Wikipedia article on Atom

From Wikipedia, the free encyclopedia[Jump to navigation]([Jump to search](*For other uses, see* [*Atom (disambiguation)*](*.Not to be confused with* [*Atum*](*.***Atom**📷An illustration of the [helium]( atom, depicting the [nucleus]( (pink) and the [electron cloud]( distribution (black). The nucleus (upper right) in helium-4 is in reality spherically symmetric and closely resembles the electron cloud, although for more complicated nuclei this is not always the case. The black bar is one [angstrom]( (10−10 m or 100 [pm]( recognized division of a chemical elementProperties[Mass range](×10−27 to 4.52×10−25 kg[Electric charge]( (neutral), or [ion]( charge[Diameter]( range62 pm ([He]( to 520 pm ([Cs]( ([data page]([Components]([Electrons]( and a compact [nucleus]( of [protons]( and [neutrons](

An **atom** is the smallest unit of ordinary [matter]( that forms a [chemical element]([[1]]( Every [solid](, [liquid](, [gas](, and [plasma]( is composed of neutral or [ionized]( atoms. Atoms are extremely small, typically around 100 [picometers]( across. They are so small that accurately predicting their behavior using [classical physics](, as if they were [tennis balls]( for example, is not possible due to [quantum effects](

Every atom is composed of a [nucleus]( and one or more [electrons]( bound to the nucleus. The nucleus is made of one or more [protons]( and a number of [neutrons]( Only the most common variety of [hydrogen]( has no neutrons. More than 99.94% of an atom’s [mass]( is in the nucleus. The protons have a positive [electric charge](, the electrons have a negative electric charge, and the neutrons have no electric charge. If the number of [protons]( and electrons are equal, then the atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively – such atoms are called [ions](

The electrons of an atom are attracted to the protons in an atomic nucleus by the [electromagnetic force]( The protons and neutrons in the nucleus are attracted to each other by the [nuclear force]( This force is usually stronger than the electromagnetic force that repels the positively charged protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus [splits]( and [leaves behind different elements]( This is a form of [nuclear decay](

The number of protons in the nucleus is the [*atomic number*]( and it defines to which chemical element the atom belongs. For example, any atom that contains 29 protons is [copper]( The number of [neutrons]( defines the [isotope]( of the element. For example, a copper atom with 34 neutrons is copper-63 (29+34), and with 36 neutrons is copper-65; natural copper is about 70% Cu-63 and the rest is Cu-65.

Atoms can attach to one or more other atoms by [chemical bonds]( to form [chemical compounds]( such as [molecules]( or [crystals]( For example, [New York City](’s [Statue of Liberty]( was originally made of pure copper, but over the years, the surface combined with [oxygen](, [carbon]( and [sulfur]( atoms to make a green [patina]( on the copper. The ability of atoms to [attach and detach]( is responsible for most of the physical changes observed in nature. [Chemistry]( is the discipline that studies these changes.

History of atomic theory

*Main article:* [*Atomic theory*](

## In philosophy

*Main article:* [*Atomism*](

The basic idea that matter is made up of tiny, indivisible particles is an old idea that appeared in many ancient cultures such as those of [Greece]( and [India]( The word *atom* is derived from the ancient Greek word *atomos*,[[a]]( which means “uncuttable”. This ancient idea was based in philosophical reasoning rather than scientific reasoning; modern atomic theory is not based on these old concepts.[[2]]([[3]]( In the early 19th century, the scientist [John Dalton]( noticed that chemical elements seemed to combine with each other by basic units of weight, and he decided to use the word “atom” to refer to these units on the assumption that these were the fundamental particles of matter. About a century later it was discovered that Dalton’s atoms are not actually indivisible, but the term stuck.

## Dalton’s law of multiple proportions

📷Atoms and molecules as depicted in [John Dalton](’s *A New System of Chemical Philosophy* vol. 1 (1808)

In the early 1800s, the English chemist [John Dalton]( compiled experimental data gathered by himself and other scientists and discovered a pattern now known as the “[law of multiple proportions](”. He noticed that in chemical compounds which contain a particular chemical element, the content of that element in these compounds will differ in weight by ratios of small whole numbers. This pattern suggested to Dalton that each chemical element combines with other elements by a basic and consistent unit of weight, and he decided to call these units “atoms”.

