Isotopes are distinct nuclear species (or nuclides ) of the same chemical element . They have the same atomic number (number of protons in their nuclei ) and position in the periodic table (and hence belong to the same chemical element), but different nucleon numbers ( mass numbers ) due to different numbers of neutrons in their nuclei. While all isotopes of a given element have similar chemical properties, they have different atomic masses and physical properties.
183-534: The term isotope is derived from the Greek roots isos ( ἴσος "equal") and topos ( τόπος "place"), meaning "the same place"; thus, the meaning behind the name is that different isotopes of a single element occupy the same position on the periodic table . It was coined by Scottish doctor and writer Margaret Todd in a 1913 suggestion to the British chemist Frederick Soddy , who popularized the term. The number of protons within
366-455: A = m 1 x 1 + m 2 x 2 + . . . + m N x N {\displaystyle {\overline {m}}_{a}=m_{1}x_{1}+m_{2}x_{2}+...+m_{N}x_{N}} where m 1 , m 2 , ..., m N are the atomic masses of each individual isotope, and x 1 , ..., x N are the relative abundances of these isotopes. Several applications exist that capitalize on
549-454: A spin isomer when the formation of an intermediate excited state has a spin far different from that of the ground state. Gamma-ray emission is hindered if the spin of the post-emission state differs greatly from that of the emitting state, especially if the excitation energy is low. The excited state in this situation is a good candidate to be metastable if there are no other states of intermediate spin with excitation energies less than that of
732-470: A capacity of 2×1 + 2×3 + 2×5 = 18. The fourth shell contains one 4s orbital, three 4p orbitals, five 4d orbitals, and seven 4f orbitals, thus leading to a capacity of 2×1 + 2×3 + 2×5 + 2×7 = 32. Higher shells contain more types of orbitals that continue the pattern, but such types of orbitals are not filled in the ground states of known elements. The subshell types are characterized by the quantum numbers . Four numbers describe an orbital in an atom completely:
915-467: A characteristic abundance, naturally occurring elements have well-defined atomic weights , defined as the average mass of a naturally occurring atom of that element. All elements have multiple isotopes , variants with the same number of protons but different numbers of neutrons . For example, carbon has three naturally occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, and
1098-486: A chromium atom to have a [Ar] 3d 4s configuration than an [Ar] 3d 4s one. A similar anomaly occurs at copper , whose atom has a [Ar] 3d 4s configuration rather than the expected [Ar] 3d 4s . These are violations of the Madelung rule. Such anomalies, however, do not have any chemical significance: most chemistry is not about isolated gaseous atoms, and the various configurations are so close in energy to each other that
1281-424: A device with adjacent layers of P-type and N-type silicon . Ionizing radiation directly penetrates the junction and creates electron–hole pairs . Nuclear isomers could replace other isotopes, and with further development, it may be possible to turn them on and off by triggering decay as needed. Current candidates for such use include Ag , Ho , Lu , and Am . As of 2004, the only successfully triggered isomer
1464-419: A different mass number. For example, carbon-12 , carbon-13 , and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13, and 14, respectively. The atomic number of carbon is 6, which means that every carbon atom has 6 protons so that the neutron numbers of these isotopes are 6, 7, and 8 respectively. A nuclide is a species of an atom with a specific number of protons and neutrons in
1647-505: A double pairing of 2 protons and 2 neutrons prevents any nuclides containing five ( 2 He , 3 Li ) or eight ( 4 Be ) nucleons from existing long enough to serve as platforms for the buildup of heavier elements via nuclear fusion in stars (see triple alpha process ). Only five stable nuclides contain both an odd number of protons and an odd number of neutrons. The first four "odd-odd" nuclides occur in low mass nuclides, for which changing
1830-451: A gamma ray from an excited nuclear state allows the nucleus to lose energy and reach a lower-energy state, sometimes its ground state . In certain cases, the excited nuclear state following a nuclear reaction or other type of radioactive decay can become a metastable nuclear excited state. Some nuclei are able to stay in this metastable excited state for minutes, hours, days, or occasionally far longer. The process of isomeric transition
2013-419: A given element all have the same number of electrons and share a similar electronic structure. Because the chemical behaviour of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behaviour. The main exception to this is the kinetic isotope effect : due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of
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#17328554246562196-407: A glowing patch on the plate at the point it struck. Thomson observed two separate parabolic patches of light on the photographic plate (see image), which suggested two species of nuclei with different mass-to-charge ratios. He wrote "There can, therefore, I think, be little doubt that what has been called neon is not a simple gas but a mixture of two gases, one of which has an atomic weight about 20 and
2379-580: A good fit for either group: hydrogen is neither highly oxidizing nor highly reducing and is not reactive with water. Hydrogen thus has properties corresponding to both those of the alkali metals and the halogens, but matches neither group perfectly, and is thus difficult to place by its chemistry. Therefore, while the electronic placement of hydrogen in group 1 predominates, some rarer arrangements show either hydrogen in group 17, duplicate hydrogen in both groups 1 and 17, or float it separately from all groups. This last option has nonetheless been criticized by
2562-419: A half-life of 4,570 years, is more stable. 90 Th has a remarkably low-lying metastable isomer only 8.355 733 554 021 (8) eV above the ground state. This low energy produces "gamma rays" at a wavelength of 148.382 182 8827 (15) nm , in the far ultraviolet , which allows for direct nuclear laser spectroscopy . Such ultra-precise spectroscopy, however, could not begin without
2745-417: A kainosymmetric first element of a group. The group 18 placement of helium nonetheless remains near-universal due to its extreme inertness. Additionally, tables that float both hydrogen and helium outside all groups may rarely be encountered. In many periodic tables, the f-block is shifted one element to the right, so that lanthanum and actinium become d-block elements in group 3, and Ce–Lu and Th–Lr form
2928-420: A metastable isomeric state. These fragments are usually produced in a highly excited state, in terms of energy and angular momentum , and go through a prompt de-excitation. At the end of this process, the nuclei can populate both the ground and the isomeric states. If the half-life of the isomers is long enough, it is possible to measure their production rate and compare it to that of the ground state, calculating
3111-468: A nonoptimal number of neutrons or protons decay by beta decay (including positron emission ), electron capture , or other less common decay modes such as spontaneous fission and cluster decay . Most stable nuclides are even-proton-even-neutron, where all numbers Z , N , and A are even. The odd- A stable nuclides are divided (roughly evenly) into odd-proton-even-neutron, and even-proton-odd-neutron nuclides. Stable odd-proton-odd-neutron nuclides are
3294-427: A nucleus. As the number of protons increases, so does the ratio of neutrons to protons necessary to ensure a stable nucleus (see graph at right). For example, although the neutron:proton ratio of 2 He is 1:2, the neutron:proton ratio of 92 U is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 ( Z = N ). The nuclide 20 Ca (calcium-40)
3477-430: A one- or two-letter chemical symbol ; those for hydrogen, helium, and lithium are respectively H, He, and Li. Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element. Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with
3660-568: A prerequisite to their use in such weapons, is disputed. Nonetheless a 12-member Hafnium Isomer Production Panel (HIPP) was created in 2003 to assess means of mass-producing the isotope. Technetium isomers 43 Tc (with a half-life of 6.01 hours) and 43 Tc (with a half-life of 61 days) are used in medical and industrial applications. Nuclear batteries use small amounts (milligrams and microcuries ) of radioisotopes with high energy densities. In one betavoltaic device design, radioactive material sits atop
3843-405: A product of stellar nucleosynthesis or another type of nucleosynthesis such as cosmic ray spallation , and have persisted down to the present because their rate of decay is very slow (e.g. uranium-238 and potassium-40 ). Post-primordial isotopes were created by cosmic ray bombardment as cosmogenic nuclides (e.g., tritium , carbon-14 ), or by the decay of a radioactive primordial isotope to
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#17328554246564026-505: A proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio ( 1 H , 3 Li , 5 B , and 7 N ; spins 1, 1, 3, 1). The only other entirely "stable" odd-odd nuclide, 73 Ta (spin 9), is thought to be the rarest of the 251 stable nuclides, and is the only primordial nuclear isomer , which has not yet been observed to decay despite experimental attempts. Many odd-odd radionuclides (such as
4209-489: A quantity known as spin , conventionally labelled "up" or "down". In a cold atom (one in its ground state), electrons arrange themselves in such a way that the total energy they have is minimized by occupying the lowest-energy orbitals available. Only the outermost electrons (so-called valence electrons ) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called core electrons . Elements are known with up to
