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1History of atomic theory

Toggle History of atomic theory subsection

1.1In philosophy

1.2Dalton's law of multiple proportions

1.3Discovery of the electron

1.4Discovery of the nucleus

1.5Bohr model

1.6Discovery of protons and neutrons

1.7The Schroedinger model

2Structure

Toggle Structure subsection

2.1Subatomic particles

2.2Nucleus

2.3Electron cloud

3Properties

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3.1Nuclear properties

3.2Mass

3.3Shape and size

3.4Radioactive decay

3.5Magnetic moment

3.6Energy levels

3.7Valence and bonding behavior

3.8States

4Identification

5Origin and current state

Toggle Origin and current state subsection

5.1Formation

5.2Earth

5.3Rare and theoretical forms

5.3.1Superheavy elements

5.3.2Exotic matter

6See also

7Notes

8References

9Bibliography

10Further reading

11External links

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From Wikipedia, the free encyclopedia

Smallest unit of a chemical element

For other uses, see Atom (disambiguation).

AtomAn 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).ClassificationSmallest recognized division of a chemical elementPropertiesMass range1.67×10−27 to 4.52×10−25 kgElectric chargezero (neutral), or ion chargeDiameter range62 pm (He) to 520 pm (Cs) (data page)ComponentsElectrons and a compact nucleus of protons and neutrons

Atoms are the basic particles of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons but a different number of neutrons are called isotopes of the same element.

Atoms are extremely small, typically around 100 picometers across. A human hair is about a million carbon atoms wide. This is smaller than the shortest wavelength of visible light, which means humans cannot see atoms with conventional microscopes. Atoms are so small that accurately predicting their behavior using classical physics is not possible due to quantum effects.

More than 99.94% of an atom's mass is in the nucleus. Protons have a positive electric charge and neutrons have no charge, so the nucleus is positively charged. The electrons are negatively charged, and this opposing charge is what binds them to the nucleus. If the numbers of protons and electrons are equal, as they normally are, then the atom is electrically neutral as a whole. If an atom has more electrons than protons, then it has an overall negative charge, and is called a negative ion (or anion). Conversely, if it has more protons than electrons, it has a positive charge, and is called a positive ion (or cation).

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.

Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to attach and detach from each other is responsible for most of the physical changes observed in nature. Chemistry is the science 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. 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.[1][2] In the early 19th century, the scientist John Dalton noticed that chemical substances seemed to combine with each other by a basic unit of weight that is consistent for each element, and he decided to use the word atom to refer to these unit.[3]

Dalton's law of multiple proportions

Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808)

In the early 1800s, the English chemist John Dalton compiled experimental data gathered by him and other scientists and discovered a pattern now known as the "law of multiple proportions". He noticed that in chemical compounds which contain two particular chemical elements, the amount of Element A per measure of Element B will differ across these compounds by ratios of small whole numbers. This pattern suggested that the elements combine with each other by basic units of weight which are consistent for each element, and Dalton decided to call these units "atoms".[4]

For example, there are two types of tin oxide: one is a grey 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 grey powder there is about 13.5 g of oxygen for every 100 g of tin, and in the white powder 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).[5][6]

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 powder there is about 28 g of oxygen for every 100 g of iron, and in the red powder there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. Dalton concluded that in these oxides, for every two atoms of iron, there are two or three atoms of oxygen respectively (Fe2O2 and Fe2O3).[b][7][8]

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.[9][10]

Discovery of the electron

In 1897, J. J. Thomson discovered that cathode rays are not electromagnetic waves but made of particles with mass because they can be deflected by electric and magnetic fields. He measured these particles to be 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.[11] It was quickly recognized that electrons are the particles that carry electric currents in metal wires.[12] Thomson concluded that these electrons emerged from the very atoms of the cathode in his instruments, which meant that atoms are not indivisible as their name suggests.

Discovery of the nucleus

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.

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.[13] This model is sometimes known as the plum pudding model.

Between 1908 and 1913, Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden performed a series of experiments in which they bombarded thin foils of metal with alpha particles. They did this to measure the scattering patterns of the alpha particles. They spotted alpha particles being deflected by angles greater than 90°. This shouldn't have been possible according to the Thomson model of the atom, whose charges were too diffuse to produce a sufficiently strong electric field. 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.[14]

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.[15] This quantization was used to explain why the electrons' orbits are stable (given that in classical physics, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation) and why elements absorb and emit electromagnetic radiation in discrete spectra.[16]

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.[17]

Chemical bonds between atoms were explained by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons.[18] As the chemical properties of the elements were known to largely repeat themselves according to the periodic law,[19] 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.[20]

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.

Discovery of protons and neutrons

Main articles: Atomic nucleus and Discovery of the neutron

In 1917 Rutherford bombarded nitrogen gas with alpha particles and observed hydrogen nuclei being emitted from the gas (Rutherford recognized these, because he had previously obtained them bombarding hydrogen with alpha particles, and observing hydrogen nuclei in the products). Rutherford concluded that the hydrogen nuclei emerged from the nuclei of the nitrogen atoms themselves (in effect, he had split a nitrogen).[21]

From his own work and the work of his students Bohr and Henry Moseley, Rutherford knew that the positive charge of any atom could always be equated to that of an integer number of hydrogen nuclei. This, coupled with the atomic mass of many elements being roughly equivalent to an integer number of hydrogen atoms - then assumed to be the lightest particles - led him to conclude that hydrogen nuclei were singular particles and a basic constituent of all atomic nuclei. He named such particles protons. Further experimentation by Rutherford found that the nuclear mass of most atoms exceeded that of the protons it possessed; he speculated that this surplus mass was composed of previously unknown neutrally charged particles, which were tentatively dubbed "neutrons".

In 1928, Walter Bothe observed that beryllium emitted a highly penetrating, electrically neutral radiation when bombarded with alpha particles. It was later discovered that this radiation could knock hydrogen atoms out of paraffin wax. Initially it was thought to be high-energy gamma radiation, since gamma radiation had a similar effect on electrons in metals, but James Chadwick found that the ionization effect was too strong for it to be due to electromagnetic radiation, so long as energy and momentum were conserved in the interaction. In 1932, Chadwick exposed various elements, such as hydrogen and nitrogen, to the mysterious "beryllium radiation", and by measuring the energies of the recoiling charged particles, he deduced that the radiation was actually composed of electrically neutral particles which could not be massless like the gamma ray, but instead were required to have a mass similar to that of a proton. Chadwick now claimed these particles as Rutherford's neutrons.[22] For his discovery of the neutron, Chadwick received the Nobel Prize in 1935.[23]

The discovery of the neutron explained the existence of isotopes, which are atoms of the same element which have slightly different masses, due to them having different numbers of neutrons but the same number of protons.

The Schroedinger model

The modern model of atomic orbitals draws zones where an electron is most likely to found at any moment.

In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (matrix mechanics).[17] One year earlier, Louis de Broglie had proposed that all particles behave like waves to some extent,[24] and in 1926 Erwin Schroedinger used this idea to develop the Schroedinger equation, a mathematical model of the atom that described the electrons as three-dimensional waveforms rather than points in space.[25]

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.[17] 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.[26]

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.[27][28]

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 the least massive of these particles by four orders of magnitude at 9.11×10−31 kg, with a negative electrical charge and a size that is too small to be measured using available techniques.[29] 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 of 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 mass of 1.6749×10−27 kg.[30][31] 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.[32] 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.[33][34]

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.[33][34]

Nucleus

Main article: Atomic nucleus

The binding energy needed for a nucleon to escape the nucleus, for various isotopes

All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to

1.07

{\displaystyle 1.07{\sqrt[{3}]{A}}}

femtometres, where

{\displaystyle A}

is the total number of nucleons.[35] This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other.[36]

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 total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay.[37]

The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud.[38]

A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus.[38]

Illustration of a nuclear fusion process that forms a deuterium nucleus, consisting of a proton and a neutron, from two protons. A positron (e+)—an antimatter electron—is emitted along with an electron neutrino.

