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The periodic table of the chemical elements

A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus. The term is also used to refer to a pure chemical substance composed of atoms with the same number of protons.[1] Common examples of elements are iron, copper, silver, gold, hydrogen, carbon, nitrogen, and oxygen. In total, 118 elements have been observed as of March 2010, of which 94 occur naturally on Earth. 80 elements have stable isotopes, namely all elements with atomic numbers 1 to 82, except elements 43 and 61 (technetium and promethium). Elements with atomic numbers 83 or higher (bismuth and above) are inherently unstable, and undergo radioactive decay. The elements from atomic number 83 to 94 have no stable nuclei, but are nevertheless found in nature, either surviving as remnants of the primordial stellar nucleosynthesis that produced the elements in the solar system, or else produced as short-lived daughter-isotopes through the natural decay of uranium and thorium.[2]

All chemical matter consists of these elements. New elements of higher atomic number are discovered from time to time, as products of artificial nuclear reactions.

Contents

History

Several old philosophies used a set of archetypal classical elements to explain patterns in nature. The term 'element' was originally used to refer to a state of matter (solid/earth, liquid/water, gas/air, and plasma/fire) or a phase of matter (as in the Chinese Wu Xing), rather than the chemical elements of modern science. The Greek, Indian (Tattva and Mahābhūta) and Japanese (go dai) traditions essentially had the same five elements: Air, Earth, Fire, Water and Aether.

The term 'elements' (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE, in his dialogue Timaeus, which includes a discussion of the composition of inorganic and organic bodies and is a rudimentary treatise on chemistry. Plato assumed that the minute particle of each element corresponded to one of the regular polyhedra: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).[citation needed][3]

Adding to the four elements of the Greek philosopher Empedocles, in about 350 BC, Aristotle also used the term "element" and conceived of a fifth element called "quintessence", which formed the heavens. Aristotle defined an element as:

Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.[4]

Building on the theory in circa 790, Arab/Persian chemist and alchemist, Jabir ibn Hayyan (Geber), postulated that metals were formed out of two elements: sulfur, ‘the stone that burns’, which characterized the principle of combustibility, and mercury, which contained the idealized principle of metallic properties.[5] Shortly thereafter, this evolved into the Arabic concept of the three principles: sulfur giving flammability or combustion, mercury giving volatility and stability, and salt giving solidity.

In the 10th century, Persian physician and alchemist Muhammad ibn Zakarīya Rāzi (Rhazes) wrote the Doubts concerning Galen, in which he refuted both the Galenic medical theory of four humours and the underlying ancient concept of four classical elements. He carried out an experiment that upset these theories by inserting a liquid with a different temperature into a body resulting in an increase or decrease of bodily heat, which resembled the temperature of that particular fluid. Razi noted particularly that a warm drink would heat up the body to a degree much higher than its own natural temperature, thus the drink would trigger a response from the body, rather than transferring only its own warmth or coldness to it. Razi's chemical experiments suggested other qualities of matter, such as "oiliness" and "sulfurousness", or inflammability and salinity, which were not readily explained by the traditional fire, water, earth and air division of elements.[6]

In 1524, Swiss chemist Paracelsus adopted Aristotle’s four element theory, but reasoned that they appeared in bodies as Geber’s three principles. Paracelsus saw these principles as fundamental, and justified them by recourse to the description of how wood burns in fire. Mercury included the cohesive principle, so that when it left in smoke the wood fell apart. Smoke represented the volatility (the mercury principle), the heat-giving flames represented flammability (sulfur), and the remnant ash represented solidity (salt).[5]

In 1669, German physician and chemist Johann Becher published his Physica Subterranea. In modification on the ideas of Paracelsus, he argued that the constituents of bodies are air, water, and three types of earth: terra fluida, the mercurial element, which contributes fluidity and volatility; terra lapida, the solidifying element, which produces fusibility or the binding quality; and terra pinguis, the fatty element, which gives material substance its oily and combustible qualities.[7] These three earths correspond with Geber’s three principles. A piece of wood, for example, according to Becher, is composed of ash and terra pinguis; when the wood is burnt, the terra pinguis is released, leaving the ash. In other words, in combustion the fatty earth burns away.

In 1661, Robert Boyle showed that there were more than just four classical elements as the ancients had assumed.[8] The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric.[9] By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that cannot be decomposed into any simpler substance.[8] Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. In 1913, Henry Moseley discovered that the physical basis of the atomic number of the atom was its nuclear charge, which eventually led to the current definition. The current definition also avoids some ambiguities due to isotopes and allotropes.

By 1919, there were seventy-two known elements.[10] In 1955, element 101 was discovered and named mendelevium in honor of Mendeleev, the first to arrange the elements in a periodic manner. In October 2006, the synthesis of element 118 was reported; however, element 117 has not yet been created in the laboratory.

Description

The lightest elements are hydrogen and helium, both theoretically created by Big Bang nucleosynthesis during the first 20 minutes of the universe[11] in a ratio of around 3:1 by mass (approximately 12:1 by number of atoms). Almost all other elements found in nature, including some further hydrogen and helium created since then, were made by various natural or (at times) artificial methods of nucleosynthesis, including occasionally breakdown activities such as nuclear fission, alpha decay, cluster decay, and cosmic ray spallation.

As of 2010, there are 118 known elements (in this context, "known" means observed well enough, even from just a few decay products, to have been differentiated from any other element).[12][13] Of these 118 elements, 94 occur naturally on Earth. Six of these occur in extreme trace quantities: technetium, atomic number 43; promethium, number 61; astatine, number 85; francium, number 87; neptunium, number 93; and plutonium, number 94. These 94 elements, and also possibly element 98 californium, have been detected in the universe at large, in the spectra of stars and also supernovae, where short-lived radioactive elements are newly being made.

