The simple voltaic (or galvanic) electrical cell, developed in the early 1800s, is named (depending on which moniker you use) either after Italian physician Luigi Galvani or after Galvani's contemporary and compatriot, Alessandro Volta. Galvani was the first person to stumble across electrochemistry, the phenomenon described below, thanks to some dead frogs. Volta was the first person to understand what Galvani had discovered and put it to use in the voltaic pile, the first battery (batteries are comprised of numerous cells).
What you see below is about as simple as a battery gets, and demonstrates the chemistry behind Volta's device and its successors. The arrangement consists of two different metals partially submerged in an electrolyte (a solution containing ions, particles that carry either a negative or positive charge). Completing the picture is a wire connecting the dry ends of the two metals.
This setup turns chemistry into usable electricity (though not very much of it) courtesy of the processes of oxidation and reduction (together referred to as redox) and of the electrons (negatively charged particles) that obligingly carry the generated current.
To understand this better, let’s take a closer look at the applet below.
You see two strips of metal: zinc on the right and copper on the left. Now, any metal you stick in your electrolyte (in this case sulfuric acid) will react by dissolving into positive ions. The difference is in how quickly they dissolve. Zinc dissolves more quickly than copper. So there are more positive ions (depicted as gray particles) in the solution around the zinc than around the copper, and more electrons (depicted here as yellow particles) left behind inside the zinc than in the copper. As a result, there is a potential (voltage) difference between the two.
Now, without a circuit, that potential difference doesn’t mean too much. But if you create a circuit by connecting these metals with a wire, as seen above, the extra electrons (depicted here as yellow particles) that have built up in the zinc side will be drawn to the copper, which, by comparison, is electron deficient. This flow is what produces electricity (by convention, the electricity flows in the opposite direction of the electrons).
The circuit continues in the solution, of course, where things get a bit more complicated. But in case you’re wondering, what happens is this:
- The sulfuric acid diluted in water breaks down into positive hydrogen ions (2H+, depicted as red particles) and negative sulfate ions (SO-24, illustrated as blue and orange molecules).
- Zinc molecules, after each losing two electrons that get drawn into the wire, become positive ions (Zn2+ – our little round gray friends) that dissolve in the acid by joining with the negative sulfate ions; in other words, the zinc corrodes.
- The zinc ions also work to nudge (repel) the hydrogen ions toward the copper, where they become attached to the incoming electrons and form molecules of hydrogen gas (appearing as light-colored bubbles in our applet) on the surface of the copper. This build-up interferes with the conduction of electricity (this is called polarization) and can halt the flow entirely, if the zinc doesn't corrode away or the solution get used up first.
This migration of negative and positive within the solution is the ionic current that balances out the electric current flowing through the wires. Use the applet speed slider to study the process in slow motion or to speed it up.
Ever hear of a “lemon battery”? It is basically a variation of what is pictured here.
In this setup, the zinc is the anode – the “negative” electrode where oxidation (giving away those electrons) occurs; the copper is the cathode – the “positive” electrode at which reduction (accepting those donated electrons) occurs.
Not long after the invention of this cell (which, as opposed to the "dry" voltaic pile, is a kind of "wet cell"), an improvement known as the Daniell cell was devised.