Pot. = V Mg will oxidize more readily than Fe Ox. Pot. = V 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

iron O2 e– e– Water droplet H2O magnesium Mg2+ Mg
A magnesium atom will lose 2 electrons as it oxidizes to an Mg2+ ion. These electrons flow into the iron.

iron O2 e– e– H2O magnesium Mg2+
The magnesium ion will leave the metal and dissolve in the water

iron O2 e– Mg(s)  Mg2+ + 2e– e– Mg2+ H2O magnesium
This oxidation half-reaction taking place on the magnesium is (click) Mg(s)  Mg2+ + 2e–

iron Anode O2 e– Mg(s)  Mg2+ + 2e– e– Mg2+ H2O magnesium
Because oxidation of magnesium occurs instead of iron, (click) magnesium is now the anode instead of the iron.

iron Anode O2 e– Mg(s)  Mg2+ + 2e– e– Mg2+ H2O magnesium
Oxygen from the air will move to the spot where the iron metal, the water, and air all meet.

iron Anode O2 e– Mg(s)  Mg2+ + 2e– e– Mg2+ H2O magnesium OH– OH–
½O2(g) + H2O + 2e–  2OH– iron Oxygen in the presence of water will gain the electrons lost by the magnesium and undergo reduction, to form hydroxide ions.

iron Anode Mg(s)  Mg2+ + 2e– Mg2+ magnesium OH– OH–
½O2(g) + H2O + 2e–  2OH– iron The hydroxide ions formed will dissolve in the water

iron Anode Cathode Mg(s)  Mg2+ + 2e– Mg2+ OH– magnesium OH–
½O2(g) + H2O + 2e–  2OH– iron Because reduction occurs here on the iron, this is the location of the cathode

iron Anode Cathode No Fe2+ ions Mg(s)  Mg2+ + 2e– Mg2+ OH– magnesium
½O2(g) + H2O + 2e–  2OH– iron Because iron did not oxidize (click), there are no Fe2+ ions in solution, so no rust can form.

iron Anode Cathode Mg(s)  Mg2+ + 2e– e– magnesium
½O2(g) + H2O + 2e–  2OH– iron As magnesium is oxidized, the electrons it loses travel through the iron toward the cathode, where they are used up as oxygen is reduced.

½O2(g) + H2O + 2e–  2OH– iron 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.

Called a “Sacrificial” Anode
Cathode magnesium Mg(s)  Mg2+ + 2e– ½O2(g) + H2O + 2e–  2OH– iron Because magnesium is consumed in order to save the iron from rusting, it is sometimes called a sacrificial anode.

iron Anode Cathode Mg(s)  Mg2+ + 2e– magnesium
½O2(g) + H2O + 2e–  2OH– iron 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.

The process is called Cathodic Protection
Anode Cathode magnesium Mg(s)  Mg2+ + 2e– ½O2(g) + H2O + 2e–  2OH– iron That’s why the process is called Cathodic Protection.

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

An Electric Hot Water Tank Temperature and Pressure Relief Valve
Cold Water Supply Hot Water Outlet Temperature and Pressure Relief Valve Upper Heating Element Sacrificial Anode Rod Lower Heating Element Drain Value Discharge Pipe 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.

Galvanized Steel is steel coated with zinc
Galvanized Steel is steel coated with zinc. The object may be dipped in molten zinc or the zinc could be electroplated on the object. This work has been released into the public domain by its author, Splarka at the English Wikipedia project. 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.

Galvanized Steel is steel coated with zinc. The object may be dipped in molten zinc or the zinc could be electroplated on 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 This work has been released into the public domain by its author, Splarka at the English Wikipedia project. The zinc protects the steel in two ways: (click) first it coats the steel to prevent oxygen and water from reaching its surface.

Galvanized Steel is steel coated with zinc. The object may be dipped in molten zinc or the zinc could be electroplated on 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 This work has been released into the public domain by its author, Splarka at the English Wikipedia project. 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

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

Stainless steel is an alloy with anywhere from 11% to 26% chromium
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. License: CC0 Public Domain / FAQ Free for commercial use / No attribution required Thank you to Taken for sharing this image on Pixabay. 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

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

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

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

Rectifier + – Anode 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 + – Anode 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 + – e– The rectifier pulls electrons from the anode, and pushes them onto the pipe,

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– (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 O2 e– e– e– e– e– e– OH– ½O2(g) + H2O + 2e–  2OH–
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– OH– ½O2(g) + H2O + 2e–  2OH– OH–
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.

Two Types of Cathodic Protection: .
Using a Sacrificial Anode (attaching a metal that is more easily oxidized than Fe—on lower on the right side of the table) . Using a rectifier or DC power supply to create an impressed current (keeping an excess on electrons on the object being protected) So we can summarize by saying that there are two main types of cathodic protection.

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

½O2(g) + H2O + 2e–  2OH– iron For example, if magnesium is attached to iron, the magnesium gets oxidized rather than the iron, thus protecting the iron from oxidation.

Using a Sacrificial Anode (attaching a metal that is more easily oxidized than Fe—or lower on the right side of the table) . Using a rectifier or DC power supply to create an impressed current (keeping an excess on electrons on the object being protected) 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.

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

– – + + Rectifier O2 e– e– e– e– e– e– OH– ½O2(g) + H2O + 2e–  2OH–
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.

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

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

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Presentation on theme: "Corrosion Prevention Corrosion of steel, which consists mainly of iron, is a major problem in our society. But steps can be taken to prevent it."— Presentation transcript:

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