The remaining 22 elements, not found on Earth or in astronomical spectra, have been derived artificially. All of the elements that are derived solely through artificial means are radioactive with very short half-lives; if any atoms of these elements were present at the formation of Earth, they are extremely likely to have already decayed, and if present in novae, have been in quantities too small to have been noted. Technetium was the first purportedly non-naturally occurring element to be synthesized, in 1937, although trace amounts of technetium have since been found in nature, and the element may have been discovered naturally in 1925. This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally occurring trace elements.

Lists of the elements are available by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as Ionization energies of the elements. The most convenient presentation of the elements is in the periodic table, which groups elements with similar chemical properties together.

Atomic number

The atomic number of an element, Z, is equal to the number of protons that defines the element. For example, all carbon atoms contain 6 protons in their nucleus; so the atomic number "Z" of carbon is 6. Carbon atoms may have different numbers of neutrons, which are known as isotopes of the element.

The number of protons in the atomic nucleus also determines its electric charge, which in turn determines the electrons of the atom in its non-ionized state. This in turn (by means of the Pauli exclusion principle) determines the atom's various chemical properties. So all carbon atoms, for example, ultimately have identical chemical properties because they all have the same number of protons in their nucleus, and therefore have the same atomic number. It is for this reason that atomic number rather than mass number (or atomic weight) is considered the identifying characteristic of an element.

Atomic mass

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a super-index on the left hand side of the atomic symbol (e.g., 238U).

The relative atomic mass of an element is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit (u). This number may be a fraction that is not close to a whole number, due to the averaging process. On the other hand, the atomic mass of a pure isotope is quite close to its mass number. Whereas the mass number is a natural (or whole) number, the atomic mass of a single isotope is a real number that is close to a natural number. In general, it differs slightly from the mass number as the mass of the protons and neutrons is not exactly 1 u, the electrons also contribute slightly to the atomic mass, and because of the nuclear binding energy. For example, the mass of 19F is 18.9984032 u. The only exception to the atomic mass of an isotope not being a natural number is 12C, which has a mass of exactly 12, due to the definition of u (it is fixed as 1/12th of the mass of a free carbon-12 atom, exactly).

Isotopes

Isotopes are atoms of the same element (that is, with the same number of protons in their atomic nucleus), but having different numbers of neutrons. Most (66 of 94) naturally occurring elements have more than one stable isotope. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons in the nucleus, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, the three isotopes of carbon are known as carbon-12, carbon-13, and carbon-14, often abbreviated to 12C, 13C, and 14C. Carbon in everyday life and in chemistry is a mixture of 12C, 13C, and 14C atoms.

Except in the case of the isotopes of hydrogen (which differ greatly from each other in relative mass—enough to cause chemical effects), the isotopes of the various elements are typically chemically nearly indistinguishable from each other. For example, the three naturally-occurring isotopes of carbon have essentially the same chemical properties, but different nuclear properties. In this example, carbon-12 and carbon-13 are stable atoms, but carbon-14 is unstable; it is radioactive, undergoing beta decay into nitrogen-14.

As illustrated by carbon, all of the elements have some isotopes that are radioactive (radioisotopes), which decay into other elements upon radiating an alpha or beta particle. Certain elements only have radioactive isotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one stable isotope, and the mean number of stable isotopes for the 80 stable elements is 3.1 stable isotopes per element. The largest number of stable isotopes that occur for an element is 10 (for tin, element 50).

Allotropes

Atoms of pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple structures (spacial arrangements of atoms), known as allotropes, which differ in their properties. For example, carbon can be found as diamond, which has a tetrahedral structure around each carbon atom; graphite, which has layers of carbon atoms with a hexagonal structure stacked on top of each other; graphene, which is a single layer of graphite that is incredibly strong; fullerenes, which have nearly spherical shapes; and carbon nanotubes, which are tubes with a hexagonal structure (even these may differ from each other in electrical properties). The ability for an element to exist in one of many structural forms is known as 'allotropy'.

Standard state

The standard state, or reference state, of an element is defined as its thermodynamically most stable state at 1 bar at a given temperature (typically at 298.15 K). In thermochemistry, an element is defined to have an enthalpy of formation of zero in its standard state. For example, the reference state for carbon is graphite, because it is more stable than the other allotropes.

Nomenclature

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When it was learned, existing names (e.g., gold, mercury, iron) were kept in most countries, and national differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For example, the Germans use "Wasserstoff" for "hydrogen", "Sauerstoff" for "oxygen" and "Stickstoff" for "nitrogen", while English and some romance languages use "sodium" for "natrium" and "potassium" for "kalium", and the French, Italians, Greeks, Portuguese and Poles prefer "azote/azot/azoto" for "nitrogen".

But for international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that "aluminium" and "caesium" take the place of the US spellings "aluminum" and "cesium", while the US "sulfur" takes the place of the British "sulphur". Chemicals that are practical to sell in bulk in many countries, however, still have national names, and those that do not use the Latin alphabet cannot be expected to use the IUPAC name.

According to IUPAC, the full name of an element is not capitalized, even if it is derived from a proper noun such as the elements californium or einsteinium (unless it would be capitalized by some other rule, such as to begin a sentence). Isotopes of chemical elements are also uncapitalized if written out: carbon-12 or uranium-235. Symbols of chemical elements, however, are capitalized: thus the symbols for the elements just discussed are Cf and Es; C-12 and U-235.