The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3 to 10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus.[39] Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element.[40][41]

If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass–energy equivalence formula, e=mc2, where m is the mass loss and c is the speed of light. This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate.[42]

The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together.[43] It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon begins to decrease. That means that a fusion process producing a nucleus that has an atomic number higher than about 26, and a mass number higher than about 60, is an endothermic process. Thus, more massive nuclei cannot undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star.[38]

Electron cloud

Main articles: Electron configuration, Electron shell, and Atomic orbitalSee also: ElectronegativityA potential well, showing, according to classical mechanics, the minimum energy V(x) needed to reach each position x. Classically, a particle with energy E is constrained to a range of positions between x1 and x2.

The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations.

Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured.[44] Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form.[45] Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation.[46]

3D views of some hydrogen-like atomic orbitals showing probability density and phase (g orbitals and higher are not shown)

Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines.[45]

The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom,[47] compared to 2.23 million eV for splitting a deuterium nucleus.[48] Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest 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.[49]

Properties

Nuclear properties

Main articles: Isotope, Stable isotope, List of nuclides, and List of elements by stability of isotopes

By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form,[50] also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single-proton element hydrogen up to the 118-proton element oganesson.[51] All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible.[52][53]

About 339 nuclides occur naturally on Earth,[54] of which 251 (about 74%) have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stable theoretically, while another 161 (bringing the total to 251) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 35 radioactive nuclides have half-lives longer than 100 million years, and are long-lived enough to have been present since the birth of the Solar System. This collection of 286 nuclides are known as primordial nuclides. Finally, an additional 53 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14).[55][note 1]

For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.1 stable isotopes per element. Twenty-six "monoisotopic elements" have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes.[56]: 1–12

Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 251 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10, and nitrogen-14. (Tantalum-180m is odd-odd and observationally stable, but is predicted to decay with a very long half-life.) Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138, and lutetium-176. Most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.[57]

Mass

Main articles: Atomic mass and mass number

The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles (called "nucleons") in a given atom is called the mass number. It is a positive integer and dimensionless (instead of having dimension of mass), because it expresses a count. An example of use of a mass number is "carbon-12," which has 12 nucleons (six protons and six neutrons).

The actual mass of an atom at rest is often expressed in daltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately 1.66×10−27 kg.[58] Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da.[59] The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 Da), but this number will not be exactly an integer except (by definition) in the case of carbon-12.[60] The heaviest stable atom is lead-208,[52] with a mass of 207.9766521 Da.[61]

As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms (about 6.022×1023). This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 Da, and so a mole of carbon-12 atoms weighs exactly 0.012 kg.[58]

Shape and size

Main article: Atomic radius

Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus.[62] This assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin.[63] On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right).[64] Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm.[65]

When subjected to external forces, like electrical fields, the shape of an atom may deviate from spherical symmetry. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by group-theoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystal-electrical fields may occur at low-symmetry lattice sites.[66][67] Significant ellipsoidal deformations have been shown to occur for sulfur ions[68] and chalcogen ions[69] in pyrite-type compounds.

Atomic dimensions are thousands of times smaller than the wavelengths of light (400–700 nm) so they cannot be viewed using an optical microscope, although individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width.[70] A single drop of water contains about 2 sextillion (2×1021) atoms of oxygen, and twice the number of hydrogen atoms.[71] A single carat diamond with a mass of 2×10−4 kg contains about 10 sextillion (1022) atoms of carbon.[note 2] If an apple were magnified to the size of the Earth, then the atoms in the apple would be approximately the size of the original apple.[72]

Radioactive decay

Main article: Radioactive decay

This diagram shows the half-life (T1⁄2) of various isotopes with Z protons and N neutrons.

Every element has one or more isotopes that have unstable nuclei that are subject to radioactive decay, causing the nucleus to emit particles or electromagnetic radiation. Radioactivity can occur when the radius of a nucleus is large compared with the radius of the strong force, which only acts over distances on the order of 1 fm.[73]

The most common forms of radioactive decay are:[74][75]

Alpha decay: this process is caused when the nucleus emits an alpha particle, which is a helium nucleus consisting of two protons and two neutrons. The result of the emission is a new element with a lower atomic number.

Beta decay (and electron capture): these processes are regulated by the weak force, and result from a transformation of a neutron into a proton, or a proton into a neutron. The neutron to proton transition is accompanied by the emission of an electron and an antineutrino, while proton to neutron transition (except in electron capture) causes the emission of a positron and a neutrino. The electron or positron emissions are called beta particles. Beta decay either increases or decreases the atomic number of the nucleus by one. Electron capture is more common than positron emission, because it requires less energy. In this type of decay, an electron is absorbed by the nucleus, rather than a positron emitted from the nucleus. A neutrino is still emitted in this process, and a proton changes to a neutron.

Gamma decay: this process results from a change in the energy level of the nucleus to a lower state, resulting in the emission of electromagnetic radiation. The excited state of a nucleus which results in gamma emission usually occurs following the emission of an alpha or a beta particle. Thus, gamma decay usually follows alpha or beta decay.

Other more rare types of radioactive decay include ejection of neutrons or protons or clusters of nucleons from a nucleus, or more than one beta particle. An analog of gamma emission which allows excited nuclei to lose energy in a different way, is internal conversion—a process that produces high-speed electrons that are not beta rays, followed by production of high-energy photons that are not gamma rays. A few large nuclei explode into two or more charged fragments of varying masses plus several neutrons, in a decay called spontaneous nuclear fission.

Each radioactive isotope has a characteristic decay time period—the half-life—that is determined by the amount of time needed for half of a sample to decay. This is an exponential decay process that steadily decreases the proportion of the remaining isotope by 50% every half-life. Hence after two half-lives have passed only 25% of the isotope is present, and so forth.[73]

Magnetic moment

Main articles: Electron magnetic moment and Nuclear magnetic moment

Elementary particles possess an intrinsic quantum mechanical property known as spin. This is analogous to the angular momentum of an object that is spinning around its center of mass, although strictly speaking these particles are believed to be point-like and cannot be said to be rotating. Spin is measured in units of the reduced Planck constant (ħ), with electrons, protons and neutrons all having spin 1⁄2 ħ, or "spin-1⁄2". In an atom, electrons in motion around the nucleus possess orbital angular momentum in addition to their spin, while the nucleus itself possesses angular momentum due to its nuclear spin.[76]

The magnetic field produced by an atom—its magnetic moment—is determined by these various forms of angular momentum, just as a rotating charged object classically produces a magnetic field, but the most dominant contribution comes from electron spin. Due to the nature of electrons to obey the Pauli exclusion principle, in which no two electrons may be found in the same quantum state, bound electrons pair up with each other, with one member of each pair in a spin up state and the other in the opposite, spin down state. Thus these spins cancel each other out, reducing the total magnetic dipole moment to zero in some atoms with even number of electrons.[77]

In ferromagnetic elements such as iron, cobalt and nickel, an odd number of electrons leads to an unpaired electron and a net overall magnetic moment. The orbitals of neighboring atoms overlap and a lower energy state is achieved when the spins of unpaired electrons are aligned with each other, a spontaneous process known as an exchange interaction. When the magnetic moments of ferromagnetic atoms are lined up, the material can produce a measurable macroscopic field. Paramagnetic materials have atoms with magnetic moments that line up in random directions when no magnetic field is present, but the magnetic moments of the individual atoms line up in the presence of a field.[77][78]

The nucleus of an atom will have no spin when it has even numbers of both neutrons and protons, but for other cases of odd numbers, the nucleus may have a spin. Normally nuclei with spin are aligned in random directions because of thermal equilibrium, but for certain elements (such as xenon-129) it is possible to polarize a significant proportion of the nuclear spin states so that they are aligned in the same direction—a condition called hyperpolarization. This has important applications in magnetic resonance imaging.[79][80]

Energy levels

These electron's energy levels (not to scale) are sufficient for ground states of atoms up to cadmium (5s2 4d10) inclusively. Do not forget that even the top of the diagram is lower than an unbound electron state.

