Galvanic corrosion occurs of an active metal (i.e. iron or zinc) in contact with a passive metal (i.e. copper) in a conductive solution. A famous example is the corrosion of iron girders in contact with copper in the statue of liberty.
Although the difference in potential can account for the protection of copper, it does not explain the girders in had a greatly accelerated corrosion rate. In a zinc/copper battery there is a solution of copper sulfate that gets reduced at the copper electrode. However, if all the copper is in metallic form in the first place, no galvanic reaction can occur.
It seems galvanic corrosion in real-life, does not increase total corrosion, but rather concentrates it (at lest for pairs of metals with similar non-galvanic corrosion rates). Copper sheets have a large surface area that initially reacted non-galvanically to form the green color we see. Some of these ions are then reduced at the expense of the oxidizing the girders. The statue of liberty and other infamous examples of corrosion involve large amounts of inactive metals connected to small amounts of active (iron/steel) girders/fasteners/etc.
This logic predicts that the attaching a copper bar to an iron bar of equal surface area will only double it's corrosion rate. That being said, predicting the corrosion rate is not easy so there could be some other effect I am overlooking. Is the surface area logic correct or would there a large increase in the total corrosion?