History
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).
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.
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.
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. 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. 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.
Description
The lightest elements are hydrogen and helium, both theoretically
created by Big Bang nucleosynthesis during the first 20 minutes of the
universe 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). 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 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 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; 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
12
C, which has a mass of exactly 12, because u is defined as 1/12th of the
mass of a free carbon-12 atom.
