Periodic table

thumb|Periodic table of the chemical elements showing the most or more commonly named [[Names for sets of chemical elements|sets of elements (in periodic tables), and a traditional dividing line between metals and nonmetals. The f-block actually fits between groups 2 and 3; it is usually shown at the foot of the table to save horizontal space.|upright=1.4]]

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"). An icon of chemistry, the periodic table 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 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 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 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 have become 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; elements beyond that can only be synthesized in 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 the patterns of the known part of the table. Some scientific discussion also continues regarding whether some elements are correctly positioned in the 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.

Structure

thumb|upright=1.5|3D views of some [[Hydrogen-like atom|hydrogen-like atomic orbitals showing probability density and phase (g orbitals and higher are not shown)]]

Each chemical element has a unique atomic number (Z for "Zahl", German for "number") representing 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 a one- or two-letter chemical symbol; those for hydrogen, helium, and lithium are respectively H, He, and Li.

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 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, 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 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, the first 94 of which are known to occur naturally on Earth. Elements up to 99 (einsteinium) have been observed in Przybylski's Star. Elements up to 100 (fermium) probably occurred in the natural nuclear fission reactor at Oklo Mine, Gabon, but they have long since decayed away. Even heavier elements may be produced in the r-process via supernovae or neutron star mergers, but this has not been confirmed. It is not clear how far they would extend past 100 and how long they would last: calculations suggest that nuclides of mass number around 280 to 290 are formed in the r-process, but quickly beta decay to nuclides that suffer spontaneous fission, so that 99.9% of the produced superheavy nuclides would decay within a month. If instead they were sufficiently long-lived, they might similarly be brought to Earth via cosmic rays, but again none have been found. Of the 94 natural elements, eighty have a stable isotope and one more (bismuth) has an almost-stable isotope (with a half-life of 2.01×1019 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 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. Groups can also be named by their first element, e.g. the "scandium group" for group 3.

Presentation forms

32 columns

18 columns

</div>

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. 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. Tables may include extra information such as the names and atomic numbers of the elements, their blocks, natural occurrences, standard atomic weight, states of matter, melting and boiling points, densities, as well as provide different classifications of the elements.

Electron configurations

The periodic table is a graphic description of the periodic law, Elements are placed in the periodic table according to their electron configurations, 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 (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.

{| class="wikitable" style="float:right; margin:0.5em; text-align:center;"

! style="text-align:right;" |ℓ =

! 0

! 1

! 2

! 3

! 4

! 5

! 6

! rowspan=2 | Shell capacity (2n2)

! style="text-align:right;" | Orbital

! s

! p

! d

! f

! g

! h

! i

! n = 1

| bgcolor="" | 1s

| colspan=6 |

| 2

! n = 2

| bgcolor="" | 2s

| bgcolor="" | 2p

| colspan=5 |

| 8

! n = 3

| bgcolor="" | 3s

| bgcolor="" | 3p

| bgcolor="" | 3d

| colspan=4 |

| 18

! n = 4

| bgcolor="" | 4s

| bgcolor="" | 4p

| bgcolor="" | 4d

| bgcolor="" | 4f

| colspan=3 |

| 32

! n = 5

| bgcolor="" | 5s

| bgcolor="" | 5p

| bgcolor="" | 5d

| bgcolor="" | 5f

| bgcolor="" | 5g

| colspan=2 |

| 50

! n = 6

| bgcolor="" | 6s

| bgcolor="" | 6p

| bgcolor="" | 6d

| bgcolor="" | 6f

| bgcolor="" | 6g

| bgcolor="" | 6h

| 72

! n = 7

| bgcolor="" | 7s

| bgcolor="" | 7p

| bgcolor="" | 7d

| bgcolor="" | 7f

| bgcolor="" | 7g

| bgcolor="" | 7h

| bgcolor="" | 7i

| 98

! Subshell capacity (4ℓ+2)

| 2

| 6

| 10

| 14

| 18

| 22

| 26

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.