The potential energy of an electron in an atom is negative relative to when the distance from the nucleus goes to infinity; its dependence on the electron's position reaches the minimum inside the nucleus, roughly in inverse proportion to the distance. In the quantum-mechanical model, a bound electron can occupy only a set of states centered on the nucleus, and each state corresponds to a specific energy level; see time-independent Schrödinger equation for a theoretical explanation. An energy level can be measured by the amount of energy needed to unbind the electron from the atom, and is usually given in units of electronvolts (eV). The lowest energy state of a bound electron is called the ground state, i.e. stationary state, while an electron transition to a higher level results in an excited state.[81] The electron's energy increases along with n because the (average) distance to the nucleus increases. Dependence of the energy on ℓ is caused not by the electrostatic potential of the nucleus, but by interaction between electrons.

For an electron to transition between two different states, e.g. ground state to first excited state, it must absorb or emit a photon at an energy matching the difference in the potential energy of those levels, according to the Niels Bohr model, what can be precisely calculated by the Schrödinger equation.

Electrons jump between orbitals in a particle-like fashion. For example, if a single photon strikes the electrons, only a single electron changes states in response to the photon; see Electron properties.

The energy of an emitted photon is proportional to its frequency, so these specific energy levels appear as distinct bands in the electromagnetic spectrum.[82] Each element has a characteristic spectrum that can depend on the nuclear charge, subshells filled by electrons, the electromagnetic interactions between the electrons and other factors.[83]

An example of absorption lines in a spectrum

When a continuous spectrum of energy is passed through a gas or plasma, some of the photons are absorbed by atoms, causing electrons to change their energy level. Those excited electrons that remain bound to their atom spontaneously emit this energy as a photon, traveling in a random direction, and so drop back to lower energy levels. Thus the atoms behave like a filter that forms a series of dark absorption bands in the energy output. (An observer viewing the atoms from a view that does not include the continuous spectrum in the background, instead sees a series of emission lines from the photons emitted by the atoms.) Spectroscopic measurements of the strength and width of atomic spectral lines allow the composition and physical properties of a substance to be determined.[84]

Close examination of the spectral lines reveals that some display a fine structure splitting. This occurs because of spin–orbit coupling, which is an interaction between the spin and motion of the outermost electron.[85] When an atom is in an external magnetic field, spectral lines become split into three or more components; a phenomenon called the Zeeman effect. This is caused by the interaction of the magnetic field with the magnetic moment of the atom and its electrons. Some atoms can have multiple electron configurations with the same energy level, which thus appear as a single spectral line. The interaction of the magnetic field with the atom shifts these electron configurations to slightly different energy levels, resulting in multiple spectral lines.[86] The presence of an external electric field can cause a comparable splitting and shifting of spectral lines by modifying the electron energy levels, a phenomenon called the Stark effect.[87]

If a bound electron is in an excited state, an interacting photon with the proper energy can cause stimulated emission of a photon with a matching energy level. For this to occur, the electron must drop to a lower energy state that has an energy difference matching the energy of the interacting photon. The emitted photon and the interacting photon then move off in parallel and with matching phases. That is, the wave patterns of the two photons are synchronized. This physical property is used to make lasers, which can emit a coherent beam of light energy in a narrow frequency band.[88]

Valence and bonding behavior

Main articles: Valence (chemistry) and Chemical bond

Valency is the combining power of an element. It is determined by the number of bonds it can form to other atoms or groups.[89] The outermost electron shell of an atom in its uncombined state is known as the valence shell, and the electrons in

that shell are called valence electrons. The number of valence electrons determines the bonding

behavior with other atoms. Atoms tend to chemically react with each other in a manner that fills (or empties) their outer valence shells.[90] For example, a transfer of a single electron between atoms is a useful approximation for bonds that form between atoms with one-electron more than a filled shell, and others that are one-electron short of a full shell, such as occurs in the compound sodium chloride and other chemical ionic salts. Many elements display multiple valences, or tendencies to share differing numbers of electrons in different compounds. Thus, chemical bonding between these elements takes many forms of electron-sharing that are more than simple electron transfers. Examples include the element carbon and the organic compounds.[91]

The chemical elements are often displayed in a periodic table that is laid out to display recurring chemical properties, and elements with the same number of valence electrons form a group that is aligned in the same column of the table. (The horizontal rows correspond to the filling of a quantum shell of electrons.) The elements at the far right of the table have their outer shell completely filled with electrons, which results in chemically inert elements known as the noble gases.[92][93]

States

Main articles: State of matter and Phase (matter)

Graphic illustrating the formation of a Bose–Einstein condensate

Quantities of atoms are found in different states of matter that depend on the physical conditions, such as temperature and pressure. By varying the conditions, materials can transition between solids, liquids, gases and plasmas.[94] Within a state, a material can also exist in different allotropes. An example of this is solid carbon, which can exist as graphite or diamond.[95] Gaseous allotropes exist as well, such as dioxygen and ozone.

At temperatures close to absolute zero, atoms can form a Bose–Einstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale.[96][97] This super-cooled collection of atoms

then behaves as a single super atom, which may allow fundamental checks of quantum mechanical behavior.[98]

Identification

Scanning tunneling microscope image showing the individual atoms making up this gold (100) surface. The surface atoms deviate from the bulk crystal structure and arrange in columns several atoms wide with pits between them (See surface reconstruction).

While atoms are too small to be seen, devices such as the scanning tunneling microscope (STM) enable their visualization at the surfaces of solids. The microscope uses the quantum tunneling phenomenon, which allows particles to pass through a barrier that would be insurmountable in the classical perspective. Electrons tunnel through the vacuum between two biased electrodes, providing a tunneling current that is exponentially dependent on their separation. One electrode is a sharp tip ideally ending with a single atom. At each point of the scan of the surface the tip's height is adjusted so as to keep the tunneling current at a set value. How much the tip moves to and away from the surface is interpreted as the height profile. For low bias, the microscope images the averaged electron orbitals across closely packed energy levels—the local density of the electronic states near the Fermi level.[99][100] Because of the distances involved, both electrodes need to be extremely stable; only then periodicities can be observed that correspond to individual atoms. The method alone is not chemically specific, and cannot identify the atomic species present at the surface.

Atoms can be easily identified by their mass. If an atom is ionized by removing one of its electrons, its trajectory when it passes through a magnetic field will bend. The radius by which the trajectory of a moving ion is turned by the magnetic field is determined by the mass of the atom. The mass spectrometer uses this principle to measure the mass-to-charge ratio of ions. If a sample contains multiple isotopes, the mass spectrometer can determine the proportion of each isotope in the sample by measuring the intensity of the different beams of ions. Techniques to vaporize atoms include inductively coupled plasma atomic emission spectroscopy and inductively coupled plasma mass spectrometry, both of which use a plasma to vaporize samples for analysis.[101]

The atom-probe tomograph has sub-nanometer resolution in 3-D and can chemically identify individual atoms using time-of-flight mass spectrometry.[102]

Electron emission techniques such as X-ray photoelectron spectroscopy (XPS) and Auger electron spectroscopy (AES), which measure the binding energies of the core electrons, are used to identify the atomic species present in a sample in a non-destructive way. With proper focusing both can be made area-specific. Another such method is electron energy loss spectroscopy (EELS), which measures the energy loss of an electron beam within a transmission electron microscope when it interacts with a portion of a sample.

Spectra of excited states can be used to analyze the atomic composition of distant stars. Specific light wavelengths contained in the observed light from stars can be separated out and related to the quantized transitions in free gas atoms. These colors can be replicated using a gas-discharge lamp containing the same element.[103] Helium was discovered in this way in the spectrum of the Sun 23 years before it was found on Earth.[104]

Origin and current state

Baryonic matter forms about 4% of the total energy density of the observable universe, with an average density of about 0.25 particles/m3 (mostly protons and electrons).[105] Within a galaxy such as the Milky Way, particles have a much higher concentration, with the density of matter in the interstellar medium (ISM) ranging from 105 to 109 atoms/m3.[106] The Sun is believed to be inside the Local Bubble, so the density in the solar neighborhood is only about 103 atoms/m3.[107] Stars form from dense clouds in the ISM, and the evolutionary processes of stars result in the steady enrichment of the ISM with elements more massive than hydrogen and helium.

