05 Corrosion Concept Overview
CORROSION | CONCEPT OVERVIEW The topic of CORROSION can be referenced on page 60 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
CONCEPT INTRO: When analyzing materials, it is important that we study how materials can be damaged, and the impacts various processes and reactions can have on the integrity of a material. Materials can be damaged by reacting chemically with their environments. Factors such as moisture and the amount of oxygen, can reactions such as OXIDATION or CORROSION to occur. Corrosion is the deterioration of a metal by chemical or electrochemical reactions with the environment. Corrosion is affected by six major factors: 1. Acidity – Acids form hydrogen ions in water solution. The hydrogen ions tend to be replaced by metal ions, freeing hydrogen gas. 2. Oxidizing Characteristics – Oxidation causes an increase in positive valence electrons. Therefore, any metal whose valence is zero would cause a change to a compound having a positive valence. This acts in the same direction as the acid reaction. The two, acid reaction and oxidation, frequently act together. 3. Electrolysis - In ELECTROLYSIS, the positive electrode called the ANODE will take up or absorb electrons. The metal in it, minus the electrons, will go into solution as a positive ion. At the negative electrode, called the CATHODE, hydrogen ions take on the electrons and become hydrogen gas. The result on the metal of this electrolysis is a pitting or roughening of the metal surface.
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4. Film Formation – A thick film forms a surface layer which inhibits any further corrosion. A film such as rust can deny free access to the corroding medium such as water. 5. Agitation or Rate of Movement – Agitation will break down layers of concentration between the metal and the corroding medium. This causes the reaction to proceed more rapidly than it would have if the layers of concentration were not disturbed. 6. Temperature – Temperature is one measure of heat energy. Adding energy to a system will increase its temperature. The highest the temperature, the faster the rate of corrosion. The EQUATION FOR THE ANODE REACTION (OXIDATION) OF A TYPICAL METAL can be referenced under the topic of CORROSION on page 60 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. OXIDATION is the removal of electrons from an atom and is the process by which metals lose mass and corrode over time. In this process, metal atoms lose one or more electrons and become metal ions as expressed in the chemical reaction. The Anode Reaction (Oxidation) of a Typical Metal “𝑀” is: 𝑀° → 𝑀$% + 𝑛𝑒 ) The reaction above is a OXIDATION REACTION occurring at the ANODE. Thus, the anode corrodes. The reaction stops unless the electrons are removed.
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Some additional examples of oxidation reactions occurring at the anode are: 𝐶𝑢 → 𝐶𝑢$% + 2𝑒 ) 𝑁𝑖 → 𝑁 % + 𝑒 ) 𝐴𝑙 → 𝐴𝑙 1% + 3𝑒 ) 𝐻4 → 2𝐻 % + 2𝑒 ) REDUCTION is the process by which metals conserve electrons. In the chemical reaction below, electrons are consumed: 𝑀$% + 𝑛 → 𝑀 Since the electrodes pass from the anode to the cathode by electrical conduction, the anode and the cathode must be electrically connected. All metals are subject to the oxidation reaction, depending on their oxidation potential. If ions of a metal with a low oxidation potential are present, they can be reduced, consuming electrons from preceding corrosion reaction. Electroplating use this reaction to deposit metals from a solution by the addition of electrons. Corrosion will stop if the electrical connection interrupted, the cathode reactants are depleted, or if the anode products are saturated. The EQUATIONS FOR POSSIBLE CATHODE REACTIONS (REDUCTION) OF A TYPICAL METAL can be referenced under the topic of CORROSION on page 60 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
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Some additional examples of reduction reactions occurring at the cathode are: 5 4 5 4
𝑂4 + 2𝑒 ) + 𝐻4 𝑂 → 2𝑂𝐻 ) 𝑂4 + 2𝑒 ) + 2𝐻1 𝑂% → 3𝐻4 𝑂
2𝑒 ) + 2𝐻1 𝑂% → 2𝐻4 𝑂 + 𝐻4 𝐶𝑢4% + 2𝑒 ) → 𝐶𝑢 𝑁𝑖 % + 𝑒 ) → 𝑁𝑖 𝐴𝑙 1% + 3𝑒 ) → 𝐴𝑙 2𝐻 % + 2𝑒 ) → 𝐻4 RUSTING: The TOPIC OF RUSTING is not provided in the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. Therefore, we must memorize these concepts and understand their application independent of the NCEES Supplied Reference Handbook. RUSTING is the corrosion of iron or iron-based alloys. Therefore, it is important to realize that nonferrous metals may corrode but will 𝑛𝑒𝑣𝑒𝑟 rust. Corrosion is a process that occurs in the presence of moisture, whether it be in the air or a surrounding fluid. In order for corrosion to occur, moisture and some medium must be present.
