An illustration of thehelium atom, depicting thenucleus (pink) and theelectron 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 oneangstrom (10−10 m or100 pm).
Classification
Smallest recognized division of a chemical element
Atoms are the basic particles of thechemical elements and the fundamental building blocks ofmatter. An atom consists of anucleus ofprotons and generallyneutrons, surrounded by an electromagnetically bound swarm ofelectrons. 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 issodium, and any atom that contains 29 protons iscopper. Atoms with the same number of protons but a different number of neutrons are calledisotopes of the same element.
Atoms are extremely small, typically around 100 picometers across. A human hair is about a millioncarbon atoms wide. Atoms are smaller than the shortest wavelength of visible light, which means humans cannot see atoms with conventional microscopes. They are so small that accurately predicting their behavior usingclassical physics is not possible due toquantum effects.
More than 99.94%[1] of an atom'smass is in the nucleus. Protons have a positiveelectric 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 ofprotons and electrons are equal, as they normally are, then the atom is electrically neutral as a whole. A charged atom is called anion. 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 theelectromagnetic force. The protons and neutrons in the nucleus are attracted to each other by thenuclear 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 nucleussplits and leaves behind different elements. This is a form ofnuclear decay.
Atoms can attach to one or more other atoms bychemical bonds to formchemical compounds such asmolecules orcrystals. 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.
The basic idea that matter is made up of tiny indivisible particles is an old idea that appeared in many ancient cultures. The wordatom is derived from theancient Greek wordatomos,[a] which means "uncuttable". But this ancient idea was based in philosophical reasoning rather than scientific reasoning. Modern atomic theory is not based on these old concepts.[2][3] In the early 19th century, the scientistJohn Dalton found evidence that matter really is composed of discrete units, and so applied the wordatom to those units.[4]
Dalton's law of multiple proportions
Various atoms and molecules fromA New System of Chemical Philosophy (John Dalton 1808).
In the early 1800s, 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 any group of chemical compounds which all contain two particular chemical elements, the amount of Element A per measure of Element B will differ across these compounds by ratios of smallwhole numbers. This pattern suggested that each element combines with other elements in multiples of a basic unit of weight, with each element having a unit of unique weight. Dalton decided to call these units "atoms".[5]
For example, there are two types oftin 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 the grey oxide there is one atom of oxygen for every atom of tin, and in the white oxide there are two atoms of oxygen for every atom of tin (SnO andSnO2).[6][7]
Dalton also analyzediron 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. These substances are known today asiron(II) oxide andiron(III) oxide, and their formulas are FeO and Fe2O3 respectively. Iron(II) oxide's formula is normally written as FeO, but since it is a crystalline substance we could alternately write it as Fe2O2, and when we contrast that with Fe2O3, the 2:3 ratio for the oxygen is plain to see.[8][9]
As a final example:nitrous oxide is 63.3%nitrogen and 36.7% oxygen,nitric oxide is 44.05% nitrogen and 55.95% oxygen, andnitrogen 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 areN2O,NO, andNO2.[10][11]
Discovery of the electron
In 1897,J. J. Thomson discovered thatcathode rays can be deflected by electric and magnetic fields, which meant that cathode rays are not a form of light but made of electrically charged particles, and their charge was negative given the direction the particles were deflected in.[12] He measured these particles to be 1,700 times lighter thanhydrogen (the lightest atom).[13] He called these new particlescorpuscles but they were later renamedelectrons since these are the particles that carry electricity.[14] Thomson also showed that electrons were identical to particles given off byphotoelectric and radioactive materials.[15] Thomson explained that an electric current is the passing of electrons from one atom to the next, and when there was no current the electrons embedded themselves in the atoms. This in turn meant that atoms were not indivisible as scientists thought. The atom was composed of electrons whose negative charge was balanced out by some source of positive charge to create an electrically neutral atom. Ions, Thomson explained, must be atoms which have an excess or shortage of electrons.[16]
Discovery of the nucleus
TheRutherford scattering experiments: The extreme scattering of some alpha particles suggested the existence of a nucleus of concentrated charge.
