   #copyright

Atom

2007 Schools Wikipedia Selection. Related subjects: General Chemistry

                                                                       Atom

     Helium atom ground state

     An accurate depiction of the atomic structure of the helium atom. The
           darkness of the electron cloud corresponds to the line-of-sight
    integral over the probability function of the 1s electron orbital. The
    magnified nucleus is schematic, showing protons in pink and neutrons in
      purple. In reality, the nucleus (and the wavefunction of each of the
    nucleons) is also spherically symmetric. (For more complex nuclei this
                                                          is not the case.)
                                                             Classification


   Smallest recognised division of a chemical element

                                                                 Properties


         Mass :               ≈ 1.67 × 10^-27 to 4.52 × 10^-25 kg
    Electric charge : zero(if the number of electrons equal of protons in
                                           an atom)
   Diameter : (approx.
        100 pm =
       1 Angstrom)                 100 pm(He) to 670 pm(Cs)

   In chemistry and physics, an atom ( Greek ἄτομος or átomos meaning
   "indivisible") is the smallest particle of a chemical element that
   retains its chemical properties. (Since until the advent of quantum
   mechanics dividing a material object was invariably equated with
   cutting it, átomos is usually translated as "indivisible".) Where as
   the word atom originally denoted a particle that cannot be cut into
   smaller particles, the atoms of modern parlance are composed of
   subatomic particles:
     * electrons, which have a negative charge and are smallest of the
       three;
     * protons, which have a positive charge and are about 1836 times
       bigger than electrons; and
     * neutrons, which have no charge and are about 1839 times bigger than
       electrons.

   Protons and neutrons make up a dense, massive atomic nucleus, and are
   collectively called nucleons. The electrons form the much larger
   electron cloud surrounding the nucleus.

   Atoms can differ in the number of each of the subatomic particles they
   contain. Atoms of the same element have the same number of protons
   (called the atomic number). Within a single element, the number of
   neutrons may vary, determining the isotope of that element. The number
   of electrons associated with an atom is most easily changed, due to the
   lower energy of binding of electrons. The number of protons (and
   neutrons) in the atomic nucleus may also change, via nuclear fusion,
   nuclear fission or radioactive decay, in which case the atom is no
   longer the same element it was.

   Atoms are electrically neutral if they have an equal number of protons
   and electrons. Atoms which have either a deficit or a surplus of
   electrons are called ions. Electrons that are furthest from the nucleus
   may be transferred to other nearby atoms or shared between atoms. By
   this mechanism atoms are able to bond into molecules and other types of
   chemical compounds like ionic and covalent network crystals.

   Atoms are the fundamental building blocks of chemistry, and are
   conserved in chemical reactions.

Atoms and molecules

   For gases and certain molecular liquids and solids (such as water and
   sugar), molecules are the smallest division of matter which retains
   chemical properties; however, there are also many solids and liquids
   which are made of atoms, but do not contain discrete molecules (such as
   salts, rocks, and liquid and solid metals). Thus, while molecules are
   common on Earth (making up all of the atmosphere and most of the
   oceans), most of the mass of the Earth (much of the crust, and all of
   the mantle and core) is not made of identifiable molecules, but rather
   represents atomic matter in other arrangements, all of which lack the
   particular type of small-scale order that is associated with molecules.

   Most molecules are made up of multiple atoms; for example, a molecule
   of water is a combination of two hydrogen atoms and one oxygen atom.
   The term "molecule" in gases has been used as a synonym for the
   fundamental particles of the gas, whatever their structure. This
   definition results in a few types of gases (for example inert elements
   that do not form compounds, such as helium), having "molecules"
   consisting of only a single atom.

History of atomic theory and discovery of atomic structure

   Philosophical atomic ruminations date back to the ancient Greeks and
   Indians in the fifth and sixth centuries BCE. It was the Greeks (
   Democritus; see below) who coined the term atomos, which meant
   "uncuttable".

