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History of the battery

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This article is about history of the development of electrochemical cells. For the history of artillery batteries, see Artillery battery.
A voltaic pile, the first chemical battery

Batteries provided the main source of electricity before the development of electric generators and electrical grids around the end of the 19th century. Successive improvements in battery technology facilitated major electrical advances, from early scientific studies to the rise of telegraphs and telephones, eventually leading to portable computers, mobile phones, electric cars, and many other electrical devices.

Scientists and engineers developed several commercially important types of battery. "Wet cells" were open containers that held liquid electrolyte and metallic electrodes. When the electrodes were completely consumed, the wet cell was renewed by replacing the electrodes and electrolyte. Open containers are unsuitable for mobile or portable use. Early electric cars used semi-sealed wet cells.

"Primary" batteries could produce current as soon as assembled, but once the active elements were consumed, they could not be electrically recharged. The development of the lead-acid battery and subsequent "secondary" or "rechargeable" types allowed energy to be restored to the cell, extending the life of permanently assembled cells.

Early electric experiments and birth of the term battery

A battery of linked glass capacitors (Leyden jars)

In 1749 Benjamin Franklin, the U.S. polymath and founding father, first used the term "battery" to describe a set of linked capacitors he used for his experiments with electricity. These capacitors were panels of glass coated with metal on each surface.[1] These capacitors were charged with a static generator and discharged by touching metal to their electrode. Linking them together in a "battery" gave a stronger discharge. Originally having the generic meaning of "a group of two or more similar objects functioning together", as in an artillery battery, the term came to be used for voltaic piles and similar devices in which many electrochemical cells were connected together in the manner of Franklin's capacitors. Today even a single electrochemical cell, e.g. a dry cell, is commonly called a battery.

Invention of the battery

The trough battery, which was in essence a Voltaic Pile laid down to prevent electrolyte leakage

In 1780, Luigi Galvani was dissecting a frog affixed to a brass hook. When he touched its leg with his iron scalpel, the leg twitched. Galvani believed the energy that drove this contraction came from the leg itself, and called it "animal electricity".

However, Alessandro Volta, a friend and fellow scientist, disagreed, believing this phenomenon was caused by two different metals joined together by a moist intermediary. He verified this hypothesis through experiment, and published the results in 1791. In 1800, Volta invented the first true battery, which came to be known as the voltaic pile. The voltaic pile consisted of pairs of copper and zinc discs piled on top of each other, separated by a layer of cloth or cardboard soaked in brine (i.e., the electrolyte). Unlike the Leyden jar, the voltaic pile produced a continuous electricity and stable current, and lost little charge over time when not in use, though his early models could not produce a voltage strong enough to produce sparks.[2] He experimented with various metals and found that zinc and silver gave the best results.

Volta believed the current was the result of two different materials simply touching each other—an obsolete scientific theory known as contact tension—and not the result of chemical reactions. As a consequence, he regarded the corrosion of the zinc plates as an unrelated flaw that could perhaps be fixed by changing the materials somehow. However, no scientist ever succeeded in preventing this corrosion. In fact, it was observed that the corrosion was faster when a higher current was drawn. This suggested that the corrosion was actually integral to the battery's ability to produce a current. This, in part, led to the rejection of Volta's contact tension theory in favor of electrochemical theory. Volta's illustrations of his Crown of Cups and voltaic pile have extra metal disks, now known to be unnecessary, on both the top and bottom. The figure associated with this section, of the zinc-copper voltaic pile, has the modern design, an indication that "contact tension" is not the source of electromotive force for the voltaic pile.

Volta's original pile models had some technical flaws, one of them involving the electrolyte leaking and causing short-circuits due to the weight of the discs compressing the brine-soaked cloth. A Scotsman named William Cruickshank solved this problem by laying the elements in a box instead of piling them in a stack. This was known as the trough battery.[3] Volta himself invented a variant that consisted of a chain of cups filled with a salt solution, linked together by metallic arcs dipped into the liquid. This was known as the Crown of Cups. These arcs were made of two different metals (e.g., zinc and copper) soldered together. This model also proved to be more efficient than his original piles,[4] though it did not prove as popular.