Up to 95% of the Milky Way's baryonic matter are concentrated inside stars, where conditions are unfavorable for atomic matter. The total baryonic mass is about 10% of the mass of the galaxy;[108] the remainder of the mass is an unknown dark matter.[109] High temperature inside stars makes most "atoms" fully ionized, that is, separates all electrons from the nuclei. In stellar remnants—with exception of their surface layers—an immense pressure make electron shells impossible.

Formation

Main article: Nucleosynthesis

Periodic table showing the origin of each element. Elements from carbon up to sulfur may be made in small stars by the alpha process. Elements beyond iron are made in large stars with slow neutron capture (s-process). Elements heavier than iron may be made in neutron star mergers or supernovae after the r-process.

Electrons are thought to exist in the Universe since early stages of the Big Bang. Atomic nuclei forms in nucleosynthesis reactions. In about three minutes Big Bang nucleosynthesis produced most of the helium, lithium, and deuterium in the Universe, and perhaps some of the beryllium and boron.[110][111][112]

Ubiquitousness and stability of atoms relies on their binding energy, which means that an atom has a lower energy than an unbound system of the nucleus and electrons. Where the temperature is much higher than ionization potential, the matter exists in the form of plasma—a gas of positively charged ions (possibly, bare nuclei) and electrons. When the temperature drops below the ionization potential, atoms become statistically favorable. Atoms (complete with bound electrons) became to dominate over charged particles 380,000 years after the Big Bang—an epoch called recombination, when the expanding Universe cooled enough to allow electrons to become attached to nuclei.[113]

Since the Big Bang, which produced no carbon or heavier elements, atomic nuclei have been combined in stars through the process of nuclear fusion to produce more of the element helium, and (via the triple alpha process) the sequence of elements from carbon up to iron;[114] see stellar nucleosynthesis for details.

Isotopes such as lithium-6, as well as some beryllium and boron are generated in space through cosmic ray spallation.[115] This occurs when a high-energy proton strikes an atomic nucleus, causing large numbers of nucleons to be ejected.

Elements heavier than iron were produced in supernovae and colliding neutron stars through the r-process, and in AGB stars through the s-process, both of which involve the capture of neutrons by atomic nuclei.[116] Elements such as lead formed largely through the radioactive decay of heavier elements.[117]

Earth

Most of the atoms that make up the Earth and its inhabitants were present in their current form in the nebula that collapsed out of a molecular cloud to form the Solar System. The rest are the result of radioactive decay, and their relative proportion can be used to determine the age of the Earth through radiometric dating.[118][119] Most of the helium in the crust of the Earth (about 99% of the helium from gas wells, as shown by its lower abundance of helium-3) is a product of alpha decay.[120]

There are a few trace atoms on Earth that were not present at the beginning (i.e., not "primordial"), nor are results of radioactive decay. Carbon-14 is continuously generated by cosmic rays in the atmosphere.[121] Some atoms on Earth have been artificially generated either deliberately or as by-products of nuclear reactors or explosions.[122][123] Of the transuranic elements—those with atomic numbers greater than 92—only plutonium and neptunium occur naturally on Earth.[124][125] Transuranic elements have radioactive lifetimes shorter than the current age of the Earth[126] and thus identifiable quantities of these elements have long since decayed, with the exception of traces of plutonium-244 possibly deposited by cosmic dust.[118] Natural deposits of plutonium and neptunium are produced by neutron capture in uranium ore.[127]

The Earth contains approximately 1.33×1050 atoms.[128] Although small numbers of independent atoms of noble gases exist, such as argon, neon, and helium, 99% of the atmosphere is bound in the form of molecules, including carbon dioxide and diatomic oxygen and nitrogen. At the surface of the Earth, an overwhelming majority of atoms combine to form various compounds, including water, salt, silicates and oxides. Atoms can also combine to create materials that do not consist of discrete molecules, including crystals and liquid or solid metals.[129][130] This atomic matter forms networked arrangements that lack the particular type of small-scale interrupted order associated with molecular matter.[131]

Rare and theoretical forms

Superheavy elements

Main article: Superheavy element

All nuclides with atomic numbers higher than 82 (lead) are known to be radioactive. No nuclide with an atomic number exceeding 92 (uranium) exists on Earth as a primordial nuclide, and heavier elements generally have shorter half-lives. Nevertheless, an "island of stability" encompassing relatively long-lived isotopes of superheavy elements[132] with atomic numbers 110 to 114 might exist.[133] Predictions for the half-life of the most stable nuclide on the island range from a few minutes to millions of years.[134] In any case, superheavy elements (with Z > 104) would not exist due to increasing Coulomb repulsion (which results in spontaneous fission with increasingly short half-lives) in the absence of any stabilizing effects.[135]

Exotic matter

Main article: Exotic matter

Each particle of matter has a corresponding antimatter particle with the opposite electrical charge. Thus, the positron is a positively charged antielectron and the antiproton is a negatively charged equivalent of a proton. When a matter and corresponding antimatter particle meet, they annihilate each other. Because of this, along with an imbalance between the number of matter and antimatter particles, the latter are rare in the universe. The first causes of this imbalance are not yet fully understood, although theories of baryogenesis may offer an explanation. As a result, no antimatter atoms have been discovered in nature.[136][137] In 1996, the antimatter counterpart of the hydrogen atom (antihydrogen) was synthesized at the CERN laboratory in Geneva.[138][139]

Other exotic atoms have been created by replacing one of the protons, neutrons or electrons with other particles that have the same charge. For example, an electron can be replaced by a more massive muon, forming a muonic atom. These types of atoms can be used to test fundamental predictions of physics.[140][141][142]

Atom | Definition, Structure, History, Examples, Diagram, & Facts | Britannica

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atom

Table of Contents

atom

Table of Contents

Introduction & Top QuestionsAtomic modelBasic propertiesAtomic numberAtomic mass and isotopesThe electronCharge, mass, and spinOrbits and energy levelsElectron shellsAtomic bondsConductors and insulatorsMagnetic propertiesThe nucleusNuclear forcesNuclear shell modelRadioactive decayNuclear energyDevelopment of atomic theoryThe atomic philosophy of the early GreeksThe emergence of experimental scienceThe beginnings of modern atomic theoryExperimental foundation of atomic chemistryAtomic weights and the periodic tableKinetic theory of gasesStudies of the properties of atomsSize of atomsElectric properties of atomsLight and spectral linesDiscovery of electronsIdentification of positive ionsDiscovery of radioactivityModels of atomic structureRutherford’s nuclear modelMoseley’s X-ray studiesBohr’s shell modelThe laws of quantum mechanicsSchrödinger’s wave equationAntiparticles and the electron’s spinAdvances in nuclear and subatomic physicsStructure of the nucleusQuantum field theory and the standard model