In the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have a half life too short for them to remain in any appreciable amounts. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question that delayed naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy).

Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century. For example, lutetium was named in reference to Paris, France. The Germans were reluctant to relinquish naming rights to the French, often calling it cassiopeium. The British discoverer of niobium originally named it columbium, in reference to the New World. It was used extensively as such by American publications prior to international standardization.

Chemical symbols

For the listing of current and not used Chemical symbols, and other symbols that look like chemical symbols, please see List of elements by symbol.

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules.

The current system of chemical notation was invented by Berzelius. In this typographical system chemical symbols are not used as mere abbreviations - though each consists of letters of the Latin alphabet - they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully universal; since Latin was the common language of science at that time, they were abbreviations based on the Latin names of metals - Cu comes from Cuprum, Fe comes from Ferrum, Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Later chemical elements were also assigned unique chemical symbols, based on the name of the element, but not necessarily in English. For example, sodium has the chemical symbol 'Na' after the Latin natrium. The same applies to "W" (wolfram) for tungsten, "Hg" (hydrargyrum) for mercury, "Sn" (stannum) for tin, "K" (kalium) for potassium, "Au" (aurum) for gold, "Pb" (plumbum) for lead, and "Sb" (stibium) for antimony.

Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not be confused with a roman numeral.

The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always lower case (small letters).

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical, meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

Isotope symbols

The three main isotopes of the element hydrogen are often written as H for protium, D for deuterium and T for tritium. This is in order to make it easier to use them in chemical equations, as it replaces the need to write out the mass number for each atom. E.g. the formula for heavy water may be written D2O instead of ²H2O.

The periodic table

The periodic table of the chemical elements is a tabular method of displaying the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869. Mendeleev intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.

The periodic table is now ubiquitous within the academic discipline of chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behavior. The table has also found wide application in physics, biology, engineering, and industry. The current standard table contains 117 confirmed elements as of March 10, 2009. (While element 118 has been synthesized, element 117 has not as of August 14, 2009.)

Abundance

During the early phases of the Big Bang, nucleosynthesis of hydrogen nuclei resulted in the production of hydrogen and helium isotopes, as well as very minuscule amounts (on the order of 10−10) of lithium and beryllium. There is argument about whether or not some boron was produced in the Big Bang, as it has been observed in some very young stars,[14] but no heavier elements than boron were produced. As a result, the primordial abundance of atoms consisted of roughly 75% 1H, 25% 4He, and 0.01% deuterium.[15] Subsequent enrichment of galactic halos occurred due to Stellar nucleosynthesis and Supernova nucleosynthesis.[16] However intergalactic space can still closely resemble the primordial abundance, unless it has been enriched by some means.

The following graph (note log scale) shows abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass.[17] Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. As physical laws and processes appear common throughout the visible universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

Abundances of the chemical elements in the Solar system.
Element Parts per million
by mass
Hydrogen 739,000
Helium 240,000
Oxygen 10,400
Carbon 4,600
Neon 1,340
Iron 1,090
Nitrogen 960
Silicon 650
Magnesium 580
Sulfur 440
Potassium 210
Nickel 100

Recently discovered elements

The first transuranium element (element with atomic number greater than 92) discovered was neptunium in 1940. As of February 2010, only the elements up to 112, copernicium, have been confirmed as discovered by IUPAC, while more or less reliable claims have been made for synthesis of elements 113, 114, 115, 116 and 118. The discovery of element 112 was acknowledged in 2009, and the name 'copernicium' and the atomic symbol 'Cn' were suggested for it.[18] The name and symbol were officially endorsed by IUPAC on February 19, 2010.[19] The heaviest element that is believed to have been synthesized to date is element 118, ununoctium, on October 9, 2006, by the Flerov Laboratory of Nuclear Reactions in Dubna, Russia.[20][21]

Element 117, ununseptium, has been synthesised[22], and its place in the periodic table is preestablished.

On April 24, 2008, Amnon Marinov and six other researchers claimed that element 122 has been detected in purified natural thorium.[23][24] If confirmed, this would be the first naturally occurring heavy element in more than 50 years.[citation needed] It has yet to be proved as it is still under confirmation by the university but could be a major development as previously all transuranic elements were artificial. The claim of Marinov et al. was criticized by a part of the scientific community, and Marinov says he has submitted the article to the journals Nature and Nature Physics but both turned it down without sending it for peer review.[25]

