How Does Corrosion Happen?
Corrosion, or rust as it is commonly referred to, is the degradation of metal. In the most simple expression most people understand it to be a one-way process:
metal → Rust
That is, over time metal turns to rust.
Now as you suspected, it is a little more complicated and you want to understand how this process really happens.
Broadly defined, corrosion falls under two categories: UNIFORM corrosion and LOCALIZED corrosion
Uniform corrosion is what you see when most of the metal is evenly coated in corrosion.
Localized corrosion, on the other hand, tends to happen in a small area and it also tends to be more destructive. The "rust hole" you see on the car door panel is an example of localized corrosion. There are many types of localized corrosion which will be discussed later.
UNIFORM CORROSION EXPLAINED
The process of uniform corrosion is fairly simple so we will start with it. Uniform corrosion begins with a metal, for example, steel exposed to oxygen and moisture in the air. As you probably know, steel primarily consists of the element Iron, represented by the symbol Fe on the periodic table. Air, as you know, primarily consists of the elements nitrogen and oxygen.
Now, this steel plate exists because someone processed iron ore, melted the metal and added other ingredients to the iron to make the steel plate. A lot of energy goes into making steel. The problem with a product, like steel, that requires energy input, is that on our planet earth there is a law called the Second Law of Thermodynamics which states that systems tend to equilibrium. Therefore, products like steel, with higher states of energy will want to resort to lower states of energy.
The way a system gives up energy is by letting go of electrons.
The flow of electrons happens to be the first step in the corrosion process.
Step 1: Steel meets Water
Imagine that the line below is the surface of a steel plate. On the surface, you would have Fe atoms.
The corrosion process begins when the steel plate is exposed to water and air. When water comes into contact with the metal, the Fe atom releases two electrons into the water, dissolves in the water itself and itself and in the process become Fe++ ion, or ferrous ions as they are known. In corrosion lingo, the solution with dissolved ions is called an electrolyte. So now you have an electrolyte (solution) with a ferrous ion and 2 elections. As Fe atoms leave the metal surface it begins to degrade or corrode. Now for a few technical terms. Whenever a metal GIVES UP electrons as the iron just did, we call that metal an ANODE. The process of giving up electrons is called OXIDATION. Also, the generic term that is used to describe a metal that supplies electrons to an electrolyte is called an ELECTRODE. In the above example, the iron is an electrode.
Second Step: Free electrons react with Oxygen and Water
The oxidation process just described in step one NEVER happens on its own. It is always part of another reaction called REDUCTION. You can think of a reduction reaction as the opposite of an oxidation reaction. The two reactions must be going on at the same time otherwise neither reaction would happen. So how are reduction reactions part of the corrosion process?
The two electrons from the Fe atom would not have released from the atom had it not been for oxygen (O2) in the water which is willing to take the electrons. The 4 electrons react with oxygen in water to form to form 4 negative hydroxide molecules (4OH-)
O2 + 2H2O + 4e ----> 4OH-
At this point the ferrous ions (Fe2+) react with the hydroxide molecules to form ferrous Hydroxide, also known as Iron(II) Hydroxide
2Fe2+ + 4OH- -------> 2Fe(OH)2
Ferrous Hydroxide then further reacts with oxygen to form Ferric Hydroxide
2Fe(OH)2 + O2 + H2O ---------> 4Fe(OH)3
In a final step, Ferric Hydroxide further reacts with oxygen in water to form Hydrated Ferric Oxide which we commonly call rust.
4Fe(OH)3 + O2 ----------> 2FeO3.2H2O
As was mentioned, this is an oxidation and reduction process. The Fe atoms oxidize (loose) electrons and the oxygen reduces (gains) electrons. This oxidation/reduction reaction is shorted to phrase REDOX reaction.
In the first step it was also mentioned that iron in the corrosion process is called an Anode because it is the electrode giving up electrons. A substance that takes electrons is called a Cathode. In the example above, oxygen is the Cathode because it takes the electrons given up by the iron. These terms are important to remember because the corroding metal will always be referred to as the anode or cathode.
