Chapter 20: Problem 86
An iron object is plated with a coating of cobalt to protect against corrosion. Does the cobalt protect iron by cathodic protection?
Short Answer
Expert verified
Cobalt does not provide cathodic protection to the iron object. This is because the standard reduction potential for iron (\(E^0 = -0.44 V\)) is more negative than that of cobalt (\(E^0 = -0.28 V\)), meaning that iron is more easily oxidized. Hence, iron would act as a sacrificial anode for the cobalt, which is not the intended purpose.
Step by step solution
01
Understanding Cathodic Protection
Cathodic protection is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. This is achieved by placing a more easily oxidized metal (having a more negative standard reduction potential) in contact with the metal we want to protect, forming a galvanic cell. The more easily oxidized metal (called the sacrificial anode) will be oxidized, and the protected metal will be reduced, preventing its corrosion.
02
Checking Standard Reduction Potentials
We need to compare the standard reduction potentials of iron and cobalt to determine if cobalt can act as a sacrificial anode for iron.
The standard reduction potential for iron (Fe):
\(Fe^{2+}(aq) + 2e^- \rightarrow Fe(s)\) \(E^0 = -0.44 V\)
The standard reduction potential for cobalt (Co):
\(Co^{2+}(aq) + 2e^- \rightarrow Co(s)\) \(E^0 = -0.28 V\)
03
Comparing Standard Reduction Potentials
Now that we have the standard reduction potentials for both iron and cobalt, we can compare them:
Iron: \(E^0 = -0.44 V\)
Cobalt: \(E^0 = -0.28 V\)
Since the standard reduction potential for iron is more negative than that of cobalt, iron is more easily oxidized.
04
Determining Cathodic Protection
As iron has a more negative standard reduction potential, it is more likely to be oxidized than cobalt. Therefore, cobalt cannot provide cathodic protection to the iron object, as it is not more easily oxidized. Instead, the iron would act as a sacrificial anode for the cobalt, which is not the desired outcome.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Standard Reduction Potential
Standard reduction potential is a measure of the tendency of a chemical species to gain electrons and be reduced. In electrochemistry, it is typically represented as the voltage difference when a standard reference electrode is coupled with another electrode. Its value lets us predict the direction of electron flow in an electrochemical cell.
- The more negative the standard reduction potential value, the more likely the species is to lose electrons and be oxidized.- Conversely, a more positive value indicates a higher likelihood of gaining electrons and being reduced.
For iron, the standard reduction potential is \(E^0 = -0.44 \, ext{V}\). This implies it readily undergoes oxidation. Meanwhile, cobalt has a standard reduction potential of \(E^0 = -0.28 \, ext{V}\), indicating it is less prone to oxidation than iron. Comparing these potentials reveals critical insights into which metal can protect the other from corrosion effects.
- The more negative the standard reduction potential value, the more likely the species is to lose electrons and be oxidized.- Conversely, a more positive value indicates a higher likelihood of gaining electrons and being reduced.
For iron, the standard reduction potential is \(E^0 = -0.44 \, ext{V}\). This implies it readily undergoes oxidation. Meanwhile, cobalt has a standard reduction potential of \(E^0 = -0.28 \, ext{V}\), indicating it is less prone to oxidation than iron. Comparing these potentials reveals critical insights into which metal can protect the other from corrosion effects.
Corrosion Prevention
Corrosion prevention involves using various methods to protect metal surfaces from deteriorating due to chemical reactions in their environment. Corrosion typically manifests as rusting, tarnishing, or similar processes, which can weaken the metal and cause structural damage.
- One common strategy is to prevent direct contact with corrosive substances like water and oxygen. - Another strategy is cathodic protection, where a more easily oxidized metal sacrifices itself to protect the base metal. - Coatings, galvanizing, and use of inhibitors are additional preventive measures.
To prevent the corrosion of iron using a coating of cobalt, it is essential to consider the electrochemical potentials of the metals involved. Unfortunately, as seen in our example, cobalt is not suitable for cathodic protection of iron since it is not more easily oxidized than iron.
- One common strategy is to prevent direct contact with corrosive substances like water and oxygen. - Another strategy is cathodic protection, where a more easily oxidized metal sacrifices itself to protect the base metal. - Coatings, galvanizing, and use of inhibitors are additional preventive measures.
To prevent the corrosion of iron using a coating of cobalt, it is essential to consider the electrochemical potentials of the metals involved. Unfortunately, as seen in our example, cobalt is not suitable for cathodic protection of iron since it is not more easily oxidized than iron.
Galvanic Cell
A galvanic cell is an electrochemical cell that derives electrical energy from spontaneous redox reactions taking place within the cell. It consists of two half-cells, where oxidation occurs at the anode and reduction happens at the cathode. This energy conversion is fundamental in how batteries work and has practical applications like cathodic protection.
- **Anode:** The electrode where oxidation happens. - **Cathode:** The electrode where reduction happens. - **Salt Bridge:** It completes the circuit by allowing ion flow, preventing charge buildup.
In the context of cathodic protection, the more easily oxidized metal acts as the anode, gradually corroding and thereby sparing the protected metal at the cathode. This dynamic keeps the protected metal (e.g., iron) free from rust as long as the anode metal continues to corrode.
- **Anode:** The electrode where oxidation happens. - **Cathode:** The electrode where reduction happens. - **Salt Bridge:** It completes the circuit by allowing ion flow, preventing charge buildup.
In the context of cathodic protection, the more easily oxidized metal acts as the anode, gradually corroding and thereby sparing the protected metal at the cathode. This dynamic keeps the protected metal (e.g., iron) free from rust as long as the anode metal continues to corrode.
Sacrificial Anode
The sacrificial anode is a pivotal part of cathodic protection systems. It is a more reactive metal that willingly corrodes, thereby protecting a less reactive metal. When the anode corrodes, it releases electrons, which flow to the less reactive metal.
- This process keeps the protected metal in a reduced state, minimizing its corrosion. - Common materials for sacrificial anodes include zinc, magnesium, and aluminum, all of which have more negative standard reduction potentials compared to the metal they protect.
In the exercise with iron and cobalt, iron cannot be protected by cobalt since iron has a more negative potential. Therefore, if cathodic protection is the goal, one must choose a material with an even higher likelihood to oxidize than iron to serve as the sacrificial anode.
- This process keeps the protected metal in a reduced state, minimizing its corrosion. - Common materials for sacrificial anodes include zinc, magnesium, and aluminum, all of which have more negative standard reduction potentials compared to the metal they protect.
In the exercise with iron and cobalt, iron cannot be protected by cobalt since iron has a more negative potential. Therefore, if cathodic protection is the goal, one must choose a material with an even higher likelihood to oxidize than iron to serve as the sacrificial anode.