Galvanic Cells

A galvanic cell converts a chemical reaction into electricity. These cells are self-contained and portable, so they are used as batteries and fuel cells.

Galvanic cells were first described in 1790 by the Italian scientist Luigi Galvani. In Galvani's experiments, a frog was dissected to expose the nerves in the lower half of a frog. A copper wire was attached to the exposed nerve and a zinc wire was attached to the leg muscle. When the two types of metal were touched together, the muscle would contract, resulting in the half-frog performing a macabre dance. An illustration of Galvani's experiment. [1]

Galvani concluded incorrectly that the electricity was coming from the tissues of the animal. Another Italian scientist, Alessandro Volta, would later establish that the source of the electricity was electron flow between the two metals, and that the frog's body fluids were just a conducting medium. (For this reason, galvanic cells are often referred to as voltaic cells as well). However, Galvani did uncover the electric nature of nerve impulses. His frog experiment is considered the beginning of electrophysiology and the understanding of how bodies convert information between chemical and electrical messaging systems [2].

There are two types of electrochemical cells: galvanic cells, which generate electricity, and electrolytic cells, which reverse the process and use an external electric current to power a chemical reaction.

Consider the redox reaction Galvani was facilitating with his frog legs:

$$Zn+Cu^{+2} \to Zn^{2+}+Cu$$

This is a spontaneous redox reaction with a negative Gibbs free energy ($\Delta G < 0$). It can be split into two half-reactions:

$$Zn \to Zn^{2+}+2e^-$$

$$Cu^{2+}+2e^- \to Cu$$

Each electrochemical cell is made of two electrodes. The cathode is positively charged and is the place where the reduction half-reaction occurs. The oxidation half-reaction occurs and the anode, the negatively charged electrode.

Zinc loses 2 electrons to become the ion $Zn^{2+}$ while $Cu^{+2}$ accepts the electrons to become elemental copper. If these two half-reactions occur in two separate containers connected by a conductive wire, an electric current will form from the transfer of the electrons from one container to the other. Zinc and copper electrodes are submerged in aqueous solutions of their salts, such as $ZnSO_4$ and $CuSO_4$, in their respective containers. These two electrodes are connected by wire and the solutions are connected by a medium or bridge which allows ions to transfer but does not allow the solutions to mix directly.

The salt bridge serves two purposes: it carries electrical charge to complete the circuit, and it ensures that both solutions remain electrically neutral. The identity of the salt in the bridge does not matter, as long as it does not interfere with the redox reaction by being oxidized or reduced or by forming a precipitate that would remove the ions involved in the cell from the solution.

What would happen if there was no salt bridge in the galvanic cell illustrated above?

The redox reaction would start the same way, but would soon stop because the solutions would not maintain their electric neutrality.

Primary cells are disposable. The reaction in the electrode is irreversible.

The dry cell is the most common type of battery used to power small household devices, such as flashlights, radios, and calculators. Despite its name, these cells are built in a water-based paste containing $MnO_{2}$ and $Zn$. The chemical reactions used in a dry cell can be modified to work in acidic or alkaline solutions. Alkaline batteries are more often available commercially.

Mercury cells are generally smaller in size and are built around a different chemical reaction than dry cell batteries. These batteries are useful in cameras, hearing aids, and similar devices that require batteries that are both small and reliable. Mercury cell batteries are often more expensive than alkaline batteries, and since they contain a heavy metal, they may pose environmental hazards if broken open or with improper disposal.

The overall reaction for a mercury cell is as follows: $Zn(s) + HgO(s) \to Hg(l) + ZnO (s)$

What is the cathode half-reaction for this cell?

Hint: consider adding liquid water, electrons, protons, or hydroxide to balance your redox half-reactions.

The cathode involves mercury ions gaining electrons to form elemental mercury.

$HgO(s) +H_{2}O(l) +2e^{-} \to Hg(l) + 2OH^{-}(aq)$

Secondary cells are rechargeable. The reaction in the electrode can be reversed by applying an electrical potential to the cell. The recharging process temporarily converts a galvanic cell to an electrolytic cell. Examples include nickel-cadmium batteries found in rechargeable power tools and the lead storage batteries found in cars.

Corrosion is a galvanic process. When metals are exposed to air or water, many of them react with oxygen to form a metal oxide. Rust, for example, is an iron oxide $Fe_2O_3$. The tarnish of silver or the blue-green color that copper develops are other examples of corrosion. Rust in particular can be a major industrial problem, because iron is so prevalent in manufacturing and because the reddish product flakes off, exposing even more metal to be oxidized. Metal products ranging from automobiles to faucets to soda cans are often coated with a thin layer of plastic, metal oxide, or a nonreactive metal (such as chromium) in order to prevent corrosion. The anodic index between the bolts and the plate allows for galvanic corrosion. [3] Dissimilar metals can also corrode each other. This process, sometimes called galvanic corrosion occurs when two metals are in contact and have different electrical properties. The anodic index can be used to determine the voltage potential between two metals. Metals with a large difference in anodic index should be kept insulated from each other to discourage corrosion.

Just as the energy from a spontaneous chemical reaction can be harnessed to create electricity, electricity can be used to drive a non-spontaneous chemical reaction. Illustration of copper electroplating. Applying electricity allows the copper ions in solution to attach to the metal, forming a thin outer layer. [4]

Electrolytic cells have many commercial functions. One example is electroplating, a process where a thin layer of one metal is deposited on top of another metal. This process is used in jewelry production and other industries where manufacturers want the qualities of a precious or non-tarnishing metal as the face of their products, but also want to sell their products at a price point lower than solid gold or platinum would require.

[1] Wells, D. A. (1857). The science of common things: A familiar explanation of the first principles of physical science. For schools, families, and young students. New York: Ivison, Blakeman, Taylor. Accessed from https://archive.org/details/sciencecommonth00wellgoog February 15, 2016.

[2] Haas, L.F. J Neurol Neurosurg Psychiatry. 1993 Oct; 56(10): 1084. Accessed from http://jnnp.bmj.com/content/56/10/1084.long February 15, 2016.

[3] Image from https://commons.wikimedia.org/wiki/File:Stainless-steel-mild-steel.jpg under Creative Commons licensing for reuse and modification.

[4] Image from https://commons.wikimedia.org/wiki/File:Copper_electroplating.svg under Creative Commons licensing for reuse and modification.