4392-478: A radioactive radiogenic nuclide daughter (e.g. uranium to radium ). A few isotopes are naturally synthesized as nucleogenic nuclides, by some other natural nuclear reaction , such as when neutrons from natural nuclear fission are absorbed by another atom. As discussed above, only 80 elements have any stable isotopes, and 26 of these have only one stable isotope. Thus, about two-thirds of stable elements occur naturally on Earth in multiple stable isotopes, with
4575-509: A series of ten transition elements ( lutetium through mercury ) follows, and finally six main-group elements ( thallium through radon ) complete the period. From lutetium onwards the 4f orbitals are in the core, and from thallium onwards so are the 5d orbitals. The seventh row is analogous to the sixth row: 7s fills ( francium and radium ), then 5f ( actinium to nobelium ), then 6d ( lawrencium to copernicium ), and finally 7p ( nihonium to oganesson ). Starting from lawrencium
4758-423: A single stable isotope (of these, 19 are so-called mononuclidic elements , having a single primordial stable isotope that dominates and fixes the atomic weight of the natural element to high precision; 3 radioactive mononuclidic elements occur as well). In total, there are 251 nuclides that have not been observed to decay. For the 80 elements that have one or more stable isotopes, the average number of stable isotopes
4941-462: A small increase especially at the end of each transition series. As metal atoms tend to lose electrons in chemical reactions, ionisation energy is generally correlated with chemical reactivity, although there are other factors involved as well. The opposite property to ionisation energy is the electron affinity , which is the energy released when adding an electron to the atom. A passing electron will be more readily attracted to an atom if it feels
5124-402: A specialized branch of relativistic quantum mechanics focusing on the properties of superheavy elements , the project's opinion was that such interest-dependent concerns should not have any bearing on how the periodic table is presented to "the general chemical and scientific community". Other authors focusing on superheavy elements since clarified that the "15th entry of the f-block represents
5307-401: A spin of 1. Similarly, 43 Tc has a spin of 1/2 and must gamma-decay to 43 Tc with a spin of 9/2. While most metastable isomers decay through gamma-ray emission, they can also decay through internal conversion . During internal conversion, energy of nuclear de-excitation is not emitted as a gamma ray, but is instead used to accelerate one of the inner electrons of
5490-491: A stable (non-radioactive) element was found by J. J. Thomson in 1912 as part of his exploration into the composition of canal rays (positive ions). Thomson channelled streams of neon ions through parallel magnetic and electric fields, measured their deflection by placing a photographic plate in their path, and computed their mass to charge ratio using a method that became known as the Thomson's parabola method. Each stream created
5673-428: A sufficiently precise initial estimate of the wavelength, something that was only achieved in 2024 after two decades of effort. The energy is so low that the ionization state of the atom affects its half-life. Neutral 90 Th decays by internal conversion with a half-life of 7 ± 1 μs , but because the isomeric energy is less than thorium's second ionization energy of 11.5 eV , this channel
Isotope - Misplaced Pages Continue
5856-621: A total 30 + 2(9) = 48 stable odd-even isotopes. There are also five primordial long-lived radioactive odd-even isotopes, 37 Rb , 49 In , 75 Re , 63 Eu , and 83 Bi . The last two were only recently found to decay, with half-lives greater than 10 years. Actinides with odd neutron number are generally fissile (with thermal neutrons ), whereas those with even neutron number are generally not, though they are fissionable with fast neutrons . All observationally stable odd-odd nuclides have nonzero integer spin. This
6039-418: A very small fraction have eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under a single element. When atomic mass is shown, it is usually the weighted average of naturally occurring isotopes; but if no isotopes occur naturally in significant quantities, the mass of the most stable isotope usually appears, often in parentheses. In the standard periodic table,
6222-433: Is 43 Tc , which decays with a half-life of about 6 hours by emitting a gamma ray of 140 keV of energy; this is close to the energy of medical diagnostic X-rays. Nuclear isomers have long half-lives because their gamma decay is "forbidden" from the large change in nuclear spin needed to emit a gamma ray. For example, 73 Ta has a spin of 9 and must gamma-decay to 73 Ta with
6405-556: Is 73 Ta , which is present in all tantalum samples at about 1 part in 8,300. Its half-life is at least 10 years, markedly longer than the age of the universe . The low excitation energy of the isomeric state causes both gamma de-excitation to the Ta ground state (which itself is radioactive by beta decay, with a half-life of only 8 hours) and direct electron capture to hafnium or beta decay to tungsten to be suppressed due to spin mismatches. The origin of this isomer
6588-681: Is aluminium-26 , which is not naturally found on Earth but is found in abundance on an astronomical scale. The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. Before the discovery of isotopes, empirically determined noninteger values of atomic mass confounded scientists. For example, a sample of chlorine contains 75.8% chlorine-35 and 24.2% chlorine-37 , giving an average atomic mass of 35.5 atomic mass units . According to generally accepted cosmology theory , only isotopes of hydrogen and helium, traces of some isotopes of lithium and beryllium, and perhaps some boron, were created at
6771-547: Is 251/80 ≈ 3.14 isotopes per element. The proton:neutron ratio is not the only factor affecting nuclear stability. It depends also on evenness or oddness of its atomic number Z , neutron number N and, consequently, of their sum, the mass number A . Oddness of both Z and N tends to lower the nuclear binding energy , making odd nuclei, generally, less stable. This remarkable difference of nuclear binding energy between neighbouring nuclei, especially of odd- A isobars , has important consequences: unstable isotopes with
6954-530: Is a metastable state of an atomic nucleus , in which one or more nucleons (protons or neutrons) occupy excited state levels (higher energy levels). "Metastable" describes nuclei whose excited states have half-lives 100 to 1000 times longer than the half-lives of the excited nuclear states that decay with a "prompt" half life (ordinarily on the order of 10 seconds). The term "metastable" is usually restricted to isomers with half-lives of 10 seconds or longer. Some references recommend 5 × 10 seconds to distinguish
7137-633: Is a radioactive form of carbon, whereas C and C are stable isotopes. There are about 339 naturally occurring nuclides on Earth, of which 286 are primordial nuclides , meaning that they have existed since the Solar System 's formation. Primordial nuclides include 35 nuclides with very long half-lives (over 100 million years) and 251 that are formally considered as " stable nuclides ", because they have not been observed to decay. In most cases, for obvious reasons, if an element has stable isotopes, those isotopes predominate in
7320-415: Is also present in ionic radii , though it is more difficult to examine because the most common ions of consecutive elements normally differ in charge. Ions with the same electron configuration decrease in size as their atomic number rises, due to increased attraction from the more positively charged nucleus: thus for example ionic radii decrease in the series Se , Br , Rb , Sr , Y , Zr , Nb , Mo , Tc . Ions of
7503-470: Is an icon of chemistry and is widely used in physics and other sciences. It is a depiction of the periodic law , which states that when the elements are arranged in order of their atomic numbers an approximate recurrence of their properties is evident. The table is divided into four roughly rectangular areas called blocks . Elements in the same group tend to show similar chemical characteristics. Vertical, horizontal and diagonal trends characterize
Isotope - Misplaced Pages Continue
7686-411: Is an s-block element, whereas all other noble gases are p-block elements. However it is unreactive at standard conditions, and has a full outer shell: these properties are like the noble gases in group 18, but not at all like the reactive alkaline earth metals of group 2. For these reasons helium is nearly universally placed in group 18 which its properties best match; a proposal to move helium to group 2
7869-663: Is because the single unpaired neutron and unpaired proton have a larger nuclear force attraction to each other if their spins are aligned (producing a total spin of at least 1 unit), instead of anti-aligned. See deuterium for the simplest case of this nuclear behavior. Only 78 Pt , 4 Be , and 7 N have odd neutron number and are the most naturally abundant isotope of their element. Elements are composed either of one nuclide ( mononuclidic elements ), or of more than one naturally occurring isotopes. The unstable (radioactive) isotopes are either primordial or postprimordial. Primordial isotopes were
8052-435: Is denoted with symbols "u" (for unified atomic mass unit) or "Da" (for dalton ). The atomic masses of naturally occurring isotopes of an element determine the standard atomic weight of the element. When the element contains N isotopes, the expression below is applied for the average atomic mass m ¯ a {\displaystyle {\overline {m}}_{a}} : m ¯
