The periodic table, also known as the periodic table of the ( chemical) elements, is a rows and columns arrangement of the . It is widely used in chemistry, physics, and other sciences, and is generally seen as an Cultural icon of chemistry. It is a graphic formulation of the Periodic trends, which states that the properties of the chemical elements exhibit an approximate periodic dependence on their . The table is divided into four roughly rectangular areas called blocks. The rows of the table are called periods, and the columns are called groups. Elements from the same group of the periodic table show similar chemical characteristics. Periodic trends run through the periodic table, with character (keeping their own electrons) increasing from left to right across a period, and from down to up across a group, and character (surrendering electrons to other atoms) increasing in the opposite direction. The underlying reason for these trends is electron configurations of atoms. The periodic table exclusively lists electrically neutral atoms that have an equal number of positively charged protons and negatively charged electrons and puts (atoms with the same number of protons but different numbers of neutrons) at the same place. Other atoms, like and , are graphically collected in other tables like the tables of nuclides (often called Segrè charts).
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. Because not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used the periodic law to predict properties of some of the missing elements. The periodic law was recognized as a fundamental discovery in the late 19th century, and it was explained with the discovery of the atomic number and pioneering work in quantum mechanics of the early 20th century that illuminated the internal structure of the atom. With Glenn T. Seaborg's 1945 discovery that the were in fact f-block rather than d-block elements, a recognisably modern form of the table was reached. 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 synthesise new elements in the laboratory. Today, all the first 118 elements are known, completing the first seven rows of the table, but chemical characterisation is still needed for the heaviest elements to confirm that their properties match their positions. It is not yet known how far the table will stretch beyond these seven rows and whether the patterns of the known part of the table will continue into this unknown region. 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.
Atoms can be subdivided into different types based on the number of protons (and thus also electrons) they have. This is called the atomic number, often symbolised Z (for "Zahl" — German for "number"). Each distinct atomic number therefore corresponds to a class of atom: these classes are called the . The chemical elements are what the periodic table classifies and organises. 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 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 of the same chemical element. Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined , defined as the average mass of a naturally occurring atom of that element.
Today, 118 elements are known, the first 94 of which are known to occur naturally on Earth at present. Of the 94 natural elements, eighty have a stable isotope and one more (bismuth) has an almost-stable isotope (with a half-life 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 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.
An electron can be thought of as inhabiting an atomic orbital, which characterises the probability it can be found in any particular region of 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 a quantity known as spin, conventionally labeled "up" or "down".Petrucci et al., p. 323 In a cold atom (one in its ground state), electrons arrange themselves in such a way that the total energy they have is minimised by occupying the lowest-energy orbitals available.Petrucci et al., p. 322 Only the outermost electrons (so-called ) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called .
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 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 1s1, 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 1s2.
Starting from the third element, lithium, the first shell is full, so its third electron occupies a 2s orbital, giving a 1s2 2s1 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. Thus the filled first shell is called a "Core electron" for this and all heavier elements. The 2s subshell is completed by the next element beryllium (1s2 2s2). The following elements then proceed to fill the 2p subshell. Boron (1s2 2s2 2p1) puts its new electron in a 2p orbital; carbon (1s2 2s2 2p2) fills a second 2p orbital; and with nitrogen (1s2 2s2 2p3) 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 the same type before filling them with the second electron. Oxygen (1s2 2s2 2p4), fluorine (1s2 2s2 2p5), and neon (1s2 2s2 2p6) 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 electron for this and all heavier elements. The eleventh electron begins the filling of the third shell by occupying a 3s orbital, giving a configuration of 1s2 2s2 2p6 3s1 for sodium. This configuration is abbreviated Ne 3s1, where Ne represents neon's configuration. Magnesium (Ne 3s2) finishes this 3s orbital, and the following six elements aluminium, silicon, phosphorus, sulfur, chlorine, and argon fill the three 3p orbitals (Ne 3s2 3p1 through Ne 3s2 3p6). This creates an analogous series in which 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.Scerri, p. 17
The first eighteen elements can thus be arranged as the start of a periodic table. Elements in 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 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.
Starting 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 4s1), and calcium then completes it (Ar 4s2). However, starting from scandium (Ar 3d1 4s2) the 3d subshell becomes the next highest in energy. The 4s and 3d subshells have approximately the same energy and they compete for filling 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 to have a Ar 3d5 4s1 configuration than an Ar 3d4 4s2 one. A similar anomaly occurs at copper. These are violations of the Madelung rule. Such anomalies however do not have any chemical significance, as the various configurations are so close in energy to each otherPetrucci et al., p. 328 that the presence of a nearby atom can shift the balance. The periodic table therefore ignores these and considers only idealised configurations.