A zinc-copper voltaic pile

Another problem with Volta's batteries was short battery life (an hour's worth at best), which was caused by two phenomena. The first was that the current produced electrolysed the electrolyte solution, resulting in a film of hydrogen bubbles forming on the copper, which steadily increased the internal resistance of the battery (this effect, called polarization, is counteracted in modern cells by additional measures). The other was a phenomenon called local action, wherein minute short-circuits would form around impurities in the zinc, causing the zinc to degrade. The latter problem was solved in 1835 by William Sturgeon, who found that amalgamated zinc, whose surface had been treated with some mercury, didn't suffer from local action.[5]

Despite its flaws, Volta's batteries provided a steadier current than Leyden jars, and made possible many new experiments and discoveries, such as the first electrolysis of water by Anthony Carlisle and William Nicholson.

See also: Leyden jar

First practical batteries

Daniell cell

Schematic representation of Daniell's original cell

A British chemist named John Frederic Daniell found a way to solve the hydrogen bubble problem in the Voltaic Pile by using a second electrolyte to consume the hydrogen produced by the first. In 1836 he invented the Daniell cell, which consisted of a copper pot filled with a copper sulfate solution, in which was immersed an unglazed earthenware container filled with sulfuric acid and a zinc electrode. The earthenware barrier was porous, which allowed ions to pass through but kept the solutions from mixing. Without this barrier, when no current was drawn the copper ions would drift to the zinc anode and undergo reduction without producing a current, which would destroy the battery's life.[6]

Over time, copper buildup would block the pores in the earthenware barrier and cut short the battery's life. Nevertheless, the Daniell cell was a great improvement over the existing technology used in the early days of battery development and was the first practical source of electricity. It provided a longer and more reliable current than the Voltaic cell because the electrolyte deposited copper (a conductor) rather than hydrogen (an insulator) on the cathode. It was also safer and less corrosive. It had an operating voltage of roughly 1.1 volts. It soon became the industry standard for use, especially with the new telegraph networks.[5]

The Daniell cell is also the historical basis for the contemporary definition of the volt, which is the unit of electromotive force in the International System of Units. The definitions of electrical units that were proposed at the 1881 International Conference of Electricians were designed so that the electromotive force of the Daniell cell would be about 1.0 volts.[7][8] With contemporary definitions, the standard potential of the Daniell cell at 25 C is actually 1.10 V.[9]

Bird's cell

A version of the Daniell cell was invented in 1837 by the Guy's hospital physician Golding Bird who used a plaster of Paris barrier to keep the solutions separate. Bird's experiments with this cell were of some importance to the new discipline of electrometallurgy.[10][11]

Porous pot cell

Porous pot cell

The porous pot version of the Daniell cell was invented by John Dancer, a Liverpool instrument maker, in 1838.[10] It consists of a central zinc anode dipped into a porous earthenware pot containing a zinc sulfate solution. The porous pot is, in turn, immersed in a solution of copper sulfate contained in a copper can, which acts as the cell's cathode. The use of a porous barrier allows ions to pass through but keeps the solutions from mixing. Without this barrier, when no current was drawn the copper ions would drift to the zinc anode and undergo reduction without producing a current, which would destroy the battery's life.[12]

Gravity cell

A 1919 illustration of a gravity cell. This particular variant is also known as a crowfoot cell due to distinctive shape of the electrodes

Sometime during the 1860s, a Frenchman by the name of Callaud invented a variant of the Daniell cell called the gravity cell.[5] This simpler version dispensed with the porous barrier. This reduced the internal resistance of the system and, thus, the battery yielded a stronger current. It quickly became the battery of choice for the American and British telegraph networks, and was used until the 1950s.[13] In the telegraph industry, this battery was often assembled on site by the telegraph workers themselves, and when it ran down it could be renewed by replacing the consumed components.[14] Even after most telegraph lines started being powered by motor-generators the gravity battery continued to be used in way stations to power the local circuit at least into the 1950s.[15]