The electrons in the atom logically had to be balanced out by a commensurate amount of positive charge, but Thomson had no idea where this positive charge came from, so he tentatively proposed that it was everywhere in the atom, the atom being in the shape of a sphere. This was the mathematically simplest hypothesis to fit the available evidence, or lack thereof. Following from this, Thomson imagined that the balance of electrostatic forces would distribute the electrons throughout the sphere in a more or less even manner.[17] Thomson's model is popularly known as theplum pudding model, though neither Thomson nor his colleagues used this analogy.[18] Thomson's model was incomplete, it was unable to predict any other properties of the elements such asemission spectra andvalencies. It was soon rendered obsolete by the discovery of theatomic nucleus.
Between 1908 and 1913,Ernest Rutherford and his colleaguesHans Geiger andErnest Marsden performed a series of experiments in which they bombarded thin foils of metal with a beam ofalpha particles. They did this to measure the scattering patterns of the alpha particles. They spotted a small number of 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. The deflections should have all been negligible. Rutherford proposed that the positive charge of the atom is concentrated in a tiny volume at the center of the atom and that the electrons surround this nucleus in a diffuse cloud. This nucleus carried almost all of the atom's mass. Only such an intense concentration of charge, anchored by its high mass, could produce an electric field that could deflect the alpha particles so strongly.[19]
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.
A problem in classical mechanics is that an accelerating charged particle radiateselectromagnetic radiation, causing the particle to losekinetic energy. Circular motion counts as acceleration, which means that an electron orbiting a central charge should spiral down into that nucleus as it loses speed. In 1913, the physicistNiels Bohr proposed a new 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 aphoton.[20] This quantization was used to explain why the electrons' orbits are stable and why elements absorb and emit electromagnetic radiation in discrete spectra.[21] Bohr's model could only predict the emission spectra of hydrogen, not atoms with more than one electron.
Back in 1815,William Prout observed that the atomic weights of many elements were multiples of hydrogen's atomic weight, which is in fact true for all of them if one takesisotopes into account. In 1898,J. J. Thomson found that the positive charge of a hydrogen ion is equal to the negative charge of an electron, and these were then the smallest known charged particles.[22] Thomson later found that the positive charge in an atom is a positive multiple of an electron's negative charge.[23] In 1913,Henry Moseley discovered that the frequencies of X-ray emissions from anexcited atom were a mathematical function of itsatomic number and hydrogen's nuclear charge. In 1919,Rutherford bombardednitrogen gas withalpha particles and detectedhydrogen ions being emitted from the gas, and concluded that they were produced by alpha particles hitting and splitting the nuclei of the nitrogen atoms.[24]
These observations led Rutherford to conclude that the hydrogen nucleus is a singular particle with a positive charge equal to the electron's negative charge.[25] He named this particle "proton" in 1920.[26] The number of protons in an atom (which Rutherford called the "atomic number"[27][28]) was found to be equal to the element's ordinal number on theperiodic table and therefore provided a simple and clear-cut way of distinguishing the elements from each other. The atomic weight of each element is higher than its proton number, so Rutherford hypothesized that the surplus weight was carried by unknown particles with no electric charge and a mass equal to that of the proton.
In 1928,Walter Bothe observed thatberyllium emitted a highly penetrating, electrically neutral radiation when bombarded with alpha particles. It was later discovered that this radiation could knock hydrogen atoms out ofparaffin wax. Initially it was thought to be high-energygamma radiation, since gamma radiation had a similar effect on electrons in metals, butJames Chadwick found that theionization 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.[29]
The current consensus model
The modern model of atomic orbitals draws zones where an electron is most likely to be found at any moment.
In 1925,Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (matrix mechanics).[30] One year earlier,Louis de Broglie had proposed that all particles behave like waves to some extent,[31] and in 1926Erwin Schrödinger used this idea to develop theSchrödinger equation, which describes electrons as three-dimensionalwaveforms rather than points in space.[32] A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both theposition andmomentum of a particle at a given point in time. This became known as theuncertainty principle, formulated by Werner Heisenberg in 1927.[30] 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.[33] Thus, the planetary model of the atom was discarded in favor of one that describedatomic orbital zones around the nucleus where a given electron is most likely to be found.[34][35] This model was able to explain observations of atomic behavior that previous models could not, such as certain structural andspectral patterns of atoms larger than hydrogen.