:1s ≪ 2s < 2p ≪ 3s < 3p ≪ 4s < 3d < 4p ≪ 5s < 4d < 5p ≪ 6s < 4f < 5d < 6p ≪ 7s < 5f < 6d < 7p ≪ ...

Here the sign ≪ means "much less than" as opposed to < meaning just "less than". 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.

Also, the ordering of the orbitals between each ≪ changes somewhat throughout each period. For example, the ordering in argon and potassium is 3p ≪ 4s < 4p ≪ 3d; by calcium it has become 3p ≪ 4s < 3d < 4p; from scandium to copper it is 3p ≪ 3d < 4s < 4p; and from zinc to krypton it is 3p < 3d ≪ 4s < 4p as the d orbitals fall into the core at gallium. Today the notion of valence has been extended by that of the oxidation state, which is the formal charge left on an element when all other elements in a compound have been removed as their ions. However, towards the right side of the d- and f-blocks, the theoretical maximum corresponding to using all valence electrons is not achievable at all; the same situation affects oxygen, fluorine, and the light noble gases up to krypton.

{| class="wikitable" style="margin:auto;text-align:center;"

|+ Number of valence electrons

! 1

! 2

! colspan=14 |

! 3

! 4

! 5

! 6

! 7

! 8

! 9

! 10

! 11

! 12

! 13

! 14

! 15

! 16

! 17

! 18

! 1

| bgcolor="" | H 1

| colspan=30 style="border-width:0" |

| bgcolor="" | He 2

! 2

| bgcolor="" | Li 1

| bgcolor="" | Be 2

| colspan=24 style="border-width:0" |

| bgcolor="" | B 3

| bgcolor="" | C 4

| bgcolor="" | N 5

| bgcolor="" | O 6

| bgcolor="" | F 7

| bgcolor="" | Ne 8

! 3

| bgcolor="" | Na 1

| bgcolor="" | Mg 2

| colspan=24 style="border-width:0" |

| bgcolor="" | Al 3

| bgcolor="" | Si 4

| bgcolor="" | P 5

| bgcolor="" | S 6

| bgcolor="" | Cl 7

| bgcolor="" | Ar 8

! 4

| bgcolor="" | K 1

| bgcolor="" | Ca 2

| colspan=14 style="border-width:0" |

| bgcolor="" | Sc 3

| bgcolor="" | Ti 4

| bgcolor="" | V 5

| bgcolor="" | Cr 6

| bgcolor="" | Mn 7

| bgcolor="" | Fe 8

| bgcolor="" | Co 9

| bgcolor="" | Ni 10

| bgcolor="" | Cu 11

| bgcolor="" | Zn 12

| bgcolor="" | Ga 3

| bgcolor="" | Ge 4

| bgcolor="" | As 5

| bgcolor="" | Se 6

| bgcolor="" | Br 7

| bgcolor="" | Kr 8

! 5

| bgcolor="" | Rb 1

| bgcolor="" | Sr 2

| colspan=14 style="border-width:0" |

| bgcolor="" | Y 3

| bgcolor="" | Zr 4

| bgcolor="" | Nb 5

| bgcolor="" | Mo 6

| bgcolor="" | Tc 7

| bgcolor="" | Ru 8

| bgcolor="" | Rh 9

| bgcolor="" | Pd 10

| bgcolor="" | Ag 11

| bgcolor="" | Cd 12

| bgcolor="" | In 3

| bgcolor="" | Sn 4

| bgcolor="" | Sb 5

| bgcolor="" | Te 6

| bgcolor="" | I 7

| bgcolor="" | Xe 8

! 6

| bgcolor="" | Cs 1

| bgcolor="" | Ba 2

| bgcolor="" | La 3

| bgcolor="" | Ce 4

| bgcolor="" | Pr 5

| bgcolor="" | Nd 6

| bgcolor="" | Pm 7

| bgcolor="" | Sm 8

| bgcolor="" | Eu 9

| bgcolor="" | Gd 10

| bgcolor="" | Tb 11

| bgcolor="" | Dy 12

| bgcolor="" | Ho 13

| bgcolor="" | Er 14

| bgcolor="" | Tm 15

| bgcolor="" | Yb 16

| bgcolor="" | Lu 3

| bgcolor="" | Hf 4

| bgcolor="" | Ta 5

| bgcolor="" | W 6

| bgcolor="" | Re 7

| bgcolor="" | Os 8

| bgcolor="" | Ir 9

| bgcolor="" | Pt 10

| bgcolor="" | Au 11

| bgcolor="" | Hg 12

| bgcolor="" | Tl 3

| bgcolor="" | Pb 4

| bgcolor="" | Bi 5

| bgcolor="" | Po 6

| bgcolor="" | At 7

| bgcolor="" | Rn 8

! 7

| bgcolor="" | Fr 1