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The chemical equation for when irons rusts due to exposure to moisture in the air is represented by the equation: 𝐹𝑒1 𝑂1 ∙ 𝑋 𝐻4 𝑂 Where: • 𝐹𝑒1 𝑂1 is Iron (III) Oxide also known as Ferric Oxide • “𝑋” represents the number of moles or amount of water present in the air The amount of water present also determines the color of rust, which may vary from black to yellow to orange brown. ELECTROCHEMISTRY: The TOPIC OF ELECTROCHEMISTRY OF A TYPICAL METAL can be referenced under the subject of CHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. GALVANIC ACTION occurs when two different metals having varying properties are put into a conductive solution that is electrically connected. The metal having the higher electro-chemical potential will act as the anode, and will corrode as electrons leave the metal. The metal having the lower electro-chemical potential will be the cathode and gain electrons. For corrosion to occur, there must be an anode and a cathode in electrical contact within the presence of an electrolyte.
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An ELECTROLYE is a medium that is typically an AQUEOUS SOLUTION where ions accumulate during a corrosion reaction. An electrolyte is necessary for corrosion to occur. A common electrolyte is WATER or AIR. In the presence of an electrolyte, the use of dissimilar metals will result in GALVANIC CORROSION. Electrochemistry problems will commonly ask for the electrical potential, “𝐸> ”, of a corrosion reaction in units of Volts (𝑉). To answer these questions, we simply reference the table in the Reference Handbook, and look up the corrosion reactions to find the corresponding electrical potential, “𝐸> ” in units of volts. There is no need to do any calculations or over analyze these type of questions. If Iron is placed in water, it will rapidly produce electrons and ions. The presence of negatively charged electrons and positively charged ions results in an electrical potential. Hydrogen atoms will also dissolve in water, producing an electron and a positively charged hydrogen ion. The potential difference between a source of hydrogen and plate of iron in water will promote an electric current from iron to the hydrogen source. This pair consisting of the hydrogen source and the iron plate is called a GALVANIC COUPLE. The OXIDATION REACTION occurs at the anode while the reduction reaction occurs at the cathode. For metallic corrosion, there must be an anode, an electrolyte, and a cathode, to form a combination that is called the GALVANIC CELL. The TABLE LISTING THE STANDARD ELECTROMOTIVE POTENTIALS OF METALS can be referenced under the topic of ELECTROCHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
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Anodic and cathodic reactions are reversible. The reaction 𝐶𝑢 → 𝐶𝑢4% + 2𝑒 ) is anodic and causes corrosion. The reaction 𝐶𝑢4% + 2𝑒 ) → 𝐶𝑢 is cathodic and causes ELECTROPLATING. The formation of rust is a very complex process, which is thought to begin with the oxidation of Iron to Ferrous (Iron +2) Ions. The dissolving of Iron in the water to produce Iron ions is referred to as SOLUTION CORROSION. OXIDATION POTENTIAL: The TABLE OF STANDARD OXIDATION POTENTIALS FOR CORROSION REACTIONS can be referenced under the topic of ELECTROCHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. The OXIDATION POTENTIAL of a given oxidation reaction is the tendency of materials to be anodic or cathodic. The PERIODIC TABLE OF ELEMENTS can be referenced under the topic of CHEMISTRY on page 55 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. From our studies of chemistry, we know that by looking at the periodic table we can see the trends for oxidation potentials. As we move down the periodic table, the oxidation potentials increase, such that the materials are increasingly anodic towards the bottom of the table. On the other hand, the materials towards the top of the table are cathodic and have higher corrosion resistance.