The gravity cell consisted of a glass jar, in which a copper cathode sat on the bottom and a zinc anode was suspended beneath the rim. Copper sulfate crystals would be scattered around the cathode and then the jar would be filled with distilled water. As the current was drawn, a layer of zinc sulfate solution would form at the top around the anode. This top layer was kept separate from the bottom copper sulfate layer by its lower density and by the polarity of the cell.

The zinc sulfate layer was clear in contrast to the deep blue copper sulfate layer, which allowed a technician to measure the battery life with a glance. On the other hand, this setup meant the battery could be used only in a stationary appliance, else the solutions would mix or spill. Another disadvantage was that a current had to be continually drawn to keep the two solutions from mixing by diffusion, so it was unsuitable for intermittent use.

Poggendorff cell

The German scientist Johann Christian Poggendorff overcame the problems with separating the electrolyte and the depolariser using a porous earthenware pot in 1842.[16] In the Poggendorff cell, sometimes called Grenet Cell due to the works of Eugene Grenet around 1859, the electrolyte was dilute sulphuric acid and the depolariser was chromic acid. The two acids were physically mixed together eliminating the porous pot. The positive electrode (cathode) was two carbon plates, with a zinc plate (negative or anode) positioned between them. Because of the tendency of the acid mixture to react with the zinc, a mechanism was provided to raise the zinc electrode clear of the acids.

The cell provided 1.9 volts. It proved popular with experimenters for many years due to its relatively high voltage; greater ability to produce a consistent current and lack of any fumes, but the relative fragility of its thin glass enclosure and the necessity of having to raise the zinc plate when the cell was not in use eventually saw it fall out of favour. The cell was also known as the 'chromic acid cell', but principally as the 'bichromate cell'. This latter name came from the practice of producing the chromic acid by adding sulphuric acid to potassium bichromate (the old name for potassium dichromate), even though the cell itself contained no bichromate.

The Fuller cell was developed from the Poggendorff cell. Although the chemistry was principally the same, the two acids were once again separated by a porous container and the zinc was treated with mercury to form an amalgam. This substantially reduced the 'local action' mainly responsible for the consumption of the zinc, but the presence of the porous pot reintroduced many of the problems that the Poggendorff cell had solved. The practice of treating zinc with mercury survived well into the 20th century until environmental considerations forced its abandonment.

Grove cell

The Grove cell was invented by Welshman William Robert Grove in 1839. It consisted of a zinc anode dipped in sulfuric acid and a platinum cathode dipped in nitric acid, separated by porous earthenware. The Grove cell provided a high current and nearly twice the voltage of the Daniell cell, which made it the favoured cell of the American telegraph networks for a time. However, it gave off poisonous nitric oxide fumes when operated.[5] The voltage also dropped sharply as the charge diminished, which became a liability as telegraph networks grew more complex. Platinum was also very expensive. (The 1841 Bunsen cell used carbon). The Grove cell was replaced by the cheaper, safer and better performing gravity cell in the 1860s.

Dun cell

Alfred Dun 1885, nitro-muriatic acid (aqua regis) - Iron & Carbon[17]

In the new element there can be used advantageously as exciting-liquid in the first case such solutions as have in a concentrated condition great depolarizing-power, which effect the whole depolarization chemically without necessitating the mechanical expedient of increased carbon surface. It is preferred to use iron as the positive electrode, and as exciting-liquid nitro muriatic acid, (aqua regis,) the mixture consisting of muriatic and nitric acids. The nitro-muriatic acid, as explained above, serves for filling both cells. For the carbon-cells it is used strong or very slightly diluted, but for the other cells very diluted, (about one-twentieth, or at the most one-tenth.) The element containing in one cell carbon and concentrated nitro-muriatic acid and in the other cell iron and dilute nitro-muriatic acid remains constant for at least twenty hours when employed for electric incandescent lighting. (p. 80 at Google Books)

