Atomic Structure & Bonding

Definition: A substance that cannot be broken down by chemical means. Elements are defined by the number of protons they possess.

Examples: copper, cesium, iron, neon

Another definition by wiki
A chemical element is a type of atom that is distinguished by its atomic number; that is, by 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 hydrogen(the first element), carbon, nitrogen, and oxygen. In total, 117 elements have been observed as of 2007, of which 94 occur naturally on Earth. Elements with atomic numbers 83 or higher (bismuth and above) are inherently unstable, and undergo radioactive decay. Of the first 82 elements, 80 have stable isotopes. Elements 43 and 61 (technetium and promethium) have no stable isotopes, and decay. The elements from 83 to atomic number 94 that have no stable nuclei, are nevertheless found in nature, either surviving as remnants of the primordial stellar nucleosynthesis which produced the elements in the solar system, or else as produced newly as short-lived daughter-isotopes in 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.

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 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 which 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 which 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.

All three of the isotopes of carbon have the same chemical properties. But they have different nuclear properties. In this example, carbon-12 and carbon-13 are stable atoms, but carbon-14 is unstable; it is slightly radioactive, decaying over time into other elements.

Like carbon, some isotopes of various elements are radioactive and decay into other elements upon radiating an alpha or beta particle. For certain elements, all their isotopes are 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 a stable isotope, 16 have only one stable isotope, and the mean number of stable isotopes for the 80 stable elements is 3.4 stable isotopes per element. The largest number of stable isotopes that occur for an element is 10 (for tin, element 50).

Allotropes

Some elements can be found as multiple elementary substances, known as allotropes, which differ in their structure and 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, and fullerenes, which have nearly spherical shapes. 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.
Recently discovered element claims

The first transuranium element (element with atomic number greater than 92) discovered was neptunium in 1940. As of August 2007, only the elements up to 111, Roentgenium, have been confirmed as valid by IUPAC, while more or less reliable claims have been made for synthesis of elements 112, 113, 114, 115, 116 and 118. 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.[17][18]

Element 117, ununseptium, has not yet been created or discovered, although its place in the periodic table is preestablished.

According to Amnon Marinov and six other researchers element 122 has been detected naturally occurring in a thorium deposit.[19], this is the first Naturally occurring heavy element in more than 50 years. 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.