Editing Periodic table (section)

== Periodic trends ==
{{Main|Periodic trends}}
As chemical reactions involve the valence electrons,<ref name="cartoon" /> elements with similar outer electron configurations may be expected to react similarly and form compounds with similar proportions of elements in them.<ref name="Greenwood27" /> Such elements are placed in the same group, and thus there tend to be clear similarities and trends in chemical behaviour as one proceeds down a group.<ref>{{cite book |last=Messler|first=R. W.|title=The essence of materials for engineers|year=2010|publisher=Jones & Bartlett Publishers|location=Sudbury, MA|isbn=978-0-7637-7833-0|page=32}}</ref> As analogous configurations occur at regular intervals, the properties of the elements thus exhibit periodic recurrences, hence the name of the periodic table and the periodic law. These periodic recurrences were noticed well before the underlying theory that explains them was developed.<ref name="Myers">{{cite book |last=Myers|first=R.|title=The basics of chemistry|url=https://archive.org/details/basicschemistry00myer_641|url-access=limited|year=2003|pages=[https://archive.org/details/basicschemistry00myer_641/page/n74 61]–67|publisher=Greenwood Publishing Group|location=Westport, CT|isbn=978-0-313-31664-7}}</ref><ref name="chang2">{{cite book|last=Chang|first=R.|title=Chemistry|url=https://archive.org/details/riimchemistry00chan/page/289|url-access=registration|year=2002|publisher=McGraw-Hill|location=New York|edition=7|isbn=978-0-07-112072-2|pages=[https://archive.org/details/riimchemistry00chan/page/289 289–310, 340–42]}}</ref>

=== Atomic radius ===
Historically, the physical size of atoms was unknown until the early 20th century. The first calculated estimate of the atomic radius of hydrogen was published by physicist [[Arthur Erich Haas|Arthur Haas]] in 1910 to within an order of magnitude (a factor of 10) of the accepted value, the [[Bohr radius]] (~0.529 Å). In his model, Haas used a single-electron configuration based on the classical atomic model proposed by [[J. J. Thomson]] in 1904, often called the [[plum-pudding model]].<ref>Haas, Arthur Erich (1884–1941) Uber die elektrodynamische Bedeutung des Planckschen Strahlungsgesetzes und uber eine neue Bestimmung des elektrischen Elementarquantums und der dimension des wasserstoffatoms. Sitzungsberichte der kaiserlichen Akademie der Wissenschaften in Wien. 2a, 119 pp 119–144 (1910). Haas AE. Die Entwicklungsgeschichte des Satzes von der Erhaltung der Kraft. Habilitation Thesis, Vienna, 1909. Hermann, A. Arthur Erich Haas, Der erste Quantenansatz für das Atom. Stuttgart, 1965 [contains a reprint]</ref>

[[Atomic radius|Atomic radii]] (the size of atoms) are dependent on the sizes of their outermost orbitals.<ref name=SB23>Siekierski and Burgess, pp. 23–26</ref> 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.<ref name="cartoon" /><ref name="chemguidear">{{cite web |url=https://www.chemguide.co.uk/atoms/properties/atradius.html |title=Atomic and Ionic Radius |last=Clark |first=Jim |date=2012 |website=Chemguide |access-date=30 March 2021 |archive-date=14 November 2020 |archive-url=https://web.archive.org/web/20201114002613/https://www.chemguide.co.uk/atoms/properties/atradius.html |url-status=live }}</ref> The first row of each block is abnormally small, due to an effect called [[kainosymmetry]] or primogenic repulsion:<ref>{{cite journal |last1=Cao |first1=Chang-Su |last2=Hu |first2=Han-Shi |last3=Li |first3=Jun |last4=Schwarz |first4=W. H. Eugen |date=2019 |title=Physical origin of chemical periodicities in the system of elements |journal=Pure and Applied Chemistry |volume=91 |issue=12 |pages=1969–1999 |doi=10.1515/pac-2019-0901 |s2cid=208868546 |doi-access=free }}</ref> the 1s, 2p, 3d, and 4f subshells have no inner analogues. For example, the 2p orbitals do not experience strong repulsion from the 1s and 2s orbitals, which have quite different angular charge distributions, and hence are not very large; but the 3p orbitals experience strong repulsion from the 2p orbitals, which have similar angular charge distributions. Thus higher s-, p-, d-, and f-subshells experience strong repulsion from their inner analogues, which have approximately the same angular distribution of charge, and must expand to avoid this. This makes significant differences arise between the small 2p elements, which prefer [[multiple bond]]ing, and the larger 3p and higher p-elements, which do not.<ref name=SB23/> Similar anomalies arise for the 1s, 2p, 3d, 4f, and the hypothetical {{Not a typo|5g}} elements:<ref name="Kaupp">{{cite journal |last=Kaupp |first=Martin |date=1 December 2006 |title=The role of radial nodes of atomic orbitals for chemical bonding and the periodic table |journal=Journal of Computational Chemistry |volume=28 |issue=1 |pages=320–25 |doi=10.1002/jcc.20522 |pmid=17143872 |s2cid=12677737 |doi-access=free }}</ref> 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.<ref name="PTSS2" />

