Galvanic cell

A galvanic cell, also known as a voltaic cell, is a device that generates an electric current. It does this by using spontaneous oxidation-reduction reactions, often called redox reactions. These reactions release chemical energy that the cell converts into electrical energy.

While these reactions happen naturally, a galvanic cell is specifically designed to harness that energy for useful work. Without this design, the energy released by the chemicals would simply escape into the surroundings as heat.

To understand how it works, we must look at the two separate parts called half-cells. Each half-cell contains a solid metal, known as an electrode, submerged in a liquid called an electrolyte. The electrolyte contains ions, which are charged particles that help carry electricity through the liquid. In a complete cell, these two half-cells are connected by an external conductor, like a metal wire, and a salt bridge or a porous membrane. The salt bridge is vital because it allows ions to move between the solutions. This movement balances the electrical charges in the liquids so the reaction can continue.

The mechanism relies on the transfer of electrons between the two metals. In one half-cell, a process called oxidation occurs at the anode. During oxidation, metal atoms lose electrons and become positive ions, which then dissolve into the electrolyte. These released electrons travel through the external wire toward the other half-cell. In the second half-cell, a process called reduction occurs at the cathode. Here, positive metal ions in the solution gain electrons from the electrode. This causes them to turn into solid metal and stick to the surface, a process known as plating.

There are two specific roles for the electrodes based on these chemical changes. The anode is the negative electrode where oxidation takes place. Because electrons are being lost there, they leave the electrode and flow away. The cathode is the positive electrode where reduction happens. It attracts positive ions, or cations, from the solution to receive electrons. This distinction is important because the flow of electrons through the external circuit is what provides the electrical power we use in technology.

The history of this discovery is a series of scientific debates. In 1780, Luigi Galvani observed a frog's leg contract when touched by two different metals. He believed this was caused by "animal electricity" stored in the biological tissue. However, in 1790, Alessandro Volta demonstrated that the frog was not the source of the electricity. Volta used brine-soaked paper to show that metals and liquids alone could create a current. In 1799, he invented the voltaic pile, which was a stack of many individual galvanic cells. This was the world's first electrical battery.

It took several more decades to fully understand the science behind these devices. About forty years after Volta's work, Michael Faraday showed that the process was truly chemical in nature. Faraday introduced the modern terminology we use today, including the terms anode, cathode, electrolyte, and ion. He correctly identified that the electromotive force, or emf, comes from the chemical reactions at the interfaces where the metal meets the liquid. This was a major step forward from earlier theories that focused only on the physical properties of the metals.

We see the application of galvanic cells in many modern systems. A single galvanic cell is often just one part of a larger battery. For example, a standard 12 V lead-acid battery is actually made of six galvanic cells connected in a series. These cells use lead and lead dioxide in a sulfuric acid electrolyte. In large telephone exchanges, many cells are connected in both series and parallel to provide enough power for many users at once. Even the Weston cell, which was used as an international standard for voltage in 1911, relies on these precise chemical principles.