Voltaic pile

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Schematic diagram of a copper-zinc voltaic pile. Each copper-zinc pair had a spacer in the middle, made of cardboard or felt soaked in salt water (the electrolyte). Volta's original piles contained an additional zinc disk at the bottom, and an additional copper disk at the top; these were later shown to be unnecessary. Voltaic pile.svg
Schematic diagram of a copper–zinc voltaic pile. Each copper–zinc pair had a spacer in the middle, made of cardboard or felt soaked in salt water (the electrolyte). Volta's original piles contained an additional zinc disk at the bottom, and an additional copper disk at the top; these were later shown to be unnecessary.
A voltaic pile on display in the Tempio Voltiano (the Volta Temple) near Volta's home in Como, Italy VoltaBattery.JPG
A voltaic pile on display in the Tempio Voltiano (the Volta Temple) near Volta's home in Como, Italy
Voltaic pile, University History Museum of the University of Pavia. Pila di volta.jpg
Voltaic pile, University History Museum of the University of Pavia.

The voltaic pile was the first electrical battery that could continuously provide an electric current to a circuit. [1] It was invented by Italian chemist Alessandro Volta, who published his experiments in 1799. [2] Its invention can be traced back to an argument between Volta and Luigi Galvani, Volta's fellow Italian scientist who had conducted experiments on frogs' legs. [3] Use of the voltaic pile enabled a rapid series of other discoveries, including the electrical decomposition (electrolysis) of water into oxygen and hydrogen by William Nicholson and Anthony Carlisle (1800), and the discovery or isolation of the chemical elements sodium (1807), potassium (1807), calcium (1808), boron (1808), barium (1808), strontium (1808), and magnesium (1808) by Humphry Davy. [4] [5]

Contents

The entire 19th-century electrical industry was powered by batteries related to Volta's (e.g. the Daniell cell and Grove cell) until the advent of the dynamo (the electrical generator) in the 1870s. [6]

History

Volta's invention was built on Luigi Galvani's 1780s discovery that a circuit of two metals and a frog's leg can cause the frog's leg to respond. [1] Volta demonstrated in 1794 that when two metals and brine-soaked cloth or cardboard are arranged in a circuit they too produce an electric current. In 1800, Volta stacked several pairs of alternating copper (or silver) and zinc discs (electrodes) separated by cloth or cardboard soaked in brine, which increased the total electromotive force. [7] [8] When the top and bottom contacts were connected by a wire, an electric current flowed through the voltaic pile and the connecting wire. This was the first "true" battery, that gave off continuous charge. [9]

Many scientific instruments that belonged to Alessandro Volta are preserved in the University History Museum of the University of Pavia, where Volta taught from 1778 to 1819; the piles on display, unfortunately, are not original, as the ones preserved in Pavia were lent on the occasion of the centenary of the invention and subsequently lost in a fire. [10]

Applications

Drawing of the voltaic pile in different configurations, from the letter sent from Alessandro Volta to Joseph Banks. Volta batteries.jpg
Drawing of the voltaic pile in different configurations, from the letter sent from Alessandro Volta to Joseph Banks.

On 20 March 1800, Alessandro Volta wrote to the London Royal Society to describe the technique for producing electric current using his device. [11] On learning of the voltaic pile, William Nicholson and Anthony Carlisle used it to discover the electrolysis of water. Humphry Davy showed that the electromotive force, which drives the electric current through a circuit containing a single voltaic cell, was caused by a chemical reaction, not by the voltage difference between the two metals. He also used the voltaic pile to decompose chemicals and to produce new chemicals. William Hyde Wollaston showed that electricity from voltaic piles had identical effects to those of electricity produced by friction. In 1802 Vasily Petrov used voltaic piles in the discovery and research of electric arc effects.

Humphry Davy and Andrew Crosse were among the first to develop large voltaic piles. [12] Davy used a 2000-pair pile made for the Royal Institution in 1808 to demonstrate carbon arc discharge [13] and isolate five new elements: barium, calcium, boron, strontium and magnesium. [14]

Electrochemistry

Because Volta believed that the electromotive force occurred at the contact between the two metals, Volta's piles had a different design than the modern design illustrated on this page. His piles had one extra disc of copper at the top, in contact with the zinc, and one extra disc of zinc at the bottom, in contact with the copper. [15] Expanding on Volta's work and the electro-magnetism work of his mentor Humphry Davy, Michael Faraday utilized both magnets and the voltaic pile in his experiments with electricity. Faraday believed that all "electricities" being studied at the time (voltaic, magnetic, thermal, and animal) were one and the same. His work to prove this theory led him to propose two laws of electrochemistry which stood in direct conflict with the current scientific beliefs of the day as laid down by Volta thirty years earlier. [16] Because of their contributions to the understanding of this field of study, Faraday and Volta are both considered to be among the fathers of electrochemistry. [17] The words "electrode" and "electrolyte", used above to describe Volta's work, are due to Faraday. [18]