| bgcolor="" | Ra 2

| bgcolor="" | Ac 3

| bgcolor="" | Th 4

| bgcolor="" | Pa 5

| bgcolor="" | U 6

| bgcolor="" | Np 7

| bgcolor="" | Pu 8

| bgcolor="" | Am 9

| bgcolor="" | Cm 10

| bgcolor="" | Bk 11

| bgcolor="" | Cf 12

| bgcolor="" | Es 13

| bgcolor="" | Fm 14

| bgcolor="" | Md 15

| bgcolor="" | No 16

| bgcolor="" | Lr 3

| bgcolor="" | Rf 4

| bgcolor="" | Db 5

| bgcolor="" | Sg 6

| bgcolor="" | Bh 7

| bgcolor="" | Hs 8

| bgcolor="" | Mt 9

| bgcolor="" | Ds 10

| bgcolor="" | Rg 11

| bgcolor="" | Cn 12

| bgcolor="" | Nh 3

| bgcolor="" | Fl 4

| bgcolor="" | Mc 5

| bgcolor="" | Lv 6

| bgcolor="" | Ts 7

| bgcolor="" | Og 8

A full explanation requires considering the energy that would be released in forming compounds with different valences rather than simply considering electron configurations alone. 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 ionizing 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 hybridization can be used for the less electronegative p-block elements.

frame|center|Oxidation states of the transition metals. The solid dots show common oxidation states, and the hollow dots show possible but unlikely states.

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. Because ionizing the transition metals further does not uncover any new inner subshells, their oxidation states tend to vary by steps of 1 instead. The very last actinides go further than the lanthanides towards low oxidation states: mendelevium is more easily reduced to the +2 state than thulium or even europium (the lanthanide with the most stable +2 state, on account of its half-filled f-shell), and nobelium outright favours +2 over +3, in contrast to ytterbium. This often leads to similarities in maximum and minimum oxidation states (e.g. sulfur and selenium in group 16 both have maximum oxidation state +6, as in SO3 and SeO3, and minimum oxidation state −2, as in sulfides and selenides); but not always (e.g. oxygen is not known to form oxidation state +6, despite being in the same group as sulfur and selenium).

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. 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.

Metallicity

thumb|right|The diamond-cubic structure, a giant covalent structure adopted by carbon (as diamond), as well as by silicon, germanium, and (grey) tin, all in group 14. (In grey tin, the band gap vanishes and metallization occurs. Tin has another allotrope, white tin, whose structure is even more metallic.)

A simple substance is a substance formed from atoms of one chemical element. The simple substances of the more electronegative atoms tend to share electrons (form covalent bonds) with each other. They form either small molecules (like hydrogen or oxygen) or giant structures stretching indefinitely (like carbon or silicon). The noble gases simply stay as single atoms, as they already have a full shell.

thumb|right|Graphite and diamond, two allotropes of carbon

The more electropositive atoms tend to instead lose electrons, creating a "sea" of electrons engulfing cations. This negatively charged "sea" pulls on all the ions and keeps them together in a metallic bond. Elements forming such bonds are often called metals; those which do not are often called nonmetals.