Though the wordatom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of varioussubatomic particles. The constituent particles of an atom are theelectron, theproton, and theneutron.
The electron is the least massive of these particles by four orders of magnitude at9.11×10−31 kg, with a negativeelectrical charge and a size that is too small to be measured using available techniques.[36] It was the lightest particle with a positive rest mass measured, until the discovery ofneutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. Electrons have been known since the late 19th century, mostly thanks toJ.J. Thomson; seehistory of subatomic physics for details.
Protons have a positive charge and a mass of1.6726×10−27 kg. The number of protons in an atom is called itsatomic 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 itproton.
Neutrons have no electrical charge and have a mass of1.6749×10−27 kg.[37][38] Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by thenuclear binding energy. Neutrons and protons (collectively known asnucleons) have comparable dimensions—on the order of2.5×10−15 m—although the 'surface' of these particles is not sharply defined.[39] The neutron was discovered in 1932 by the English physicistJames Chadwick.
In theStandard Model of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed ofelementary particles calledquarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of twoup quarks (each with charge +2/3) and onedown 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.[40][41]
The quarks are held together by thestrong interaction (or strong force), which is mediated bygluons. The protons and neutrons, in turn, are held to each other in the nucleus by thenuclear 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 ofgauge bosons, which are elementary particles that mediate physical forces.[40][41]
Thebinding energy needed for a nucleon to escape the nucleus, for various isotopes
All the bound protons and neutrons in an atom make up a tinyatomic nucleus, and are collectively callednucleons. The radius of a nucleus is approximately equal tofemtometres, where is the total number of nucleons.[42] 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 theresidual strong force. At distances smaller than 2.5 fm this force is much more powerful than theelectrostatic force that causes positively charged protons to repel each other.[43]
Atoms of the sameelement have the same number of protons, called theatomic number. Within a single element, the number of neutrons may vary, determining theisotope of that element. The total number of protons and neutrons determine thenuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoingradioactive decay.[44]
The proton, the electron, and the neutron are classified asfermions. Fermions obey thePauli exclusion principle which prohibitsidentical 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.[45]
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.[45]
Illustration of a nuclear fusion process that forms a deuterium nucleus, consisting of a proton and a neutron, from two protons. Apositron (e+)—anantimatter electron—is emitted along with an electronneutrino.
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—thecoulomb barrier—and fuse together into a single nucleus.[46]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.[47][48]
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 agamma ray, or the kinetic energy of abeta particle), as described byAlbert Einstein'smass–energy equivalence formula,E = mc2, wherem is the mass loss andc is thespeed of light. This deficit is part of thebinding 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.[49]
The fusion of two nuclei that create larger nuclei with lower atomic numbers thaniron andnickel—a total nucleon number of about 60—is usually anexothermic process that releases more energy than is required to bring them together.[50] It is this energy-releasing process that makes nuclear fusion instars a self-sustaining reaction. For heavier nuclei, the binding energy pernucleon begins to decrease. That means that a fusion process producing a nucleus that has an atomic number higher than about 26, and amass number higher than about 60, is anendothermic process. Thus, more massive nuclei cannot undergo an energy-producing fusion reaction that can sustain thehydrostatic equilibrium of a star.[45]
A potential well, showing, according toclassical mechanics, the minimum energyV(x) needed to reach each positionx. Classically, a particle with energyE is constrained to a range of positions betweenx1 andx2.