At zinc (Ar 3d10 4s2), 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 3d10 4s2 4p1) through krypton (Ar 3d10 4s2 4p6), in a manner analogous to 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).Petrucci et al., pp. 326–7
The next eighteen elements fill 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 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, 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 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:Petrucci et al., p. 331 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 was incomplete as most of its elements do not occur in nature. The missing elements beyond uranium started to be synthesised in the laboratory in 1940, when neptunium was made. The row was completed with the synthesis of tennessine in 2010 (the last element oganesson had already been made in 2002), and 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.
32 columns
18 columns
Both forms represent 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.Scerri, p. 375 The form chosen is an editorial choice, and does not imply any change of scientific claim or statement. For example, when discussing 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 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.
Like the group 1 metals, hydrogen has one electron in its outermost shellGray, p. 12 and typically loses its only electron in chemical reactions. It has some metal-like chemical properties, being able to displace some metals from their salts.
Helium is an unreactive noble gas 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. Therefore, helium is nearly universally placed in group 18 which its properties best match: a proposal to move helium to group 2 was rejected by IUPAC in 1988 for these reasons. However, helium only has two electrons in its outermost shell, whereas the other noble gases have eight; and it is an s-block element, whereas all other noble gases are p-block elements. Also, 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. In these ways helium better matches the alkaline earth metals
Some arguments in favour of Sc-Y-La-Ac can still be encountered in the literature, but many authors challenge them as being logically inconsistent. For example, it has been argued that lanthanum and actinium cannot be f-block elements because their atoms 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 the f-shells complete filling at ytterbium and nobelium (matching the Sc-Y-Lu-Lr form), not at lutetium and lawrencium (as in Sc-Y-La-Ac). Such exceptions to the Madelung rule have in any case never been considered as relevant for positioning any other elements on the periodic table. 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. 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.
Atomic radius (the size of atoms) are dependent on the sizes of their outermost orbitals.Siekierski and Burgess, pp. 23–26 They generally decrease going left to right along the main-group elements, because the nuclear charge increases but 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 to which they would be orthogonality. 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 the small 2p elements, which prefer , and the larger 3p and higher p-elements, which do not. Similar anomalies arise for the 1s, 2p, 3d, 4f, and the hypothetical 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 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,Greenwood and Earnshaw, p. 29 because the added core 3d and 4f subshells provide only incomplete shielding of 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 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.
Thallium and lead atoms are about 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 the periodic table. Spin–orbit interaction splits the p-subshell: one p-orbital is relativistically stabilised and shrunken (it fills in thallium and lead), but the other two (filling in bismuth through radon) are relativistically destabilised 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 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 112 (copernicium), 113 (nihonium), and 114 (flerovium), so the chemical characterisation of the heaviest elements remains a topic of current research.
In 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 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.
Some atoms, like the noble gases, have no electron affinity: they cannot form stable gas-phase anions. The noble gases, having high ionisation energies and no electron affinity, have little inclination towards gaining or losing electrons and are generally unreactive.
Some exceptions to the trends occur: oxygen and fluorine have lower electron affinities than their heavier homologues sulfur and chlorine, because they are small atoms and hence the newly added electron would experience significant repulsion from the already present ones. For the nonmetallic elements, electron affinity likewise somewhat correlates with reactivity, but not perfectly since other factors are involved. For example, fluorine has a lower electron affinity than chlorine, but is more reactive.
The electron configuration suggests a ready explanation from the number of electrons available for bonding, although a full explanation requires considering the energy that would be released in forming compounds with different valences rather than simply considering electron configurations alone.Greenwood and Earnshaw, p. 113 For example, magnesium forms Mg2+ rather than Mg+ cations when dissolved in water, because the latter would spontaneously disproportionate into Mg0 and Mg2+ cations. This is because the enthalpy of hydration (surrounding the cation with water molecules) increases in magnitude with the charge and radius of the ion. In Mg+, the outermost orbital (which determines ionic radius) is still 3s, so the hydration enthalpy is small and insufficient to compensate the energy required to remove the electron; but ionising again to Mg2+ uncovers the core 2p subshell, making the hydration enthalpy large enough to allow magnesium(II) compounds to form. For similar reasons, the common oxidation states of the heavier p-block elements (where the ns electrons become lower in energy than the np) tend to vary by steps of 2, because that is necessary to uncover an inner subshell and decrease the ionic radius (e.g. Tl+ uncovers 6s, and Tl3+ uncovers 5d, so once thallium loses two electrons it tends to lose the third one as well). Analogous arguments based on orbital hybridisation can be used for the less electronegative p-block elements.Siekierski and Burgess, pp. 45–54