Rechargeable batteries and dry cells


19th-century illustration of Planté's original lead-acid cell

Up to this point, all existing batteries would be permanently drained when all their chemical reactions were spent. In 1859, Gaston Planté invented the lead–acid battery, the first-ever battery that could be recharged by passing a reverse current through it. A lead acid cell consists of a lead anode and a lead dioxide cathode immersed in sulphuric acid. Both electrodes react with the acid to produce lead sulfate, but the reaction at the lead anode releases electrons whilst the reaction at the lead dioxide consumes them, thus producing a current. These chemical reactions can be reversed by passing a reverse current through the battery, thereby recharging it.

Planté's first model consisted of two lead sheets separated by rubber strips and rolled into a spiral.[18] His batteries were first used to power the lights in train carriages while stopped at a station[citation needed]. In 1881, Camille Alphonse Faure invented an improved version that consisted of a lead grid lattice into which a lead oxide paste was pressed, forming a plate. Multiple plates could be stacked for greater performance. This design was easier to mass-produce.

Compared to other batteries, Planté's was rather heavy and bulky for the amount of energy it could hold. However, it could produce remarkably large currents in surges. It also had very low internal resistance, meaning that a single battery could be used to power multiple circuits.[5]

The lead-acid battery is still used today in automobiles and other applications where weight is not a big factor. The basic principle has not changed since 1859. In the early 1930s, a gel electrolyte (instead of a liquid) produced by adding silica to a charged cell was used in the LT battery of portable vacuum-tube radios. In the 1970s, "sealed" versions became common (commonly known as a "gel cell" or "SLA"), allowing the battery to be used in different positions without failure or leakage.

Today cells are classified as "primary" if they produce a current only until their chemical reactants are exhausted, and "secondary" if the chemical reactions can be reversed by recharging the cell. The lead-acid cell was the first "secondary" cell.

Leclanché cell

A 1912 illustration of a Leclanché cell

In 1866, Georges Leclanché invented a battery that consisted of a zinc anode and a manganese dioxide cathode wrapped in a porous material, dipped in a jar of ammonium chloride solution. The manganese dioxide cathode had a little carbon mixed into it as well, which improved conductivity and absorption.[19] It provided a voltage of 1.4 volts.[20] This cell achieved very quick success in telegraphy, signalling and electric bell work.

The dry cell form was used to power early telephones—usually from an adjacent wooden box affixed to the wall—before telephones could draw power from the telephone line itself. The Leclanché cell could not provide a sustained current for very long. In lengthy conversations, the battery would run down, rendering the conversation inaudible.[21] This was because certain chemical reactions in the cell increased the internal resistance and, thus, lowered the voltage. These reactions reversed themselves when the battery was left idle, so it was good only for intermittent use.[5]

Zinc-carbon cell, the first dry cell

Many experimenters tried to immobilize the electrolyte of an electrochemical cell to make it more convenient to use. The Zamboni pile of 1812 was a high-voltage dry battery but capable of delivering only minute currents. Various experiments were made with cellulose, sawdust, spun glass, asbestos fibers, and gelatine.[22]

In 1886, Carl Gassner obtained a German patent[23] on a variant of the Leclanché cell, which came to be known as the dry cell because it did not have a free liquid electrolyte. Instead, the ammonium chloride was mixed with Plaster of Paris to create a paste, with a small amount of zinc chloride added in to extend the shelf life. The manganese dioxide cathode was dipped in this paste, and both were sealed in a zinc shell, which also acted as the anode. In November 1887, he obtained U.S. Patent 373,064 for the same device.