References & Edit History

Top Questions

What is an atom?An atom is the basic building block of chemistry. It is the smallest unit into which matter can be divided without the release of electrically charged particles. It also is the smallest unit of matter that has the characteristic properties of a chemical element.Are all atoms the same size?All atoms are roughly the same size, whether they have 3 or 90 electrons. Approximately 50 million atoms of solid matter lined up in a row would measure 1 cm (0.4 inches). A convenient unit of length for measuring atomic sizes is the angstrom, defined as 10−10 meters.What does the mass of an atom consist of?The mass of an atom consists of the mass of the nucleus plus that of the electrons. That means the atomic mass unit is not exactly the same as the mass of the proton or neutron.How is the atomic number of an atom defined?The single most important characteristic of an atom is its atomic number (usually denoted by the letter Z), which is defined as the number of units of positive charge (protons) in the nucleus. For example, if an atom has a Z of 6, it is carbon, while a Z of 92 corresponds to uranium.atomsHow atoms can be seen.(more)See all videos for this articleatom, the basic building block of all matter and chemistry. Atoms can combine with other atoms to form molecules but cannot be divided into smaller parts by ordinary chemical processes.Explore an atom's interior to discover the layout of its nucleus, protons, and electronsSee all videos for this articleMost of the atom is empty space. The rest consists of three basic types of subatomic particles: protons, neutrons, and electrons. The protons and neutrons form the atom’s central nucleus. (The ordinary hydrogen atom is an exception; it contains one proton but no neutrons.) As their names suggest, protons have a positive electrical charge, while neutrons are electrically neutral—they carry no charge; overall, then, the nucleus has a positive charge. Circling the nucleus is a cloud of electrons, which are negatively charged. Like opposite ends of a magnet that attract one another, the negative electrons are attracted to a positive force, which binds them to the nucleus. The nucleus is small and dense compared with the electrons, which are the lightest charged particles in nature. The electrons circle the nucleus in orbital paths called shells, each of which holds only a certain number of electrons.Investigate varying electron configurations in electron shells around an atom's nucleusAtomic model of electron configurations.(more)See all videos for this articleAn ordinary, neutral atom has an equal number of protons (in the nucleus) and electrons (surrounding the nucleus). Thus the positive and negative charges are balanced. Some atoms, however, lose or gain electrons in chemical reactions or in collisions with other particles. Ordinary atoms that either gain or lose electrons are called ions. If a neutral atom loses an electron, it becomes a positive ion. If it gains an electron, it becomes a negative ion. These basic subatomic particles—protons, neutrons, and electrons—are themselves made up of smaller substances, such as quarks and leptons.More than 90 types of atoms exist in nature, and each kind of atom forms a different chemical element. Chemical elements are made up of only one type of atom—gold contains only gold atoms, and neon contains only neon atoms--and they are ranked in order of their atomic number (the total number of protons in its nucleus) in a chart called the periodic table. Accordingly, because an atom of iron has 26 protons in its nucleus, its atomic number is 26 and its ranking on the periodic table of chemical elements is 26. Because an ordinary atom has the same number of electrons as protons, an element’s atomic number also tells how many electrons its atoms have, and it is the number and arrangement of the electrons in their orbiting shells that determines how one atom interacts with another. The key shell is the outermost one, called the valence shell. If this outermost shell is complete, or filled with the maximum number of electrons for that shell, the atom is stable, with little or no tendency to interact with other atoms. But atoms with incomplete outer shells seek to fill or to empty such shells by gaining or losing electrons or by sharing electrons with other atoms. This is the basis of an atom’s chemical activity. Atoms that have the same number of electrons in the outer shell have similar chemical properties.

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Facts You Should Know: The Periodic Table Quiz

This article opens with a broad overview of the fundamental properties of the atom and its constituent particles and forces. Following this overview is a historical survey of the most influential concepts about the atom that have been formulated through the centuries. Atomic model Most matter consists of an agglomeration of molecules, which can be separated relatively easily. Molecules, in turn, are composed of atoms joined by chemical bonds that are more difficult to break. Each individual atom consists of smaller particles—namely, electrons and nuclei. These particles are electrically charged, and the electric forces on the charge are responsible for holding the atom together. Attempts to separate these smaller constituent particles require ever-increasing amounts of energy and result in the creation of new subatomic particles, many of which are charged.

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As noted in the introduction to this article, an atom consists largely of empty space. The nucleus is the positively charged centre of an atom and contains most of its mass. It is composed of protons, which have a positive charge, and neutrons, which have no charge. Protons, neutrons, and the electrons surrounding them are long-lived particles present in all ordinary, naturally occurring atoms. Other subatomic particles may be found in association with these three types of particles. They can be created only with the addition of enormous amounts of energy, however, and are very short-lived. All atoms are roughly the same size, whether they have 3 or 90 electrons. Approximately 50 million atoms of solid matter lined up in a row would measure 1 cm (0.4 inch). A convenient unit of length for measuring atomic sizes is the angstrom (Å), defined as 10−10 metre. The radius of an atom measures 1–2 Å. Compared with the overall size of the atom, the nucleus is even more minute. It is in the same proportion to the atom as a marble is to a football field. In volume the nucleus takes up only 10−14 metres of the space in the atom—i.e., 1 part in 100,000. A convenient unit of length for measuring nuclear sizes is the femtometre (fm), which equals 10−15 metre. The diameter of a nucleus depends on the number of particles it contains and ranges from about 4 fm for a light nucleus such as carbon to 15 fm for a heavy nucleus such as lead. In spite of the small size of the nucleus, virtually all the mass of the atom is concentrated there. The protons are massive, positively charged particles, whereas the neutrons have no charge and are slightly more massive than the protons. The fact that nuclei can have anywhere from 1 to nearly 300 protons and neutrons accounts for their wide variation in mass. The lightest nucleus, that of hydrogen, is 1,836 times more massive than an electron, while heavy nuclei are nearly 500,000 times more massive.

Basic properties Atomic number The single most important characteristic of an atom is its atomic number (usually denoted by the letter Z), which is defined as the number of units of positive charge (protons) in the nucleus. For example, if an atom has a Z of 6, it is carbon, while a Z of 92 corresponds to uranium. A neutral atom has an equal number of protons and electrons so that the positive and negative charges exactly balance. Since it is the electrons that determine how one atom interacts with another, in the end it is the number of protons in the nucleus that determines the chemical properties of an atom.

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The classical “planetary” model of an atom. The protons and neutrons in the nucleus are circled by electrons in “orbit” around the nucleus. The number of protons determines which element is represented, the number of electrons determines its charge, and the number of neutrons determines which isotope of the element is represented.atom, Smallest unit into which matter can be divided and still retain the characteristic properties of an element. The word derives from the Greek atomos (“indivisible”), and the atom was believed to be indivisible until the early 20th century, when electrons and the nucleus were discovered. It is now known that an atom has a positively charged nucleus that makes up more than 99.9% of the atom’s mass but only about 1/100,000 of its volume. The nucleus is composed of positively charged protons and electrically neutral neutrons, each about 2,000 times as massive as an electron. Most of the atom’s volume consists of a cloud of electrons that have very small mass and negative charge. The electron cloud is bound to the nucleus by the attraction of opposite charges. In a neutral atom, the protons in the nucleus are balanced by the electrons. An atom that has gained or lost electrons becomes negatively or positively charged and is called an ion.

quantum chromodynamics Summary

Quantum chromodynamics (QCD), in physics, the theory that describes the action of the strong force. QCD was constructed in analogy to quantum electrodynamics (QED), the quantum field theory of the electromagnetic force. In QED the electromagnetic interactions of charged particles are described

neutrino Summary

Neutrino, elementary subatomic particle with no electric charge, very little mass, and 12 unit of spin. Neutrinos belong to the family of particles called leptons, which are not subject to the strong force. Rather, neutrinos are subject to the weak force that underlies certain processes of

magnetic resonance imaging Summary

Magnetic resonance imaging (MRI), three-dimensional diagnostic imaging technique used to visualize organs and structures inside the body without the need for X-rays or other radiation. MRI is valuable for providing detailed anatomical images and can reveal minute changes that occur over time. It

transuranium element Summary

Transuranium element, any of the chemical elements that lie beyond uranium in the periodic table—i.e., those with atomic numbers greater than 92. Twenty-six of these elements have been discovered and named or are awaiting confirmation of their discovery. Eleven of them, from neptunium through