Electromotive force

The strength of the pile is expressed in terms of its electromotive force, or emf, given in volts. Alessandro Volta's theory of contact tension considered that the emf, which drives the electric current through a circuit containing a voltaic cell, occurs at the contact between the two metals. Volta did not consider the electrolyte, which was typically brine in his experiments, to be significant. However, chemists soon realized that water in the electrolyte was involved in the pile's chemical reactions, and led to the evolution of hydrogen gas from the copper or silver electrode. [4] [19] [20] [21]

The modern, atomistic understanding of a cell with zinc and copper electrodes separated by an electrolyte is the following. When the cell is providing an electrical current through an external circuit, the metallic zinc at the surface of the zinc anode is oxidized and dissolves into the electrolyte as electrically charged ions (Zn2+), leaving two negatively charged electrons (
e−
) behind in the metal:

anode (oxidation): Zn → Zn2+ + 2
e−

This reaction is called oxidation. While zinc is entering the electrolyte, two positively charged hydrogen ions (H+) from the electrolyte accept two electrons at the copper cathode surface, become reduced and form an uncharged hydrogen molecule (H2):

cathode (reduction): 2 H+ + 2
e−
 → H2

This reaction is called reduction. The electrons used from the copper to form the molecules of hydrogen are made up by an external wire or circuit that connects it to the zinc. The hydrogen molecules formed on the surface of the copper by the reduction reaction ultimately bubble away as hydrogen gas.

One will observe that the global electro-chemical reaction does not immediately involve the electrochemical couple Cu2+/Cu (Ox/Red) corresponding to the copper cathode. The copper metal disk thus only serves here as a "chemically inert" noble metallic conductor for the transport of electrons in the circuit and does not chemically participate in the reaction in the aqueous phase. Copper does act as a catalyst for the hydrogen-evolution reaction, which otherwise could occur equally well directly at the zinc electrode without current flow through the external circuit. The copper electrode could be replaced in the system by any sufficiently noble/inert and catalytically active metallic conductor (Ag, Pt, stainless steel, graphite, ...). The global reaction can be written as follows:

Zn + 2H+ → Zn2+ + H2

This is usefully stylized by means of the electro-chemical chain notation:

(anode: oxidation) Zn | Zn2+ || 2H+ | H2 | Cu (cathode: reduction)

in which a vertical bar each time represents an interface. The double vertical bar represents the interfaces corresponding to the electrolyte impregnating the porous cardboard disk.

When no current is drawn from the pile, each cell, consisting of zinc/electrolyte/copper, generates 0.76 V with a brine electrolyte. The voltages from the cells in the pile add, so the six cells in the diagram above generate 4.56 V of electromotive force.

Dry piles

A number of high-voltage dry piles were invented between 1800 and the 1830s in an attempt to determine the source of electricity of the wet voltaic pile, and specifically to support Volta's hypothesis of contact tension. Indeed, Volta himself experimented with a pile whose cardboard discs had dried out, most likely accidentally.

The first to publish the discovery of a dry pile that produced a current was Johann Wilhelm Ritter in 1802, albeit in an obscure journal; over the next decade, it was announced repeatedly as a new discovery. One form of dry pile is the Zamboni pile. Francis Ronalds in 1814 was one of the first to realize that dry piles also worked through chemical reaction rather than metal-to-metal contact, even though corrosion was not visible due to the very small currents generated. [22] [23]

The dry pile could be referred to as the ancestor of the modern dry cell.[ original research? ]

See also

Related Research Articles

<span class="mw-page-title-main">Alessandro Volta</span> Italian physicist and chemist (1745–1827)

Alessandro Giuseppe Antonio Anastasio Volta was an Italian physicist and chemist who was a pioneer of electricity and power, and is credited as the inventor of the electric battery and the discoverer of methane. He invented the voltaic pile in 1799, and reported the results of his experiments in a two-part letter to the president of the Royal Society, which was published in 1800. With this invention, Volta proved that electricity could be generated chemically and debunked the prevalent theory that electricity was generated solely by living beings. Volta's invention sparked a great amount of scientific excitement and led others to conduct similar experiments, which eventually led to the development of the field of electrochemistry.