Unlike previous wet cells, Gassner's dry cell was more solid, did not require maintenance, did not spill, and could be used in any orientation. It provided a potential of 1.5 volts. The first mass-produced model was the Columbia dry cell, first marketed by the National Carbon Company in 1896.[24] The NCC improved Gassner's model by replacing the plaster of Paris with coiled cardboard, an innovation that left more space for the cathode and made the battery easier to assemble. It was the first convenient battery for the masses and made portable electrical devices practical, and led directly to the invention of the flashlight.

The zinc–carbon battery (as it came to be known) is still manufactured today.

In parallel, in 1887 Wilhelm Hellesen developed his own dry cell design. It has been claimed that Hellesen's design preceded that of Gassner.[25]

In 1887, a dry-battery was developed by Yai Sakizō (屋井 先蔵) of Japan, then patented in 1892.[26][27] In 1893, Yai Sakizō's dry-battery was exhibited in World's Columbian Exposition and commanded considerable international attention.

NiCd, the first alkaline battery

In 1899, a Swedish scientist named Waldemar Jungner invented the nickel–cadmium battery, a rechargeable battery that had nickel and cadmium electrodes in a potassium hydroxide solution; the first battery to use an alkaline electrolyte. It was commercialized in Sweden in 1910 and reached the United States in 1946. The first models were robust and had significantly better energy density than lead-acid batteries, but were much more expensive.

20th century: new technologies and ubiquity


Nickel-iron batteries manufactured between 1972 and 1975 under the "Exide" brand, originally developed in 1901 by Thomas Edison.
A set of modern batteries

Jungner had invented a nickel–iron battery the same year as his Ni-Cad battery, but found it to be inferior to its cadmium counterpart and, as a consequence, never bothered patenting it. It produced a lot more hydrogen gas when being charged, meaning it could not be sealed, and the charging process was less efficient (it was, however, cheaper). However, Thomas Edison picked up Jungner's nickel-iron battery design, patented it himself and sold it in 1903. Edison wanted to commercialise a more lightweight and durable substitute for the lead-acid battery that powered some early automobiles, and hoped that by doing so electric cars would become the standard, with his firm as its main battery vendor. However, customers found his first model to be prone to leakage and short battery life, and it did not outperform the lead-acid cell by much either. Although Edison was able to produce a more reliable and powerful model seven years later, by this time the inexpensive and reliable Model T Ford had made gasoline engine cars the standard. Nevertheless, Edison's battery achieved great success in other applications such as electric and diesel-electric rail vehicles, providing backup power for railroad crossing signals, or to provide power for the lamps used in mines.[28][29][30]

Common alkaline batteries

Until the late 1950s the zinc–carbon battery continued to be a popular primary cell battery, but its relatively low battery life hampered sales. In 1955, an engineer working for Union Carbide at the National Carbon Company Parma Research Laboratory named Lewis Urry was tasked with finding a way to extend the life of zinc-carbon batteries, but Urry decided instead that alkaline batteries held more promise. Until then, longer-lasting alkaline batteries were unfeasibly expensive. Urry's battery consisted of a manganese dioxide cathode and a powdered zinc anode with an alkaline electrolyte. Using powdered zinc gave the anode a greater surface area. These batteries hit the market in 1959.

Nickel-hydrogen and nickel metal-hydride

The nickel–hydrogen battery entered the market as an energy-storage subsystem for commercial communication satellites.[31][32]

The first consumer grade nickel–metal hydride batteries (NiMH) for smaller applications appeared on the market in 1989 as a variation of the 1970s nickel–hydrogen battery.[33] NIMH batteries tend to have longer lifespans than NiCd batteries (and their lifespans continue to increase as manufacturers experiment with new alloys) and, since cadmium is toxic, NiMH batteries are less damaging to the environment.