See also

References

  1. ^ International Union of Pure and Applied Chemistry. "[1]". Compendium of Chemical Terminology Internet edition.
  2. ^ A. Earnshaw, Norman Greenwood (1997). Chemistry of the Elements (2 ed.). Butterworth-Heinemann. 
  3. ^ Hillar, Marian (2004). "The Problem of the Soul in Aristotle's De anima". NASA WMAP. http://www.socinian.org/aristotles_de_anima.html. Retrieved 2006-08-10. 
  4. ^ Partington, J.R. (1937). A Short History of Chemistry. New York: Dover Publications, Inc.. ISBN 0486659771. 
  5. ^ a b Strathern, Paul. (2000). Mendeleyev’s Dream – the Quest for the Elements. New York: Berkley Books.
  6. ^ G. Stolyarov II (2002), "Rhazes: The Thinking Western Physician", The Rational Argumentator, Issue VI
  7. ^ Partington, J.R. (1937). A Short History of Chemistry. New York: Dover Publications, Inc.
  8. ^ a b Boyle, Robert (1661). The Sceptical Chymist. London. 
  9. ^ Lavoisier, Antoine Laurent (1790). Elements of chemistry translated by Robert Kerr. Edinburgh. pp. 175–176. http://books.google.de/books?hl=de&lr=&id=4BzAjCpEK4gC&pg=PA175. 
  10. ^ Carey, George, W. (1914). The Chemistry of Human Life. Los Angeles. 
  11. ^ Gaitskell, R; Utyonkov, V. K.; Lobanov, Yu. V.; Abdullin, F. Sh.; Polyakov, A. N.; Sagaidak, R. N.; Shirokovsky, I. V.; Tsyganov, Yu. S. et al. (2006). "Evidence for Dark Matter" (PDF). Physical Review C 74: timeline on page 10. doi:10.1103/PhysRevC.74.044602. http://gaitskell.brown.edu/physics/talks/0408_SLAC_SummerSchool/Gaitskell_DMEvidence_v16.pdf. Retrieved 2008-10-08. 
  12. ^ Sanderson, Katherine (17 October 2006). "Heaviest element made - again". nature@news.com. Nature. http://www.nature.com/news/2006/061016/full/061016-4.html. Retrieved 2006-10-19. 
  13. ^ Phil Schewe and Ben Stein (17 October 2006). "Elements 116 and 118 Are Discovered". Physics News Update. American Institute of Physics. http://www.aip.org/pnu/2006/797.html. Retrieved 2006-10-19. 
  14. ^ Hubble Observations Bring Some Surprises - New York Times
  15. ^ Wright, Edward L. (September 12, 2004). "Big Bang Nucleosynthesis". UCLA Division of Astronomy. http://www.astro.ucla.edu/~wright/BBNS.html. Retrieved 2007-02-22. 
  16. ^ G. Wallerstein, I. Iben Jr., P. Parker, A. M. Boesgaard, G. M. Hale, A. E. Champagne, C. A. Barnes, F. KM-dppeler, V. V. Smith, R. D. Hoffman, F. X. Timmes, C. Sneden, R.N. Boyd, B.S. Meyer, D.L. Lambert (1999). "Synthesis of the elements in stars: forty years of progress" (pdf). Reviews of Modern Physics 69 (4): 995–1084. doi:10.1103/RevModPhys.69.995. http://www.cococubed.com/papers/wallerstein97.pdf. Retrieved 2006-08-04. 
  17. ^ Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. http://kencroswell.com/alchemy.html. 
  18. ^ "IUPAC Announces Start of the Name Approval Process for the Element of Atomic Number 112". 20 July 2009. http://media.iupac.org/news/112_Naming_Process_20090720.pdf. Retrieved 2009-08-27. 
  19. ^ IUPAC (International Union of Pure and Applied Chemistry): Element 112 is Named Copernicium
  20. ^ Phil Schewe and Ben Stein (17 October 2006). "Elements 116 and 118 Are Discovered". Physics News Update. American Institute of Physics. http://www.aip.org/pnu/2006/797.html. Retrieved 2006-10-19. 
  21. ^ Oganessian, Yu. Ts. et al. (2006-10-09). "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245Cm+48Ca fusion reactions". Physical Review C 74 (4): 044602. doi:10.1103/PhysRevC.74.044602. 
  22. ^ http://www.jinr.ru/img_sections/PAC/NP/31/PAK_NP_31_recom_eng.pdf
  23. ^ Marinov, A.; Rodushkin, I.; Kolb, D.; Pape, A.; Kashiv, Y.; Brandt, R.; Gentry, R. V.; Miller, H. W. (2008). "Evidence for a long-lived superheavy nucleus with atomic mass number A=292 and atomic number Z=~122 in natural Th". ArXiv.org 74: 044602. doi:10.1103/PhysRevC.74.044602. http://arxiv.org/abs/0804.3869. Retrieved 2008-04-28. 
  24. ^ Battersby, Stephen (2008-05-02). "Has the heaviest element been found?". NewScientist. http://www.newscientist.com/article/dn13828-has-the-heaviest-element-been-found.html. Retrieved 2009-05-01. 
  25. ^ Van Noorden, Richard (2008-05-02). "Heaviest element claim criticised". Chemistry World (Royal Society of Chemistry). http://www.rsc.org/chemistryworld/News/2008/May/02050802.asp. Retrieved 2009-05-01. 

Further reading

  • E.R. Scerri, The Periodic Table, Its Story and Its Significance, Oxford University Press, NY, 2007.

External links


A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus.[1] Common examples of elements are iron, copper, silver, gold, hydrogen, carbon, nitrogen, and oxygen. All chemical matter consists of these elements.

Save for the hydrogen and helium in the universe, which are thought to have been mostly produced in the Big Bang, most chemical elements are thought to have been produced by later processes. These processes are divided into cosmic ray spallation (important for lithium, beryllium and boron, though some of these may have formed in the Big Bang), and stellar nucleosynthesis which produces all elements heavier than boron (with carbon being the first of this series). The very heaviest elements (those beyond element 94, plutonium) decay with half lives too short to allow them to be observed naturally on Earth.

In total, 118 elements have been characterized, as of March, 2010, and new elements of higher atomic number are "discovered" from time to time, as new synthetic products of artificial nuclear reactions. Of the known elements, the first 94 occur naturally on Earth. Of these, oxygen is the most abundant element in the Earth's crust. About 80 elements have stable isotopes: namely all elements with atomic numbers 1 to 82, except elements 43 and 61 (technetium and promethium). About half of the 80 stable elements are expected to be radioactive with such long half lives that their decay happens only in theory, and is too slow to have yet been detected by experiment (see the list of nuclides). These elements (such as bismuth, only recently measured as unstable) have half lives at least 100 million to 1000 million times longer than the estimated age of the universe.