The electrons in an atom are attracted to the protons in the nucleus by theelectromagnetic force. This force binds the electrons inside anelectrostaticpotential 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 aparticle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensionalstanding wave—a wave form that does not move relative to the nucleus. This behavior is defined by anatomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured.[51] Only a discrete (orquantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form.[52] Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation.[53]
3D views of somehydrogen-like atomic orbitals showing probability density and phase (g orbitals and higher are not shown)
Each atomic orbital corresponds to a particularenergy level of the electron. The electron can change its state to a higher energy level by absorbing aphoton with sufficient energy to boost it into the new quantum state. Likewise, throughspontaneous 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 foratomic spectral lines.[52]
The amount of energy needed to remove or add an electron—theelectron binding energy—is far less than thebinding energy of nucleons. For example, it requires only 13.6 eV to strip aground-state electron from a hydrogen atom,[54] compared to 2.23 million eV for splitting adeuterium nucleus.[55] Atoms areelectrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are calledions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able tobond intomolecules and other types ofchemical compounds likeionic andcovalent networkcrystals.[56]
By definition, any two atoms with an identical number ofprotons in their nuclei belong to the samechemical element. Atoms with equal numbers of protons but a different number ofneutrons 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,[57] also called protium), one neutron (deuterium), two neutrons (tritium) andmore than two neutrons. The known elements form a set of atomic numbers, from the single-proton elementhydrogen up to the 118-proton elementoganesson.[58] 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.[59][60]
About 339 nuclides occur naturally onEarth,[61] of which 251 (about 74%) have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stabletheoretically, 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 theSolar System. This collection of 286 nuclides are known asprimordial nuclides. Finally, an additional 53 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such asradium fromuranium), or as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14).[62][note 1]
For 80 of the chemical elements, at least onestable 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 elementtin. Elements43,61, and all elements numbered83 or higher have no stable isotopes.[63]: 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 theshell 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 protonsand odd number of neutrons:hydrogen-2 (deuterium),lithium-6,boron-10, andnitrogen-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, andlutetium-176. Most odd-odd nuclei are highly unstable with respect tobeta decay, because the decay products are even-even, and are therefore more strongly bound, due tonuclear pairing effects.[64]
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 themass 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 actualmass of an atom at rest is often expressed indaltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom ofcarbon-12, which is approximately1.66×10−27 kg.[65]Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da.[66] The value of this number is called theatomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the dalton (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.[67] The heavieststable atom is lead-208,[59] with a mass of207.9766521 Da.[68]
As even the most massive atoms are far too light to work with directly, chemists instead use the unit ofmoles. One mole of atoms of any element always has the same number of atoms (about6.022×1023). 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 thedalton, 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.[65]
Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of anatomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus.[69] 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 achemical 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 aquantum mechanical property known asspin.[70] On theperiodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right).[71] Consequently, the smallest atom ishelium with a radius of 32 pm, while one of the largest iscaesium at 225 pm.[72]
When subjected to external forces, likeelectrical fields, the shape of an atom may deviate fromspherical symmetry. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown bygroup-theoretical considerations. Aspherical deviations might be elicited for instance incrystals, where large crystal-electrical fields may occur atlow-symmetry lattice sites.[73][74] Significantellipsoidal deformations have been shown to occur for sulfur ions[75] andchalcogen ions[76] inpyrite-type compounds.
Atomic dimensions are thousands of times smaller than the wavelengths oflight (400–700 nm) so they cannot be viewed using anoptical microscope, although individual atoms can be observed using ascanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width.[77] A single drop of water contains about 2 sextillion (2×1021) atoms of oxygen, and twice the number of hydrogen atoms.[78] A singlecaratdiamond with a mass of2×10−4 kg contains about 10 sextillion (1022) atoms ofcarbon.[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.[79]
This diagram shows thehalf-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.[80]
The most common forms of radioactive decay are:[81][82]
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 loweratomic number.
Beta decay (andelectron capture): these processes are regulated by theweak 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 anantineutrino, while proton to neutron transition (except in electron capture) causes the emission of apositron and aneutrino. 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 ofradioactive decay include ejection of neutrons or protons or clusters ofnucleons from a nucleus, or more than onebeta particle. An analog of gamma emission which allows excited nuclei to lose energy in a different way, isinternal 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 calledspontaneous nuclear fission.