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 delocalized 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.

thumb|upright=1.6|Graph of carbon atoms being brought together to form a diamond crystal, demonstrating formation of the electronic band structure and band gap. The right graph shows the energy levels as a function of the spacing between atoms. When far apart (right side of graph) all the atoms have discrete valence orbitals p and s with the same energies. However, when the atoms come closer (left side), their electron orbitals begin to spatially overlap. The orbitals [[Hybridization (chemistry)|hybridize into N molecular orbitals each with a different energy, where N is the number of atoms in the crystal. Since N is such a large number, adjacent orbitals are extremely close together in energy so the orbitals can be considered a continuous energy band. At the actual diamond crystal cell size (denoted by a), two bands are formed, called the valence and conduction bands, separated by a 5.5 eV band gap. (Here only the valence 2s and 2p electrons have been illustrated; the 1s orbitals do not significantly overlap, so the bands formed from them are much narrower.)]]

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 band gaps and are semiconductors at ambient conditions: electrons can cross the gap when thermally excited. (Boron is also a semiconductor at ambient conditions.) The band gap disappears in tin, so that tin and lead become metals. See metallization pressure for values for all nonmetals.

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 metalloids.

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 network 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 characterized in this way. Theoretical considerations and current experimental evidence suggest that all of those elements would metallise if they could form condensed phases,

File:Iron electrolytic and 1cm3 cube.jpg|Iron, a metal

Sulfur - El Desierto mine, San Pablo de Napa, Daniel Campos Province, Potosí, Bolivia.jpg|Sulfur, a nonmetal

Arsen 1a.jpg|Arsenic, an element often called a semi-metal or metalloid

</gallery>

Generally, metals are shiny and dense. 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. antimony, bismuth, and uranium are brittle (not an exhaustive list); gallium, rubidium, caesium, and mercury are liquid at or close to room temperature; and noble metals 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. Near the borderline, band gaps are small and thus many elements in that region are semiconductors, such as silicon, germanium,

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. and that included in the Encyclopædia Britannica does not refer to metalloids or semi-metals at all.|name=metalloids For example, unlike all the other elements generally considered metalloids or nonmetals, antimony's only stable form has metallic conductivity. Moreover, the element resembles bismuth and, more generally, the other p-block metals in its physical and chemical behaviour. On this basis some authors have argued that it is better classified as a metal than as a metalloid. On the other hand, selenium has some semiconducting properties in its most stable form (though it also has insulating allotropes) and it has been argued that it should be considered a metalloid Some similarities can also be found between the main groups and the transition metal groups, or between the early actinides and early transition metals, when the elements have the same number of valence electrons. Thus uranium somewhat resembles chromium and tungsten in group 6,

Classification of elements

[[File:Simple Periodic Table Chart-en.svg|upright=1.5|thumb|right|A periodic table colour-coded to show some commonly used sets of similar elements. The categories and their boundaries differ somewhat between sources. Lutetium and lawrencium in group 3 are also transition metals. Some divide the p-block elements from groups 13 to 16 by metallicity, and IUPAC does not presently mention it as allowable in its Principles of Chemical Nomenclature.

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 transactinides or superheavy elements are the short-lived elements beyond the actinides, starting at lawrencium or rutherfordium (depending on where the actinides are taken to end).

Many more categorizations 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 metal", which has been given such a wide range of definitions that it has been criticized 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 specialize 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.