Elements with atomic numbers 83 or higher (bismuth and above) are unstable to the point that their instability has been detected, and they undergo radioactive decay. The elements from atomic number 83 to 94 are composed entirely of radioactive isotopes. However, along with unstable elements 43 and 61, they are nevertheless found on Earth, though sometimes in very small amounts. Some of these elements, notably uranium and thorium, have one or more isotopes with half lives long enough to survive as remnants of the primordial explosive stellar nucleosynthesis that produced the heavy elements before formation of our solar system (see primordial nuclide). Other radioactive elements continue to be produced in natural processes, such as production by cosmic rays, or as shorter-lived daughter nuclides or transmutation products from natural decay of longer-lived radioactives.[2]

of the chemical elements]]

When two distinct elements are chemically combined, with the atoms held together by chemical bonds, the result is termed a compound. Chemical compounds may result in elements combined in exact whole number ratios of atoms (a familiar example is water). The term "compound" does not always imply an exact combination ratio, however, inasmuch as chemical bonding of many types of elements results in crystalline solids and metallic alloys, for which exact formulas do not exist. Most of the solid substance of the Earth is of this latter type: the atoms are present in the substance of the Earth's crust, mantle, and core are combined into chemical compounds of many compositions, but these do have precise formulas.

Contents

History

's 1869 periodic table]]

Ancient philosophy posited a set of classical elements to explain patterns in nature. Elements originally referred to earth, water, air and fire rather than the chemical elements of modern science.

The term 'elements' (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE, in his dialogue Timaeus, which includes a discussion of the composition of inorganic and organic bodies and is a speculative treatise on chemistry. Plato believed the elements introduced a century earlier by Empedocles were composed of small polyhedral forms: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).[3][4]

Aristotle, c. 350 BCE, also used the term stoicheia and added a fifth element called aether, which formed the heavens. Aristotle defined an element as:

Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.[5]

In 1661, Robert Boyle showed that there were more than just four classical elements as the ancients had assumed.[6] The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric.[7] By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that cannot be decomposed into any simpler substance.[6] Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. In 1913, Henry Moseley discovered that the physical basis of the atomic number of the atom was its nuclear charge, which eventually led to the current definition. The current definition also avoids some ambiguities due to isotopes and allotropes.

By 1919, there were seventy-two known elements.[8] In 1955, element 101 was discovered and named mendelevium in honor of Mendeleev, the first to arrange the elements in a periodic manner. In October 2006, the synthesis of element 118 was reported; the synthesis of element 117 was reported in April 2010.[9]

Description

The lightest elements are hydrogen and helium, both theoretically created by Big Bang nucleosynthesis during the first 20 minutes of the universe[10] in a ratio of around 3:1 by mass (approximately 12:1 by number of atoms). Almost all other elements found in nature, including some further hydrogen and helium created since then, were made by various natural or (at times) artificial methods of nucleosynthesis, including occasionally breakdown activities such as nuclear fission, alpha decay, cluster decay, and cosmic ray spallation.

As of 2010, there are 118 known elements (in this context, "known" means observed well enough, even from just a few decay products, to have been differentiated from any other element).[11][12] Of these 118 elements, 94 occur naturally on Earth[citation needed]. Six of these occur in extreme trace quantities: technetium, atomic number 43; promethium, number 61; astatine, number 85; francium, number 87; neptunium, number 93; and plutonium, number 94. These 94 elements, and also possibly element 98 californium, have been detected in the universe at large, in the spectra of stars and also supernovae, where short-lived radioactive elements are newly being made. The first 94 elements have been detected directly on Earth as naturally-occurring fission or transmutation products of uranium and thorium. Some californium may be present on Earth, but at present, natural californium is only known from supernovae spectra.

The remaining 24 elements, not found on Earth or in astronomical spectra, have been derived artificially. All of the elements that are derived solely through artificial means are radioactive with very short half-lives; if any atoms of these elements were present at the formation of Earth, they are extremely likely to have already decayed, and if present in novae, have been in quantities too small to have been noted. Technetium was the first purportedly non-naturally occurring element to be synthesized, in 1937, although trace amounts of technetium have since been found in nature, and the element may have been discovered naturally in 1925. This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally occurring trace elements.

Lists of the elements are available by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as Ionization energies of the elements. The nuclides of stable and radioactive elements are also available as a list of nuclides, sorted by length of half life for those that are unstable. One of the most convenient, and certainly the most traditional presentation of the elements, is in form of periodic table, which groups elements with similar chemical properties (and usually also similar electronic structures) together.

Atomic number

The atomic number of an element, Z, is equal to the number of protons that defines the element. For example, all carbon atoms contain 6 protons in their nucleus; so the atomic number "Z" of carbon is 6. Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known as isotopes of the element.

The number of protons in the atomic nucleus also determines its electric charge, which in turn determines the electrons of the atom in its non-ionized state. This in turn (by means of the Pauli exclusion principle) determines the atom's various chemical properties. So all carbon atoms, for example, ultimately have identical chemical properties because they all have the same number of protons in their nucleus, and therefore have the same atomic number. It is for this reason that atomic number rather than mass number (or atomic weight) is considered the identifying characteristic of an element.

Atomic mass

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a super-index on the left hand side of the atomic symbol (e.g., 238U).