Eachradioactive isotope has a characteristic decay time period—thehalf-life—that is determined by the amount of time needed for half of a sample to decay. This is anexponential 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.[80]
Elementary particles possess an intrinsic quantum mechanical property known asspin. This is analogous to theangular momentum of an object that is spinning around itscenter 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 reducedPlanck constant (ħ), with electrons, protons and neutrons all having spin1⁄2 ħ, or "spin-1⁄2". In an atom, electrons in motion around thenucleus possess orbitalangular momentum in addition to their spin, while the nucleus itself possesses angular momentum due to its nuclear spin.[83]
Themagnetic field produced by an atom—itsmagnetic 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 thePauli exclusion principle, in which no two electrons may be found in the samequantum 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.[84]
Inferromagnetic 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 anexchange 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.[84][85]
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 ofthermal equilibrium, but for certain elements (such asxenon-129) it is possible topolarize a significant proportion of the nuclear spin states so that they are aligned in the same direction—a condition calledhyperpolarization. This has important applications inmagnetic resonance imaging.[86][87]
Energy levels
These electron's energy levels (not to scale) are sufficient for ground states of atoms up tocadmium (5s2 4d10) inclusively. The top of the diagram is lower than an unbound electron state.
Thepotential energy of an electron in an atom isnegative relative to when thedistance from the nucleusgoes to infinity; its dependence on the electron'sposition reaches theminimum inside the nucleus, roughly ininverse proportion to the distance. In the quantum-mechanical model, a bound electron can occupy only a set ofstates centered on the nucleus, and each state corresponds to a specificenergy level; seetime-independent Schrödinger equation for a theoretical explanation. An energy level can be measured by theamount of energy needed to unbind the electron from the atom, and is usually given in units ofelectronvolts (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.[88] The electron's energy increases along withn because the (average) distance to the nucleus increases. Dependence of the energy onℓ is caused not by theelectrostatic potential of the nucleus, but by interaction between electrons.
For an electron totransition between two different states, e.g.ground state to firstexcited state, it must absorb or emit aphoton at an energy matching the difference in the potential energy of those levels, according to theNiels Bohr model, what can be precisely calculated by theSchrödinger 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; seeElectron properties.
The energy of an emitted photon is proportional to itsfrequency, so these specific energy levels appear as distinct bands in theelectromagnetic spectrum.[89] 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.[90]
An example of absorption lines in a spectrum
When a continuousspectrum 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 darkabsorption 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 ofemission lines from the photons emitted by the atoms.Spectroscopic measurements of the strength and width ofatomic spectral lines allow the composition and physical properties of a substance to be determined.[91]
Close examination of the spectral lines reveals that some display afine structure splitting. This occurs because ofspin–orbit coupling, which is an interaction between the spin and motion of the outermost electron.[92] When an atom is in an external magnetic field, spectral lines become split into three or more components; a phenomenon called theZeeman 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 multipleelectron 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.[93] The presence of an externalelectric field can cause a comparable splitting and shifting of spectral lines by modifying the electron energy levels, a phenomenon called theStark effect.[94]
If a bound electron is in an excited state, an interacting photon with the proper energy can causestimulated 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 makelasers, which can emit a coherent beam of light energy in a narrow frequency band.[95]
Valency is the combining power of an element. It is determined by the number of bonds it can form to other atoms or groups.[96] The outermost electron shell of an atom in its uncombined state is known as thevalence shell, and the electrons inthat shell are calledvalence electrons. The number of valence electrons determines thebondingbehavior with other atoms. Atoms tend tochemically react with each other in a manner that fills (or empties) their outer valence shells.[97] 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 compoundsodium 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 theorganic compounds.[98]
Thechemical elements are often displayed in aperiodic 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 thenoble gases.[99][100]
Quantities of atoms are found in different states of matter that depend on the physical conditions, such astemperature andpressure. By varying the conditions, materials can transition betweensolids,liquids,gases, andplasmas.[101] Within a state, a material can also exist in differentallotropes. An example of this is solid carbon, which can exist asgraphite ordiamond.[102] Gaseous allotropes exist as well, such asdioxygen andozone.