For transition metals, common oxidation states are nearly always at least +2 for similar reasons (uncovering the next subshell); this holds even for the metals with anomalous dx+1s1 or dx+2s0 configurations (except for silver), because repulsion between d-electrons means that the movement of the second electron from the s- to the d-subshell does not appreciably change its ionisation energy.Siekierski and Burgess, pp. 134–137 Because ionising the transition metals further does not uncover any new inner subshells, their oxidation states tend to vary by steps of 1 instead. The lanthanides and late actinides generally show a stable +3 oxidation state, removing the outer s-electrons and then (usually) one electron from the (n−2)f-orbitals, that are similar in energy to ns. The common and maximum oxidation states of the d- and f-block elements tend to depend on the ionisation energies. As the energy difference between the (n−1)d and ns orbitals rises along each transition series, it becomes less energetically favourable to ionise further electrons. Thus, the early transition metal groups tend to prefer higher oxidation states, but the +2 oxidation state becomes more stable for the late transition metal groups. The highest formal oxidation state thus increases from +3 at the beginning of each d-block row, to +7 or +8 in the middle (e.g. osmium tetroxide), and then to +2 at the end. The lanthanides and late actinides usually have high fourth ionisation energies and hence rarely surpass the +3 oxidation states, whereas early actinides have low fourth ionisation energies and so for example neptunium and plutonium can reach +7.Siekierski and Burgess, pp. 178–180
As elements in the same group share the same valence configurations, they usually exhibit similar chemical behaviour. For example, the in the first group all have one valence electron, and form a very homogeneous class of elements: they are all soft and reactive metals. However, there are many factors involved, and groups can often be rather hetereogeneous. For instance, hydrogen also has one valence electron and is in the same group as the alkali metals, but its chemical behaviour is quite different. The stable elements of carbon group comprise a nonmetal (carbon), two semiconductors (silicon and germanium), and two metals (tin and lead); they are nonetheless united by having four valence electrons.Scerri, pp. 14–15 This often leads to similarities in maximum and minimum oxidation states (e.g. sulfur and selenium in chalcogen both have maximum oxidation state +6, as in sulfur trioxide and SeO3, and minimum oxidation state −2, as in and ); but not always (e.g. oxygen is not known to form oxidation state +6, despite being in the same group as sulfur and selenium).
Electronegativity depends on how strongly the nucleus can attract an electron pair, and so it exhibits a similar variation to the other properties already discussed: electronegativity tends to fall going up to down, and rise going left to right. The alkali and alkaline earth metals are among the most electropositive elements, while the chalcogens, halogens, and noble gases are among the most electronegative ones.
Electronegativity is generally measured on the Pauling scale, on which the most electronegative reactive atom (fluorine) is given electronegativity 4.0, and the least electronegative atom (caesium) is given electronegativity 0.79. In fact neon is the most electronegative element, but the Pauling scale cannot measure its electronegativity because it does not form covalent bonds with most elements.
An element's electronegativity varies with the identity and number of the atoms it is bonded to, as well as how many electrons it has already lost: an atom becomes more electronegative when it has lost more electrons.Greenwood and Earnshaw, pp. 25–6 This sometimes makes a large difference: lead in the +2 oxidation state has electronegativity 1.87 on the Pauling scale, while lead in the +4 oxidation state has electronegativity 2.33.
The more electropositive atoms, however, tend to instead lose electrons, creating a "sea" of electrons engulfing cations. The outer orbitals of one atom overlap to share electrons with all its neighbours, creating a giant structure of molecular orbitals extending over all the atoms. This negatively charged "sea" pulls on all the ions and keeps them together in a metallic bond. Elements forming such bonds are often called ; those which do not are often called . Some elements can form multiple simple substances with different structures: these are called . For example, diamond and graphite are two allotropes of carbon.
The metallicity of an element can be predicted from electronic properties. When atomic orbitals overlap during metallic or covalent bonding, they create both bonding and antibonding molecular orbitals of equal capacity, with the antibonding orbitals of higher energy. Net bonding character occurs when there are more electrons in the bonding orbitals than there are in the antibonding orbitals. Metallic bonding is thus possible when the number of electrons delocalised by each atom is less than twice the number of orbitals contributing to the overlap. This is the situation for elements in groups 1 through 13; they also have too few valence electrons to form giant covalent structures where all atoms take equivalent positions, and so almost all of them metallise. The exceptions are hydrogen and boron, which have too high an ionisation energy. Hydrogen thus forms a covalent H2 molecule, and boron forms a giant covalent structure based on icosahedral B12 clusters. In a metal, the bonding and antibonding orbitals have overlapping energies, creating a single band that electrons can freely flow through, allowing for electrical conduction.Siekierski and Burgess, pp. 60–66
In group 14, both metallic and covalent bonding become possible. In a diamond crystal, covalent bonds between carbon atoms are strong, because they have a small atomic radius and thus the nucleus has more of a hold on the electrons. Therefore, the bonding orbitals that result are much lower in energy than the antibonding orbitals, and there is no overlap, so electrical conduction becomes impossible: carbon is a nonmetal. However, covalent bonding becomes weaker for larger atoms and the energy gap between the bonding and antibonding orbitals decreases. Therefore, silicon and germanium have smaller and are : electrons can cross the gap when thermally excited. The band gap disappears in tin, so that tin and lead become metals.