8235-420: Is expected to show slightly less inertness than neon and to form (HeO)(LiF) 2 with a structure similar to the analogous beryllium compound (but with no expected neon analogue), have resulted in more chemists advocating a placement of helium in group 2. This relates to the electronic argument, as the reason for neon's greater inertness is repulsion from its filled p-shell that helium lacks, though realistically it
8418-479: Is forbidden in thorium cations and 90 Th decays by gamma emission with a half-life of 1740 ± 50 s . This conveniently moderate lifetime allows the development of a nuclear clock of unprecedented accuracy. The most common mechanism for suppression of gamma decay of excited nuclei, and thus the existence of a metastable isomer, is lack of a decay route for the excited state that will change nuclear angular momentum along any given direction by
8601-600: Is impossible when the nucleus begins in a zero-spin state, as such an emission would not conserve angular momentum. Hafnium isomers (mainly Hf) have been considered as weapons that could be used to circumvent the Nuclear Non-Proliferation Treaty , since it is claimed that they can be induced to emit very strong gamma radiation . This claim is generally discounted. DARPA had a program to investigate this use of both nuclear isomers. The potential to trigger an abrupt release of energy from nuclear isotopes,
8784-443: Is mysterious, though it is believed to have been formed in supernovae (as are most other heavy elements). Were it to relax to its ground state, it would release a photon with a photon energy of 75 keV . It was first reported in 1988 by C. B. Collins that theoretically Ta can be forced to release its energy by weaker X-rays, although at that time this de-excitation mechanism had never been observed. However,
8967-628: Is observationally the heaviest stable nuclide with the same number of neutrons and protons. All stable nuclides heavier than calcium-40 contain more neutrons than protons. Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin ). No element has nine or eight stable isotopes. Five elements have seven stable isotopes, eight have six stable isotopes, ten have five stable isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have two stable isotopes (counting 73 Ta as stable), and 26 elements have only
9150-453: Is similar to gamma emission from any excited nuclear state, but differs by involving excited metastable states of nuclei with longer half-lives. As with other excited states, the nucleus can be left in an isomeric state following the emission of an alpha particle , beta particle , or some other type of particle. The gamma ray may transfer its energy directly to one of the most tightly bound electrons , causing that electron to be ejected from
9333-413: Is specified by the name of the particular element (this indicates the atomic number) followed by a hyphen and the mass number (e.g. helium-3 , helium-4 , carbon-12 , carbon-14 , uranium-235 and uranium-239 ). When a chemical symbol is used, e.g. "C" for carbon, standard notation (now known as "AZE notation" because A is the mass number , Z the atomic number , and E for element ) is to indicate
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#17328554246569516-439: Is the decay of a nuclear isomer to a lower-energy nuclear state. The actual process has two types (modes): Isomers may decay into other elements, though the rate of decay may differ between isomers. For example, Lu can beta-decay to Hf with a half-life of 160.4 d, or it can undergo isomeric transition to Lu with a half-life of 160.4 d, which then beta-decays to Hf with a half-life of 6.68 d. The emission of
9699-399: Is universally accepted by chemists that these configurations are exceptional and that the d-block really ends in accordance with the Madelung rule at zinc, cadmium, and mercury. The relevant fact for placement is that lanthanum and actinium (like thorium) have valence f-orbitals that can become occupied in chemical environments, whereas lutetium and lawrencium do not: their f-shells are in
9882-419: Is unlikely that helium-containing molecules will be stable outside extreme low-temperature conditions (around 10 K ). The first-row anomaly in the periodic table has additionally been cited to support moving helium to group 2. It arises because the first orbital of any type is unusually small, since unlike its higher analogues, it does not experience interelectronic repulsion from a smaller orbital of
10065-469: The Big Bang , while all other nuclides were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously produced nuclides. (See nucleosynthesis for details of the various processes thought responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through
10248-444: The CNO cycle . The nuclides 3 Li and 5 B are minority isotopes of elements that are themselves rare compared to other light elements, whereas the other six isotopes make up only a tiny percentage of the natural abundance of their elements. 53 stable nuclides have an even number of protons and an odd number of neutrons. They are a minority in comparison to
10431-538: The Girdler sulfide process . Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project ) by a type of production mass spectrometry . Periodic table The periodic table , also known as the periodic table of the elements , is an ordered arrangement of the chemical elements into rows (" periods ") and columns (" groups "). It
10614-413: The atom's nucleus is called its atomic number and is equal to the number of electrons in the neutral (non-ionized) atom. Each atomic number identifies a specific element, but not the isotope; an atom of a given element may have a wide range in its number of neutrons . The number of nucleons (both protons and neutrons) in the nucleus is the atom's mass number , and each isotope of a given element has
10797-410: The binding energy of the nucleus (see mass defect ), the slight difference in mass between proton and neutron, and the mass of the electrons associated with the atom, the latter because the electron:nucleon ratio differs among isotopes. The mass number is a dimensionless quantity . The atomic mass, on the other hand, is measured using the atomic mass unit based on the mass of the carbon-12 atom. It
10980-465: The fissile 92 U . Because of their odd neutron numbers, the even-odd nuclides tend to have large neutron capture cross-sections, due to the energy that results from neutron-pairing effects. These stable even-proton odd-neutron nuclides tend to be uncommon by abundance in nature, generally because, to form and enter into primordial abundance, they must have escaped capturing neutrons to form yet other stable even-even isotopes, during both
11163-423: The gamma decay from a metastable state is referred to as isomeric transition, but this process typically resembles shorter-lived gamma decays in all external aspects with the exception of the long-lived nature of the meta-stable parent nuclear isomer. The longer lives of nuclear isomers' metastable states are often due to the larger degree of nuclear spin change which must be involved in their gamma emission to reach
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#173285542465611346-456: The principal quantum number n , the azimuthal quantum number ℓ (the orbital type), the orbital magnetic quantum number m ℓ , and the spin magnetic quantum number m s . The sequence in which the subshells are filled is given in most cases by the Aufbau principle , also known as the Madelung or Klechkovsky rule (after Erwin Madelung and Vsevolod Klechkovsky respectively). This rule
11529-436: The residual strong force . Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their copresence pushes protons slightly apart, reducing the electrostatic repulsion between the protons, and they exert an attractive nuclear force on each other and on protons. For this reason, one or more neutrons are necessary for two or more protons to bind into
11712-421: The s-process and r-process of neutron capture, during nucleosynthesis in stars . For this reason, only 78 Pt and 4 Be are the most naturally abundant isotopes of their element. 48 stable odd-proton-even-neutron nuclides, stabilized by their paired neutrons, form most of the stable isotopes of the odd-numbered elements; the very few odd-proton-odd-neutron nuclides comprise
11895-442: The visible range. The amount of energy released is related to bond-dissociation energy or ionization energy and is usually in the range of a few to few tens of eV per bond. However, a much stronger type of binding energy , the nuclear binding energy , is involved in nuclear processes. Due to this, most nuclear excited states decay by gamma ray emission. For example, a well-known nuclear isomer used in various medical procedures
12078-471: The 1s and 2s orbitals, which have quite different angular charge distributions, and hence are not very large; but the 3p orbitals experience strong repulsion from the 2p orbitals, which have similar angular charge distributions. Thus higher s-, p-, d-, and f-subshells experience strong repulsion from their inner analogues, which have approximately the same angular distribution of charge, and must expand to avoid this. This makes significant differences arise between
12261-401: The 5f orbitals are in the core, and probably the 6d orbitals join the core starting from nihonium. Again there are a few anomalies along the way: for example, as single atoms neither actinium nor thorium actually fills the 5f subshell, and lawrencium does not fill the 6d shell, but all these subshells can still become filled in chemical environments. For a very long time, the seventh row