   The first philosophical statements relating to an idea similar to atoms
   was developed by Democritus in Greece in the fifth century BCE around
   450 BCE. The idea was lost for centuries until scientific interest was
   rekindled during the Renaissance Period.
   Various atoms and molecules as depicted in John Dalton's A New System
   of Chemical Philosophy (1808).
   Enlarge
   Various atoms and molecules as depicted in John Dalton's A New System
   of Chemical Philosophy (1808).

   In 1803, John Dalton used the concept of atoms to explain why elements
   always reacted in simple proportions, and why certain gases dissolved
   better in water than others. He proposed that each element consists of
   atoms of a single, unique type, and that these atoms could join up to
   form compound chemicals.

   In 1897 JJ Thomson, through his work on cathode rays, discovered the
   electron and their subatomic nature, which destroyed the concept of
   atoms as being indivisible units. Thomson would also later discover the
   existence of isotopes through his work on ionized gases.

   Thomson believed that the electrons were distributed evenly throughout
   the atom, balanced by the presence of a uniform sea of positive charge.
   However, in 1909, Rutherford's gold foil experiment suggested that the
   positive charge of an atom and most of its mass was concentrated in a
   nucleus at the centre of the atom, with the electrons orbiting it like
   planets around a sun. In 1913, Niels Bohr added quantum mechanics into
   this model, which now stated that the electrons were confined to
   clearly defined orbits and could not freely spiral in or out.

   In 1926, Erwin Schrodinger proposed that electrons behave not like
   particles, but like waves. A consequence of this notion, pointed out by
   Werner Heisenberg a year later, is that it is mathematically impossible
   to obtain precise values for both the position and momentum of a
   particle at any point in time; this became known as the Uncertainty
   principle. Instead, for any given value of position you could only
   obtain a range of probable values for momentum, and vice versa. Thus,
   the planetary model of the atom was discarded in favour of one that
   described zones around the nucleus where a given electron is most
   likely to exist.

Properties of the atom in present theory

Subatomic particles

   This model of a helium atom shows the electrons (yellow), the protons
   (grey), and the neutrons (pink). Also shown are the up quarks (red),
   and the down quarks (blue) that make up the nucleons as well as the
   gluons (black) which hold the quarks together.
   Enlarge
   This model of a helium atom shows the electrons (yellow), the protons
   (grey), and the neutrons (pink). Also shown are the up quarks (red),
   and the down quarks (blue) that make up the nucleons as well as the
   gluons (black) which hold the quarks together.

   Although the name "atom" was applied at a time when atoms were thought
   to be indivisible, it is now known that the atom can be broken down
   into a number of smaller components. The first of these to be
   discovered was the negatively charged electron, which is easily ejected
   from atoms during ionization. The electrons (or more specifically,
   electron clouds) orbit a small, dense body containing all of the
   positive charge in the atom, called the atomic nucleus. This nucleus is
   itself made up of nucleons: positively charged protons and chargeless
   neutrons.

   Before 1961, the subatomic particles were thought to consist of only
   protons, neutrons and electrons. However, protons and neutrons
   themselves are now known to consist of still smaller particles called
   quarks. In addition, the electron is known to have a nearly massless
   neutral partner called a neutrino. Together, the electron and neutrino
   are both leptons.

   Ordinary atoms are composed only of quarks and leptons of the first
   generation. The proton is composed of two up quarks and one down quark,
   whereas the neutron is composed of one up quark and two down quarks.
   Although they do not occur in ordinary matter, two other heavier
   generations of quarks and leptons may be generated in high-energy
   collisions.

   The subatomic force carrying particles (called gauge bosons) are also
   important to atoms. Electrons are bound to the nucleus by photons
   carrying the electromagnetic force. Protons and neutrons are bound
   together in the nucleus by gluons carrying the strong nuclear force.