The Structure of the Atom – Introductory Chemistry

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I. Introduction to Chemistry1. Overview of ChemistryLumenLearning2. Chemistry in the Modern WorldSaylor Academy3. The Scientific MethodSaylor Academy4. MeasurementsOpenStax5. A Description of MatterSaylor Academy6. Dimensional AnalysisLumenLearningII. Particles of Matter7. History of Atomic StructureLumenLearning8. Early Ideas in Atomic TheoryOpenStax9. The Periodic TableOpenStax10. EnergyLumenLearning11. Energy BasicsOpenStax12. Introduction to ThermodynamicsLumenLearning13. Classification of Matter (Solid, Liquid, Gas)LumenLearning14. Properties of GasesLumenLearning15. Gas LawsLumenLearning16. Ideal Gas LawLumenLearning17. Kinetic Molecular TheoryLumenLearningIII. Elements of Chemistry18. Physical and Chemical Properties of MatterLumenLearning19. Classification of Matter (Elements, Compounds, Mixtures)LumenLearning20. The Periodic TableLumenLearning21. Molecular and Ionic CompoundsOpenStax22. Naming CompoundsLumenLearningIV. Subatomic Particles23. Discoveries Leading to the Nuclear Atom ModelLumenLearning24. The Structure of the AtomLumenLearning25. The Nature of LightLumenLearning26. Bohr's TheoryLumenLearning27. Quantum Numbers for ElectronsSaylor Academy28. Orbital ShapesLumenLearning29. Organization of Electrons in AtomsSaylor Academy30. Electron ConfigurationLumenLearning31. Electron Structure and the Periodic TableSaylor Academy32. Periodic TrendsSaylor AcademyV. Nuclear Chemistry33. RadioactivityLumenLearning34. Nuclear Structure and StabilityOpenStax35. Nuclear ReactionsLumenLearning36. Nuclear TransmutationLumenLearning37. Nuclear FissionLumenLearning38. Nuclear FusionLumenLearning39. Use of IsotopesLumenLearning40. Uses of RadioisotopesOpenStax41. Biological Effects of RadiationOpenStaxVI. Chemical Bonding42. Lewis Dot Symbols and Lewis Structures (Writing Lewis Symbols for Atoms)LumenLearning43. The Ionic BondLumenLearning44. Crystals and Band TheoryLumenLearning45. The Covalent BondLumenLearning46. Covalent BondingOpenStax47. Lewis Symbols and StructuresOpenStax48. ElectronegativityLumenLearning49. Molecular GeometryLumenLearning50. Molecular Structure and PolarityOpenStaxVII. Solutions51. Intermolecular ForcesLumenLearning52. SolubilityLumenLearning53. MolarityOpenStax54. ColloidsOpenStaxVIII. Chemical Reactions55. Writing and Balancing Chemical EquationsOpenStax56. Molar MassLumenLearning57. Reaction StoichiometryLumenLearning58. Bond Energy and EnthalpyLumenLearning59. Strengths of Ionic and Covalent BondsOpenStax60. EntropyLumenLearning61. Activation Energy and Temperature DependenceLumenLearning62. CatalysisLumenLearningIX. Acids and Bases63. Acids and BasesLumenLearning64. Strength of AcidsLumenLearning65. Strength of BasesLumenLearning66. Buffer SolutionsLumenLearning67. BuffersSaylor Academy68. The Chemistry of Acid RainSaylor AcademyX. Electrochemistry69. Oxidation-Reduction ReactionsLumenLearning70. Electrochemical CellsLumenLearning71. Using Electrochemistry to Generate ElectricityLumenLearning72. ElectrolysisOpenStax73. CorrosionOpenStaxXI. Organic Chemistry74. Classes of Organic CompoundsLumenLearning75. Aliphatic HydrocarbonsLumenLearning76. HydrocarbonsOpenStax77. Alkenes and AlkynesLumenLearning78. Aromatic HydrocarbonsLumenLearning79. Functional Groups Names, Properties, and ReactionsLumenLearning80. PolymersSaylor Academy81. Synthetic Organic PolymersLumenLearningXII. Biochemistry82. LipidsLumenLearning83. CarbohydratesLumenLearning84. ProteinsLumenLearning85. Protein StructureLumenLearning86. Nucleic AcidsLumenLearning87. Classes of Organic CompoundsLumenLearning

Introductory Chemistry

24 The Structure of the Atom

LumenLearning

Overview of Atomic Structure

Atoms are made up of particles called protons, neutrons, and electrons, which are responsible for the mass and charge of atoms.

LEARNING OBJECTIVES

Discuss the electronic and structural properties of an atom

KEY TAKEAWAYS

Key Points

An atom is composed of two regions: the nucleus, which is in the center of the atom and contains protons and neutrons, and the outer region of the atom, which holds its electrons in orbit around the nucleus.

Protons and neutrons have approximately the same mass, about 1.67 × 10-24 grams, which scientists define as one atomic mass unit (amu) or one Dalton.

Each electron has a negative charge (-1) equal to the positive charge of a proton (+1).

Neutrons are uncharged particles found within the nucleus.

Key Terms

atom: The smallest possible amount of matter which still retains its identity as a chemical element, consisting of a nucleus surrounded by electrons.

proton: Positively charged subatomic particle forming part of the nucleus of an atom and determining the atomic number of an element. It weighs 1 amu.

neutron: A subatomic particle forming part of the nucleus of an atom. It has no charge. It is equal in mass to a proton or it weighs 1 amu.

An atom is the smallest unit of matter that retains all of the chemical properties of an element. Atoms combine to form molecules, which then interact to form solids, gases, or liquids. For example, water is composed of hydrogen and oxygen atoms that have combined to form water molecules. Many biological processes are devoted to reassembling molecules into different, more useful molecules.

Atomic Particles

Atoms consist of three basic particles: protons, electrons, and neutrons. The nucleus (center) of the atom contains the protons (positively charged) and the neutrons (no charge). The outermost regions of the atom are called electron shells and contain the electrons (negatively charged). Atoms have different properties based on the arrangement and number of their basic particles.

Elements, such as helium, depicted here, are made up of atoms. Atoms are made up of protons and neutrons located within the nucleus, with electrons in orbitals surrounding the nucleus.

The hydrogen atom (H) contains only one proton, one electron, and no neutrons. This can be determined using the atomic number and the mass number of the element (see the concept on atomic numbers and mass numbers).

Atomic Mass

Protons and neutrons have approximately the same mass, about 1.67 × 10-24 grams. Scientists define this amount of mass as one atomic mass unit (amu) or one Dalton. Although similar in mass, protons are positively charged, while neutrons have no charge. Therefore, the number of neutrons in an atom contributes significantly to its mass, but not to its charge.

Electrons are much smaller in mass than protons, weighing only 9.11 × 10-28 grams, or about 1/1800 of an atomic mass unit. Therefore, they do not contribute much to an element’s overall atomic mass. When considering atomic mass, it is customary to ignore the mass of any electrons and calculate the atom’s mass based on the number of protons and neutrons alone.

Electrons contribute greatly to the atom’s charge, as each electron has a negative charge equal to the positive charge of a proton. Scientists define these charges as “+1” and “-1. ” In an uncharged, neutral atom, the number of electrons orbiting the nucleus is equal to the number of protons inside the nucleus. In these atoms, the positive and negative charges cancel each other out, leading to an atom with no net charge.

Both protons and neutrons have a mass of 1 amu and are found in the nucleus. However, protons have a charge of +1, and neutrons are uncharged. Electrons have a mass of approximately 0 amu, orbit the nucleus, and have a charge of -1.

Exploring Electron Properties: Compare the behavior of electrons to that of other charged particles to discover properties of electrons such as charge and mass.

Volume of Atoms

Accounting for the sizes of protons, neutrons, and electrons, most of the volume of an atom—greater than 99 percent—is, in fact, empty space. Despite all this empty space, solid objects do not just pass through one another. The electrons that surround all atoms are negatively charged and cause atoms to repel one another, preventing atoms from occupying the same space. These intermolecular forces prevent you from falling through an object like your chair.

Atomic Number and Mass Number

The atomic number is the number of protons in an element, while the mass number is the number of protons plus the number of neutrons.

LEARNING OBJECTIVES

Determine the relationship between the mass number of an atom, its atomic number, its atomic mass, and its number of subatomic particles

KEY TAKEAWAYS

Key Points

Neutral atoms of each element contain an equal number of protons and electrons.

The number of protons determines an element’s atomic number and is used to distinguish one element from another.

