Question

An iron object is plated with a coating of cobalt to protect against corrosion. Does the cobalt protect iron by cathodic protection? Explain.

Short answer

No, the cobalt coating does not provide cathodic protection to the iron object against corrosion. This is because the standard electrode potential of iron (\( E^0(Fe^{2+}/Fe) = - 0.44 V \)) is more negative than that of cobalt (\( E^0(Co^{2+}/Co) = - 0.28 V \)). As a result, iron is more easily oxidized (corroded) than cobalt, and the iron will still corrode preferentially over the cobalt coating.

Step-by-step solution

Understand 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 usually achieved by connecting the metal to be protected to a more easily corroded "sacrificial metal" that will act as the anode. The sacrificial metal then corrodes preferentially, protecting the cathode (the metal we want to preserve) from corrosion.

Electrochemical Series of Metals

To determine if cobalt can provide cathodic protection to iron, we need to look at the electrochemical series of metals. The electrochemical series is a list of metals arranged in order of their standard electrode potentials. In general, when two metals are in contact with each other in an electrolyte, the metal with the lower (more negative) electrode potential will be the anode, meaning that it will corrode preferentially.

Check Iron and Cobalt's Position on Electrochemical Series

Now, let's find the standard electrode potentials of both iron and cobalt from the electrochemical series: Iron(Fe): \[ E^0(Fe^{2+}/Fe) = - 0.44 V \] Cobalt(Co): \[ E^0(Co^{2+}/Co) = - 0.28 V \]

Determine if Cobalt can provide Cathodic Protection for Iron

From the electrochemical series, we can see that the standard electrode potential of iron (Fe) is -0.44 V, which is more negative than the standard electrode potential of cobalt (Co) which is -0.28 V. This means that in an electrochemical cell, iron is more easily oxidized (corroded) than cobalt. So, when a cobalt coating is applied on an iron object, the iron will still corrode preferentially over cobalt. Thus, the cobalt coating does not provide cathodic protection to the iron object against corrosion.

Key concepts

Cathodic Protection

Cathodic protection is a clever method used to prevent metal surfaces from corroding. Imagine your metal object as the center of a protective bubble. By making the metal of interest—the cathode—less likely to react, it stays safe from corrosion. This is typically done by connecting it to a more reactive, and more easily corroded, metal known as the "sacrificial anode."

Here's how it works:

• The sacrificial metal corrodes instead of the protected metal.
• As the name suggests, it sacrifices itself to save the more valuable metal.
• This technique is commonly used for iron structures, like pipes, by attaching metals like zinc.

In the example of cathodic protection, the choice of the sacrificial metal is crucial for the system to work effectively. The metal needs to have a more negative standard electrode potential than the metal it aims to protect. If not correctly chosen, the system can end up not providing the protection intended.

Standard Electrode Potentials

Standard electrode potentials (SEP) are a key concept in understanding which metals will corrode and which will stay protected when used together. Simply put, SEP measures the tendency of a metal to lose electrons, making it a part of the Electrochemical Series—a ranking list of metals based on how easily they oxidize.

In an electrochemical series:

• Metals with more negative SEP values are more proactive—they give up electrons easily and tend to corrode faster.
• Metals with less negative, or more positive, potentials are more passive—they hold on to their electrons and resist corrosion better.

To understand which metal can protect another, you compare their SEPs. If metal A has a SEP more negative than metal B, metal A will corrode before metal B.

For iron and cobalt, iron has a SEP of -0.44 V, which is more negative than cobalt's -0.28 V. Hence, iron will act as the anode and corrode before cobalt if both are in contact. This comparison clearly shows why cobalt cannot protect iron from corrosion by cathodic protection.

Metal Corrosion

Corrosion is a natural process where metals deteriorate due to reactions with their environment. Most commonly seen as rust on iron, corrosion can severely damage metal structures over time.

Corrosion occurs because metals tend to revert to their more stable, lower energy forms, usually as oxides or sulfides. This process is essentially a redox reaction where the metal loses electrons, gradually deteriorating:

• Moisture (water) and oxygen are primary agents fueling corrosion, especially in metals like iron.
• Electrolytes, such as salts, speed up this process by aiding electron transfer.

Preventing corrosion is critical in prolonging the lifespan of metal structures. Several strategies exist:

• Physical barriers like paints or coatings can isolate the metal from corrosive elements.
• Using cathodic protection slows down or halts the process by making the metal the cathode of an electrochemical cell.

Understanding these concepts allows engineers and scientists to design more durable and sustainable metal structures. In practice, knowing which metals corrode and which can protect others is crucial in the fight against corrosion.