Elements in groups 15 through 17 have too many electrons to form giant covalent molecules that stretch in all three dimensions. For the lighter elements, the bonds in small diatomic molecules are so strong that a condensed phase is disfavoured: thus nitrogen (N2), oxygen (O2), white phosphorus (P4), sulfur (S8), and the stable halogens (F2, Cl2, Br2, and I2) readily form covalent molecules with few atoms. The heavier ones tend to form long chains (e.g. red phosphorus, grey selenium, tellurium) or layered structures (e.g. carbon as graphite, black phosphorus, grey arsenic, grey antimony, bismuth) that only extend in one or two rather than three dimensions. As these structures do not use all their orbitals for bonding, they end up with bonding, nonbonding, and antibonding bands in order of increasing energy. Similarly to group 14, the band gaps shrink for the heavier elements and free movement of electrons between the chains or layers becomes possible. Thus for example black phosphorus, black arsenic, grey selenium, tellurium, and iodine are semiconductors; grey arsenic, grey antimony, and bismuth are (exhibiting quasi-metallic conduction, with a very small band overlap); and polonium and probably astatine are true metals. Finally, the natural group 18 elements all stay as individual atoms.
The dividing line between metals and nonmetals is roughly diagonal from top left to bottom right, with the transition series appearing to the left of this diagonal (as they have many available orbitals for overlap). This is expected, as metallicity tends to be correlated with electropositivity and the willingness to lose electrons, which increases right to left and up to down. Thus the metals greatly outnumber the nonmetals. Elements near the borderline are difficult to classify: they tend to have properties that are intermediate between those of metals and nonmetals, and may have some properties characteristic of both. They are often termed semimetals or . The term "semimetal" used in this sense should not be confused with its strict physical sense having to do with band structure: bismuth is physically a semimetal, but is generally considered a metal by chemists.
The following table considers the most stable allotropes at standard conditions. The elements coloured yellow form simple substances that are well-characterised by metallic bonding. Elements coloured light blue form giant covalent structures, whereas those coloured dark blue form small covalently bonded molecules that are held together by weaker van der Waals forces. The noble gases are coloured in violet: their molecules are single atoms and no covalent bonding occurs. Greyed-out cells are for elements which have not been prepared in sufficient quantities for their most stable allotropes to have been characterised in this way. Theoretical considerations and current experimental evidence suggest that all of those elements would metallise if they could form condensed phases, except perhaps for oganesson.
Generally, metals are shiny and dense. They usually have high melting and boiling points due to the strength of the metallic bond, and are often malleable and ductile (easily stretched and shaped) because the atoms can move relative to each other without breaking the metallic bond. They conduct electricity because their electrons are free to move in all three dimensions. Similarly, they conduct heat, which is transferred by the electrons as extra kinetic energy: they move faster. These properties persist in the liquid state, as although the crystal structure is destroyed on melting, the atoms still touch and the metallic bond persists, though it is weakened. Metals tend to be reactive towards nonmetals. Some exceptions can be found to these generalisations: for example, manganese, arsenic, antimony, and bismuth are brittle; chromium is extremely hard; gallium, rubidium, caesium, and mercury are liquid at or close to room temperature; and such as gold are chemically very inert.
Nonmetals exhibit different properties. Those forming giant covalent crystals exhibit high melting and boiling points, as it takes considerable energy to overcome the strong covalent bonds. Those forming discrete molecules are held together mostly by dispersion forces, which are more easily overcome; thus they tend to have lower melting and boiling points, and many are liquids or gases at room temperature. Nonmetals are often dull-looking. They tend to be reactive towards metals, except for the noble gases, which are inert towards most substances. They are brittle when solid as their atoms are held tightly in place. They are less dense and conduct electricity poorly, because there are no mobile electrons. Near the borderline, band gaps are small and thus many elements in that region are semiconductors, such as silicon, germanium, selenium, and tellurium. Again there are exceptions; for example, diamond has the highest thermal conductivity of all known materials, greater than any metal.
It is common to designate a class of metalloids straddling the boundary between metals and nonmetals, as elements in that region are intermediate in both physical and chemical properties. However, no consensus exists in the literature for precisely which elements should be so designated. When such a category is used, silicon, germanium, arsenic, and tellurium are almost always included, and boron and antimony usually are; but most sources include other elements as well, without agreement on which extra elements should be added, and some others subtract from this list instead.