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If the anode is small compared to the cathode, the current density,
@ABC AA D
, will be high
and corrosion will be rapid. CORROSION CONTROL: The TOPIC OF CORROSION CONTROL is not provided in the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. Therefore, we must memorize these concepts and understand their application independent of the NCEES Supplied Reference Handbook. There are several approaches for corrosion control with the strategy depending on the specific situation. The most obvious strategy is the avoidance of galvanic cells. For example, iron or steel screws should not be used in bronze marine hardware, since iron is anodic to “𝐶𝑢” and will corrode more easily. Also, single-phase alloys are more resistive to corrosion than the same alloys with multiple phases. For example, corrosion can be established in Sterling Silver (92.5% 𝐴𝑔 𝑎𝑛𝑑 7.5% 𝐶𝑢) which has phases “𝛼 𝑎𝑛𝑑 𝛽". In this case, Cu-rich 𝛽-phase acts as the anode. PROTECTIVE COATINGS are commonly used in controlling corrosion. Structural steel is commonly galvanized with zinc to prevent rusting. If the Zinc coating is punctured, the steel becomes the cathode and is protected. On the other hand, if a tin coating on steel is punctured, the steel underneath corrodes very rapidly because it is anodic to tin.
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ANODIZING is a process that uses the electrical current to force the formation of a thicket 𝐴𝑙4 𝑂1 film than occurs naturally. The iron oxide film is porous and not stable. Aluminum (𝐴𝑙) has a higher oxidation potential than Iron (𝐹𝑒). Yet, aluminum has better corrosion resistance than iron. This is because aluminum reacts readily with oxygen to form a stable coating of aluminum oxide, which protects the base metal. This phenomenon is known as PASSIVITY and also accounts for the excellent corrosion resistance of stainless steel due to the presence of chromium oxide protective film. STAINLESS STEELS are stainless because their surfaces are passivated by a thin chromium oxide film. The chromium reacts to form iron 𝐶𝑟 ions before the iron reacts. If oxygen is present, it immediately claims the released electrons and forms an extremely thin chromium-oxide film that halts all further reaction. The reaction is called PASSIVATION, and gives a high degree of protection. INHIBITORS are also used to provide another variation of a passive surface film. Inhibitors contain chromate or similar highly oxidized ions that are absorbed onto the metal surface. The protection received is similar to that afforded by the passive film on stainless steel, except for the fact that it cannot be regenerated from the metal. New inhibitor metal must be added if the film is flushed away. GALVANIC PROTECTION: The TOPIC OF GALVANIC PROTECTION is not provided in the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. We must memorize these concepts and understand their application independent of the NCEES Supplied Reference Handbook.
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GALVANIC PROTECTION is used to achieve service protection by using a DC current that feeds electrons into the metal that must be protected against corrosion. Magnesium strips are commonly used as CATHODIC PROTECTION for steel structures and steel components to prevent them from rusting. From the table on page 59, it is clear that Magnesium (𝑀𝑔) is anodic to steel (which is essentially Iron, 𝐹𝑒). Thus Magnesium becomes the anode and corrodes, and steel becomes the cathode and is protected. In this case, “𝑀𝑔” is known as a SACRIFICIAL ANODE.
CONCEPT EXAMPLE: The following problem introduces the concept reviewed within this module. Use this content as a primer for the subsequent material.
What is the electrical potential of the element lead gaining two electrons? A. −0.788 B. +0.126 C. +2.363 D. +2.925
SOLUTION: The PERIODIC TABLE OF ELEMENTS can be referenced under the topic of CHEMISTRY on page 55 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
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The first thing we must do in this problem is identify the atomic symbol used to represent the element of lead. Looking at the periodic table, we know that potassium is represented by the symbol “𝑃𝑏”. The TABLE LISTING THE STANDARD ELECTROMOTIVE POTENTIALS OF METALS can be referenced under the topic of ELECTROCHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. The next step in this problem is to look at the table in reference handbook that lists out the standard oxidation potential for corrosion reactions, and the associated electrical potential for each reaction. In the problem statement we are asked for the electrical potential of a chemical reaction when lead gains an electron. We realize that the chemical reaction must use the symbol “𝑃𝑏” to represent lead, as well as have the addition or subtraction of electrons, represented by the symbol 𝑒 ) . Looking at the table in the reference handbook, we find the chemical reaction for lead gaining two electrons is represented as: 𝑃𝑏 → 𝑃𝑏 4% + 2𝑒 )
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Looking at the associated electrical potential for this reaction, we find the electrical potential of lead gaining two electrons is +0.126 Volts.