<span class="mw-page-title-main">Anode</span> Electrode through which conventional current flows into a polarized electrical device

An anode usually is an electrode of a polarized electrical device through which conventional current enters the device. This contrasts with a cathode, which is usually an electrode of the device through which conventional current leaves the device. A common mnemonic is ACID, for "anode current into device". The direction of conventional current in a circuit is opposite to the direction of electron flow, so electrons flow from the anode of a galvanic cell, into an outside or external circuit connected to the cell. For example, the end of a household battery marked with a "+" is the cathode.

<span class="mw-page-title-main">Electrochemistry</span> Branch of chemistry

Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.

<span class="mw-page-title-main">Electrochemical cell</span> Electro-chemical device

An electrochemical cell is a device that generates electrical energy from chemical reactions. Electrical energy can also be applied to these cells to cause chemical reactions to occur. Electrochemical cells that generate an electric current are called voltaic or galvanic cells and those that generate chemical reactions, via electrolysis for example, are called electrolytic cells.

<span class="mw-page-title-main">Electrolysis</span> Technique in chemistry and manufacturing

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<span class="mw-page-title-main">Electromotive force</span> Electrical action produced by a non-electrical source

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<span class="mw-page-title-main">Lemon battery</span> Simple battery made with a lemon for educational purposes

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<span class="mw-page-title-main">Galvanic cell</span> Electrochemical device

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<span class="mw-page-title-main">Electrolytic cell</span> Cell that uses electrical energy to drive a non-spontaneous redox reaction

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<span class="mw-page-title-main">Standard electrode potential</span> Electromotive force of a half reaction cell versus standard hydrogen electrode

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In electrochemistry, a half-cell is a structure that contains a conductive electrode and a surrounding conductive electrolyte separated by a naturally occurring Helmholtz double layer. Chemical reactions within this layer momentarily pump electric charges between the electrode and the electrolyte, resulting in a potential difference between the electrode and the electrolyte. The typical anode reaction involves a metal atom in the electrode being dissolved and transported as a positive ion across the double layer, causing the electrolyte to acquire a net positive charge while the electrode acquires a net negative charge. The growing potential difference creates an intense electric field within the double layer, and the potential rises in value until the field halts the net charge-pumping reactions. This self-limiting action occurs almost instantly in an isolated half-cell; in applications two dissimilar half-cells are appropriately connected to constitute a Galvanic cell.

<span class="mw-page-title-main">Daniell cell</span> Type of electrochemical cell

The Daniell cell is a type of electrochemical cell invented in 1836 by John Frederic Daniell, a British chemist and meteorologist, and consists of a copper pot filled with a copper (II) sulfate solution, in which is immersed an unglazed earthenware container filled with sulfuric acid and a zinc electrode. He was searching for a way to eliminate the hydrogen bubble problem found in the voltaic pile, and his solution was to use a second electrolyte to consume the hydrogen produced by the first. Zinc sulfate may be substituted for the sulfuric acid. The Daniell cell was a great improvement over the existing technology used in the early days of battery development. A later variant of the Daniell cell called the gravity cell or crowfoot cell was invented in the 1860s by a Frenchman named Callaud and became a popular choice for electrical telegraphy.

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<span class="mw-page-title-main">Voltameter</span> Instrument for measuring electric charge

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<span class="mw-page-title-main">Leclanché cell</span> Battery (cell) with an anode of zinc and a cathode of manganese dioxide

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<span class="mw-page-title-main">Penny battery</span> Voltaic pile

The penny battery is a voltaic pile which uses various coinage as the metal disks (pennies) of a traditional voltaic pile. The coins are stacked with pieces of electrolyte soaked paper in between. The penny battery experiment is common during electrochemistry units in an educational setting.

<span class="mw-page-title-main">Galvanic corrosion</span> Electrochemical process

Galvanic corrosion is an electrochemical process in which one metal corrodes preferentially when it is in electrical contact with another, in the presence of an electrolyte. A similar galvanic reaction is exploited in primary cells to generate a useful electrical voltage to power portable devices. This phenomenon is named after Italian physician Luigi Galvani (1737–1798).

This article provides information on the following six methods of producing electric power.

  1. Friction: Energy produced by rubbing two material together.
  2. Heat: Energy produced by heating the junction where two unlike metals are joined.
  3. Light: Energy produced by light being absorbed by photoelectric cells, or solar power.
  4. Chemical: Energy produced by chemical reaction in a voltaic cell, such as an electric battery.
  5. Pressure: Energy produced by compressing or decompressing specific crystals.
  6. Magnetism: Energy produced in a conductor that cuts or is cut by magnetic lines of force.