The relative atomic mass of an element is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit (u). This number may be a fraction that is not close to a whole number, due to the averaging process. On the other hand, the atomic mass of a pure isotope is quite close to its mass number. Whereas the mass number is a natural (or whole) number, the atomic mass of a single isotope is a real number that is close to a natural number. In general, it differs slightly from the mass number as the mass of the protons and neutrons is not exactly 1 u, the electrons also contribute slightly to the atomic mass, and because of the nuclear binding energy. For example, the mass of 19F is 18.9984032 u. The only exception to the atomic mass of an isotope not being a natural number is 12C, which has a mass of exactly 12, because u is defined as 1/12th of the mass of a free carbon-12 atom.

Isotopes

Isotopes are atoms of the same element (that is, with the same number of protons in their atomic nucleus), but having different numbers of neutrons. Most (66 of 94) naturally occurring elements have more than one stable isotope. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons in the nucleus, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, the three isotopes of carbon are known as carbon-12, carbon-13, and carbon-14, often abbreviated to 12C, 13C, and 14C. Carbon in everyday life and in chemistry is a mixture of 12C, 13C, and 14C atoms.

Except in the case of the isotopes of hydrogen (which differ greatly from each other in relative mass—enough to cause chemical effects), the isotopes of the various elements are typically chemically nearly indistinguishable from each other. For example, the three naturally occurring isotopes of carbon have essentially the same chemical properties, but different nuclear properties. In this example, carbon-12 and carbon-13 are stable atoms, but carbon-14 is unstable; it is radioactive, undergoing beta decay into nitrogen-14.

As illustrated by carbon, all of the elements have some isotopes that are radioactive (radioisotopes), which decay into other elements upon radiating an alpha or beta particle. Certain elements only have radioactive isotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one stable isotope, and the mean number of stable isotopes for the 80 stable elements is 3.1 stable isotopes per element. The largest number of stable isotopes that occur for an element is 10 (for tin, element 50).

Allotropes

Atoms of pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple structures (spacial arrangements of atoms), known as allotropes, which differ in their properties. For example, carbon can be found as diamond, which has a tetrahedral structure around each carbon atom; graphite, which has layers of carbon atoms with a hexagonal structure stacked on top of each other; graphene, which is a single layer of graphite that is incredibly strong; fullerenes, which have nearly spherical shapes; and carbon nanotubes, which are tubes with a hexagonal structure (even these may differ from each other in electrical properties). The ability for an element to exist in one of many structural forms is known as 'allotropy'.

Standard state

The standard state, or reference state, of an element is defined as its thermodynamically most stable state at 1 bar at a given temperature (typically at 298.15 K). In thermochemistry, an element is defined to have an enthalpy of formation of zero in its standard state. For example, the reference state for carbon is graphite, because it is more stable than the other allotropes.

Nomenclature and symbols

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When these facts were learned, the existing names for anciently-known elements (e.g., gold, mercury, iron) were kept in most countries. National differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For a few illustrative examples: German speakers use "Wasserstoff" (water substance) for "hydrogen", "Sauerstoff" (acid substance) for "oxygen" and "Stickstoff" (smothering substance) for "nitrogen", while English and some romance languages use "sodium" for "natrium" and "potassium" for "kalium", and the French, Italians, Greeks, Portuguese and Poles prefer "azote/azot/azoto" (from roots meaning "no life") for "nitrogen".

For purposes of international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that "aluminium" and "caesium" take the place of the U.S. spellings "aluminum" and "cesium", while the U.S. "sulfur" takes the place of the British "sulphur". Chemicals that are practical to sell in bulk in many countries, however, still have national names. Also, those countries that do not use the Latin alphabet to write their national language, cannot be expected to use the IUPAC name for elements.

In languages such as German, where all nouns of all types are capitalized, chemical element names are capitalized. However, chemical elements are not proper nouns in English, according to IUPAC. Thus, in English, the full name of an element is not capitalized, even if it is derived from a proper noun such as the elements californium or einsteinium. Thus, these elements are only capitalized in English when as words they would be capitalized by some other rule, such as to begin a sentence, or as a heading or title. There is a tendency of some English sources to capitalize rare elements, such as uranium, but this is incorrect. Isotopes of chemical elements are also uncapitalized if written out: carbon-12 or uranium-235.

Symbols of chemical elements, however, are capitalized: thus the symbols for the elements just discussed are Cf and Es; C-12 and U-235. IUPAC prefers that isotope symbols be written in superscript notation, however: 12C and 235U.

In the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have a half life too short for them to remain in any appreciable amounts. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question that delayed naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy).

Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century. For example, lutetium was named in reference to Paris, France. The Germans were reluctant to relinquish naming rights to the French, often calling it cassiopeium. Similarly, the British discoverer of niobium originally named it columbium, in reference to the New World. It was used extensively as such by American publications prior to international standardization.

Chemical symbols

For the listing of current and not used Chemical symbols, and other symbols that look like chemical symbols, please see List of elements by symbol.

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules.

The current system of chemical notation was invented by Berzelius. In this typographical system chemical symbols are not used as mere abbreviations - though each consists of letters of the Latin alphabet - they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully universal; since Latin was the common language of science at that time, they were abbreviations based on the Latin names of metals - Cu comes from Cuprum, Fe comes from Ferrum, Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Later chemical elements were also assigned unique chemical symbols, based on the name of the element, but not necessarily in English. For example, sodium has the chemical symbol 'Na' after the Latin natrium. The same applies to "W" (wolfram) for tungsten, "Fe" (ferrum) for Iron, "Hg" (hydrargyrum) for mercury, "Sn" (stannum) for tin, "K" (kalium) for potassium, "Au" (aurum) for gold, "Ag" (argentum) for silver, "Pb" (plumbum) for lead, and "Sb" (stibium) for antimony.

Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not be confused with a roman numeral.

The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always lower case (small letters).