Therefore, the correct answer choice is B. +𝟎. 𝟏𝟐𝟔
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CONCEPT INTRO: When analyzing materials, it is important that we study how materials can be damaged, and the impacts various processes and reactions can have on the integrity of a material. Materials can be damaged by reacting chemically with their environments. Factors such as moisture and the amount of oxygen, can reactions such as OXIDATION or CORROSION to occur. Corrosion is the deterioration of a metal by chemical or electrochemical reactions with the environment. Corrosion is affected by six major factors: 1. Acidity – Acids form hydrogen ions in water solution. The hydrogen ions tend to be replaced by metal ions, freeing hydrogen gas. 2. Oxidizing Characteristics – Oxidation causes an increase in positive valence electrons. Therefore, any metal whose valence is zero would cause a change to a compound having a positive valence. This acts in the same direction as the acid reaction. The two, acid reaction and oxidation, frequently act together. 3. Electrolysis - In ELECTROLYSIS, the positive electrode called the ANODE will take up or absorb electrons. The metal in it, minus the electrons, will go into solution as a positive ion. At the negative electrode, called the CATHODE, hydrogen ions take on the electrons and become hydrogen gas. The result on the metal of this electrolysis is a pitting or roughening of the metal surface.
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4. Film Formation – A thick film forms a surface layer which inhibits any further corrosion. A film such as rust can deny free access to the corroding medium such as water. 5. Agitation or Rate of Movement – Agitation will break down layers of concentration between the metal and the corroding medium. This causes the reaction to proceed more rapidly than it would have if the layers of concentration were not disturbed. 6. Temperature – Temperature is one measure of heat energy. Adding energy to a system will increase its temperature. The highest the temperature, the faster the rate of corrosion. The EQUATION FOR THE ANODE REACTION (OXIDATION) OF A TYPICAL METAL can be referenced under the topic of CORROSION on page 60 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. OXIDATION is the removal of electrons from an atom and is the process by which metals lose mass and corrode over time. In this process, metal atoms lose one or more electrons and become metal ions as expressed in the chemical reaction. The Anode Reaction (Oxidation) of a Typical Metal “𝑀” is: 𝑀° → 𝑀$% + 𝑛𝑒 ) The reaction above is a OXIDATION REACTION occurring at the ANODE. Thus, the anode corrodes. The reaction stops unless the electrons are removed.
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Some additional examples of oxidation reactions occurring at the anode are: 𝐶𝑢 → 𝐶𝑢$% + 2𝑒 ) 𝑁𝑖 → 𝑁 % + 𝑒 ) 𝐴𝑙 → 𝐴𝑙 1% + 3𝑒 ) 𝐻4 → 2𝐻 % + 2𝑒 ) REDUCTION is the process by which metals conserve electrons. In the chemical reaction below, electrons are consumed: 𝑀$% + 𝑛 → 𝑀 Since the electrodes pass from the anode to the cathode by electrical conduction, the anode and the cathode must be electrically connected. All metals are subject to the oxidation reaction, depending on their oxidation potential. If ions of a metal with a low oxidation potential are present, they can be reduced, consuming electrons from preceding corrosion reaction. Electroplating use this reaction to deposit metals from a solution by the addition of electrons. Corrosion will stop if the electrical connection interrupted, the cathode reactants are depleted, or if the anode products are saturated. The EQUATIONS FOR POSSIBLE CATHODE REACTIONS (REDUCTION) OF A TYPICAL METAL can be referenced under the topic of CORROSION on page 60 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
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Some additional examples of reduction reactions occurring at the cathode are: 5 4 5 4
𝑂4 + 2𝑒 ) + 𝐻4 𝑂 → 2𝑂𝐻 ) 𝑂4 + 2𝑒 ) + 2𝐻1 𝑂% → 3𝐻4 𝑂
2𝑒 ) + 2𝐻1 𝑂% → 2𝐻4 𝑂 + 𝐻4 𝐶𝑢4% + 2𝑒 ) → 𝐶𝑢 𝑁𝑖 % + 𝑒 ) → 𝑁𝑖 𝐴𝑙 1% + 3𝑒 ) → 𝐴𝑙 2𝐻 % + 2𝑒 ) → 𝐻4 RUSTING: The TOPIC OF RUSTING is not provided in the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. Therefore, we must memorize these concepts and understand their application independent of the NCEES Supplied Reference Handbook. RUSTING is the corrosion of iron or iron-based alloys. Therefore, it is important to realize that nonferrous metals may corrode but will 𝑛𝑒𝑣𝑒𝑟 rust. Corrosion is a process that occurs in the presence of moisture, whether it be in the air or a surrounding fluid. In order for corrosion to occur, moisture and some medium must be present.