12444-742: The 5s orbitals ( rubidium and strontium ), then 4d ( yttrium through cadmium , again with a few anomalies along the way), and then 5p ( indium through xenon ). Again, from indium onward the 4d orbitals are in the core. Hence the fifth row has the same structure as the fourth. The sixth row of the table likewise starts with two s-block elements: caesium and barium . After this, the first f-block elements (coloured green below) begin to appear, starting with lanthanum . These are sometimes termed inner transition elements. As there are now not only 4f but also 5d and 6s subshells at similar energies, competition occurs once again with many irregular configurations; this resulted in some dispute about where exactly
12627-476: The 83 primordial elements that survived from the Earth's formation. The remaining eleven natural elements decay quickly enough that their continued trace occurrence rests primarily on being constantly regenerated as intermediate products of the decay of thorium and uranium. All 24 known artificial elements are radioactive. Under an international naming convention, the groups are numbered numerically from 1 to 18 from
12810-400: The 94 natural elements, eighty have a stable isotope and one more ( bismuth ) has an almost-stable isotope (with a half-life of 2.01×10 years, over a billion times the age of the universe ). Two more, thorium and uranium , have isotopes undergoing radioactive decay with a half-life comparable to the age of the Earth . The stable elements plus bismuth, thorium, and uranium make up
12993-578: The AZE notation is different from how it is written: 2 He is commonly pronounced as helium-four instead of four-two-helium, and 92 U as uranium two-thirty-five (American English) or uranium-two-three-five (British) instead of 235-92-uranium. Some isotopes/nuclides are radioactive , and are therefore referred to as radioisotopes or radionuclides , whereas others have never been observed to decay radioactively and are referred to as stable isotopes or stable nuclides . For example, C
13176-474: The accepted value, the Bohr radius (~0.529 Å). In his model, Haas used a single-electron configuration based on the classical atomic model proposed by J. J. Thomson in 1904, often called the plum-pudding model . Atomic radii (the size of atoms) are dependent on the sizes of their outermost orbitals. They generally decrease going left to right along the main-group elements, because the nuclear charge increases but
13359-403: The alkali metals which are reactive solid metals. This and hydrogen's formation of hydrides , in which it gains an electron, brings it close to the properties of the halogens which do the same (though it is rarer for hydrogen to form H than H ). Moreover, the lightest two halogens ( fluorine and chlorine ) are gaseous like hydrogen at standard conditions. Some properties of hydrogen are not
13542-424: The almost integral masses for the two isotopes Cl and Cl. After the discovery of the neutron by James Chadwick in 1932, the ultimate root cause for the existence of isotopes was clarified, that is, the nuclei of different isotopes for a given element have different numbers of neutrons, albeit having the same number of protons. A neutral atom has the same number of electrons as protons. Thus different isotopes of
13725-452: The atom is still determined by the outer electrons. The increasing nuclear charge across the series and the increased number of inner electrons for shielding somewhat compensate each other, so the decrease in radius is smaller. The 4p and 5d atoms, coming immediately after new types of transition series are first introduced, are smaller than would have been expected, because the added core 3d and 4f subshells provide only incomplete shielding of
13908-523: The atom. Their energies are quantised , which is to say that they can only take discrete values. Furthermore, electrons obey the Pauli exclusion principle : different electrons must always be in different states. This allows classification of the possible states an electron can take in various energy levels known as shells, divided into individual subshells, which each contain one or more orbitals. Each orbital can contain up to two electrons: they are distinguished by
14091-413: The atom. These excited electrons then leave at a high speed. This occurs because inner atomic electrons penetrate the nucleus where they are subject to the intense electric fields created when the protons of the nucleus re-arrange in a different way. In nuclei that are far from stability in energy, even more decay modes are known. After fission, several of the fission fragments that may be produced have
14274-401: The beginning of a new shell. Thus, with the exception of the first row, each period length appears twice: The overlaps get quite close at the point where the d-orbitals enter the picture, and the order can shift slightly with atomic number and atomic charge. Starting from the simplest atom, this lets us build up the periodic table one at a time in order of atomic number, by considering
14457-624: The beta decay of actinium-230 forms thorium-230. The term "isotope", Greek for "at the same place", was suggested to Soddy by Margaret Todd , a Scottish physician and family friend, during a conversation in which he explained his ideas to her. He received the 1921 Nobel Prize in Chemistry in part for his work on isotopes. In 1914 T. W. Richards found variations between the atomic weight of lead from different mineral sources, attributable to variations in isotopic composition due to different radioactive origins. The first evidence for multiple isotopes of
14640-411: The cases of single atoms. In hydrogen , there is only one electron, which must go in the lowest-energy orbital 1s. This electron configuration is written 1s , where the superscript indicates the number of electrons in the subshell. Helium adds a second electron, which also goes into 1s, completely filling the first shell and giving the configuration 1s . Starting from the third element, lithium ,
14823-470: The characterization "nuclear isomer" is usually applied only to configurations with half-lives of 10 seconds or longer. Quantum mechanics predicts that certain atomic species should possess isomers with unusually long lifetimes even by this stricter standard and have interesting properties. Some nuclear isomers are so long-lived that they are relatively stable and can be produced and observed in quantity. The most stable nuclear isomer occurring in nature
15006-411: The chemical elements are a periodic function of their atomic number . Elements are placed in the periodic table according to their electron configurations , the periodic recurrences of which explain the trends in properties across the periodic table. An electron can be thought of as inhabiting an atomic orbital , which characterizes the probability it can be found in any particular region around
15189-411: The chemist and philosopher of science Eric Scerri on the grounds that it appears to imply that hydrogen is above the periodic law altogether, unlike all the other elements. Helium is the only element that routinely occupies a position in the periodic table that is not consistent with its electronic structure. It has two electrons in its outermost shell, whereas the other noble gases have eight; and it
15372-415: The composition of group 3 , the options can be shown equally (unprejudiced) in both forms. Periodic tables usually at least show the elements' symbols; many also provide supplementary information about the elements, either via colour-coding or as data in the cells. The above table shows the names and atomic numbers of the elements, and also their blocks, natural occurrences and standard atomic weights . For
15555-443: The core, and cannot be used for chemical reactions. Thus the relationship between yttrium and lanthanum is only a secondary relationship between elements with the same number of valence electrons but different kinds of valence orbitals, such as that between chromium and uranium; whereas the relationship between yttrium and lutetium is primary, sharing both valence electron count and valence orbital type. As chemical reactions involve
15738-549: The de-excitation of Ta by resonant photo-excitation of intermediate high levels of this nucleus ( E ≈ 1 MeV) was observed in 1999 by Belic and co-workers in the Stuttgart nuclear physics group. 72 Hf is another reasonably stable nuclear isomer. It possesses a half-life of 31 years and the highest excitation energy of any comparably long-lived isomer. One gram of pure Hf contains approximately 1.33 gigajoules of energy,
15921-424: The distribution of protons and neutrons is so much further from spherical geometry that de-excitation to the nuclear ground state is strongly hindered. In general, these states either de-excite to the ground state far more slowly than a "usual" excited state, or they undergo spontaneous fission with half-lives of the order of nanoseconds or microseconds —a very short time, but many orders of magnitude longer than
16104-422: The electrons, and so the occupation is not quite consistently filling the 3d orbitals one at a time. The precise energy ordering of 3d and 4s changes along the row, and also changes depending on how many electrons are removed from the atom. For example, due to the repulsion between the 3d electrons and the 4s ones, at chromium the 4s energy level becomes slightly higher than 3d, and so it becomes more profitable for
16287-494: The elemental abundance found on Earth and in the Solar System. However, in the cases of three elements ( tellurium , indium , and rhenium ) the most abundant isotope found in nature is actually one (or two) extremely long-lived radioisotope(s) of the element, despite these elements having one or more stable isotopes. Theory predicts that many apparently "stable" nuclides are radioactive, with extremely long half-lives (discounting
16470-429: The elements are listed in order of increasing atomic number. A new row ( period ) is started when a new electron shell has its first electron . Columns ( groups ) are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. oxygen , sulfur , and selenium are in the same column because they all have four electrons in