Electron configuration

   The nucleus of an atom is surrounded by a cloud of electrons, and it is
   primarily the interaction of these clouds that govern the chemical
   behaviour of atoms. A popular concept is that the electrons orbit the
   atom in neat circles like planets around a sun, but this is an obsolete
   model that is nonetheless still taught to schoolchildren because it is
   simpler and sufficient for school-level chemistry. In the true modern
   model of the atom, the positions of the electrons around the atom's
   nucleus are described through probabilities—that is, an electron can
   theoretically be found at any arbitrary position around the nucleus,
   but is more likely to be found in certain regions than others. This
   pattern is referred to as its atomic orbital and the shape of its
   orbital depends on its energy level (or, more specifically, its quantum
   state).
   The five atomic orbitals of a neon atom, separated and arranged in
   order of increasing energy. Each orbital holds up to two electrons,
   which exist for most of the time in the zones represented by the
   bubbles.
   The five atomic orbitals of a neon atom, separated and arranged in
   order of increasing energy. Each orbital holds up to two electrons,
   which exist for most of the time in the zones represented by the
   bubbles.

   Each atomic orbital can hold up to two electrons. The orbitals are
   organized into shells and subshells, based on their overall energy and
   angular momentum. Generally speaking, higher energy shells can hold
   more electrons and are located farther from the nucleus. A shell can
   hold up to 2n^2 electrons (where n is the shell number). The electrons
   in the outermost shell, called the valence electrons, have the greatest
   influence on chemical behaviour. Core electrons (those not in the outer
   shell) play a role, but it is usually in terms of a secondary effect
   due to screening of the positive charge in the atomic nucleus.

   In the most stable ground state, an atom's electrons will fill up its
   orbitals in order of increasing energy. Under some circumstances an
   electron may be excited to a higher energy level (that is, it absorbs
   energy from an external source and leaps to a higher shell), leaving a
   space in a lower shell. An excited atom's electrons will spontaneously
   fall back to lower levels, emitting the energy it had gained as a
   photon. This behaviour is the root of a substance's absorption
   spectrum.

Nucleon properties

   The constituent protons and neutrons of the atomic nucleus are
   collectively called nucleons. The nucleons are held together in the
   nucleus by the strong nuclear force which is carried by gluons.

   Nuclei can undergo transformations that affect the number of protons
   and neutrons they contain, a process called radioactive decay. When
   nuclei transformations take place spontaneously, this process is called
   radioactivity. Radioactive transformations proceed by a wide variety of
   modes, but the most common are alpha decay (emission of a helium
   nucleus) and beta decay (emission of an electron). Decays involving
   electrons or positrons are due to the weak nuclear interaction.

   In addition, like the electrons of the atom, the nucleons of nuclei may
   be pushed into excited states of higher energy. However, these
   transitions typically require thousands of times more energy than
   electron excitations. When an excited nucleus emits a photon to return
   to the ground state, the photon has very high energy and is called a
   gamma ray.

   Nuclear transformations also take place in nuclear reactions. In
   nuclear fusion, two light nuclei come together and merge into a single
   heavier nucleus. In nuclear fission, a single large nucleus is divided
   into two or more smaller nuclei.

Atom size and speed

   depiction of a hydrogen atom showing the Van der Waals radius. (Image
   not to scale)
   Enlarge
   depiction of a hydrogen atom showing the Van der Waals radius. (Image
   not to scale)

   Atoms are much smaller than the wavelengths of light that human vision
   can detect, so atoms cannot be seen in any kind of optical microscope.
   However, there are ways of detecting the positions of atoms on the
   surface of a solid or a thin film so as to obtain images. These
   include: electron microscopes (such as in scanning tunneling microscopy
   (STM)), atomic force microscopy (AFM), nuclear magnetic resonance (NMR)
   and x-ray microscopy.