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The chemical equation for when irons rusts due to exposure to moisture in the air is represented by the equation: 𝐹𝑒1 𝑂1 ∙ 𝑋 𝐻4 𝑂 Where: • 𝐹𝑒1 𝑂1 is Iron (III) Oxide also known as Ferric Oxide • “𝑋” represents the number of moles or amount of water present in the air The amount of water present also determines the color of rust, which may vary from black to yellow to orange brown. ELECTROCHEMISTRY: The TOPIC OF ELECTROCHEMISTRY OF A TYPICAL METAL can be referenced under the subject of CHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. GALVANIC ACTION occurs when two different metals having varying properties are put into a conductive solution that is electrically connected. The metal having the higher electro-chemical potential will act as the anode, and will corrode as electrons leave the metal. The metal having the lower electro-chemical potential will be the cathode and gain electrons. For corrosion to occur, there must be an anode and a cathode in electrical contact within the presence of an electrolyte.
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An ELECTROLYE is a medium that is typically an AQUEOUS SOLUTION where ions accumulate during a corrosion reaction. An electrolyte is necessary for corrosion to occur. A common electrolyte is WATER or AIR. In the presence of an electrolyte, the use of dissimilar metals will result in GALVANIC CORROSION. Electrochemistry problems will commonly ask for the electrical potential, “𝐸> ”, of a corrosion reaction in units of Volts (𝑉). To answer these questions, we simply reference the table in the Reference Handbook, and look up the corrosion reactions to find the corresponding electrical potential, “𝐸> ” in units of volts. There is no need to do any calculations or over analyze these type of questions. If Iron is placed in water, it will rapidly produce electrons and ions. The presence of negatively charged electrons and positively charged ions results in an electrical potential. Hydrogen atoms will also dissolve in water, producing an electron and a positively charged hydrogen ion. The potential difference between a source of hydrogen and plate of iron in water will promote an electric current from iron to the hydrogen source. This pair consisting of the hydrogen source and the iron plate is called a GALVANIC COUPLE. The OXIDATION REACTION occurs at the anode while the reduction reaction occurs at the cathode. For metallic corrosion, there must be an anode, an electrolyte, and a cathode, to form a combination that is called the GALVANIC CELL. The TABLE LISTING THE STANDARD ELECTROMOTIVE POTENTIALS OF METALS can be referenced under the topic of ELECTROCHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
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Anodic and cathodic reactions are reversible. The reaction 𝐶𝑢 → 𝐶𝑢4% + 2𝑒 ) is anodic and causes corrosion. The reaction 𝐶𝑢4% + 2𝑒 ) → 𝐶𝑢 is cathodic and causes ELECTROPLATING. The formation of rust is a very complex process, which is thought to begin with the oxidation of Iron to Ferrous (Iron +2) Ions. The dissolving of Iron in the water to produce Iron ions is referred to as SOLUTION CORROSION. OXIDATION POTENTIAL: The TABLE OF STANDARD OXIDATION POTENTIALS FOR CORROSION REACTIONS can be referenced under the topic of ELECTROCHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. The OXIDATION POTENTIAL of a given oxidation reaction is the tendency of materials to be anodic or cathodic. The PERIODIC TABLE OF ELEMENTS can be referenced under the topic of CHEMISTRY on page 55 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. From our studies of chemistry, we know that by looking at the periodic table we can see the trends for oxidation potentials. As we move down the periodic table, the oxidation potentials increase, such that the materials are increasingly anodic towards the bottom of the table. On the other hand, the materials towards the top of the table are cathodic and have higher corrosion resistance.