Many other physical properties of the elements exhibit periodic variation in accordance with the periodic law, such as , , heats of fusion, heats of vaporisation, atomisation energy, and so on. Similar periodic variations appear for the compounds of the elements, which can be observed by comparing hydrides, oxides, sulfides, halides, and so on. Chemical properties are more difficult to describe quantitatively, but likewise exhibit their own periodicities. Examples include the variation in the and basic properties of the elements and their compounds, the stabilities of compounds, and methods of isolating the elements. Periodicity is and has been used very widely to predict the properties of unknown new elements and new compounds, and is central to modern chemistry.Greenwood and Earnshaw, pp. 29–31
The lanthanides are considered to be the elements La–Lu, which are all very similar to each other: historically they included only Ce–Lu, but lanthanum became included by common usage. The rare earth elements (or rare earth metals) add scandium and yttrium to the lanthanides. Analogously, the actinides are considered to be the elements Ac–Lr (historically Th–Lr), although variation of properties in this set is much greater than within the lanthanides. IUPAC recommends the names lanthanoids and actinoids to avoid ambiguity, as the -ide suffix typically denotes a negative ion; however lanthanides and actinides remain common. With the increasing recognition of lutetium and lawrencium as d-block elements, some authors began to define the lanthanides as La–Yb and the actinides as Ac–No, matching the f-block.
Many more categorisations exist and are used according to certain disciplines. In astrophysics, a metal is defined as any element with atomic number greater than 2, i.e. anything except hydrogen and helium. The term "semimetal" has a different definition in physics than it does in chemistry: bismuth is a semimetal by physical definitions, but chemists generally consider it a metal. A few terms are widely used, but without any very formal definition, such as "heavy metals", which has been given such a wide range of definitions that it has been criticised as "effectively meaningless".
The scope of terms varies significantly between authors. For example, according to IUPAC, the noble gases extend to include the whole group, including the very radioactive superheavy element oganesson. However, among those who specialise in the superheavy elements, this is not often done: in this case "noble gas" is typically taken to imply the unreactive behaviour of the lighter elements of the group. Since calculations generally predict that oganesson should not be particularly inert due to relativistic effects, and may not even be a gas at room temperature if it could be produced in bulk, its status as a noble gas is often questioned in this context. Furthermore, national variations are sometimes encountered: in Japan, alkaline earth metals often do not include beryllium and magnesium as their behaviour is different from the heavier group 2 metals.
John Newlands published a letter in the Chemical News in February 1863 on the periodicity among the chemical elements. In 1864 Newlands published an article in the Chemical News showing that if the elements are arranged in the order of their atomic weights, those having consecutive numbers frequently either belong to the same group or occupy similar positions in different groups, and he pointed out that each eighth element starting from a given one is in this arrangement a kind of repetition of the first, like the eighth note of an octave in music (The Law of Octaves). However, Newlands' formulation only worked well for the main-group elements, and encountered serious problems with the others.
German chemist Lothar Meyer noted the sequences of similar chemical and physical properties repeated at periodic intervals. According to him, if the atomic weights were plotted as ordinates (i.e. vertically) and the atomic volumes as abscissas (i.e. horizontally)—the curve obtained a series of maximums and minimums—the most electropositive elements would appear at the peaks of the curve in the order of their atomic weights. In 1864, a book of his was published; it contained an early version of the periodic table containing 28 elements, and classified elements into six families by their valence—for the first time, elements had been grouped according to their valence. Works on organizing the elements by atomic weight had until then been stymied by inaccurate measurements of the atomic weights.Meyer, Julius Lothar; Die modernen Theorien der Chemie (1864); table on page 137 In 1868, he revised his table, but this revision was published as a draft only after his death.Scerri, pp. 106–108
In 1875, the French chemist Paul-Émile Lecoq de Boisbaudran, working without knowledge of Mendeleev's prediction, discovered a new element in a sample of the mineral sphalerite, and named it gallium. He isolated the element and began determining its properties. Mendeleev, reading de Boisbaudran's publication, sent a letter claiming that gallium was his predicted eka-aluminium. Although Lecoq de Boisbaudran was initially sceptical, and suspected that Mendeleev was trying to take credit for his discovery, he later admitted that Mendeleev was correct.Scerri, p. 149 In 1879, the Swedish chemist Lars Fredrik Nilson discovered a new element, which he named scandium: it turned out to be eka-boron. Eka-silicon was found in 1886 by German chemist Clemens Winkler, who named it germanium. The properties of gallium, scandium, and germanium matched what Mendeleev had predicted.Scerri, p. 151–2 In 1889, Mendeleev noted at the Faraday Lecture to the Royal Institution in London that he had not expected to live long enough "to mention their discovery to the Chemical Society of Great Britain as a confirmation of the exactitude and generality of the periodic law". Even the discovery of the noble gases at the close of the 19th century, which Mendeleev had not predicted, fitted neatly into his scheme as an eighth main group.Scerri, pp. 164–169
Mendeleev nevertheless had some trouble fitting the known lanthanides into his scheme, as they did not exhibit the periodic change in valencies that the other elements did. After much investigation, the Czech chemist Bohuslav Brauner suggested in 1902 that the lanthanides could all be placed together in one group on the periodic table. He named this the "asteroid hypothesis" as an astronomical analogy: just as there is an asteroid belt instead of a single planet between Mars and Jupiter, so the place below yttrium was thought to be occupied by all the lanthanides instead of just one element.