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical, meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

Isotope symbols

The three main isotopes of the element hydrogen are often written as H for protium, D for deuterium and T for tritium. This is in order to make it easier to use them in chemical equations, as it replaces the need to write out the mass number for each atom. E.g. the formula for heavy water may be written D2O instead of ²H2O.

The periodic table

The periodic table of the chemical elements is a tabular method of displaying the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869. Mendeleev intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.

The periodic table is now ubiquitous within the academic discipline of chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behavior. The table has also found wide application in physics, biology, engineering, and industry. The current standard table contains 118 confirmed elements as of April 10, 2010.

Abundance

During the early phases of the Big Bang, nucleosynthesis of hydrogen nuclei resulted in the production of hydrogen and helium isotopes, as well as very minuscule amounts (on the order of 10−10) of lithium and beryllium. There evidence that some boron was produced in the Big Bang, since it has been observed in some very young stars, even though carbon has not.[13] There is agreement that no heavier elements than boron were produced in the Big Bang. As a result, the primordial abundance of atoms consisted of roughly 75% 1H, 25% 4He, and 0.01% deuterium.[14] Subsequent enrichment of galactic halos occurred due to stellar nucleosynthesis and supernova nucleosynthesis.[15] However intergalactic space can still closely resemble the primordial abundance, unless it has been enriched by some means.

The following graph (note log scale) shows abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass.[16] Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. As physical laws and processes appear common throughout the visible universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

ElementParts per million
by mass
Hydrogen 739,000
Helium 240,000
Oxygen 10,400
Carbon 4,600
Neon 1,340
Iron 1,090
Nitrogen 960
Silicon 650
Magnesium 580
Sulfur 440
Potassium 210
Nickel 100

Recently discovered elements

The first transuranium element (element with atomic number greater than 92) discovered was neptunium in 1940. As of February 2010, only the elements up to 112, copernicium, have been confirmed as discovered by IUPAC, while more or less reliable claims have been made for synthesis of elements 113, 114, 115, 116, 117[17] and 118. The discovery of element 112 was acknowledged in 2009, and the name 'copernicium' and the atomic symbol 'Cn' were suggested for it.[18] The name and symbol were officially endorsed by IUPAC on February 19, 2010.[19] The heaviest element that is believed to have been synthesized to date is element 118, ununoctium, on October 9, 2006, by the Flerov Laboratory of Nuclear Reactions in Dubna, Russia.[20][21] Element 117 was the latest element claimed to be discovered, in 2009.[17]

A 2008 claim by Amnon Marinov and six other researchers for discovery of naturally existing element 122[22][23] was criticized by other scientists, and turned down without peer review by both Nature and Nature Physics.[24]

See also

References

  1. ^ International Union of Pure and Applied Chemistry. "[1]". Compendium of Chemical Terminology Internet edition.
  2. ^ A. Earnshaw, Norman Greenwood (1997). Chemistry of the Elements (2 ed.). Butterworth-Heinemann. 
  3. ^ http://books.google.com/books?id=xSjvowNydN8C&lpg=PP1&ots=eRla--Y6Ul&dq=Plato%20timaeus&pg=PA45#v=onepage&q=cube&f=false
  4. ^ Hillar, Marian (2004). "The Problem of the Soul in Aristotle's De anima". NASA WMAP. http://www.socinian.org/aristotles_de_anima.html. Retrieved 2006-08-10. 
  5. ^ Partington, J.R. (1937). A Short History of Chemistry. New York: Dover Publications, Inc.. ISBN 0486659771. 
  6. ^ a b Boyle, Robert (1661). The Sceptical Chymist. London. ISBN 0922802904. 
  7. ^ Lavoisier, Antoine Laurent (1790). Elements of chemistry translated by Robert Kerr. Edinburgh. pp. 175–176. ISBN 9780415179140. http://books.google.com/?id=4BzAjCpEK4gC&pg=PA175. 
  8. ^ Carey, George, W. (1914). The Chemistry of Human Life. Los Angeles. ISBN 0766128407. 
  9. ^ Glanz, James (2010-04-06). "Scientists Discover Heavy New Element". The New York Times. http://www.nytimes.com/2010/04/07/science/07element.html?hp. 
  10. ^ Gaitskell, R; Utyonkov, V. K.; Lobanov, Yu. V.; Abdullin, F. Sh.; Polyakov, A. N.; Sagaidak, R. N.; Shirokovsky, I. V.; Tsyganov, Yu. S. et al. (2006). "Evidence for Dark Matter" (PDF). Physical Review C 74: timeline on page 10. doi:10.1103/PhysRevC.74.044602. http://gaitskell.brown.edu/physics/talks/0408_SLAC_SummerSchool/Gaitskell_DMEvidence_v16.pdf. Retrieved 2008-10-08. 
  11. ^ Sanderson, Katherine (17 October 2006). "Heaviest element made - again". nature@news.com. Nature. http://www.nature.com/news/2006/061016/full/061016-4.html. Retrieved 2006-10-19. 
  12. ^ Phil Schewe and Ben Stein (17 October 2006). "Elements 116 and 118 Are Discovered". Physics News Update. American Institute of Physics. http://www.aip.org/pnu/2006/797.html. Retrieved 2006-10-19. 
  13. ^ Hubble Observations Bring Some Surprises - New York Times
  14. ^ Wright, Edward L. (September 12, 2004). "Big Bang Nucleosynthesis". UCLA Division of Astronomy. http://www.astro.ucla.edu/~wright/BBNS.html. Retrieved 2007-02-22. 
  15. ^ G. Wallerstein, I. Iben Jr., P. Parker, A. M. Boesgaard, G. M. Hale, A. E. Champagne, C. A. Barnes, F. KM-dppeler, V. V. Smith, R. D. Hoffman, F. X. Timmes, C. Sneden, R.N. Boyd, B.S. Meyer, D.L. Lambert (1999). "Synthesis of the elements in stars: forty years of progress" (pdf). Reviews of Modern Physics 69 (4): 995–1084. doi:10.1103/RevModPhys.69.995. http://www.cococubed.com/papers/wallerstein97.pdf. Retrieved 2006-08-04. 
  16. ^ Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. http://kencroswell.com/alchemy.html. 
  17. ^ a b http://www.jinr.ru/img_sections/PAC/NP/31/PAK_NP_31_recom_eng.pdf
  18. ^ "IUPAC Announces Start of the Name Approval Process for the Element of Atomic Number 112". 20 July 2009. http://media.iupac.org/news/112_Naming_Process_20090720.pdf. Retrieved 2009-08-27. 
  19. ^ IUPAC (International Union of Pure and Applied Chemistry): Element 112 is Named Copernicium
  20. ^ Phil Schewe and Ben Stein (17 October 2006). "Elements 116 and 118 Are Discovered". Physics News Update. American Institute of Physics. http://www.aip.org/pnu/2006/797.html. Retrieved 2006-10-19. 
  21. ^ Oganessian, Yu. Ts. et al.; Utyonkov, V.; Lobanov, Yu.; Abdullin, F.; Polyakov, A.; Sagaidak, R.; Shirokovsky, I.; Tsyganov, Yu. et al. (2006-10-09). [Expression error: Unexpected < operator "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245Cm+48Ca fusion reactions"]. Physical Review C 74 (4): 044602. doi:10.1103/PhysRevC.74.044602. 
  22. ^ Marinov, A.; Rodushkin, I.; Kolb, D.; Pape, A.; Kashiv, Y.; Brandt, R.; Gentry, R. V.; Miller, H. W. (2008). "Evidence for a long-lived superheavy nucleus with atomic mass number A=292 and atomic number Z=~122 in natural Th". ArXiv.org 74: 044602. doi:10.1103/PhysRevC.74.044602. http://arxiv.org/abs/0804.3869. Retrieved 2008-04-28. 
  23. ^ Battersby, Stephen (2008-05-02). "Has the heaviest element been found?". NewScientist. http://www.newscientist.com/article/dn13828-has-the-heaviest-element-been-found.html. Retrieved 2009-05-01. 
  24. ^ Van Noorden, Richard (2008-05-02). "Heaviest element claim criticised". Chemistry World (Royal Society of Chemistry). http://www.rsc.org/chemistryworld/News/2008/May/02050802.asp. Retrieved 2009-05-01. 