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If the anode is small compared to the cathode, the current density,
@ABC AA D
, will be high
and corrosion will be rapid. CORROSION CONTROL: The TOPIC OF CORROSION CONTROL is not provided in the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. Therefore, we must memorize these concepts and understand their application independent of the NCEES Supplied Reference Handbook. There are several approaches for corrosion control with the strategy depending on the specific situation. The most obvious strategy is the avoidance of galvanic cells. For example, iron or steel screws should not be used in bronze marine hardware, since iron is anodic to “𝐶𝑢” and will corrode more easily. Also, single-phase alloys are more resistive to corrosion than the same alloys with multiple phases. For example, corrosion can be established in Sterling Silver (92.5% 𝐴𝑔 𝑎𝑛𝑑 7.5% 𝐶𝑢) which has phases “𝛼 𝑎𝑛𝑑 𝛽". In this case, Cu-rich 𝛽-phase acts as the anode. PROTECTIVE COATINGS are commonly used in controlling corrosion. Structural steel is commonly galvanized with zinc to prevent rusting. If the Zinc coating is punctured, the steel becomes the cathode and is protected. On the other hand, if a tin coating on steel is punctured, the steel underneath corrodes very rapidly because it is anodic to tin.
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ANODIZING is a process that uses the electrical current to force the formation of a thicket 𝐴𝑙4 𝑂1 film than occurs naturally. The iron oxide film is porous and not stable. Aluminum (𝐴𝑙) has a higher oxidation potential than Iron (𝐹𝑒). Yet, aluminum has better corrosion resistance than iron. This is because aluminum reacts readily with oxygen to form a stable coating of aluminum oxide, which protects the base metal. This phenomenon is known as PASSIVITY and also accounts for the excellent corrosion resistance of stainless steel due to the presence of chromium oxide protective film. STAINLESS STEELS are stainless because their surfaces are passivated by a thin chromium oxide film. The chromium reacts to form iron 𝐶𝑟 ions before the iron reacts. If oxygen is present, it immediately claims the released electrons and forms an extremely thin chromium-oxide film that halts all further reaction. The reaction is called PASSIVATION, and gives a high degree of protection. INHIBITORS are also used to provide another variation of a passive surface film. Inhibitors contain chromate or similar highly oxidized ions that are absorbed onto the metal surface. The protection received is similar to that afforded by the passive film on stainless steel, except for the fact that it cannot be regenerated from the metal. New inhibitor metal must be added if the film is flushed away. GALVANIC PROTECTION: The TOPIC OF GALVANIC PROTECTION is not provided in the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. We must memorize these concepts and understand their application independent of the NCEES Supplied Reference Handbook.
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GALVANIC PROTECTION is used to achieve service protection by using a DC current that feeds electrons into the metal that must be protected against corrosion. Magnesium strips are commonly used as CATHODIC PROTECTION for steel structures and steel components to prevent them from rusting. From the table on page 59, it is clear that Magnesium (𝑀𝑔) is anodic to steel (which is essentially Iron, 𝐹𝑒). Thus Magnesium becomes the anode and corrodes, and steel becomes the cathode and is protected. In this case, “𝑀𝑔” is known as a SACRIFICIAL ANODE.
CONCEPT EXAMPLE: The following problem introduces the concept reviewed within this module. Use this content as a primer for the subsequent material.
What is the electrical potential of the element lead gaining two electrons? A. −0.788 B. +0.126 C. +2.363 D. +2.925
SOLUTION: The PERIODIC TABLE OF ELEMENTS can be referenced under the topic of CHEMISTRY on page 55 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing.
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The first thing we must do in this problem is identify the atomic symbol used to represent the element of lead. Looking at the periodic table, we know that potassium is represented by the symbol “𝑃𝑏”. The TABLE LISTING THE STANDARD ELECTROMOTIVE POTENTIALS OF METALS can be referenced under the topic of ELECTROCHEMISTRY on page 59 of the NCEES Supplied Reference Handbook, Version 9.4 for Computer Based Testing. The next step in this problem is to look at the table in reference handbook that lists out the standard oxidation potential for corrosion reactions, and the associated electrical potential for each reaction. In the problem statement we are asked for the electrical potential of a chemical reaction when lead gains an electron. We realize that the chemical reaction must use the symbol “𝑃𝑏” to represent lead, as well as have the addition or subtraction of electrons, represented by the symbol 𝑒 ) . Looking at the table in the reference handbook, we find the chemical reaction for lead gaining two electrons is represented as: 𝑃𝑏 → 𝑃𝑏 4% + 2𝑒 )
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Looking at the associated electrical potential for this reaction, we find the electrical potential of lead gaining two electrons is +0.126 Volts.
Therefore, the correct answer choice is B. +𝟎. 𝟏𝟐𝟔
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