A chemical element, often called simply an element, is a substance that cannot be decomposed or transformed into other chemical substances by ordinary chemical processes. All matter fundamentally consists of these elements and as of 2006, 116 unique elements have been discovered or artificially created. The smallest particle of such an element is an atom, which consists of electrons centered about a nucleus of protons and neutrons.

Chemistry terminology

Earlier an element or pure element was defined as a substance which "cannot be further broken down into another compound with different chemical properties" -- which should be taken to mean it consists of atoms of one element. However, due to allotropy, the isotope effect, and the confusion with the more useful term referring to the general class of atoms (irrespective of what compound it may be in), this usage is in disfavor amongst contemporary chemists, and sees restricted, mostly historical, use. This definition was motivated by the observation that these elements could not be dissociated by chemical means into other compounds. For example, water could be converted into hydrogen and oxygen, but hydrogen and oxygen could not be further decomposed, thus "elemental". There are also many counterexamples (for example "elemental oxygen" (O2) can be decomposed by solely chemical means into oxygen ions and atoms which have drastically different chemical properties). This article will concern itself with the latter definition.

Description

The lightest elements are hydrogen and helium. Hydrogen is thought to have been the first element to appear after the Big Bang. All the heavier elements are made, both naturally and artificially, through various methods of nucleosynthesis. As of 2005, there are 116 known elements: 93 occur naturally on earth (including technetium and plutonium), and 94 (including promethium) have been detected in the universe at large. The 23 elements not found on earth are derived artificially; 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. All artificially derived elements are radioactive with short half-lives, so if any atoms of these elements were present at the formation of Earth are extremely likely to have already decayed.

Lists of the elements by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as Ionization energies of the elements are available. 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 which defines the element. For example, all carbon atoms contain 6 protons in their nucleus, so for carbon Z=6. These atoms may have different amounts of neutrons, and are known as isotopes of the element. The atomic mass of an element, A, is measured in unified atomic mass units (u) is the average mass of all the atoms of the element in an environment of interest (usually the earth's crust and atmosphere). Since electrons are of negligible mass, and neutrons are barely more than the mass of the proton, this usually corresponds to the sum of the protons and neutrons in the nucleus of the most abundant isotope, though this is not always the case (notably chlorine, which is about three-quarters 35Cl and a quarter 37Cl).


Atomic mass

The atomic masses that are given on the periodic table are calculated by the following method. As an example, assume there exists three isotopes of element X and their respective atomic masses are 10, 20 and 30 AMU for sake of demonstration. Now also assume that 50% of the isotopes of element X are the 10 AMU version and the two heavier isotopes each account for 25% of the total number of atoms (particles) of this hypothetical element. As a result 10 * 0.5 = 5 AMU and 20 * 0.25= 5 AMU and 30 * 0.25 = 7.5 AMU. The average atomic mass that results is 17.5 AMU. The reason is because the method to calculate the average mass takes into account the relative abundance of all of the isotopes of an element, which is multiplied against their individual masses.


Isotopes

Some isotopes are radioactive and decay into other elements upon radiating an alpha or beta particle. Some elements have no nonradioactive isotopes, in particular all elements with atomic numbers greater than 82.
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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" and "Sauerstoff" for "oxygen," while English and some romance languages use "sodium" for "natrium" and "potassium" for "kalium," and the French prefer the term "azote" for "nitrogen." This is also used by the Greeks.

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." But chemicals which are practicable to be sold in bulk within many countries, however, still have national names, and those which 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 (unless it would be capitalized by some other rule, for instance if it begins a sentence).

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 which delayed the 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.