The same year, English physicist Henry Moseley using X-ray spectroscopy confirmed van den Broek's proposal experimentally. Moseley determined the value of the nuclear charge of each element from aluminium to gold and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge. Nuclear charge is identical to proton count and determines the value of the atomic number ( Z) of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley's research immediately resolved discrepancies between atomic weight and chemical properties; these were cases such as tellurium and iodine, where atomic number increases but atomic weight decreases. Although Moseley was soon killed in World War I, the Swedish physicist Manne Siegbahn continued his work up to uranium, and established that it was the element with the highest atomic number then known (92). Based on Moseley and Siegbahn's research, it was also known which atomic numbers corresponded to missing elements yet to be found.
Bohr called his electron shells "rings" in 1913: atomic orbitals within shells did not exist at the time of his planetary model. Bohr explains in Part 3 of his famous 1913 paper that the maximum electrons in a shell is eight, writing, "We see, further, that a ring of n electrons cannot rotate in a single ring round a nucleus of charge ne unless n < 8." For smaller atoms, the electron shells would be filled as follows: "rings of electrons will only join if they contain equal numbers of electrons; and that accordingly the numbers of electrons on inner rings will only be 2, 4, 8." However, in larger atoms the innermost shell would contain eight electrons: "on the other hand, the periodic system of the elements strongly suggests that already in neon N = 10 an inner ring of eight electrons will occur." His proposed electron configurations for the atoms (shown to the right) mostly do not accord with those now known.Niels Bohr, "On the Constitution of Atoms and Molecules, Part III, Systems containing several nuclei" Philosophical Magazine 26:857--875 (1913)Kragh, Helge. "Niels Bohr's Second Atomic Theory."
The first one to systematically expand and correct the chemical potentials of Bohr's atomic theory was Walther Kossel in 1914 and in 1916. Kossel explained that in the periodic table new elements would be created as electrons were added to the outer shell. In Kossel's paper, he writes: "This leads to the conclusion that the electrons, which are added further, should be put into concentric rings or shells, on each of which ... only a certain number of electrons—namely, eight in our case—should be arranged. As soon as one ring or shell is completed, a new one has to be started for the next element; the number of electrons, which are most easily accessible, and lie at the outermost periphery, increases again from element to element and, therefore, in the formation of each new shell the chemical periodicity is repeated."W. Kossel, "Über Molekülbildung als Folge des Atom- baues," Ann. Phys., 1916, 49, 229-362 (237).Translated in Helge Kragh, Aarhus, LARS VEGARD, ATOMIC STRUCTURE, AND THE PERIODIC SYSTEM, Bull. Hist. Chem., VOLUME 37, Number 1 (2012), p.43.
In a 1919 paper, Irving Langmuir postulated the existence of "cells" which we now call orbitals, which could each only contain two electrons each, and these were arranged in "equidistant layers" which we now call shells. He made an exception for the first shell to only contain two electrons. The chemist Charles Rugeley Bury suggested in 1921 that eight and eighteen electrons in a shell form stable configurations. Bury proposed that the electron configurations in transitional elements depended upon the valence electrons in their outer shell. He introduced the word transition to describe the elements now known as or transition elements.
Prompted by Bohr, Wolfgang Pauli took up the problem of electron configurations in 1923. Pauli extended Bohr's scheme to use four , and formulated his exclusion principle which stated that no two electrons could have the same four quantum numbers. This explained the lengths of the periods in the periodic table (2, 8, 18, and 32), which corresponded to the number of electrons that each shell could occupy.Scerri, pp. 218–23 In 1925, Friedrich Hund arrived at configurations close to the modern ones. As a result of these advances, periodicity became based on the number of chemically active or valence electrons rather than by the valences of the elements. The Aufbau principle that describes the electron configurations of the elements was first empirically observed by Erwin Madelung in 1926, though the first to publish it was Vladimir Karapetoff in 1930. In 1961, Vsevolod Klechkovsky derived the first part of the Madelung rule (that orbitals fill in order of increasing n + ℓ) from the Thomas–Fermi model; the complete rule was derived from a similar potential in 1971 by Yury N. Demkov and Valentin N. Ostrovsky.
The quantum theory clarified the transition metals and lanthanides as forming their own separate groups, transitional between the main groups, although some chemists had already proposed tables showing them this way before then: the English chemist Henry Bassett did so in 1892, the Danish chemist Julius Thomsen in 1895, and the Swiss chemist Alfred Werner in 1905. Bohr used Thomsen's form in his 1922 Nobel Lecture; Werner's form is very similar to the modern 32-column form. In particular, this supplanted Brauner's asteroidal hypothesis.