Uniform corrosion is the uniform degradation of metal. Oxygen in contact with iron will not initiate corrosion but a solvent, such as water, is needed so that the iron can dissolve and form a ferrous ion solution. In corrosion lingo, the solution is called an Electrolyte. In other words, dissolved ions are necessary for corrosion and so we need a solvent, such as water, to dissolve the metal. We also need a means for the electrons to travel from the metal to the oxygen. For uniform corrosion, the electrolyte acts both as a medium for dissolved ions AND as a conducting pathway for the electrons to migrate from the metal to the dissolved oxygen. To summarize, there are 4 necessary components to corrosion:
1) Anode (supplies electrons)
2) Cathode (receives electrons)
3) Electrical conductor (path for electrons to move from the anode to cathode)
4) Electrolyte (solvent that dissolves metal into ions)
Note: The word electrolyte can be confusing because it has two correct meanings. It can be used to describe the solution with dissolved ions, or it can be used to describe the dissolved ions. When talking about electrolytes here, we are referring to dissolved ions that are in a solution.
LOCALIZED CORROSION EXPLAINED
Engineers are usually more concerned with the types of corrosion that fall under the "localized corrosion". Like an iceberg, a small area at the surface can represent a lot more corrosion beneath. For this reason, localized corrosion is responsible for the majority of metal failures due to corrosion. Some examples of localized corrosion include galvanic corrosion, pitting corrosion and crevice corrosion. In this section we will discuss one of the most common forms of localized corrosion: galvanic corrosion.
GALVANIC CORROSION EXPLAINED
Galvanic corrosion occurs when two dissimilar metals are in contact with each other in the presence of an electrolyte .
How can corrosion simply be the result of two dissimilar metals that are in contact with each other?
The fact that you now understand uniform corrosion will help in understanding galvanic corrosion. Remember that the anode in our example was the iron metal that gave up electrons and the cathode was oxygen. When two dissimilar metals (electrodes) are in contact with each other the same electron flow can happen except they now happen between two metals that have different electrode potentials. Of course, you will need to understand electrode potential before we can move on.
The electrode (or redox) potential of an electrode refers to the ease with which an electrode (metal) gives up electrons to a solution (electrolyte). In other words, how easy is it for iron to give up two electrons and become a dissolved ferrous ion. Every metal reacts differently with an electrolyte, some metals are more reactive than others. The Electrochemical Series is a list of metals and their Nobility (how reactive they are). A more noble metal is considered less reactive and therefore has a more difficult time giving up electrons. A less noble metal is more reactive and gives up electrons more easily.
As stated earlier, galvanic corrosion is induced when two metals with different electrode potentials are connected to each other when they are submersed in an electrolyte.
Here is an example to further explain electrode potential and galvanic corrosion.
Based on the definition of galvanic corrosion, we are going to take two metal plates and submerse them both in a glass filled with water. The water is our electrolyte and the metal plates are separated.
Now remember from the recap in the section on uniform corrosion we stated 4 conditions for corrosion. Those 4 conditions also apply for galvanic corrosion. So we need (1) Anode, (2) Cathode, (3) path for electrons to travel from the anode to the cathode and (4) an electrolyte for the metal to dissolve into.
Lets start with the anode and cathode. As mentioned, galvanic corrosion is when two metals are in contact with each other. In the uniform corrosion example, the iron was the anode and the oxygen acted as the cathode. In this example one metal will act as the anode and the other metal will act as the cathode. So the anode metal will give electrons and the cathode metal will take electrons.
You might be wondering how does one know which metal is the anode and which metal is the cathode. The simple answer is that if you refer to the Electrochemical Series chart and randomly choose two metals, the metal that is less noble (more reactive) will be the anode and the more noble metal will act as the cathode. This is intuative since the more noble metal is more resistant to giving up electrons.
At this point you may be wondering what makes a metal more or less noble, we'll take a quick break from explaining galvanic corrosion to explain.
What makes a metal more or less reactive?
We are randomly going to pick two metals: Iron and Zinc.
Based on the Electrochemical series chart we know that iron is more noble than zinc. When we place the two metal plates in the electrolyte we know that zinc will be the anode (give up electrons) and iron will be the cathode (take electrons). The following illustration explains how and why the electrons transfer and why the zinc found itself to be less noble than iron.
