A galvanic cell is an electrochemical cell that stores chemical energy and makes it available in an electrical form, and a battery is a string of two or more cells in series. Other types of electrochemical cells include electrolytic cells, fuel cells, flow cells, or voltaic cells.[1]
Though an early form of battery may have been used in antiquity (the Baghdad Battery), the development of modern batteries started with the Voltaic pile, invented by the Italian physicist Alessandro Volta in 1800.[2] According to a 2005 estimate, the worldwide battery industry generates US$48 billion in sales annually.[3]
Formally, an electrical "battery" is a series-connected array of similar voltaic cells ("cells"). However, in many contexts it is common to call a single cell a battery.[4]
The two types of batteries, primary and secondary, both convert chemical energy to electrical energy. However, primary batteries can only be used once, as they use up their chemicals in an irreversible reaction. Secondary batteries can be recharged because the chemical reactions they use are reversible; they are recharged by running a current parallel to the battery, with an orientation opposite to the original discharge.[5]


The earliest known artifacts that may have served as batteries are the Baghdad Batteries, which existed some time between 250 BC and 640 AD. However, it is not known what electrical function they may have served, and if they were in fact batteries at all. Scientists have developed several theories about its use, including medicine (as a painkiller) and electroplating jewelry.[6]
The story of the modern battery begins in the 1780s with the discovery of "animal electricity" by Luigi Galvani, which he published in 1791.[7] He created an electric circuit consisting of two different metals, with one touching a frog's leg and the other touching both the leg and the first metal, thus closing the circuit. In modern terms, the frog's leg served as both electrolyte and detector, and the metals served as electrodes. He noticed that even though the frog was dead, its legs would twitch when he touched them with the metals.[8]
By 1791, Alessandro Volta realized that the frog could be replaced by cardboard soaked in salt water, employing another form of detection. Having already studied the electrostatic phenomenon of capacitance, Volta was able to quantitatively measure the electromotive force (emf) associated with each electrode-electrolyte interface (voltage) in volts, which were named after him. Such a device is called a voltaic cell, or cell for short. In 1799, Volta invented the modern battery by placing many galvanic cells in series, literally piling them one above the other. This Voltaic Pile gave a greatly enhanced net emf for the combination,[9] with a voltage of about 50 volts for a 32-cell pile.[10] In many parts of Europe batteries continue to be called piles. Unfortunately, Volta did not appreciate that the voltage was due to chemical reactions. He thought that his cells were an inexhaustible source of energy, and that the associated chemical effects (e.g., corrosion) were a mere nuisance -- rather than, as Michael Faraday showed around 1830, an unavoidable by-product of their operation.
Later, researchers placed galvanic cells in series. Such banks of cells are called batteries, presumably after the earlier use by Benjamin Franklin to describe Leyden jars (capacitors) in series and in parallel.[11]
Although early batteries were of great value for experimental purposes, their limitations made them impractical for large current drain. Later, batteries, starting with the Daniell cell in 1836, provided more reliable currents and were adopted by industry for use in stationary devices, particularly in telegraph networks where they were the only practical source of electricity, since electrical distribution networks did not exist then.[12] These wet cells used liquid electrolytes, which were prone to leaks and spillage if not handled correctly. Many used glass jars to hold their components, which made them fragile. These characteristics made wet cells unsuitable for portable appliances. Near the end of the 19th century, the invention of dry cell batteries, which replaced liquid electrolyte with a paste made portable electrical devices practical.

[edit] How batteries work
Main article: Electrochemical cell

One example of a galvanic cell, element of a battery
A battery is a device that converts chemical energy directly to electrical energy.[13] It consists of one or more voltaic cells. Each voltaic cell consists of two half cells connected in series by a conductive electrolyte. Each cell has a positive electrode (cathode), and a negative electrode (anode). These do not touch each other but are immersed in a solid or liquid electrolyte.[14] In a practical cell the materials are enclosed in a container, and a separator between the electrodes prevents the electrodes from coming into contact.
Each half cell has a net electromotive force (or emf), with the net emf of the battery being the difference between the emfs of the half-cells, a fact first recognized by Volta. Thus, if the electrodes have emfs , then the net emf is . (Hence, two identical electrodes and a common electrolyte give zero net emf.) Each half cell emf is due to a charge-transferring (or faradaic) chemical reaction at the electrode-electrolyte interface, which transfers charge across the interface. The reaction stops when the charge transfer is enough to cancel out the tendency of the reaction to occur. Non-charge-transferring, or nonfaradaic, reactions can also occur at the interface. These are undesirable, using up the chemicals without producing current (which is the rate of charge transfer). Additional, but relatively ineffective, faradaic reactions (also called parasitic or "side-reactions") can also occur.
The electrical potential difference across the terminals of a battery is known as its terminal voltage, measured in volts. The terminal voltage of a battery that is neither charging nor discharging is called the open-circuit voltage, and gives the emf of the battery. The terminal voltage of a battery that is discharging is smaller in magnitude than the open-circuit voltage, and the terminal voltage of a battery being charged is greater than the open-circuit voltage. [15]
The voltage developed across a cell's terminals depends on the chemicals used in it and their concentrations. For example, alkaline and carbon-zinc cells both measure about 1.5 volts, due to the energy release of the associated chemical reactions. Because of the high electrochemical potential changes in the reactions of lithium compounds, lithium cells can provide as much as 3 volts or more.