At temperatures close toabsolute zero, atoms can form aBose–Einstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale.[103][104] This super-cooled collection of atoms then behaves as a singlesuper atom, which may allow fundamental checks of quantum mechanical behavior.[105]
While atoms are too small to be seen, devices such as thescanning tunneling microscope (STM) enable their visualization at the surfaces of solids. The microscope uses thequantum tunneling phenomenon, which allows particles to pass through a barrier that would be insurmountable in the classical perspective. Electrons tunnel through the vacuum between twobiased 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 localdensity of the electronic states near theFermi level.[106][107] 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 isionized by removing one of its electrons, its trajectory when it passes through amagnetic 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. Themass spectrometer uses this principle to measure themass-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 includeinductively coupled plasma atomic emission spectroscopy andinductively coupled plasma mass spectrometry, both of which use a plasma to vaporize samples for analysis.[108]
Spectra ofexcited states can be used to analyze the atomic composition of distantstars. Specific lightwavelengths 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 agas-discharge lamp containing the same element.[110]Helium was discovered in this way in the spectrum of the Sun 23 years before it was found on Earth.[111]
Origin and current state
Baryonic matter forms about 4% of the total energy density of theobservable universe, with an average density of about 0.25 particles/m3 (mostlyprotons and electrons).[112] Within a galaxy such as theMilky Way, particles have a much higher concentration, with the density of matter in theinterstellar medium (ISM) ranging from 105 to 109 atoms/m3.[113] The Sun is believed to be inside theLocal Bubble, so the density in thesolar neighborhood is only about 103 atoms/m3.[114] 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;[115] the remainder of the mass is an unknowndark matter.[116] Hightemperature inside stars makes most "atoms" fully ionized, that is, separatesall electrons from the nuclei. Instellar remnants—with exception of their surface layers—an immensepressure make electron shells impossible.
Periodic table showing the origin of each element. Elements from carbon up to sulfur may be made in small stars by thealpha 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 ther-process.
Ubiquitousness and stability of atoms relies on theirbinding energy, which means that an atom has a lower energy than an unbound system of the nucleus and electrons. Where thetemperature is much higher thanionization potential, the matter exists in the form ofplasma—a gas of positively charged ions (possibly, bare nuclei) and electrons. When the temperature drops below the ionization potential, atoms becomestatistically favorable. Atoms (complete with bound electrons) became to dominate overcharged particles 380,000 years after the Big Bang—an epoch calledrecombination, when the expanding Universe cooled enough to allow electrons to become attached to nuclei.[120]
Isotopes such as lithium-6, as well as some beryllium and boron are generated in space throughcosmic ray spallation.[122] 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 insupernovae and collidingneutron stars through ther-process, and inAGB stars through thes-process, both of which involve the capture of neutrons by atomic nuclei.[123] Elements such aslead formed largely through the radioactive decay of heavier elements.[124]
Earth
Most of the atoms that make up theEarth and its inhabitants were present in their current form in thenebula that collapsed out of amolecular cloud to form theSolar System. The rest are the result of radioactive decay, and their relative proportion can be used to determine theage of the Earth throughradiometric dating.[125][126] Most of thehelium in the crust of the Earth (about 99% of the helium from gas wells, as shown by its lower abundance ofhelium-3) is a product ofalpha decay.[127]
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.[128] Some atoms on Earth have been artificially generated either deliberately or as by-products of nuclear reactors or explosions.[129][130] Of thetransuranic elements—those with atomic numbers greater than 92—onlyplutonium andneptunium occur naturally on Earth.[131][132] Transuranic elements have radioactive lifetimes shorter than the current age of the Earth[133] and thus identifiable quantities of these elements have long since decayed, with the exception of traces ofplutonium-244 possibly deposited by cosmic dust.[125] Natural deposits of plutonium and neptunium are produced byneutron capture in uranium ore.[134]
The Earth contains approximately1.33×1050 atoms.[135] Although small numbers of independent atoms ofnoble gases exist, such asargon,neon, andhelium, 99% ofthe atmosphere is bound in the form of molecules, includingcarbon dioxide anddiatomicoxygen andnitrogen. At the surface of the Earth, an overwhelming majority of atoms combine to form various compounds, includingwater,salt,silicates, andoxides. Atoms can also combine to create materials that do not consist of discrete molecules, includingcrystals and liquid or solidmetals.[136][137] This atomic matter forms networked arrangements that lack the particular type of small-scale interrupted order associated with molecular matter.[138]