When the iron is placed into the water some of the iron atoms will release electrons and dissolve into the solution. The electrons will stay behind and remain on the iron plate. As more iron atoms dissolve and become positive ions, negatively charged electrons build up on the plate. As this process continues, some positively charged irons ions will be attracted back to the plate, take on electrons and become iron. At a certain point a dynamic equilibrium is established of iron leaving the plate and iron coming back.
The difference between iron and zinc is that zinc is more reactive and more easily gives up electrons. When the zinc plate reaches equilibrium, more of the zinc ions will remain in solution and there will be a greater number of electrons left on the zinc plate.
The essence of electrode potential is assigning a number value to that state of equilibrium. Since we are talking about how many electrons are left on the metal, you may have already guessed that measurement is going to be related to electrical current.
And you are correct to assume that. What we are going to do is examine the difference in potential between the negativity of the metal and the positivity of the water with its positive ions.
This potential difference could be measured as voltage but the problem is that it is impossible to measure the difference between the water and the metal. It is easy to attach a volt meter to the metal (electrode) but attaching it to the water is impossible.
What we need is a reference metal; another metal that we can compare all other metals to. This metal would be O and any other metal would be either positive or negative relative to the reference metal. Just like on earth we choose sea level as a reference level, anything above is positive anything below is negative, or below sea level. Of course the sea level is just an arbitrary but useful reference.
This reference electrode I am talking about is called the STANDARD HYDROGEN ELECTRODE and it has a value of zero at a atmospheric pressure of 1 bar and a temperature of 298 Kelvin (25 Celsius).
To determine the electrode potential of iron, a standard hydrogen electrode is placed in the electrolyte with the iron electrode (note - in reality this experiment is set up a bit differently but that does not matter here).
Both the electrodes will begin to react with the electrolyte and an equilibrium will be established. On the electrochemical series chart it gives iron a value of -0.44 volts. Here's what that means.
With both electrodes reaching equilibrium, the fact that iron has a negative value (-0.44 volts) means that the iron electrode is more negative than the standard hydrogen electrode. This means that after both electrodes have reached equilibrium, the iron electrode will have a higher number of negatively charged electrons than the hydrogen electrode.
But now we can measure the difference by using a volt meter. When the volt meter is attached to the hydrogen electrode and the iron electrode, the voltmeter can measure the Electrode Potential difference between the two electrodes, in other words how negative or positive the one electrode is in reference to the other electrode. The voltmeter will give a number in volts which happens to be -0.44 volts for iron. When copper, on the other hand, is put into the electrolyte, the copper is less reactive than the hydrogen electrode and electrons will flow from the hydrogen electrode to the copper electrode and the voltmeter will give a postive value (+0.33 volts). By convention and to keep negative and positive signs consistent, the negative terminal of the voltmeter is always attached to the reference cell which allows, for example, one to easily determine that a negative reading is an indication that the electrode being measured is relatively negative (relative to the reference electrode) making it the anode.
BACK TO GALVANIC CORROSION
While reading the section on electrode potential you may have already concluded where the corrosion is taking place. Like uniform corrosion, the dissolving of the anode is corrosion. What happens during galvanic corrosion is that when ever two dissimilar metals are connected electrically and through an electrolyte, the more noble metal will always act as the cathode and steal electrons from the less noble metal which becomes the anode. This redox reactions can cause very rapid corrosion in the case where the electrode potential difference is large.
GALVANIC CORROSION PROTECTION
Engineers have used their knowledge of galvanic corrosion to help with corrosion control. If engineers are trying to protect an underground steel pipe (steel is made from iron), they will connect a metal to the pipe that is less noble and that will act as an anode. For example, zinc is a more reactive metal and when attached to steel will donate electrons to the steel and in the process make the steel more noble and less likely to corrode. In the process, of course, the zinc will corrode away. Galvanized metal is a metal with a zinc coating that works on the same principle. It is important to remember that galvanic protection only works when there is an electrolyte such as water or soil.
Ben Heikoop (2015)