Lithium and lithium-ion batteries

Lithium is the metal with lowest density and with the greatest electrochemical potential and energy-to-weight ratio. The low atomic weight and small size of its ions also speeds its diffusion. So theory suggests it would make an ideal material for batteries. Experimentation with lithium batteries began in 1912 under G.N. Lewis, and in the 1970s the first lithium batteries came onto the market. Three volt lithium primary cells such as the CR123A type and three volt button cells are still widely used, especially in cameras and very small devices. Energizer and other brands market 1.5 volt lithium primary cells, with advantages for some applications over alkaline cells.

Three important developments marked the 1980s. In 1980 an American chemist, John B. Goodenough, disclosed the LiCoO2 cathode (positive lead) and a Moroccan research scientist, Rachid Yazami, discovered the graphite anode (negative lead) with the solid electrolyte. In 1981, Japanese chemist Tokio Yamabe and late Shjzukuni Yata discovered novel nano-carbonacious-PAS, [34] and found that it was very effective for the anode in the conventional liquid electrolyte. [35] [36] This led a research team managed by Akira Yoshino of Asahi Chemical, Japan, to build the first lithium-ion battery prototype in 1985, a rechargeable and more stable version of the lithium battery; Sony commercialized the lithium-ion battery in 1991. [37]

In 1997 the lithium polymer battery was released by Sony and Asahi Kasei. These batteries hold their electrolyte in a solid polymer composite instead of in a liquid solvent, and the electrodes and separators are laminated to each other. The latter difference allows the battery to be encased in a flexible wrapping instead of in a rigid metal casing, which means such batteries can be specifically shaped to fit a particular device. This advantage has favored lithium polymer batteries in the design of portable electronic devices such as mobile phones and personal digital assistants, and of radio-controlled aircraft, as such batteries allow for more flexible and compact design. They generally have a lower energy density than normal lithium-ion batteries.