For example, there are two types of [tin oxide]( one is a black powder that is 88.1% tin and 11.9% oxygen, and the other is a white powder that is 78.7% tin and 21.3% oxygen. Adjusting these figures, in the black oxide there is about 13.5 g of oxygen for every 100 g of tin, and in the white oxide there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 1:2. Dalton concluded that in these oxides, for every tin atom there are one or two oxygen atoms respectively ([SnO]( and [SnO2]([[4]]([[5]](

Dalton also analyzed [iron oxides]( There is one type of iron oxide that is a black powder which is 78.1% iron and 21.9% oxygen; and there is another iron oxide that is a red powder which is 70.4% iron and 29.6% oxygen. Adjusting these figures, in the black oxide there is about 28 g of oxygen for every 100 g of iron, and in the red oxide there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. In these respective oxides, for every two atoms of iron, there are two or three atoms of oxygen ([Fe2O2]( and [Fe2O3]([[b]]([[6]]([[7]](

As a final example: [nitrous oxide]( is 63.3% nitrogen and 36.7% oxygen, [nitric oxide]( is 44.05% nitrogen and 55.95% oxygen, and [nitrogen dioxide]( is 29.5% nitrogen and 70.5% oxygen. Adjusting these figures, in nitrous oxide there is 80 g of oxygen for every 140 g of nitrogen, in nitric oxide there is about 160 g of oxygen for every 140 g of nitrogen, and in nitrogen dioxide there is 320 g of oxygen for every 140 g of nitrogen. 80, 160, and 320 form a ratio of 1:2:4. The respective formulas for these oxides are [N2O](, [NO](, and [NO2]([[8]]([[9]](

## Kinetic theory of gases

*Main article:* [*Kinetic theory of gases*](

In 1738 [Daniel Bernoulli]( [[10]]( and a number of other scientists found that they could better explain the behavior of gases by describing them as collections of sub-microscopic particles and modelling their behavior using [statistics]( and [probability]( Unlike Dalton’s atomic theory, the kinetic theory of gases describes not how gases react chemically with each other to form compounds, but how they behave physically: diffusion, viscosity, conductivity, pressure, etc.

## Brownian motion

In 1827, [botanist]( [Robert Brown](,_born_1773)) used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as “[Brownian motion](”. This was thought to be caused by water molecules knocking the grains about. In 1905, [Albert Einstein]( proved the reality of these molecules and their motions by producing the first [statistical physics]( analysis of [Brownian motion]([[11]]([[12]]([[13]]( French physicist [Jean Perrin]( used Einstein’s work to experimentally determine the mass and dimensions of molecules, thereby providing physical evidence for the particle nature of matter.[[14]](

## Discovery of the electron

📷The [Geiger–Marsden experiment](*Left:* Expected results: alpha particles passing through the plum pudding model of the atom with negligible deflection.*Right:* Observed results: a small portion of the particles were deflected by the concentrated positive charge of the nucleus.

In 1897, [J. J. Thomson]( discovered that [cathode rays]( are not electromagnetic waves but made of particles that are 1,800 times lighter than [hydrogen]( (the lightest atom). Thomson concluded that these particles came from the atoms within the cathode — they were *subatomic* particles. He called these new particles *corpuscles* but they were later renamed [*electrons*]( Thomson also showed that electrons were identical to particles given off by [photoelectric]( and radioactive materials.[[15]]( It was quickly recognized that electrons are the particles that carry [electric currents]( in metal wires.[[16]]( Thomson concluded that these electrons emerged from the very atoms of the cathode in his instruments, which meant that atoms are not indivisible as the name *atomos* suggests.