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 aprimordial nuclide, and heavier elements generally have shorter half-lives. Nevertheless, an "island of stability" encompassing relatively long-lived isotopes of superheavy elements[139] with atomic numbers110 to114 might exist.[140] Predictions for the half-life of the most stable nuclide on the island range from a few minutes to millions of years.[141] In any case, superheavy elements (withZ > 104) would not exist due to increasingCoulomb repulsion (which results inspontaneous fission with increasingly short half-lives) in the absence of any stabilizing effects.[142]
Each particle of matter has a correspondingantimatter particle with the opposite electrical charge. Thus, thepositron is a positively chargedantielectron and theantiproton is a negatively charged equivalent of aproton. 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 ofbaryogenesis may offer an explanation. As a result, no antimatter atoms have been discovered in nature.[143][144] In 1996, the antimatter counterpart of the hydrogen atom (antihydrogen) was synthesized at theCERN laboratory inGeneva.[145][146]
Otherexotic 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 massivemuon, forming amuonic atom. These types of atoms can be used to test fundamental predictions of physics.[147][148][149]
^Pullman (1998).The Atom in the History of Human Thought, p. 199: "The constant ratios, expressible in terms of integers, of the weights of the constituents in composite bodies could be construed as evidence on a macroscopic scale of interactions at the microscopic level between basic units with fixed weights. For Dalton, this agreement strongly suggested a corpuscular structure of matter, even though it did not constitute definite proof."
^Thomson, J.J. (August 1901)."On bodies smaller than atoms".The Popular Science Monthly:323–335.Archived from the original on 1 December 2016. Retrieved21 June 2009.
^J. J. Thomson (1907).On the Corpuscular Theory of Matter, p. 26: "The simplest interpretation of these results is that the positive ions are the atoms or groups of atoms of various elements from which one or more corpuscles have been removed [...] while the negative electrified body is one with more corpuscles than the unelectrified one."
^J. J. Thomson (1907).The Corpuscular Theory of Matter, p. 103: "In default of exact knowledge of the nature of the way in which positive electricity occurs in the atom, we shall consider a case in which the positive electricity is distributed in the way most amenable to mathematical calculation, i.e., when it occurs as a sphere of uniform density, throughout which the corpuscles are distributed."
^J. J. Thomson (1907).The Corpuscular Theory of Matter. p. 26–27: "In an unelectrified atom there are as many units of positive electricity as there are of negative; an atom with a unit of positive charge is a neutral atom which has lost one corpuscle, while an atom with a unit of negative charge is a neutral atom to which an additional corpuscle has been attached."
^The Development of the Theory of Atomic Structure (Rutherford 1936). Reprinted inBackground to Modern Science: Ten Lectures at Cambridge arranged by the History of Science Committee 1936: "In 1919 I showed that when light atoms were bombarded by α-particles they could be broken up with the emission of a proton, or hydrogen nucleus. We therefore presumed that a proton must be one of the units of which the nuclei of other atoms were composed..."
^Orme Masson (1921)."The Constitution of Atoms".The London, Edinburgh, and Dublin Philosophical Magazine and Journal of Science.41 (242):281–285.doi:10.1080/14786442108636219. Footnote by Ernest Rutherford: 'At the time of writing this paper in Australia, Professor Orme Masson was not aware that the name "proton" had already been suggested as a suitable name for the unit of mass nearly 1, in terms of oxygen 16, that appears to enter into the nuclear structure of atoms. The question of a suitable name for this unit was discussed at an informal meeting of a number of members of Section A of the British Association at Cardiff this year. The name "baron" suggested by Professor Masson was mentioned, but was considered unsuitable on account of the existing variety of meanings. Finally the name "proton" met with general approval, particularly as it suggests the original term "protyle " given by Prout in his well-known hypothesis that all atoms are built up of hydrogen. The need of a special name for the nuclear unit of mass 1 was drawn attention to by Sir Oliver Lodge at the Sectional meeting, and the writer then suggested the name "proton."'
^Eric Scerri (2020).The Periodic Table: Its Story and Its Significance, p. 185
^Helge Kragh (2012).Niels Bohr and the Quantum Atom, p. 33
^Shultis, J. Kenneth; Faw, Richard E. (2002).Fundamentals of Nuclear Science and Engineering. CRC Press. pp. 10–17.ISBN978-0-8247-0834-4.OCLC123346507.