The exact position of the lanthanides, and thus the composition of group 3, remained under dispute for decades longer because their electron configurations were initially measured incorrectly.Scerri, pp. 392−401 On chemical grounds Bassett, Werner, and Bury grouped scandium and yttrium with lutetium rather than lanthanum (the former two left an empty space below yttrium as lutetium had not yet been discovered). Hund assumed in 1927 that all the lanthanide atoms had configuration Xe4f0−145d16s2, on account of their prevailing trivalency (it is now known that the relationship between chemistry and electron configuration is more complicated than that).
In 1988, IUPAC released a report supporting this composition of group 3, a decision that was reaffirmed in 2021. Variation can still be found in textbooks on the composition of group 3, and some argumentation against this format is still published today, but chemists and physicists who have considered the matter largely agree on group 3 containing scandium, yttrium, lutetium, and lawrencium and challenge the counterarguments as being inconsistent.
A significant controversy arose with elements 102 through 106 in the 1960s and 1970s, as competition arose between the LBNL team (now led by Albert Ghiorso) and a team of Soviet scientists at the Joint Institute for Nuclear Research (JINR) led by Georgy Flyorov. Each team claimed discovery, and in some cases each proposed their own name for the element, creating an element naming controversy that lasted decades. These elements were made by bombardment of actinides with light ions.Scerri, pp. 356–9 IUPAC at first adopted a hands-off approach, preferring to wait and see if a consensus would be forthcoming. Unfortunately, it was also the height of the Cold War, and it became clear after some time that this would not happen. As such, IUPAC and the International Union of Pure and Applied Physics (IUPAP) created a Transfermium Working Group (TWG, fermium being element 100) in 1985 to set out criteria for discovery, which were published in 1991. After some further controversy, these elements received their final names in 1997, including seaborgium (106) in honour of Seaborg.
The TWG's criteria were used to arbitrate later element discovery claims from LBNL and JINR, as well as from research institutes in Germany (GSI) and Japan (Riken). Currently, consideration of discovery claims is performed by a IUPAC/IUPAP Joint Working Party. After priority was assigned, the elements were officially added to the periodic table, and the discoverers were invited to propose their names. By 2016, this had occurred for all elements up to 118, therefore completing the periodic table's first seven rows. The discoveries of elements beyond 106 were made possible by techniques devised by Yuri Oganessian at the JINR: cold fusion (bombardment of lead and bismuth by heavy ions) made possible the 1981–2004 discoveries of elements 107 through 112 at GSI and 113 at Riken, and he led the JINR team (in collaboration with American scientists) to discover elements 114 through 118 using hot fusion (bombardment of actinides by calcium ions) in 1998–2010.Scerri, pp. 356–363() The heaviest known element, oganesson (118), is named in Oganessian's honour. Element 114 is named flerovium in honour of his predecessor and mentor Flyorov.
In celebration of the periodic table's 150th anniversary, the United Nations declared the year 2019 as the International Year of the Periodic Table, celebrating "one of the most significant achievements in science". The discovery criteria set down by the TWG were updated in 2020 in response to experimental and theoretical progress that had not been foreseen in 1991. Today, the periodic table is among the most recognisable icons of chemistry. IUPAC is involved today with many processes relating to the periodic table: the recognition and naming of new elements, recommending group numbers and collective names, and the updating of atomic weights.
If the eighth period follows the pattern set by the earlier periods, then it would contain fifty elements, filling the 8s, , 6f, 7d, and finally 8p subshells in that order. But there is present discussion regarding whether this would truly be the case, as calculations predict that by this point relativistic effects should result in significant deviations from the Madelung rule. Various different models have been suggested. All agree that the eighth period should begin like the previous ones with two 8s elements, and that there should then follow a new series of g-block elements filling up the orbitals, but the precise configurations calculated for these elements vary widely between sources. Beyond this series, calculations do not agree on what exactly should follow. Filling of the , 6f, 7d, and 8p shells is indeed expected to occur in approximately that order, but they are likely to be intermingled with each other and with the 9s and 9p subshells, so that it is not clear which elements should go in which groups anymore. Eric Scerri has raised the question of whether an extended periodic table should take into account the failure of the Madelung rule in this region, or if such exceptions should be ignored. The shell structure may also be fairly formal at this point: already the electron distribution in an oganesson atom is expected to be rather uniform, with no discernible shell structure.