## Discovery of the nucleus

*Main article:* [*Geiger–Marsden experiment*](

[J. J. Thomson]( thought that the negatively-charged electrons were distributed throughout the atom in a sea of positive charge that was distributed across the whole volume of the atom.[[17]]( This model is sometimes known as the [plum pudding model](

[Ernest Rutherford]( and his colleagues [Hans Geiger]( and [Ernest Marsden]( came to have doubts about the Thomson model after they encountered difficulties when they tried to build an instrument to measure the charge-to-mass ratio of [alpha particles]( (these are positively-charged particles emitted by certain radioactive substances such as [radium]( The alpha particles were being scattered by the air in the detection chamber, which made the measurements unreliable. Thomson had encountered a similar problem in his work on cathode rays, which he solved by creating a near-perfect vacuum in his instruments. Rutherford didn’t think he’d run into this same problem because alpha particles are much heavier than electrons. According to Thomson’s model of the atom, the positive charge in the atom is not concentrated enough to produce an electric field strong enough to deflect an alpha particle, and the electrons are so lightweight they should be pushed aside effortlessly by the much heavier alpha particles. Yet there was scattering, so Rutherford and his colleagues decided to investigate this scattering carefully.[[18]](

Between 1908 and 1913, Rutherford and his colleagues performed a series of experiments in which they bombarded thin foils of metal with alpha particles. They spotted alpha particles being deflected by angles greater than 90°. To explain this, Rutherford proposed that the positive charge of the atom is not distributed throughout the atom’s volume as Thomson believed, but is concentrated in a tiny nucleus at the center. Only such an intense concentration of charge could produce an electric field strong enough to deflect the alpha particles as observed.[[18]](

## Discovery of isotopes

While experimenting with the products of [radioactive decay](, in 1913 [radiochemist]( [Frederick Soddy]( discovered that there appeared to be more than one type of atom at each position on the [periodic table]([[19]]( The term [isotope]( was coined by [Margaret Todd]( as a suitable name for different atoms that belong to the same element. J. J. Thomson created a technique for [isotope separation]( through his work on [ionized gases](, which subsequently led to the discovery of [stable isotopes]([[20]](

## Bohr model

📷The Bohr model of the atom, with an electron making instantaneous “quantum leaps” from one orbit to another with gain or loss of energy. This model of electrons in orbits is obsolete.*Main article:* [*Bohr model*](

In 1913, the physicist [Niels Bohr]( proposed a model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon.[[21]]( This quantization was used to explain why the electrons’ orbits are stable (given that normally, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation, see [*synchrotron radiation*]( and why elements absorb and emit electromagnetic radiation in discrete spectra.[[22]](

Later in the same year [Henry Moseley]( provided additional experimental evidence in favor of [Niels Bohr’s theory]( These results refined [Ernest Rutherford](’s and [Antonius van den Broek](’s model, which proposed that the atom contains in its [nucleus]( a number of positive [nuclear charges]( that is equal to its (atomic) number in the periodic table. Until these experiments, [atomic number]( was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today.[[23]](

[Chemical bonds]( between atoms were explained by [Gilbert Newton Lewis]( in 1916, as the interactions between their constituent electrons.[[24]]( As the [chemical properties]( of the elements were known to largely repeat themselves according to the [periodic law](,[[25]]( in 1919 the American chemist [Irving Langmuir]( suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of [electron shells]( about the nucleus.[[26]](

The Bohr model of the atom was the first complete physical model of the atom. It described the overall structure of the atom, how atoms bond to each other, and predicted the spectral lines of hydrogen. Bohr’s model was not perfect and was soon superseded by the more accurate Schrödinger model, but it was sufficient to evaporate any remaining doubts that matter is composed of atoms. For chemists, the idea of the atom had been a useful heuristic tool, but physicists had doubts as to whether matter really is made up of atoms as nobody had yet developed a complete physical model of the atom.