   Since the electron cloud does not have a sharp cutoff, the size of an
   atom is not easily defined. For atoms that can form solid crystal
   lattices, the distance between the centers of adjacent atoms can be
   easily determined by x-ray diffraction, giving an estimate of the
   atoms' size. For any atom, one might use the radius at which the
   electrons of the valence shell are most likely to be found. As an
   example, the size of a hydrogen atom is estimated to be approximately
   1.06×10^-10 m (twice the Bohr radius). Compare this to the size of the
   proton (the only particle in the nucleus of the hydrogen atom), which
   is approximately 10^-15 m. So the ratio of the size of the hydrogen
   atom to its nucleus is about 100,000:1. If an atom were the size of a
   stadium, the nucleus would be the size of a marble. If an atom were the
   size of the United States, an electron would be 3cm long and wide.
   Nearly all the mass of an atom is in its nucleus, yet almost all the
   space in an atom is occupied by its electrons.

   Atoms of different elements do vary in size, but the sizes (volumes) do
   not scale well with the mass of the atom. Heavier atoms do tend
   generally to be more dense. The diameters of atoms are roughly the same
   to within a factor of less than three for the heavier atoms, and the
   most noticeable effect on size with atomic mass is a reverse one:
   atomic size actually shrinks with increasing mass in each periodic
   table row . The reason for these effects is that heavy elements have
   large positive charge on their nuclei, which strongly attract the
   electrons to the centre of the atom. This contracts the size of the
   electron shells, so that more electrons may fit into a smaller volume.
   These effects may be striking: for example, atoms of the densest
   element iridium (atomic weight about 192) are about the same size as
   aluminium atoms (atomic weight about 27), and this contributes greatly
   to the density ratio of more than eight between these metals.

   The temperature of a collection of atoms is a measure of the average
   energy of motion of those atoms above the minimum zero-point energy
   demanded by quantum mechanics; at 0 kelvins (absolute zero) atoms would
   have no extra energy above the minimum. As the temperature of the
   system is increased, the kinetic energy of the particles in the system
   is increased, and their speed of motion increases. At room temperature,
   atoms making up gases in the air move at an average speed of 500 m/s
   (about 1100 mph or 1800 km/h).

Elements, isotopes and ions

   Atoms with the same atomic number Z share a wide variety of physical
   properties and exhibit almost identical chemical properties (for the
   closest instance to an exception to this principle, see deuterium and
   heavy water). Atoms are classified into chemical elements by their
   atomic number Z, which corresponds to the number of protons in the
   atom. For example, all atoms containing six protons (Z = 6) are
   classified as carbon. The elements may be sorted according to the
   periodic table in order of increasing atomic number. When this is done,
   certain repeating cycles of regularities in chemical and physical
   properties are evident.

   The mass number A, or nucleon number of an element is the total number
   of protons and neutrons in an atom of that element, so-called because
   each proton and neutron has a mass of about 1  amu. A particular
   collection of a certain number of protons Z, and neutrons A-Z, is
   called a nuclide.

   Each element can have numerous different nuclides with the same Z
   (number of protons and electrons) but varying numbers of neutrons. Such
   a family of nuclides are called the isotopes of the element (isotope =
   "same place", because these nuclides share the same chemical symbol and
   place on the periodic table). When writing the name of a particular
   nuclide, the element name (which specifies the Z) is preceded by the
   mass number if written as a superscript, or else followed by the mass
   number if not a superscript. For example, the nuclide carbon-14, which
   may also be written ^14C, is one of the isotopes of carbon, and it
   contains 6 protons and 8 neutrons in each atom, for a total mass number
   of 14. For a complete table of known nuclides, including radioactive
   and stable nuclides, see isotope table (divided).

   The atomic mass listed for each element in the periodic table is an
   average of the isotope masses found in nature, weighted by their
   abundance.