The number of neutrons is variable, resulting in isotopes, which are different forms of the same atom that vary only in the number of neutrons they possess.

Together, the number of protons and the number of neutrons determine an element’s mass number.

Since an element’s isotopes have slightly different mass numbers, the atomic mass is calculated by obtaining the mean of the mass numbers for its isotopes.

Key Terms

mass number: The sum of the number of protons and the number of neutrons in an atom.

atomic number: The number of protons in an atom.

atomic mass: The average mass of an atom, taking into account all its naturally occurring isotopes.

Atomic Number

Neutral atoms of an element contain an equal number of protons and electrons. The number of protons determines an element’s atomic number (Z) and distinguishes one element from another. For example, carbon’s atomic number (Z) is 6 because it has 6 protons. The number of neutrons can vary to produce isotopes, which are atoms of the same element that have different numbers of neutrons. The number of electrons can also be different in atoms of the same element, thus producing ions (charged atoms). For instance, iron, Fe, can exist in its neutral state, or in the +2 and +3 ionic states.

Mass Number

An element’s mass number (A) is the sum of the number of protons and the number of neutrons. The small contribution of mass from electrons is disregarded in calculating the mass number. This approximation of mass can be used to easily calculate how many neutrons an element has by simply subtracting the number of protons from the mass number. Protons and neutrons both weigh about one atomic mass unit (amu). Isotopes of the same element will have the same atomic number but different mass numbers.

Carbon has an atomic number of six, and two stable isotopes with mass numbers of twelve and thirteen, respectively. Its relative atomic mass is 12.011.

Scientists determine the atomic mass by calculating the mean of the mass numbers for its naturally-occurring isotopes. Often, the resulting number contains a decimal. For example, the atomic mass of chlorine (Cl) is 35.45 amu because chlorine is composed of several isotopes, some (the majority) with an atomic mass of 35 amu (17 protons and 18 neutrons) and some with an atomic mass of 37 amu (17 protons and 20 neutrons).

Given an atomic number (Z) and mass number (A), you can find the number of protons, neutrons, and electrons in a neutral atom. For example, a lithium atom (Z=3, A=7 amu) contains three protons (found from Z), three electrons (as the number of protons is equal to the number of electrons in an atom), and four neutrons (7 – 3 = 4).

Isotopes

Isotopes are various forms of an element that have the same number of protons, but a different number of neutrons.

LEARNING OBJECTIVES

Discuss the properties of isotopes and their use in radiometric dating

KEY TAKEAWAYS

Key Points

Isotopes are atoms of the same element that contain an identical number of protons, but a different number of neutrons.

Despite having different numbers of neutrons, isotopes of the same element have very similar physical properties.

Some isotopes are unstable and will undergo radioactive decay to become other elements.

The predictable half-life of different decaying isotopes allows scientists to date material based on its isotopic composition, such as with Carbon-14 dating.

Key Terms

isotope: Any of two or more forms of an element where the atoms have the same number of protons, but a different number of neutrons within their nuclei.

half-life: The time it takes for half of the original concentration of an isotope to decay back to its more stable form.

radioactive isotopes: an atom with an unstable nucleus, characterized by excess energy available that undergoes radioactive decay and creates most commonly gamma rays, alpha or beta particles.

radiocarbon dating: Determining the age of an object by comparing the ratio of the [latex]^{14}\text{C}[/latex] concentration found in it to the amount of [latex]^{14}\text{C}[/latex] in the atmosphere.

What is an Isotope?

Isotopes are various forms of an element that have the same number of protons but a different number of neutrons. Some elements, such as carbon, potassium, and uranium, have multiple naturally-occurring isotopes. Isotopes are defined first by their element and then by the sum of the protons and neutrons present.

Carbon-12 (or [latex]^{12}\text{C}[/latex] ) contains six protons, six neutrons, and six electrons; therefore, it has a mass number of 12 amu (six protons and six neutrons).

Carbon-14 (or [latex]^{14}\text{C}[/latex]) contains six protons, eight neutrons, and six electrons; its atomic mass is 14 amu (six protons and eight neutrons).

While the mass of individual isotopes is different, their physical and chemical properties remain mostly unchanged.

Radiocarbon Dating

Carbon is normally present in the atmosphere in the form of gaseous compounds like carbon dioxide and methane. Carbon-14 ([latex]^{14}\text{C}[/latex]) is a naturally-occurring radioisotope that is created from atmospheric [latex]^{14}\text{N}[/latex] (nitrogen) by the addition of a neutron and the loss of a proton, which is caused by cosmic rays. This is a continuous process so more [latex]^{14}\text{C}[/latex] is always being created in the atmosphere. Once produced, the [latex]^{14}\text{C}[/latex] often combines with the oxygen in the atmosphere to form carbon dioxide. Carbon dioxide produced in this way diffuses in the atmosphere, is dissolved in the ocean, and is incorporated by plants via photosynthesis. Animals eat the plants and, ultimately, the radiocarbon is distributed throughout the biosphere.

In living organisms, the relative amount of [latex]^{14}\text{C}[/latex] in their body is approximately equal to the concentration of [latex]^{14}\text{C}[/latex] in the atmosphere. When an organism dies, it is no longer ingesting [latex]^{14}\text{C}[/latex], so the ratio between [latex]^{14}\text{C}[/latex] and [latex]^{12}\text{C}[/latex] will decline as [latex]^{14}\text{C}[/latex] gradually decays back to [latex]^{14}\text{N}[/latex]. This slow process, which is called beta decay, releases energy through the emission of electrons from the nucleus or positrons.

After approximately 5,730 years, half of the starting concentration of [latex]^{14}\text{C}[/latex] will have been converted back to [latex]^{14}\text{N}[/latex]. This is referred to as its half-life, or the time it takes for half of the original concentration of an isotope to decay back to its more stable form. Because the half-life of [latex]^{14}\text{C}[/latex] is long, it is used to date formerly-living objects such as old bones or wood. Comparing the ratio of the [latex]^{14}\text{C}[/latex] concentration found in an object to the amount of [latex]^{14}\text{C}[/latex] in the atmosphere, the amount of the isotope that has not yet decayed can be determined. On the basis of this amount, the age of the material can be accurately calculated, as long as the material is believed to be less than 50,000 years old. This technique is called radiocarbon dating, or carbon dating for short.

The age of carbon-containing remains that are less than about 50,000 years old, such as this pygmy mammoth, can be determined using carbon dating. (credit: Bill Faulkner, NPS)

Other elements have isotopes with different half lives. For example, [latex]^{40}\text{K}[/latex] (potassium-40) has a half-life of 1.25 billion years, and [latex]^{235}\text{U}[/latex] (uranium-235) has a half-life of about 700 million years. Scientists often use these other radioactive elements to date objects that are older than 50,000 years (the limit of carbon dating). Through the use of radiometric dating, scientists can study the age of fossils or other remains of extinct organisms.

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Chemistry

What is an atom? Facts about the building blocks of the universe

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By Tim Sharp, Daisy Dobrijevic published 15 December 2021

Atoms are made up of a nucleus, protons and electrons.

Atoms consist of a nucleus made of protons and neutrons orbited by electrons.

