We already know that corrosion involves the flow of electrons. Areas that lose electrons corrode. Areas that the electrons flow to do not corode. So from this we can deduce the direction of electron flow in each of the metal pairs that we saw in experiment 3 (using the evidence of corrosion).

When the nail is attached to aluminium, the nail does not rust, but there are signs of corrosion on the aluminium. The electrons must be flowing from the aluminium (corroding) and into the nail (not corroding).

When the nail is attached to silver solder, the nail rusts, but there are no signs of corrosion on the solder. The electrons must be flowing from the nail (corroding) and into the solder (not corroding).

When the nail is attached to copper, the nail rusts, but there are no signs of corrosion on the copper. The electrons must be flowing from the nail (corroding) and into the copper (not corroding).

When the solder is attached to copper, the copper corrodes, but there are no signs of corrosion on the solder. The electrons must be flowing from the copper (corroding) and into the solder (not corroding).

We know that metals want to oxidise and get back to their natural state, and the above results show that some metals are "stronger" than others. In the first example, both the iron nail and the aluminium want to throw off ions into the water (corrode) and get rid of electrons. The aluminium is stronger - and corrodes - losing atoms and forcing the iron nail to accept its spare electrons. Because the electron current is flowing into the iron nail - it does not corrode. The iron nail is the protected cathode (electrons flow to the cathode), the aluminium is the corroding anode (electrons flow away from the anode).
In the second example, the flow is reversed. This time the iron nail wins the battle for corrosion, forcing the solder to accept its spare electrons. The electron current is flowing into the solder - so it does not corrode. The iron nail is the corroding anode (electrons flow away from the anode), the solder is the protected cathode (electrons flow to the cathode).
Following the same logic, in the third example, iron wins the battle against copper, and the electrons must be flowing from the corroding iron nail into the copper. In the last example, the corroding copper wins, and the electrons must flow from the copper into the solder.

So we can rank these metals in order of reactivity: Aluminium beats iron, iron beats solder, iron beats copper, and copper beats solder:

1) Aluminium
2) Iron
3) Copper
4) Solder

So we can confidently predict that if aluminium and copper were joined together in the solution, the aluminium would corrode, and the copper would be protected. This is another form of corrosion protection. By examining the metal that we want to protect, we can attach a piece of a stronger (more reactive) metal to it, and the weaker metal is protected (will not corrode), but the stronger metal will become the anode and will corrode instead.

For example, if we had an iron cage with steel mesh that we wanted to put in the sea to catch lobsters, we would soon find that it had rusted away. In order to prevent this, we could attach a piece of aluminium to the iron cage and mesh. Just like the aluminium foil attached to the nail, the cage would not rust, but the aluminium would corrode, pushing electrons into the iron cage and mesh and preventing it from corroding.

In reality, a metal called zinc would be used for this purpose (as it is even stronger than Aluminium). Sometimes a metal called magnesium is used (as this is stronger still). These pieces of metal attached to structures to prevent rusting are called sacrificial anodes - because they sacrifice themselves and corrode away to protect the structure that they are attached to. The metals used must be carefully selected - the anode must always be a stronger metal than the thing it is protecting, otherwise the electrons would flow in the opposite direction and the structure that you are trying to protect would corrode even faster! When the sacrificial anode has corroded away completely, it is simply replaced with a new one.

Initial conclusion: some metals are more reactive than others. When they are joined together, the stronger metal will corrode, and force the weaker metal to accept the spare electrons produced. The weaker metal will be prevented from corroding. This effect can be used to protect structures from corrosion through the use of sacrifical anodes.