Further reading

  • E.R. Scerri, The Periodic Table, Its Story and Its Significance, Oxford University Press, NY, 2007.

External links


Study guide

Up to date as of January 14, 2010

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A chemical element is a substance that cannot be changed into any other substance by a normal chemical reaction.

An element is made up of atoms containing electrons, protons and neutrons. Each atom in an element will contain the same number of protons and electrons, for example every carbon atom has six protons and six electrons.

The periodic table of the elements contains lists all of the elements in order of their number of protons and is laid out so that the elements in each group (vertical column) will have similar properties.

Just because two substances contain the same atoms does not always mean that they will have the same properties. Graphite and diamonds are both made completely from carbon atoms, but the way in which they join with each other means that the have different properties.

Atoms of the same element may also contain different numbers of neutrons. Atoms of the same element with different numbers of neutrons are called isotopes.


Simple English

A chemical element is a substance that cannot be divided into smaller chemical parts. An element can be a solid, liquid or gas. The smallest particle of such an element is an atom. Atoms are made up of protons, neutrons, and electrons.

A pure element contains only one kind of atom. The number of protons in the atoms of an element is known as that element's atomic number. For example, all atoms with 6 protons in their nuclei are of the chemical element carbon, and all atoms with 92 protons in their nuclei are of the element uranium.

117 different chemical elements are known to modern chemistry. 90 of these elements can be found in nature, and the others can only be made in laboratories. The last natural element discovered was uranium, in 1789.[1] The first man-made element was Technetium in 1937. All man-made elements are radioactive and unstable.

Chemical elements are commonly arranged in the periodic table.

Atoms of the same element, whose nuclei contain a different number of neutrons, are said to be different isotopes of the element.

Chemical elements are also given a unique chemical symbol. Chemical symbols of elements come from their English or Latin names (For example, carbon has the chemical symbol 'C', and sodium has chemical symbol 'Na', after the Latin natrium). Tungsten is called "W" after its German name, wolfram.[1] Others were named after famous people, like einsteinium was named after Albert Einstein.[1]

Elements can join (react) to form pure compounds (such as water, salts, oxides, and organic compounds). In many cases, these compounds have a fixed composition and their own structure and properties.

Some elements, particularly metal elements mix together in any proportion to form new structures. Such new structures are not compounds. They are called mixtures.

= Classification

= Elements can be classified based on physical states. At room temperature and pressure, most elements are solids, only 11 are gases and 2 are liquids.

Elements can also be classified into metals and non-metals. There are many more metals than non-metals.

However, a few elements have properties in between those of metals and non-metals. These elements are called semimetals (or metalloids).

Other pages

References

  1. 1.0 1.1 1.2 Devlin, Jacinta; Cochrane, Helen (2005). "1" (in English). Science Links 2. Coffey, Rhonda (VELS Edition ed.). Heinemann. pp. 5. ISBN I740815203. 
rue:Хемічный елемент







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