(Image credit: Rost-9D via Getty Images)

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History of the atom

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Atoms are the basic units of matter. Everything in the universe apart from energy is made of matter therefore atoms make up everything in the universe, according to Northwestern University. The term "atom" comes from the Greek word for indivisible, because it was once thought that atoms were the smallest things in the universe and could not be divided. We now know that atoms are made up of three particles known as subatomic particles: protons, neutrons and electrons — which are composed of even smaller particles, such as quarks.Atoms were created after the Big Bang 13.7 billion years ago. As the hot, dense new universe cooled, conditions became suitable for quarks and electrons to form. Quarks came together to form protons and neutrons, and these particles combined into nuclei. This all took place within the first few minutes of the universe's existence, according to CERN.It took 380,000 years for the universe to cool enough to slow down the electrons so that the nuclei could capture them to form the first atoms. The earliest atoms were primarily hydrogen and helium, which are still the most abundant elements in the universe, according to Jefferson Lab. Gravity eventually caused clouds of gas to coalesce and form stars, and heavier atoms were (and still are) created within the stars and sent throughout the universe when the star exploded (supernova).Related: What is antimatter, how is it made and is it dangerous? Subatomic particlesProtons and neutrons are heavier than electrons and reside in the nucleus at the center of the atom. Electrons are extremely lightweight and exist in a cloud orbiting the nucleus. The electron cloud has a radius 10,000 times greater than the nucleus, according to the Los Alamos National Laboratory.Protons and neutrons have approximately the same mass. However, one proton is about 1,835 times more massive than an electron. Atoms always have an equal number of protons and electrons, and the number of protons and neutrons is usually the same as well. Adding a proton to an atom makes a new element, while adding a neutron makes an isotope, or heavier version, of that atom.What does the nucleus do?The nucleus was discovered in 1911 by Ernest Rutherford, a physicist from New Zealand, according to the American Institute of Physics. In 1920, Rutherford proposed the name proton for the positively charged particles of the atom. He also theorized that there was a neutral particle within the nucleus, which James Chadwick, a British physicist and student of Rutherford's, was able to confirm in 1932.

Virtually all the mass of an atom resides in its nucleus, according to Chemistry LibreTexts. The protons and neutrons that make up the nucleus are approximately the same mass (the proton is slightly less) and have the same angular momentum, or spin.

The nucleus is held together by the strong force, one of the four basic forces in nature. This force between the protons and neutrons overcomes the repulsive electrical force that would otherwise push the protons apart, according to the rules of electricity. Some atomic nuclei are unstable because the binding force varies for different atoms based on the size of the nucleus. These atoms will then decay into other elements, such as carbon-14 decaying into nitrogen-14. What are protons?Protons are positively charged particles found within atomic nuclei. Rutherford discovered them in experiments with cathode-ray tubes that were conducted between 1911 and 1919. Protons are about 99.86% as massive as neutrons according to the Jefferson Lab.

The number of protons in an atom is unique to each element. For example, carbon atoms have six protons, hydrogen atoms have one and oxygen atoms have eight. The number of protons in an atom is referred to as the atomic number of that element. The number of protons also determines the chemical behavior of the element. Elements are arranged in the Periodic Table of the Elements in order of increasing atomic number.

Three quarks make up each proton — two "up" quarks (each with a two-thirds positive charge) and one "down" quark (with a one-third negative charge) — and they are held together by other subatomic particles called gluons, which are massless.What is an electron?Electrons are tiny compared to protons and neutrons, over 1,800 times smaller than either a proton or a neutron. Electrons are about 0.054% as massive as neutrons, according to Jefferson Lab.

Joseph John (J.J.) Thomson, a British physicist, discovered the electron in 1897, according to the Science History Institute. Originally known as "corpuscles," electrons have a negative charge and are electrically attracted to the positively charged protons. Electrons surround the atomic nucleus in pathways called orbitals, an idea that was put forth by Erwin Schrödinger, an Austrian physicist, in the 1920s. Today, this model is known as the quantum model or the electron cloud model. The inner orbitals surrounding the atom are spherical but the outer orbitals are much more complicated.

An atom's electron configuration refers to the locations of the electrons in a typical atom. Using the electron configuration and principles of physics, chemists can predict an atom's properties, such as stability, boiling point and conductivity, according to the Los Alamos National Laboratory.

Related: What is quantum entanglement?Where are neutrons located?The neutron's existence was theorized by Rutherford in 1920 and discovered by Chadwick in 1932, according to the American Physical Society. Neutrons were found during experiments when atoms were shot at a thin sheet of beryllium. Subatomic particles with no charge were released — the neutron.

Neutrons are uncharged particles found within all atomic nuclei (except for hydrogen). A neutron's mass is slightly larger than that of a proton. Like protons, neutrons are also made of quarks — one "up" quark (with a positive 2/3 charge) and two "down" quarks (each with a negative one-third charge). History of the atomThe theory of the atom dates at least as far back as 440 B.C. to Democritus, a Greek scientist and philosopher. Democritus most likely built his theory of atoms upon the work of past philosophers, according to Andrew G. Van Melsen, author of "From Atomos to Atom: The History of the Concept Atom" (Duquesne University Press, 1952). Democritus' explanation of the atom begins with a stone. A stone cut in half gives two halves of the same stone. If the stone were to be continuously cut, at some point there would exist a piece of the stone small enough that it could no longer be cut. The term "atom" comes from the Greek word for indivisible, which Democritus concluded must be the point at which a being (any form of matter) cannot be divided any more, according to educational website Lumen Learning. His explanation included the ideas that atoms exist separately from each other, that there are an infinite amount of atoms, that atoms are able to move, that they can combine together to create matter but do not merge to become a new atom, and that they cannot be divided, according to Universe Today. However, because most philosophers at the time — especially the very influential Aristotle — believed that all matter was created from earth, air, fire and water, Democritus' atomic theory was put aside.John Dalton, a British chemist, built upon Democritus' ideas in 1803 when he put forth his own atomic theory, according to the chemistry department at Purdue University. Dalton's theory included several ideas from Democritus, such as atoms are indivisible and indestructible and that different atoms form together to create all matter. Dalton's additions to the theory included the following ideas: That all atoms of a certain element were identical, that atoms of one element will have different weights and properties than atoms of another element, that atoms cannot be created or destroyed and that matter is formed by atoms combining in simple whole numbers.Thomson, the British physicist who discovered the electron in 1897, proved that atoms can be divided, according to the Chemical Heritage Foundation. He was able to determine the existence of electrons by studying the properties of electric discharge in cathode-ray tubes. According to Thomson's 1897 paper, the rays were deflected within the tube, which proved that there was something that was negatively charged within the vacuum tube. In 1899, Thomson published a description of his version of the atom, commonly known as the "plum pudding model." An excerpt of this paper is found on the Chem Team site. Thomson's model of the atom included a large number of electrons suspended in something that produced a positive charge giving the atom an overall neutral charge. His model resembled plum pudding, a popular British dessert that had raisins suspended in a round cake-like ball.The next scientist to further modify and advance the atomic model was Rutherford, who studied under Thomson, according to the chemistry department at Purdue University. In 1911, Rutherford published his version of the atom, which included a positively charged nucleus orbited by electrons. This model arose when Rutherford and his assistants fired alpha particles at thin sheets of gold. An alpha particle is made up of two protons and two neutrons, all held together by the same strong nuclear force that binds the nucleus, according to the Jefferson Lab. The scientists noticed that a small percentage of the alpha particles were scattered at very large angles to the original direction of motion while the majority passed right through hardly disturbed. Rutherford was able to approximate the size of the nucleus of the gold atom, finding it to be at least 10,000 times smaller than the size of the entire atom with much of the atom being empty space. Rutherford's model of the atom is still the basic model that is used today. Several other scientists furthered the atomic model, including Niels Bohr (built upon Rutherford's model to include properties of electrons based on the hydrogen spectrum), Erwin Schrödinger (developed the quantum model of the atom), Werner Heisenberg (stated that one cannot know both the position and velocity of an electron simultaneously), and Murray Gell-Mann and George Zweig (independently developed the theory that protons and neutrons were composed of quarks). Additional resourcesRead more about the early universe, from CERN. Learn more about the history of atomic chemistry in this video from Khan Academy. Check out this usefultive slide show about atoms from the Jefferson Lab.

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官方博客：Atom 的落幕（2022.6.8）

原文：Sunsetting Atom

当我们在 2011 年发布 Atom 的时候，我们希望给开发者们带来的是一个在深度可订制的同时又易于使用的编辑器，让更多的人能够从事软件开发。而现在，我们决定通过 Visual Studio Code 和 GitHub Codespaces 来继续实现这一目标，在云上提供快速和可靠的软件开发体验。

今天，我们正式宣布 Atom 的落幕，Atom 组织下的所有项目将会在 2022 年 12 月 1…

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