16653-465: The elements thus exhibit periodic recurrences, hence the name of the periodic table and the periodic law. These periodic recurrences were noticed well before the underlying theory that explains them was developed. Historically, the physical size of atoms was unknown until the early 20th century. The first calculated estimate of the atomic radius of hydrogen was published by physicist Arthur Haas in 1910 to within an order of magnitude (a factor of 10) of
16836-434: The emitted gamma ray must carry inhibits decay rate by about 5 orders of magnitude. The highest known spin change of 8 units occurs in the decay of Ta, which suppresses its decay by a factor of 10 from that associated with 1 unit. Instead of a natural gamma-decay half-life of 10 seconds, it has a half-life of more than 10 seconds, or at least 3 × 10 years, and thus has yet to be observed to decay. Gamma emission
17019-410: The energy is released very quickly, so that Hf can produce extremely high powers (on the order of exawatts ). Other isomers have also been investigated as possible media for gamma-ray stimulated emission . Holmium 's nuclear isomer 67 Ho has a half-life of 1,200 years, which is nearly the longest half-life of any holmium radionuclide. Only Ho , with
17202-434: The equivalent of exploding about 315 kg (700 lb) of TNT . In the natural decay of Hf , the energy is released as gamma rays with a total energy of 2.45 MeV. As with Ta , there are disputed reports that Hf can be stimulated into releasing its energy. Due to this, the substance is being studied as a possible source for gamma-ray lasers . These reports indicate that
17385-410: The even-even isotopes, which are about 3 times as numerous. Among the 41 even- Z elements that have a stable nuclide, only two elements (argon and cerium) have no even-odd stable nuclides. One element (tin) has three. There are 24 elements that have one even-odd nuclide and 13 that have two odd-even nuclides. Of 35 primordial radionuclides there exist four even-odd nuclides (see table at right), including
17568-410: The extent to which chemical or electronic properties should decide periodic table placement. Like the group 1 metals, hydrogen has one electron in its outermost shell and typically loses its only electron in chemical reactions. Hydrogen has some metal-like chemical properties, being able to displace some metals from their salts . But it forms a diatomic nonmetallic gas at standard conditions, unlike
17751-476: The f-block is supposed to begin, but most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle. Even though lanthanum does not itself fill the 4f subshell as a single atom, because of repulsion between electrons, its 4f orbitals are low enough in energy to participate in chemistry. At ytterbium , the seven 4f orbitals are completely filled with fourteen electrons; thereafter,
17934-488: The f-block. Thus the d-block is split into two very uneven portions. This is a holdover from early mistaken measurements of electron configurations; modern measurements are more consistent with the form with lutetium and lawrencium in group 3, and with La–Yb and Ac–No as the f-block. The 4f shell is completely filled at ytterbium, and for that reason Lev Landau and Evgeny Lifshitz in 1948 considered it incorrect to group lutetium as an f-block element. They did not yet take
18117-545: The f-shells complete filling at ytterbium and nobelium, matching the Sc-Y-Lu-Lr form, and not at lutetium and lawrencium as the Sc-Y-La-Ac form would have it. Not only are such exceptional configurations in the minority, but they have also in any case never been considered as relevant for positioning any other elements on the periodic table: in gaseous atoms, the d-shells complete their filling at copper, palladium, and gold, but it
18300-459: The filling of the third shell by occupying a 3s orbital, giving a configuration of 1s 2s 2p 3s for sodium. This configuration is abbreviated [Ne] 3s , where [Ne] represents neon's configuration. Magnesium ([Ne] 3s ) finishes this 3s orbital, and the following six elements aluminium , silicon , phosphorus , sulfur , chlorine , and argon fill the three 3p orbitals ([Ne] 3s 3p through [Ne] 3s 3p ). This creates an analogous series in which
18483-409: The first 94 of which are known to occur naturally on Earth at present. The remaining 24, americium to oganesson (95–118), occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are primordial and 11 occur only in decay chains of primordial elements. A few of the latter are so rare that they were not discovered in nature, but were synthesized in the laboratory before it
18666-451: The first seven shells occupied. The first shell contains only one orbital, a spherical s orbital. As it is in the first shell, this is called the 1s orbital. This can hold up to two electrons. The second shell similarly contains a 2s orbital, and it also contains three dumbbell-shaped 2p orbitals, and can thus fill up to eight electrons (2×1 + 2×3 = 8). The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals, and thus has
18849-406: The first shell is full, so its third electron occupies a 2s orbital, giving a 1s 2s configuration. The 2s electron is lithium's only valence electron, as the 1s subshell is now too tightly bound to the nucleus to participate in chemical bonding to other atoms: such a shell is called a " core shell ". The 1s subshell is a core shell for all elements from lithium onward. The 2s subshell is completed by
19032-441: The first slot of the d-block which is left vacant to indicate the place of the f-block inserts", which would imply that this form still has lutetium and lawrencium (the 15th entries in question) as d-block elements in group 3. Indeed, when IUPAC publications expand the table to 32 columns, they make this clear and place lutetium and lawrencium under yttrium in group 3. Several arguments in favour of Sc-Y-La-Ac can be encountered in
19215-524: The first) differ in the opposite direction. Thus for example many properties in the p-block show a zigzag rather than a smooth trend along the group. For example, phosphorus and antimony in odd periods of group 15 readily reach the +5 oxidation state, whereas nitrogen, arsenic, and bismuth in even periods prefer to stay at +3. A similar situation holds for the d-block, with lutetium through tungsten atoms being slightly smaller than yttrium through molybdenum atoms respectively. Thallium and lead atoms are about
19398-486: The galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the Solar System , isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites . The atomic mass ( m r ) of an isotope (nuclide) is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to
19581-401: The ground state of tantalum-180) with comparatively short half-lives are known. Usually, they beta-decay to their nearby even-even isobars that have paired protons and paired neutrons. Of the nine primordial odd-odd nuclides (five stable and four radioactive with long half-lives), only 7 N is the most common isotope of a common element. This is the case because it is a part of
19764-421: The ground state. This high spin change causes these decays to be forbidden transitions and delayed. Delays in emission are caused by low or high available decay energy. The first nuclear isomer and decay-daughter system (uranium X 2 /uranium Z, now known as 91 Pa / 91 Pa ) was discovered by Otto Hahn in 1921. The nucleus of a nuclear isomer occupies a higher energy state than
19947-474: The group was in the d-block . The Roman numerals used correspond to the last digit of today's naming convention (e.g. the group 4 elements were group IVB, and the group 14 elements were group IVA). In Europe , the lettering was similar, except that "A" was used for groups 1 through 7, and "B" was used for groups 11 through 17. In addition, groups 8, 9 and 10 used to be treated as one triple-sized group, known collectively in both notations as group VIII. In 1988,
20130-405: The half-life of a more usual nuclear excited state. Fission isomers may be denoted with a postscript or superscript "f" rather than "m", so that a fission isomer, e.g. of plutonium -240, can be denoted as plutonium-240f or 94 Pu . Most nuclear excited states are very unstable and "immediately" radiate away the extra energy after existing on the order of 10 seconds. As a result,
20313-409: The integers 20 and 22 and that neither is equal to the known molar mass (20.2) of neon gas. This is an example of Aston's whole number rule for isotopic masses, which states that large deviations of elemental molar masses from integers are primarily due to the fact that the element is a mixture of isotopes. Aston similarly showed in 1920 that the molar mass of chlorine (35.45) is a weighted average of
20496-430: The isomeric states (e.g., hafnium-178m2, or 72 Hf ). A different kind of metastable nuclear state (isomer) is the fission isomer or shape isomer . Most actinide nuclei in their ground states are not spherical, but rather prolate spheroidal , with an axis of symmetry longer than the other axes, similar to an American football or rugby ball . This geometry can result in quantum-mechanical states where
20679-460: The laboratory. By 2010, the first 118 elements were known, thereby completing the first seven rows of the table; however, chemical characterization is still needed for the heaviest elements to confirm that their properties match their positions. New discoveries will extend the table beyond these seven rows , though it is not yet known how many more elements are possible; moreover, theoretical calculations suggest that this unknown region will not follow
