# Watercooling and Electrochemistry

Tips on avoiding corrosion – Steven Bazzell and Mathew Barchok

We all know that having a copper block and an aluminum radiator is bad. Now, you can add an additive to your water to slow corrosion down to a bare minimum, but going all copper is just the better way to go.

I will attempt to explain this whole process, so that you can have a better understanding of how this works, why this works, and how to prevent it from happening.

I’m not going to fully explain the process of redox reactions and electrochemistry, so you’re going to have to rough it through.

Since we are dealing with Cu (copper) and Al (aluminum), let’s find their Standard Reduction Potentials at 25*C (298K).

Cu(+) + e(-) –> Cu(s) E(red) = +0.52v

Al(3+) + 3e(-) –> Al(s) E(red) = -1.66v

These two equations are for a copper solution and an aluminum solution going to their respective solids.

When the corrosion in your watercooling setup happens, you’re actually getting two reactions taking place.

Since we start off in a watercooling setup with the solid metals of Cu and Al, the equations will now look like this:

Cu(s) –> Cu(+) + e(-) E(red) = -0.52v

Al(s) –> Al(3+) + 3e(-) E(red) = +1.66v

NOTE: electrons move from the (-) side –> the (+) side.

Now we see that the equation for the Al(s) has a very positive standard reduction potential. In other words, it will tend to dissociate into water easier. We also see that copper has a negative standard reduction potential. This means that copper is NOT going to dissociate into water – period.

So, while you’re running your watercooling setup along, Al is just barely, ever so slightly starting to dissociate into the water, giving you those 3 electrons ( 3e(-) ) in solution. By Le Chatelier’s principle, the water (now negative) is going to want to become neutral again.

So What Happens?

Those electrons travel through the water over to the copper block and cause it to corrode itself, using up those electrons to create Cu(s). A slightly small reaction occurs converting solid copper -> solution + electrons = solid copper again. Since the reaction performs forwards, and backwards, the current it creates sums to 0v. In other words, we will ignore this reaction.

You end up with an aluminum radiator that is deteriorating and a copper block that is corroding. The addition of an additive will slow this process down an extreme amount, but it will still happen.

Since the water is warmer than 25*C (most of the time), the process occurs faster. If the water going through the tubes is a lot lower than that, the reaction is slowed down dramatically (ex: water chiller running -20*C water).

This isn’t the actual full net ionic equation, but this is the equation that sums it up and is easier for you to understand:

Al(s) –> 3e(-)

Those 3 electrons cause the equation of the copper to go in the opposite direction, thus creating Cu(s) (the corrosion you see that is normally blue/green).

3e(-) + 3Cu(+) –> 3Cu(s)

If you were to connect voltmeter to this setup, you would actually get a reading.

HOW TO SOLVE THIS?

• Convert everything to copper
• Use a sacrificial anode.

Mathew Barchok contributed the following discussion:

Sacrificial Anodes:

Sacrificial anodes have long been used to protect against corrosion. The theory is simple: Some metals are more easily oxidized than others. In the case of metals, oxidation results in the removal of electrons from the metal to form a positive ion.

In the case where two different metals are connected through a joint or a wire, electrons can flow freely between the two. If an atom of the less easily oxidized metal is oxidized, electrons will flow from the more easily oxidized side to that atom and reduce it back to the neutral state.

As a result, an atom on the more easily oxidized side will become oxidized. For this to work, it must be in an environment where the oxidized can flow to the other end. Otherwise, opposite charges will build up on the metals and stop the electrons from flowing. Water and soil do this effectively.

Good examples of the use of sacrificial anodes are the protection of iron pipes and bridges. Chunks of zinc are connected to the pipe or bridge with a cable. The zinc is destroyed in a couple years due to corrosion, but provided it is replaced, the bridge or pipe can last decandes. With a little modification, a current meter can be placed in the cable to monitor the system.

Zinc is the most commonly used sacrifical anode because it has many desirable properties, including the fact that it is more readily oxidized than most metals. Unfortunately, it will not work with aluminum. Why? Because aluminum is more readily oxidized. That means that the aluminum will act as a sacrificial anode to the zinc!

To protect aluminum using this method, there is only one metal you can use. That’s it – One. Magnesium.

Sure, there are other metals that are more easily oxidized than magnesium. Unfortunately, these other metals have a tendency to react violently with water. Even magnesium itself is not that great for many reasons, so it must be used as an alloy.

For a good disussion of this subject,