## The Schrödinger model

The [Stern–Gerlach experiment]( of 1922 provided further evidence of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom’s angular momentum, or [spin]( As this spin direction is initially random, the beam would be expected to deflect in a random direction. Instead, the beam was split into two directional components, corresponding to the atomic [spin]( being oriented up or down with respect to the magnetic field.[[27]](

In 1925, [Werner Heisenberg]( published the first consistent mathematical formulation of quantum mechanics ([matrix mechanics]([[23]]( One year earlier, [Louis de Broglie]( had proposed the [de Broglie hypothesis]( that all particles behave like waves to some extent,[[28]]( and in 1926 [Erwin Schrödinger]( used this idea to develop the [Schrödinger equation](, a mathematical model of the atom (wave mechanics) that described the electrons as three-dimensional [waveforms]( rather than point particles.[[29]](

A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the [position]( and [momentum]( of a particle at a given point in time. This became known as the [uncertainty principle](, formulated by [Werner Heisenberg]( in 1927.[[23]]( In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa.[[30]]( This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and [spectral]( patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described [atomic orbital]( zones around the nucleus where a given electron is most likely to be observed.[[31]]([[32]](

## Discovery of the neutron

The development of the [mass spectrometer]( allowed the mass of atoms to be measured with increased accuracy. The device uses a magnet to bend the trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom’s mass to its charge. The chemist [Francis William Aston]( used this instrument to show that isotopes had different masses. The [atomic mass]( of these isotopes varied by integer amounts, called the [whole number rule]([[33]]( The explanation for these different isotopes awaited the discovery of the [neutron](, an uncharged particle with a mass similar to the [proton](, by the physicist [James Chadwick]( in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus.[[34]](

## Fission, high-energy physics and condensed matter

In 1938, the German chemist [Otto Hahn](, a student of Rutherford, directed neutrons onto uranium atoms expecting to get [transuranium elements]( Instead, his chemical experiments showed [barium]( as a product.[[35]]([[36]]( A year later, [Lise Meitner]( and her nephew [Otto Frisch]( verified that Hahn’s result were the first experimental *nuclear fission*.[[37]]([[38]]( In 1944, Hahn received the [Nobel Prize in Chemistry]( Despite Hahn’s efforts, the contributions of Meitner and Frisch were not recognized.[[39]](

In the 1950s, the development of improved [particle accelerators]( and [particle detectors]( allowed scientists to study the impacts of atoms moving at high energies.[[40]]( Neutrons and protons were found to be [hadrons](, or composites of smaller particles called [quarks]( The [standard model of particle physics]( was developed that so far has successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions.[[41]](

## Structure

## Subatomic particles

*Main article:* [*Subatomic particle*](

Though the word *atom* originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various [subatomic particles]( The constituent particles of an atom are the [electron](, the [proton]( and the [neutron](

The electron is by far the least massive of these particles at 9.11×10−31 kg, with a negative [electrical charge]( and a size that is too small to be measured using available techniques.[[42]]( It was the lightest particle with a positive rest mass measured, until the discovery of [neutrino]( mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an [ion]( Electrons have been known since the late 19th century, mostly thanks to [J.J. Thomson](; see [history of subatomic physics]( for details.

Protons have a positive charge and a mass 1,836 times that of the electron, at 1.6726×10−27 kg. The number of protons in an atom is called its [atomic number]( [Ernest Rutherford]( (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it [proton](

Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the electron, or 1.6749×10−27 kg.[[43]]([[44]]( Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the [nuclear binding energy]( Neutrons and protons (collectively known as [nucleons]( have comparable dimensions—on the order of 2.5×10−15 m—although the ‘surface’ of these particles is not sharply defined.[[45]]( The neutron was discovered in 1932 by the English physicist [James Chadwick](

In the [Standard Model]( of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed of [elementary particles]( called [quarks]( There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two [up quarks]( (each with charge +2/3) and one [down quark]( (with a charge of −1/3). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles.[[46]]([[47]](

The quarks are held together by the [strong interaction]( (or strong force), which is mediated by [gluons]( The protons and neutrons, in turn, are held to each other in the nucleus by the [nuclear force](, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of [gauge bosons](, which are elementary particles that mediate physical forces.[[46]]([[47]](


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