Nuclear stability will likely prove a decisive factor constraining the number of possible elements. It depends on the balance between the electric repulsion between protons and the strong force binding protons and neutrons together. Protons and neutrons are arranged in shells, just like electrons, and so a closed shell can significantly increase stability: the known superheavy nuclei exist because of such a shell closure. They are probably close to a predicted island of stability, where superheavy nuclides should have significantly longer half-lives: predictions range from minutes or days, to millions or billions of years. However, as the number of protons increases beyond about 126, this stabilising effect should vanish as a closed shell is passed. It is not clear if any further-out shell closures exist, due to an expected smearing out of distinct nuclear shells (as is already expected for the electron shells at oganesson). Furthermore, even if later shell closures exist, it is not clear if they would allow such heavy elements to exist.Scerri, p. 386 As such, it may be that the periodic table practically ends around element 120, as elements become too short-lived to observe; the era of discovering new elements would thus be close to its end.
Alternatively, quark matter may become stable at high mass numbers, in which the nucleus is composed of freely flowing up quark and instead of binding them into protons and neutrons; this would create a continent of stability instead of an island. Other effects may come into play: for example, in very heavy elements the 1s electrons are likely to spend a significant amount of time so close to the nucleus that they are actually inside it, which would make them vulnerable to electron capture.
Even if eighth-row elements can exist, producing them is likely to be difficult, and it should become even more difficult as atomic number rises. Although the 8s elements are expected to be reachable with present means, the first few elements are expected to require new technology, if they can be produced at all. Experimentally characterising these elements chemically would also pose a great challenge.
Alternative periodic tables are often developed to highlight or emphasize chemical or physical properties of the elements that are not as apparent in traditional periodic tables, with different ones skewed more towards emphasizing chemistry or physics at either end.Scerri, pp. 402–3 The standard form, which remains by far the most common, is somewhere in the middle.
The many different forms of the periodic table have prompted the questions of whether there is an optimal or definitive form of the periodic table, and if so, what it might be. There are no current consensus answers to either question. Janet's left-step table is being increasingly discussed as a candidate for being the optimal or most fundamental form; Scerri has written in support of it, as it clarifies helium's nature as an s-block element, increases regularity by having all period lengths repeated, faithfully follows Madelung's rule by making each period correspond to one value of + , and regularises atomic number triads and the first-row anomaly trend. While he notes that its placement of helium atop the alkaline earth metals can be seen a disadvantage from a chemical perspective, he counters this by appealing to the first-row anomaly, pointing out that the periodic table "fundamentally reduces to quantum mechanics", and that it is concerned with "abstract elements" and hence atomic properties rather than macroscopic properties.
Electron configuration
Elements are known with up to 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 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 characterised by the . Four numbers describe an orbital in an atom completely: the principal quantum number n, the azimuthal quantum number ℓ (the orbital type), the magnetic quantum number mℓ, and the spin quantum number s.
The order of subshell filling
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 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 the beginning of a new shell. Thus, with the exception of the first row, each period length appears twice:
1
hydrogen helium 2×1 = 2 elements 3
lithium4
beryllium5
boron6
carbon7
nitrogen8
oxygen9
fluorine10
neon11
sodium12
magnesium13
aluminium14
silicon15
phosphorus16
sulfur17
chlorine18
argon2×(1+3) = 8 elements 1
hydrogen helium 2×1 = 2 elements 3
lithium4
beryllium 5
boron6
carbon7
nitrogen8
oxygen9
fluorine10
neon2×(1+3) = 8 elements 11
sodium12
magnesium 13
aluminium14
silicon15
phosphorus16
sulfur17
chlorine18
argon2×(1+3) = 8 elements 19
potassium20
calcium21
scandium22
titanium23
vanadium24
chromium25
manganese26
iron27
cobalt28
nickel29
copper30
zinc31
gallium32
germanium33
arsenic34
selenium35
bromine36
krypton37
rubidium38
strontium39
yttrium40
zirconium41
niobium42
molybdenum43
technetium44
ruthenium45
rhodium46
palladium47
silver48
cadmium49
indium50
tin51
antimony52
tellurium53
iodine54
xenon2×(1+3+5) = 18 elements
Electron configuration table
Group names and numbers
Presentation forms
Variations
Period 1
Group 3
Periodic trends
Atomic radius
Ionisation energy
Electron affinity
Valence and oxidation states
Electronegativity
Metallicity
Further manifestations of periodicity
Classification of elements
History
Early history
Mendeleev
Atomic number
Electron shells
.
+ Bohr's electron configurations for light elements 2,2 2,4 4,3 4,2,2 4,4,1 8,2 8,2,1 8,4,2,2 8,8,2
Synthetic elements
Future extension beyond the seventh period
Alternative periodic tables
Notes
Bibliography
Further reading
External links
target="_blank" rel="nofollow"> Digital Collections featuring select visual representations of the periodic table of the elements, with an emphasis on alternative layouts including circular, cylindrical, pyramidal, spiral, and triangular forms.
target="_blank" rel="nofollow"> Eric Scerri, leading philosopher of science specializing in the history and philosophy of the periodic table
target="_blank" rel="nofollow"> Periodic table of endangered elements
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