   The simplest atom is the hydrogen isotope protium, which has atomic
   number 1 and atomic mass number 1; it consists of one proton and one
   electron. The hydrogen isotope which also contains one neutron so is
   called deuterium or hydrogen-2; the hydrogen isotope with two neutrons
   is called tritium or hydrogen-3. Tritium is an unstable isotope which
   decays through a process called radioactivity. Many isotopes of each
   element are radioactive; the number which are stable varies greatly
   with the element (tin has 10 stable isotopes; see list of stable
   isotopes). Lead (Z = 82) is the last element which has stable isotopes.
   The elements with atomic number 83 (bismuth) and greater have no stable
   isotopes and are all radioactive.

   Virtually all elements heavier than hydrogen and helium were created
   through stellar nucleosynthesis and supernova nucleosynthesis. The
   solar system is thought to be formed of clouds of elements from many
   such supernovae, which date from more than 4.6 billion years ago. Most
   of the elements lighter than uranium (Z = 92) have either stable
   isotopes, or else radioisotopes long-lived enough to occur naturally on
   Earth. Two notable exceptions of light but short-lived radioactive
   elements are technetium Z = 43 (although some technetium has been found
   on Earth, this occurred only after the element was first synthesized
   artificially), and promethium Z = 61, which is found naturally only in
   stars where it was recently made. Several other short-lived heavy
   elements that do not occur on Earth have been found to be present in
   stars. Elements not normally found in nature have been artificially
   created by nuclear bombardment; as of 2006, elements have been created
   through atomic number 116 (given the temporary name ununhexium). These
   ultra-heavy elements are generally highly unstable and decay quickly.

   Atoms that have lost or gained electrons to become electrically
   non-neutral, are called atomic ions. Ions are divided into cations with
   positive (+) electric charge; or anions with negative (-) charge.

Valence and bonding

   The number of electrons in an atom's outermost shell (the valence
   shell) governs its bonding behaviour. Therefore, elements with the same
   number of valence electrons are grouped together in the columns of the
   periodic table of the elements. Alkali metals contain one electron on
   their outer shell; alkaline earth metals, two electrons; halogens,
   seven electrons; and various others.

   Every atom is most stable with a full valence shell. This means that
   atoms with full valence shells (the noble gases) are very unreactive.
   Conversely, atoms with few electrons in their valence shell are more
   reactive. Alkali metals are therefore very reactive, with caesium,
   rubidium, and francium being the most reactive of all metals. Also,
   atoms that need only few electrons (such as the halogens) to fill their
   valence shells are reactive. Fluorine is the most reactive of all
   elements.

   Atoms may fill their valence shells by chemical bonding. This can be
   achieved one of two ways: an atom can either share electrons with other
   atoms (a covalent bond), or it can remove electrons from (or donate
   electrons to) other atoms (an ionic bond). The formation of a bond
   causes a strong attraction between two atoms, creating molecules or
   ionic compounds. Many other types of bonds exist, including:
     * polar covalent bonds;
     * coordinate covalent bonds;
     * metallic bonds;
     * hydrogen bonds; and
     * van der Waals bonds.

Atomic spectrum

   Since each element in the periodic table consists of an atom in a
   unique configuration with different numbers of protons and electrons,
   each element can also be uniquely described by the energies of its
   atomic orbitals and the number of electrons within them. Normally, an
   atom is found in its lowest-energy ground state; states with higher
   energy are called excited states. An electron may move from a
   lower-energy orbital to a higher-energy orbital by absorbing a photon
   with energy equal to the difference between the energies of the two
   levels. An electron in a higher-energy orbital may drop to a
   lower-energy orbital by emitting a photon. Since each element has a
   unique set of energy levels, each creates its own light pattern unique
   to itself: its own spectral signature.

   If a set of atoms is heated (such as in an arc lamp), their electrons
   will move into excited states. When these atoms fall back toward the
   ground state, they will produce an emission spectrum. If a set of atoms
   is illuminated by a continuous spectrum, it will only absorb specific
   wavelengths (energies) of photon that correspond to the differences in
   its energy levels. The resulting pattern of gaps is called the
   absorption spectrum.