20862-463: The largest number of stable isotopes for an element being ten, for tin ( 50 Sn ). There are about 94 elements found naturally on Earth (up to plutonium inclusive), though some are detected only in very tiny amounts, such as plutonium-244 . Scientists estimate that the elements that occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes ( nuclides ) in total. Only 251 of these naturally occurring nuclides are stable, in
21045-409: The last elements in this seventh row were given names in 2016. This completes the modern periodic table, with all seven rows completely filled to capacity. The following table shows the electron configuration of a neutral gas-phase atom of each element. Different configurations can be favoured in different chemical environments. The main-group elements have entirely regular electron configurations;
21228-425: The late seventh period, potentially leading to a collapse of periodicity. Electron configurations are only clearly known until element 108 ( hassium ), and experimental chemistry beyond 108 has only been done for elements 112 ( copernicium ) through 115 ( moscovium ), so the chemical characterization of the heaviest elements remains a topic of current research. The trend that atomic radii decrease from left to right
21411-482: The least common. The 146 even-proton, even-neutron (EE) nuclides comprise ~58% of all stable nuclides and all have spin 0 because of pairing. There are also 24 primordial long-lived even-even nuclides. As a result, each of the 41 even-numbered elements from 2 to 82 has at least one stable isotope , and most of these elements have several primordial isotopes. Half of these even-numbered elements have six or more stable isotopes. The extreme stability of helium-4 due to
21594-505: The leftmost column (the alkali metals) to the rightmost column (the noble gases). The f-block groups are ignored in this numbering. Groups can also be named by their first element, e.g. the "scandium group" for group 3. Previously, groups were known by Roman numerals . In the United States , the Roman numerals were followed by either an "A" if the group was in the s- or p-block , or a "B" if
21777-455: The lightest elements, whose ratio of neutron number to atomic number varies the most between isotopes, it usually has only a small effect although it matters in some circumstances (for hydrogen, the lightest element, the isotope effect is large enough to affect biology strongly). The term isotopes (originally also isotopic elements , now sometimes isotopic nuclides ) is intended to imply comparison (like synonyms or isomers ). For example,
21960-399: The literature, but they have been challenged as being logically inconsistent. For example, it has been argued that lanthanum and actinium cannot be f-block elements because as individual gas-phase atoms, they have not begun to fill the f-subshells. But the same is true of thorium which is never disputed as an f-block element, and this argument overlooks the problem on the other end: that
22143-400: The longest-lived isotope), and thorium X (Ra) are impossible to separate. Attempts to place the radioelements in the periodic table led Soddy and Kazimierz Fajans independently to propose their radioactive displacement law in 1913, to the effect that alpha decay produced an element two places to the left in the periodic table, whereas beta decay emission produced an element one place to
22326-414: The mass number (number of nucleons) with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left (e.g. 2 He , 2 He , 6 C , 6 C , 92 U , and 92 U ). Because the atomic number is given by the element symbol, it is common to state only the mass number in
22509-630: The metastable half life from the normal "prompt" gamma-emission half-life. Occasionally the half-lives are far longer than this and can last minutes, hours, or years. For example, the 73 Ta nuclear isomer survives so long (at least 10 years) that it has never been observed to decay spontaneously. The half-life of a nuclear isomer can even exceed that of the ground state of the same nuclide, as shown by 73 Ta as well as 75 Re , 77 Ir , 83 Bi , 84 Po , 95 Am and multiple holmium isomers . Sometimes,
22692-521: The metastable state. Metastable isomers of a particular isotope are usually designated with an "m". This designation is placed after the mass number of the atom; for example, cobalt-58m1 is abbreviated 27 Co , where 27 is the atomic number of cobalt. For isotopes with more than one metastable isomer, "indices" are placed after the designation, and the labeling becomes m1, m2, m3, and so on. Increasing indices, m1, m2, etc., correlate with increasing levels of excitation energy stored in each of
22875-427: The most common amount of 1 quantum unit ħ in the spin angular momentum. This change is necessary to emit a gamma photon, which has a spin of 1 unit in this system. Integral changes of 2 and more units in angular momentum are possible, but the emitted photons carry off the additional angular momentum. Changes of more than 1 unit are known as forbidden transitions . Each additional unit of spin change larger than 1 that
23058-424: The new IUPAC (International Union of Pure and Applied Chemistry) naming system (1–18) was put into use, and the old group names (I–VIII) were deprecated. 32 columns 18 columns For reasons of space, the periodic table is commonly presented with the f-block elements cut out and positioned as a distinct part below the main body. This reduces the number of element columns from 32 to 18. Both forms represent
23241-402: The next element beryllium (1s 2s ). The following elements then proceed to fill the 2p subshell. Boron (1s 2s 2p ) puts its new electron in a 2p orbital; carbon (1s 2s 2p ) fills a second 2p orbital; and with nitrogen (1s 2s 2p ) all three 2p orbitals become singly occupied. This is consistent with Hund's rule , which states that atoms usually prefer to singly occupy each orbital of
23424-401: The next row, for potassium and calcium the 4s subshell is the lowest in energy, and therefore they fill it. Potassium adds one electron to the 4s shell ([Ar] 4s ), and calcium then completes it ([Ar] 4s ). However, starting from scandium ([Ar] 3d 4s ) the 3d subshell becomes the next highest in energy. The 4s and 3d subshells have approximately the same energy and they compete for filling
23607-419: The noble gases' boiling points and solubilities in water, where helium is too close to neon, and the large difference characteristic between the first two elements of a group appears only between neon and argon. Moving helium to group 2 makes this trend consistent in groups 2 and 18 as well, by making helium the first group 2 element and neon the first group 18 element: both exhibit the characteristic properties of
23790-429: The non-excited nucleus existing in the ground state . In an excited state, one or more of the protons or neutrons in a nucleus occupy a nuclear orbital of higher energy than an available nuclear orbital. These states are analogous to excited states of electrons in atoms. When excited atomic states decay, energy is released by fluorescence . In electronic transitions, this process usually involves emission of light near
23973-407: The nuclear charge for the outer electrons. Hence for example gallium atoms are slightly smaller than aluminium atoms. Together with kainosymmetry, this results in an even-odd difference between the periods (except in the s-block) that is sometimes known as secondary periodicity: elements in even periods have smaller atomic radii and prefer to lose fewer electrons, while elements in odd periods (except
24156-450: The nucleus are held more tightly and are more difficult to remove. Ionisation energy thus is minimized at the first element of each period – hydrogen and the alkali metals – and then generally rises until it reaches the noble gas at the right edge of the period. There are some exceptions to this trend, such as oxygen, where the electron being removed is paired and thus interelectronic repulsion makes it easier to remove than expected. In
24339-424: The nucleus, for example, carbon-13 with 6 protons and 7 neutrons. The nuclide concept (referring to individual nuclear species) emphasizes nuclear properties over chemical properties, whereas the isotope concept (grouping all atoms of each element) emphasizes chemical over nuclear. The neutron number greatly affects nuclear properties, but its effect on chemical properties is negligible for most elements. Even for
24522-519: The nuclides 6 C , 6 C , 6 C are isotopes (nuclides with the same atomic number but different mass numbers), but 18 Ar , 19 K , 20 Ca are isobars (nuclides with the same mass number). However, isotope is the older term and so is better known than nuclide and is still sometimes used in contexts in which nuclide might be more appropriate, such as nuclear technology and nuclear medicine . An isotope and/or nuclide
24705-403: The number of protons in its nucleus . Each distinct atomic number therefore corresponds to a class of atom: these classes are called the chemical elements . The chemical elements are what the periodic table classifies and organizes. Hydrogen is the element with atomic number 1; helium , atomic number 2; lithium , atomic number 3; and so on. Each of these names can be further abbreviated by
24888-416: The other about 22. The parabola due to the heavier gas is always much fainter than that due to the lighter, so that probably the heavier gas forms only a small percentage of the mixture." F. W. Aston subsequently discovered multiple stable isotopes for numerous elements using a mass spectrograph . In 1919 Aston studied neon with sufficient resolution to show that the two isotopic masses are very close to
25071-689: The other naturally occurring nuclides are radioactive but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. These include the afore-mentioned cosmogenic nuclides , the nucleogenic nuclides, and any radiogenic nuclides formed by ongoing decay of a primordial radioactive nuclide, such as radon and radium from uranium. An additional ~3000 radioactive nuclides not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived nuclides not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae . An example