References

  1. 1 2 "Battery: Voltaic Pile". americanhistory.si.edu. Retrieved 2024-05-12.
  2. ↑ "Alessandro Volta | Biography, Facts, Battery, & Invention | Britannica". www.britannica.com. 2024-04-15. Retrieved 2024-05-12.
  3. ↑ "The Voltaic Pile | Distinctive Collections Spotlights". libraries.mit.edu. Retrieved 2023-01-24.
  4. 1 2 Decker, Franco (January 2005). "Volta and the 'Pile'". Electrochemistry Encyclopedia. Case Western Reserve University. Archived from the original on 2012-07-16.
  5. ↑ Russell, Colin (August 2003). "Enterprise and electrolysis..." Chemistry World.
  6. ↑ "Alessandro Volta | Biography, Facts, Battery, & Invention | Britannica". www.britannica.com. 2024-04-15. Retrieved 2024-05-12.
  7. ↑ Brockman, C. J. (June 1927). "Primary cells—A brief historical sketch". Journal of Chemical Education . 4 (6): 770. doi:10.1021/ed004p770. ISSN   0021-9584.
  8. ↑ Mottelay, Paul Fleury (2008). Bibliographical History of Electricity and Magnetism (Reprint of 1892 ed.). Read Books. p. 247. ISBN   978-1-4437-2844-7.
  9. ↑ "The Voltaic Pile | Distinctive Collections Spotlights". libraries.mit.edu. Retrieved 2023-01-24.
  10. ↑ "Sala Volta". Musei Unipv. Retrieved 21 August 2022.
  11. ↑ Volta, Alessandro (1800). "On the Electricity Excited by the Mere Contact of Conducting Substances of Different Kinds". Philosophical Transactions of the Royal Society of London (in French). 90: 403–431. doi: 10.1098/rstl.1800.0018 . A partial translation of this paper is available online; see "Volta and the Battery" . Retrieved 2012-12-01. A complete translation was published in Dibner, Bern (1964). Alessandro Volta and the Electric Battery. Franklin Watts. pp. 111–131. OCLC   247967.
  12. ↑ Encyclopædia Britannica, 1911 edition, Volume V09, Page 185
  13. ↑ Tracking Down the Origin of Arc Plasma Science. II. Early Continuous Discharges
  14. ↑ Kenyon, T. K. (2008). "Science and Celebrity: Humphry Davy's Rising Star". Chemical Heritage Magazine. 26 (4): 30–35. Retrieved 22 March 2018.
  15. ↑ Cecchini, R.; Pelosi, G. (April 1992). "Alessandro Volta and his battery". IEEE Antennas and Propagation Magazine. 34 (2): 30–37. Bibcode:1992IAPM...34...30C. doi:10.1109/74.134307. S2CID   6515671.
  16. ↑ James, Frank A. J. L. (1989). "Michael Faraday's first law of electrochemistry: how context develops new knowledge". In Stock, J. T.; Orna, M. V. (eds.). Electrochemistry, past and present. Washington, DC: American Chemical Society. pp. 32–49. ISBN   9780841215726.
  17. ↑ Stock, John T. (1989). "Electrochemistry in retrospect: an overview". In Orna, Mary Virginia (ed.). Electrochemistry, past and present. Washington, DC: American Chemical Society. pp. 1–17. ISBN   9780841215726.
  18. ↑ James, F.A.J.L. (18 July 2013). "The Royal Institution of Great Britain: 200 years of scientific discovery and communication". Interdisciplinary Science Reviews. 24 (3): 225–231. doi:10.1179/030801899678777.
  19. ↑ Turner, Edward (1841). Liebig, Justus; Gregory, William (eds.). Elements of chemistry: including the actual state and prevalent doctrines of the science (7 ed.). London: Taylor and Walton. p. 102. During the action of a simple circle, as of zinc and copper, excited by dilute sulfuric acid, all of the hydrogen developed in the voltaic action is evolved at the surface of the copper.
  20. ↑ Goodisman, Jerry (2001). "Observations on Lemon Cells". Journal of Chemical Education. 78 (4): 516. Bibcode:2001JChEd..78..516G. doi:10.1021/ed078p516. Goodisman notes that many chemistry textbooks use an incorrect model for a cell with zinc and copper electrodes in an acidic electrolyte.
  21. ↑ Graham-Cumming, John (2009). "Tempio Voltiano". The Geek Atlas: 128 Places Where Science and Technology Come Alive. O'Reilly Media. p. 97. ISBN   9780596523206.
  22. ↑ Ronalds, B.F. (2016). Sir Francis Ronalds: Father of the Electric Telegraph. London: Imperial College Press. ISBN   978-1-78326-917-4.
  23. ↑ Ronalds, B.F. (July 2016). "Francis Ronalds (1788-1873): The First Electrical Engineer?". Proceedings of the IEEE. 104 (7): 1489–1498. doi:10.1109/JPROC.2016.2571358. S2CID   20662894.
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