^Matis, Howard S. (9 August 2000)."The Isotopes of Hydrogen".Guide to the Nuclear Wall Chart. Lawrence Berkeley National Lab.Archived from the original on 18 December 2007.
^Tuli, Jagdish K. (April 2005)."Nuclear Wallet Cards". National Nuclear Data Center, Brookhaven National Laboratory.Archived from the original on 3 October 2011.
^Staff (2007)."Small Miracles: Harnessing nanotechnology". Oregon State University.Archived from the original on 21 May 2011. – describes the width of a human hair as105 nm and 10 carbon atoms as spanning 1 nm.
^Padilla, Michael J.; Miaoulis, Ioannis; Cyr, Martha (2002).Prentice Hall Science Explorer: Chemical Building Blocks. Upper Saddle River, New Jersey: Prentice-Hall, Inc. p. 32.ISBN978-0-13-054091-1.OCLC47925884.There are 2,000,000,000,000,000,000,000 (that's 2 sextillion) atoms of oxygen in one drop of water—and twice as many atoms of hydrogen.
^Goebel, Greg (1 September 2007)."[4.3] Magnetic Properties of the Atom".Elementary Quantum Physics. In The Public Domain website. Archived from the original on 29 June 2011.
^Yarris, Lynn (Spring 1997)."Talking Pictures".Berkeley Lab Research Review.Archived from the original on 13 January 2008.
^Liang, Z.-P.; Haacke, E.M. (1999). Webster, J.G. (ed.).Encyclopedia of Electrical and Electronics Engineering: Magnetic Resonance Imaging. Vol. 2. John Wiley & Sons. pp. 412–426.ISBN978-0-471-13946-1.
^Zeghbroeck, Bart J. Van (1998)."Energy levels". Shippensburg University. Archived fromthe original on 15 January 2005.
^Jacox, Marilyn; Gadzuk, J. William (November 1997)."Scanning Tunneling Microscope". National Institute of Standards and Technology.Archived from the original on 7 January 2008.
^"The Nobel Prize in Physics 1986". The Nobel Foundation.Archived from the original on 17 September 2008. Retrieved11 January 2008. In particular, see the Nobel lecture by G. Binnig and H. Rohrer.
^Lochner, Jim; Gibb, Meredith; Newman, Phil (30 April 2007)."What Do Spectra Tell Us?". NASA/Goddard Space Flight Center.Archived from the original on 16 January 2008.
^Winter, Mark (2007)."Helium". WebElements.Archived from the original on 30 December 2007.
^Choppin, Gregory R.; Liljenzin, Jan-Olov; Rydberg, Jan (2001).Radiochemistry and Nuclear Chemistry. Elsevier. p. 441.ISBN978-0-7506-7463-8.OCLC162592180.
Oliver Manuel (2001).Origin of Elements in the Solar System: Implications of Post-1957 Observations. Springer.ISBN978-0-306-46562-8.OCLC228374906.
Andrew G. van Melsen (2004) [1952].From Atomos to Atom: The History of the Concept Atom. Translated by Henry J. Koren. Dover Publications.ISBN0-486-49584-1.
J.P. Millington (1906).John Dalton. J. M. Dent & Co. (London); E. P. Dutton & Co. (New York).
Charles H. Holbrow; James N. Lloyd; Joseph C. Amato; Enrique Galvez; M. Elizabeth Parks (2010).Modern Introductory Physics. Springer Science & Business Media.ISBN978-0-387-79079-4.
Eric R. Scerri (2020).The Periodic Table, Its Story and Its Significance (2nd ed.). New York: Oxford University Press.ISBN978-0-190-91436-3.
Further reading
Gangopadhyaya, Mrinalkanti (1981).Indian Atomism: History and Sources. Atlantic Highlands, New Jersey: Humanities Press.ISBN978-0-391-02177-8.OCLC10916778.
Wurtz, Charles Adolphe (1881).The Atomic Theory. New York: D. Appleton and company.ISBN978-0-559-43636-9.{{cite book}}:ISBN / Date incompatibility (help)