25254-726: The others. There are 41 odd-numbered elements with Z = 1 through 81, of which 39 have stable isotopes ( technetium ( 43 Tc ) and promethium ( 61 Pm ) have no stable isotopes). Of these 39 odd Z elements, 30 elements (including hydrogen-1 where 0 neutrons is even ) have one stable odd-even isotope, and nine elements: chlorine ( 17 Cl ), potassium ( 19 K ), copper ( 29 Cu ), gallium ( 31 Ga ), bromine ( 35 Br ), silver ( 47 Ag ), antimony ( 51 Sb ), iridium ( 77 Ir ), and thallium ( 81 Tl ), have two odd-even stable isotopes each. This makes
25437-450: The outer electrons are still in the same shell. However, going down a column, the radii generally increase, because the outermost electrons are in higher shells that are thus further away from the nucleus. The first row of each block is abnormally small, due to an effect called kainosymmetry or primogenic repulsion: the 1s, 2p, 3d, and 4f subshells have no inner analogues. For example, the 2p orbitals do not experience strong repulsion from
25620-409: The outer shell structures of sodium through argon are analogous to those of lithium through neon, and is the basis for the periodicity of chemical properties that the periodic table illustrates: at regular but changing intervals of atomic numbers, the properties of the chemical elements approximately repeat. The first eighteen elements can thus be arranged as the start of a periodic table. Elements in
25803-427: The outermost p-subshell). Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it. Today, 118 elements are known,
25986-425: The patterns of the known part of the table. Some scientific discussion also continues regarding whether some elements are correctly positioned in today's table. Many alternative representations of the periodic law exist, and there is some discussion as to whether there is an optimal form of the periodic table. 1 Each chemical element has a unique atomic number ( Z — for "Zahl", German for "number") representing
26169-432: The period 1 elements hydrogen and helium remains an open issue under discussion, and some variation can be found. Following their respective s and s electron configurations, hydrogen would be placed in group 1, and helium would be placed in group 2. The group 1 placement of hydrogen is common, but helium is almost always placed in group 18 with the other noble gases. The debate has to do with conflicting understandings of
26352-400: The periodic law to predict some properties of some of the missing elements . The periodic law was recognized as a fundamental discovery in the late 19th century. It was explained early in the 20th century, with the discovery of atomic numbers and associated pioneering work in quantum mechanics , both ideas serving to illuminate the internal structure of the atom. A recognisably modern form of
26535-572: The periodic table. Metallic character increases going down a group and from right to left across a period. Nonmetallic character increases going from the bottom left of the periodic table to the top right. The first periodic table to become generally accepted was that of the Russian chemist Dmitri Mendeleev in 1869; he formulated the periodic law as a dependence of chemical properties on atomic mass . As not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used
26718-407: The periodic table. Spin–orbit interaction splits the p-subshell: one p-orbital is relativistically stabilized and shrunken (it fills in thallium and lead), but the other two (filling in bismuth through radon) are relativistically destabilized and expanded. Relativistic effects also explain why gold is golden and mercury is a liquid at room temperature. They are expected to become very strong in
26901-415: The possibility of proton decay , which would make all nuclides ultimately unstable). Some stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay products have yet been observed, and so these isotopes are said to be "observationally stable". The predicted half-lives for these nuclides often greatly exceed the estimated age of
27084-414: The presence of a nearby atom can shift the balance. Therefore, the periodic table ignores them and considers only idealized configurations. At zinc ([Ar] 3d 4s ), the 3d orbitals are completely filled with a total of ten electrons. Next come the 4p orbitals, completing the row, which are filled progressively by gallium ([Ar] 3d 4s 4p ) through krypton ([Ar] 3d 4s 4p ), in a manner analogous to
27267-429: The previous p-block elements. From gallium onwards, the 3d orbitals form part of the electronic core, and no longer participate in chemistry. The s- and p-block elements, which fill their outer shells, are called main-group elements ; the d-block elements (coloured blue below), which fill an inner shell, are called transition elements (or transition metals, since they are all metals). The next eighteen elements fill
27450-432: The primary exceptions). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms; so different isotopologues have different sets of vibrational modes. Because vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range. Atomic nuclei consist of protons and neutrons bound together by
27633-457: The properties of the various isotopes of a given element. Isotope separation is a significant technological challenge, particularly with heavy elements such as uranium or plutonium. Lighter elements such as lithium, carbon, nitrogen, and oxygen are commonly separated by gas diffusion of their compounds such as CO and NO. The separation of hydrogen and deuterium is unusual because it is based on chemical rather than physical properties, for example in
27816-459: The pull of the nucleus more strongly, and especially if there is an available partially filled outer orbital that can accommodate it. Therefore, electron affinity tends to increase down to up and left to right. The exception is the last column, the noble gases, which have a full shell and have no room for another electron. This gives the halogens in the next-to-last column the highest electron affinities. Nuclear isomer A nuclear isomer
27999-586: The relative mass difference between isotopes is much less so that the mass-difference effects on chemistry are usually negligible. (Heavy elements also have relatively more neutrons than lighter elements, so the ratio of the nuclear mass to the collective electronic mass is slightly greater.) There is also an equilibrium isotope effect . Similarly, two molecules that differ only in the isotopes of their atoms ( isotopologues ) have identical electronic structures, and therefore almost indistinguishable physical and chemical properties (again with deuterium and tritium being
28182-446: The right. Soddy recognized that emission of an alpha particle followed by two beta particles led to the formation of an element chemically identical to the initial element but with a mass four units lighter and with different radioactive properties. Soddy proposed that several types of atoms (differing in radioactive properties) could occupy the same place in the table. For example, the alpha-decay of uranium-235 forms thorium-231, whereas
28365-433: The same column have the same number of valence electrons and have analogous valence electron configurations: these columns are called groups. The single exception is helium, which has two valence electrons like beryllium and magnesium, but is typically placed in the column of neon and argon to emphasise that its outer shell is full. (Some contemporary authors question even this single exception, preferring to consistently follow
28548-486: The same element get smaller as more electrons are removed, because the attraction from the nucleus begins to outweigh the repulsion between electrons that causes electron clouds to expand: thus for example ionic radii decrease in the series V , V , V , V . The first ionisation energy of an atom is the energy required to remove an electron from it. This varies with the atomic radius: ionisation energy increases left to right and down to up, because electrons that are closer to
28731-440: The same element. This is most pronounced by far for protium ( H ), deuterium ( H ), and tritium ( H ), because deuterium has twice the mass of protium and tritium has three times the mass of protium. These mass differences also affect the behavior of their respective chemical bonds, by changing the center of gravity ( reduced mass ) of the atomic systems. However, for heavier elements,
28914-473: The same periodic table. The form with the f-block included in the main body is sometimes called the 32-column or long form; the form with the f-block cut out the 18-column or medium-long form. The 32-column form has the advantage of showing all elements in their correct sequence, but it has the disadvantage of requiring more space. The form chosen is an editorial choice, and does not imply any change of scientific claim or statement. For example, when discussing
29097-422: The same size as indium and tin atoms respectively, but from bismuth to radon the 6p atoms are larger than the analogous 5p atoms. This happens because when atomic nuclei become highly charged, special relativity becomes needed to gauge the effect of the nucleus on the electron cloud. These relativistic effects result in heavy elements increasingly having differing properties compared to their lighter homologues in
29280-407: The same type before filling them with the second electron. Oxygen (1s 2s 2p ), fluorine (1s 2s 2p ), and neon (1s 2s 2p ) then complete the already singly filled 2p orbitals; the last of these fills the second shell completely. Starting from element 11, sodium , the second shell is full, making the second shell a core shell for this and all heavier elements. The eleventh electron begins