See also

Notes and references

  1. ^ "Benjamin Franklin et al.; Leonard W. Labaree, ed., ''The Papers of Benjamin Franklin'' (New Haven, Connecticut: Yale University Press, 1961) vol. 3, page 352: Letter to Peter Collinson, April 29, 1749. paragraph 18". Retrieved 2012-08-29. 
  2. ^ Finn, Bernard S. (September 2002). "Origin of Electrical Power". National Museum of American History. Retrieved 2012-08-29. 
  3. ^ Institute and Museum of the History of Science. "Trough Battery". Retrieved 2007-01-15. 
  4. ^ Decker, Franco (January 2005). "Volta and the 'Pile'". Electrochemistry Encyclopedia. Case Western Reserve University. Retrieved 2012-11-30. 
  5. ^ a b c d e f Calvert, James B. (2000). "The Electromagnetic Telegraph". Retrieved 2007-01-12. 
  6. ^ Carboni, Giorgio (1999-06-07). "Experiments in Electrochemistry". Fun Science Gallery. Retrieved 2012-08-29. 
  7. ^ Borvon, Gérard (September 10, 2012). "History of the electrical units". Association S-EAU-S. 
  8. ^ Hamer, Walter J. (January 15, 1965). Standard Cells: Their Construction, Maintenance, and Characteristics (PDF). National Bureau of Standards Monograph #84. US National Bureau of Standards. 
  9. ^ Spencer, James N.; Bodner, George M.; Rickard, Lyman H. (2010). Chemistry: Structure and Dynamics (Fifth Edition). John Wiley & Sons. p. 564. ISBN 9780470587119. 
  10. ^ a b Watt, Alexander; Philip, Arnold (2005). Electroplating and Electrorefining of Metals. Watchmaker Publishing. pp. 90–92. ISBN 1929148453.  Reprint of an 1889 volume.
  11. ^ Golding Bird, Report of the Seventh Meeting of the British Society for the Advancement of Science, vol.6 (1837), p.45, London: J. Murray, 1838.
  12. ^ Giorgio Carboni, Experiments in Electrochemistry; Last accessed on Jul 30, 2010.
  13. ^ Tools of Telegraphy, Telegraph Lore. Last accessed Jan 9, 2007 Archived April 5, 2009, at the Wayback Machine.
  14. ^ Gregory S. Raven, Recollections of a Narrow Gauge Lightning Slinger Archived October 31, 2007, at the Wayback Machine.
  15. ^ Tools of Telegraphy, Telegraph Lore; Last accessed Jul 30, 2010
  16. ^ [1]
  17. ^ United States. Patent Office (1886). Specifications and Drawings of Patents Relating to Electricity Issued by the U. S. pp. 80–81. In the new element there can be used advantageously as exciting-liquid in the first case such solutions as have in a concentrated condition great depolarizing-power, which effect the whole depolarization chemically without necessitating the mechanical expedient of increased carbon surface. It is preferred to use iron as the positive electrode, and as exciting-liquid nitro muriatic acid, (aqua regis,) the mixture consisting of muriatic and nitric acids. The nitro-muriatic acid, as explained above, serves for filling both cells. For the carbon-cells it is used strong or very slightly diluted, but for the other cells very diluted, (about one-twentieth, or at the most one-tenth.) The element containing in one cell carbon and concentrated nitro-muriatic acid and in the other cell iron and dilute nitro-muriatic acid remains constant for at least twenty hours when employed for electric incandescent lighting. (p. 80 at Google Books) 
  18. ^ "Gaston Planté (1834-1889)". Corrosion Doctors. Retrieved 2012-08-29. 
  19. ^ "Zinc-Carbon Batteries". Molecular Expressions. Retrieved 2012-08-29. 
  20. ^ The Boy Electrician by J.W.Simms M.I.E.E. (Page 61)
  21. ^ Battery Facts. "Leclanché Cell". Retrieved 2007-01-09. 
  22. ^ W. E. Ayrton Practical Electricity; A Laboratory and Lecture Course for First Year ... 1897, reprint Read Books, 2008 ISBN 1-4086-9150-7, page 458
  23. ^ DE patent 37758, Carl Gassner, Jr., issued 1886-04-08 
  24. ^ "The Columbia Dry Cell Battery". National Historic Chemical Landmarks. American Chemical Society. Retrieved 2014-02-21. 
  25. ^ Energi på dåse, Jytte Thorndahl. Last accessed on June 26, 2007 Archived September 28, 2007, at the Wayback Machine.
  26. ^ "1) The Yai dry-battery". The history of the battery. Battery association of Japan. Retrieved 2012-08-29. 
  27. ^ "乾電池の発明者は日本人だった 理大ゆかりの屋井先蔵". Tokyo University of Science. 2004-07-07. Retrieved 2012-08-29. 
  28. ^ IEEE History Center. "Edison's Alkaline Battery". Retrieved 2011-06-20. 
  29. ^ "Systematic design of an autonomous hybrid locomotive | EUrailmag". Retrieved 2013-04-17. 
  30. ^ "Magma #10 Project". 2012-05-15. Retrieved 2013-04-17. 
  31. ^ "A nickel/hydrogen battery for PV systems". IEEE Xplore. 2011-09-27. doi:10.1109/62.59267. Retrieved 2012-08-29. 
  32. ^ "Nickel-Hydrogen Battery Technology—Development and Status" (PDF). Archived from the original (PDF) on 2009-03-18. Retrieved 2012-08-29. 
  33. ^ Sponsored by (2008-03-06). "In search of the perfect battery". Retrieved 2012-08-29. 
  34. ^ T. Yamabe, K. Tanaka, K. Ohzeki, and S.Yata, Solid State Communications, 44,823, (1982)
  35. ^ S. Yata, U.S. Patent #4,601,849
  36. ^ Shjzukuni Yata, Kazuyoshi Tanaka and Tokio Yamabe, Polyacene (PAS) Batteries, MRS Proceedings, Volume 496,1997
  37. ^ P. Novak, K. Muller, K. S. V. Santhanam, O. Haas, Electrochemically Active Polymers for Rechargeable Batteries, Chem. Rev., 97, p.272 (1997)