   In spectroscopic analysis, scientists can use a spectrometer to study
   the atoms in stars and other distant objects. Due to the distinctive
   spectral lines that each element produces, they are able to tell the
   chemical composition of distant planets, stars and nebulae.

   Not all parts of the atomic spectrum are in visible light part of the
   electromagnetic spectrum. For example, the hyperfine transitions
   (including the important 21 cm line) produce low-energy radio waves.
   When electrons deep inside atoms of high atomic number are knocked out
   (for example by beta radiation), replacement electrons fall deep into
   the electric potential of the high-Z nucleus, producing high-energy
   x-rays.

Exotic atoms

   An exotic atom is usually made from a normal matter atom with a
   substitution from abnormal or rarely encountered matter, such as
   antimatter, muons, mesons, or other objects. A few exotic atoms (such
   as atoms of antimatter) are not made of any normal atomic constituents
   at all. All exotic atoms (save antiatoms made from antinucleons and
   positrons), are highly unstable, decaying with lifetimes of a few
   microseconds or less. The antimatter counterparts of stable particles
   are also stable, but difficult to store for more than short periods,
   since they annihilate if allowed to contact ordinary matter.

   The most familiar examples of exotic atoms are the antiatom
   antihydrogen (composed of an antiproton and positron) which has been
   produced in tiny quantities, and positronium, an analogue to the
   hydrogen atom in which a positron is substituted for the usual proton
   nucleus. Positronium is unstable; it is a common phase in the
   attraction between an electron and positron before the annihilation
   reaction in which the matter particles are destroyed and two gamma rays
   are emitted.

Atoms and the Big Bang Theory

   In models of the Big Bang, Big Bang nucleosynthesis predicts that
   within one to three minutes of the Big Bang almost all atomic material
   in the universe was created. During this process, nuclei of hydrogen
   and helium formed abundantly, but almost no elements heavier than
   lithium. Hydrogen makes up approximately 75% of the atoms in the
   universe; helium makes up 24%; and all other elements make up just 1%.
   However, although nuclei (fully- ionized atoms) were created, neutral
   atoms themselves could not form in the intense heat.

   Big Bang chronology of the atom continues to approximately 379,000
   years after the Big Bang when the cosmic temperature had dropped to
   just 3,000  K. It was then cool enough to allow the nuclei to capture
   electrons. This process is called recombination, during which the first
   neutral atoms took form. Once atoms become neutral, they only absorb
   photons of a discrete absorption spectrum. This allows most of the
   photons in the universe to travel unimpeded for billions of years.
   These photons are still detectable today in the cosmic microwave
   background.

   After Big Bang nucleosynthesis, no heavier elements could be created
   until the formation of the first stars. These stars fused heavier
   elements through stellar nucleosynthesis during their lives and through
   supernova nucleosynthesis as they died. The seeding of the interstellar
   medium by heavy elements eventually allowed the formation of
   terrestrial planets like the Earth.

Atom size comparisons

   Various analogies have been used to demonstrate the minuteness of the
   atom:
     * A human hair is about 1 million carbon atoms wide.
     * An HIV virus is the width of 800 carbon atoms and contains about
       100 million atoms total. An E. coli bacterium contains perhaps 100
       billion atoms.
     * A speck of dust might contain 3x10^12 (3 trillion) atoms.
     * The number of atoms in 12 grams of charcoal (about 6 x 10^23) is
       more than 1,400,000 times the age of the universe in seconds.

   Retrieved from " http://en.wikipedia.org/wiki/Atom"
   This reference article is mainly selected from the English Wikipedia
   with only minor checks and changes (see www.wikipedia.org for details
   of authors and sources) and is available under the GNU Free
   Documentation License. See also our Disclaimer.