29463-417: The same type. This makes the first row of elements in each block unusually small, and such elements tend to exhibit characteristic kinds of anomalies for their group. Some chemists arguing for the repositioning of helium have pointed out that helium exhibits these anomalies if it is placed in group 2, but not if it is placed in group 18: on the other hand, neon, which would be the first group 18 element if helium
29646-414: The same value of n + ℓ, the one with lower n is occupied first. In general, orbitals with the same value of n + ℓ are similar in energy, but in the case of the s-orbitals (with ℓ = 0), quantum effects raise their energy to approach that of the next n + ℓ group. Hence the periodic table is usually drawn to begin each row (often called a period) with the filling of a new s-orbital, which corresponds to
29829-406: The sense of never having been observed to decay as of the present time. An additional 35 primordial nuclides (to a total of 286 primordial nuclides), are radioactive with known half-lives, but have half-lives longer than 100 million years, allowing them to exist from the beginning of the Solar System. See list of nuclides for details. All the known stable nuclides occur naturally on Earth;
30012-412: The short-lived elements without standard atomic weights, the mass number of the most stable known isotope is used instead. Other tables may include properties such as state of matter, melting and boiling points, densities, as well as provide different classifications of the elements. The periodic table is a graphic description of the periodic law, which states that the properties and atomic structures of
30195-410: The small 2p elements, which prefer multiple bonding , and the larger 3p and higher p-elements, which do not. Similar anomalies arise for the 1s, 2p, 3d, 4f, and the hypothetical 5g elements: the degree of this first-row anomaly is highest for the s-block, is moderate for the p-block, and is less pronounced for the d- and f-blocks. In the transition elements, an inner shell is filling, but the size of
30378-418: The so-called isomeric yield ratio . Metastable isomers can be produced through nuclear fusion or other nuclear reactions . A nucleus produced this way generally starts its existence in an excited state that relaxes through the emission of one or more gamma rays or conversion electrons . Sometimes the de-excitation does not completely proceed rapidly to the nuclear ground state . This usually occurs as
30561-452: The step of removing lanthanum from the d-block as well, but Jun Kondō realized in 1963 that lanthanum's low-temperature superconductivity implied the activity of its 4f shell. In 1965, David C. Hamilton linked this observation to its position in the periodic table, and argued that the f-block should be composed of the elements La–Yb and Ac–No. Since then, physical, chemical, and electronic evidence has supported this assignment. The issue
30744-485: The superscript and leave out the atomic number subscript (e.g. He , He , C , C , U , and U ). The letter m (for metastable) is sometimes appended after the mass number to indicate a nuclear isomer , a metastable or energetically excited nuclear state (as opposed to the lowest-energy ground state ), for example 73 Ta ( tantalum-180m ). The common pronunciation of
30927-474: The table appearing on the IUPAC web site, but this creates an inconsistency with quantum mechanics by making the f-block 15 elements wide (La–Lu and Ac–Lr) even though only 14 electrons can fit in an f-subshell. There is moreover some confusion in the literature on which elements are then implied to be in group 3. While the 2021 IUPAC report noted that 15-element-wide f-blocks are supported by some practitioners of
31110-418: The table was reached in 1945 with Glenn T. Seaborg 's discovery that the actinides were in fact f-block rather than d-block elements. The periodic table and law are now a central and indispensable part of modern chemistry. The periodic table continues to evolve with the progress of science. In nature, only elements up to atomic number 94 exist; to go further, it was necessary to synthesize new elements in
31293-406: The transition and inner transition elements show twenty irregularities due to the aforementioned competition between subshells close in energy level. For the last ten elements (109–118), experimental data is lacking and therefore calculated configurations have been shown instead. Completely filled subshells have been greyed out. Although the modern periodic table is standard today, the placement of
31476-401: The transition series, the outer electrons are preferentially lost even though the inner orbitals are filling. For example, in the 3d series, the 4s electrons are lost first even though the 3d orbitals are being filled. The shielding effect of adding an extra 3d electron approximately compensates the rise in nuclear charge, and therefore the ionisation energies stay mostly constant, though there is
31659-419: The universe, and in fact, there are also 31 known radionuclides (see primordial nuclide ) with half-lives longer than the age of the universe. Adding in the radioactive nuclides that have been created artificially, there are 3,339 currently known nuclides . These include 905 nuclides that are either stable or have half-lives longer than 60 minutes. See list of nuclides for details. The existence of isotopes
31842-419: The valence configurations and place helium over beryllium.) There are eight columns in this periodic table fragment, corresponding to at most eight outer-shell electrons. A period begins when a new shell starts filling. Finally, the colouring illustrates the blocks : the elements in the s-block (coloured red) are filling s-orbitals, while those in the p-block (coloured yellow) are filling p-orbitals. Starting
32025-403: The valence electrons, elements with similar outer electron configurations may be expected to react similarly and form compounds with similar proportions of elements in them. Such elements are placed in the same group, and thus there tend to be clear similarities and trends in chemical behaviour as one proceeds down a group. As analogous configurations occur at regular intervals, the properties of
32208-489: Was Ta , which required more photon energy to trigger than was released. An isotope such as Lu releases gamma rays by decay through a series of internal energy levels within the nucleus, and it is thought that by learning the triggering cross sections with sufficient accuracy, it may be possible to create energy stores that are 10 times more concentrated than high explosive or other traditional chemical energy storage. An isomeric transition or internal transition (IT)
32391-439: Was brought to wide attention by William B. Jensen in 1982, and the reassignment of lutetium and lawrencium to group 3 was supported by IUPAC reports dating from 1988 (when the 1–18 group numbers were recommended) and 2021. The variation nonetheless still exists because most textbook writers are not aware of the issue. A third form can sometimes be encountered in which the spaces below yttrium in group 3 are left empty, such as
32574-459: Was determined that they do exist in nature after all: technetium (element 43), promethium (element 61), astatine (element 85), neptunium (element 93), and plutonium (element 94). No element heavier than einsteinium (element 99) has ever been observed in macroscopic quantities in its pure form, nor has astatine ; francium (element 87) has been only photographed in the form of light emitted from microscopic quantities (300,000 atoms). Of
32757-424: Was first observed empirically by Madelung, and Klechkovsky and later authors gave it theoretical justification. The shells overlap in energies, and the Madelung rule specifies the sequence of filling according to: Here the sign ≪ means "much less than" as opposed to < meaning just "less than". Phrased differently, electrons enter orbitals in order of increasing n + ℓ, and if two orbitals are available with
32940-507: Was first suggested in 1913 by the radiochemist Frederick Soddy , based on studies of radioactive decay chains that indicated about 40 different species referred to as radioelements (i.e. radioactive elements) between uranium and lead, although the periodic table only allowed for 11 elements between lead and uranium inclusive. Several attempts to separate these new radioelements chemically had failed. For example, Soddy had shown in 1910 that mesothorium (later shown to be Ra), radium (Ra,
33123-417: Was incomplete as most of its elements do not occur in nature. The missing elements beyond uranium started to be synthesized in the laboratory in 1940, when neptunium was made. (However, the first element to be discovered by synthesis rather than in nature was technetium in 1937.) The row was completed with the synthesis of tennessine in 2010 (the last element oganesson had already been made in 2002), and
33306-474: Was rejected by IUPAC in 1988 for these reasons. Nonetheless, helium is still occasionally placed in group 2 today, and some of its physical and chemical properties are closer to the group 2 elements and support the electronic placement. Solid helium crystallises in a hexagonal close-packed structure, which matches beryllium and magnesium in group 2, but not the other noble gases in group 18. Recent theoretical developments in noble gas chemistry, in which helium
33489-524: Was removed from that spot, does exhibit those anomalies. The relationship between helium and beryllium is then argued to resemble that between hydrogen and lithium, a placement which is much more commonly accepted. For example, because of this trend in the sizes of orbitals, a large difference in atomic radii between the first and second members of each main group is seen in groups 1 and 13–17: it exists between neon and argon, and between helium and beryllium, but not between helium and neon. This similarly affects
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