History

Early history

In 1817, German physicist Johann Wolfgang Döbereiner began one of the earliest attempts to classify the elements. In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups triads. Chlorine, bromine, and iodine formed a triad; as did calcium, strontium, and barium; lithium, sodium, and potassium; and sulfur, selenium, and tellurium. Various chemists continued his work and were able to identify more and more relationships between small groups of elements. However, they could not build one scheme that encompassed them all.

thumb|right|upright=1.5|Newlands's table of the elements in 1866.|alt=Newlands's table of the elements.

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). In 1868, he revised his table, but this revision was published as a draft only after his death.

Mendeleev

The definitive breakthrough came from the Russian chemist Dmitri Mendeleev. Although other chemists (including Meyer) had found some other versions of the periodic system at about the same time, Mendeleev was the most dedicated to developing and defending his system, and it was his system that most affected the scientific community. On 17 February 1869 (1 March 1869 in the Gregorian calendar), Mendeleev began arranging the elements and comparing them by their atomic weights. He began with a few elements, and over the course of the day his system grew until it encompassed most of the known elements. After he found a consistent arrangement, his printed table appeared in May 1869 in the journal of the Russian Chemical Society. When elements did not appear to fit in the system, he boldly predicted that either valencies or atomic weights had been measured incorrectly, or that there was a missing element yet to be discovered.

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. 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. 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.

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 New Zealand physicist Ernest Rutherford coined the word "atomic number" for this nuclear charge. In van den Broek's published article he illustrated the first electronic periodic table showing the elements arranged according to the number of their electrons.

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. Based on Moseley and Siegbahn's research, it was also known which atomic numbers corresponded to missing elements yet to be found: 43, 61, 72, 75, 85, and 87.

The dawn of atomic physics also clarified the situation of isotopes. In the decay chains of the primordial radioactive elements thorium and uranium, it soon became evident that there were many apparent new elements that had different atomic weights but exactly the same chemical properties. In 1913, Frederick Soddy coined the term "isotope" to describe this situation, and considered isotopes to merely be different forms of the same chemical element. This furthermore clarified discrepancies such as tellurium and iodine: tellurium's natural isotopic composition is weighted towards heavier isotopes than iodine's, but tellurium has a lower atomic number.

Electron shells

The Danish physicist Niels Bohr applied Max Planck's idea of quantization to the atom. He concluded that the energy levels of electrons were quantised: only a discrete set of stable energy states were allowed. Bohr then attempted to understand periodicity through electron configurations, surmising in 1913 that the outer electrons should be responsible for the chemical properties of the element. In 1913, he produced the first electronic periodic table based on a quantum atom.

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 electrons cannot rotate in a single ring round a nucleus of charge ne unless < 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 = 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. They were improved further after the work of Arnold Sommerfeld and Edmund Stoner discovered more quantum numbers.</blockquote>

In a 1919 paper, Irving Langmuir postulated the existence of "cells" which we now call orbitals, which could each only contain eight 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 transition metals or transition elements. Bohr's theory was vindicated by the discovery of element 72: Georges Urbain claimed to have discovered it as the rare earth element celtium, but Bury and Bohr had predicted that element 72 could not be a rare earth element and had to be a homologue of zirconium. Dirk Coster and Georg von Hevesy searched for the element in zirconium ores and found element 72, which they named hafnium after Bohr's hometown of Copenhagen (Hafnia in Latin). Urbain's celtium proved to be simply purified lutetium (element 71). Hafnium and rhenium thus became the last stable elements to be discovered. 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. 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. and regularises atomic number triads and the first-row anomaly trend.

Notes

References

Bibliography

Scerri, Eric R. (2020). The Periodic Table, Its Story and Its Significance (2nd ed.). New York: Oxford University Press. .

External links

Periodic Table featured topic page on Science History Institute 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.
IUPAC Periodic Table of the Elements
Dynamic periodic table, with interactive layouts
Eric Scerri, leading philosopher of science specializing in the history and philosophy of the periodic table
The Internet Database of Periodic Tables
Periodic table of endangered elements
Periodic table of samples
Periodic table of videos
WebElements
The Periodic Graphics of Elements