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.<ref name="chemguidear" /> The 4p and 5d atoms, coming immediately after new types of transition series are first introduced, are smaller than would have been expected,<ref name="Greenwood29">Greenwood and Earnshaw, p. 29</ref> 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.<ref name=SB23/> Together with kainosymmetry, this results in an even-odd difference between the periods (except in the s-block){{efn|Properties of the p-block elements nevertheless do affect the succeeding s-block elements. The 3s shell in sodium is above a kainosymmetric 2p core, but the 4s shell in potassium is above the much larger 3p core. Hence while one would have already expected potassium atoms to be larger than sodium atoms, the size difference is greater than usual.<ref name=SB23/>}} 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.<ref name="PTSS2" /><ref>{{cite journal |last1=Imyanitov |first1=Naum S. |date=2018 |title=Is the periodic table appears doubled? Two variants of division of elements into two subsets. Internal and secondary periodicity |url= |journal=Foundations of Chemistry |volume=21 |issue= |pages=255–284 |doi=10.1007/s10698-018-9321-z |s2cid=254514910 |access-date=}}</ref> A similar situation holds for the d-block, with lutetium through tungsten atoms being slightly smaller than yttrium through molybdenum atoms respectively.<ref>{{cite journal |last1=Chistyakov |first1=V. M. |date=1968 |title=Biron's Secondary Periodicity of the Side d-subgroups of Mendeleev's Short Table |url=https://archive.org/details/sim_russian-journal-of-general-chemistry_1968-02_38_2/page/212/mode/2up |journal=Journal of General Chemistry of the USSR |volume=38 |issue=2 |pages=213–214 |doi= |access-date=6 January 2024}}</ref><ref name="Calc1">{{cite journal|author1=P. Pyykkö|author2=M. Atsumi|year=2009|title=Molecular Single-Bond Covalent Radii for Elements 1-118|journal=Chemistry: A European Journal|volume=15|issue=1|pages=186–197|doi=10.1002/chem.200800987|pmid=19058281}}</ref>

[[File:Pouring liquid mercury bionerd.jpg|thumb|right|Liquid mercury. Its liquid state at standard conditions is the result of relativistic effects.<ref name=PekkaPyykko/>]]
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 quantum chemistry|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&nbsp;subshell: one p&nbsp;orbital is relativistically stabilized and shrunken (it fills in thallium and lead), but the other two (filling in bismuth through radon) are relativistically destabilized and expanded.<ref name=SB23/> Relativistic effects also explain why [[gold]] is golden and [[mercury (element)|mercury]] is a liquid at room temperature.<ref name="PekkaPyykko">{{cite journal |doi=10.1021/ar50140a002 |title=Relativity and the periodic system of elements |year=1979 |last1=Pyykkö |first1=Pekka |last2=Desclaux |first2=Jean Paul |journal=Accounts of Chemical Research |volume=12 |issue=8 |page=276}}</ref><ref name="Norrby">{{cite journal |doi=10.1021/ed068p110 |title=Why is mercury liquid? Or, why do relativistic effects not get into chemistry textbooks? |year=1991 |last1=Norrby |first1=Lars J. |journal=Journal of Chemical Education |volume=68 |issue=2 |page=110 |bibcode = 1991JChEd..68..110N}}</ref> They are expected to become very strong in the late seventh period, potentially leading to a collapse of periodicity.<ref name=actrev/> Electron configurations are only clearly known until element 108 ([[hassium]]), and experimental chemistry beyond 108 has only been done for elements 112 ([[copernicium]]) through 115 ([[moscovium]]), so the chemical characterization of the heaviest elements remains a topic of current research.<ref name="Schändel 2003 277">{{cite book|title=The Chemistry of Superheavy Elements|last=Schädel|first=M.|year=2003|publisher=Kluwer Academic Publishers|location=Dordrecht|isbn=978-1-4020-1250-1|page=277}}</ref><ref name=moscovium>{{cite journal |last1=Yakushev |first1=A. |last2=Khuyagbaatar |first2=J. |first3=Ch. E. |last3=Düllmann |first4=M. |last4=Block |first5=R. A. |last5=Cantemir |first6=D. M. |last6=Cox |first7=D. |last7=Dietzel |first8=F. |last8=Giacoppo |first9=Y. |last9=Hrabar |first10=M. |last10=Iliaš |first11=E. |last11=Jäger |first12=J. |last12=Krier |first13=D. |last13=Krupp |first14=N. |last14=Kurz |first15=L. |last15=Lens |first16=S. |last16=Löchner |first17=Ch. |last17=Mokry |first18=P. |last18=Mošať |first19=V. |last19=Pershina |first20=S. |last20=Raeder |first21=D. |last21=Rudolph |first22=J. |last22=Runke |first23=L. G. |last23=Sarmiento |first24=B. |last24=Schausten |first25=U. |last25=Scherer |first26=P. |last26=Thörle-Pospiesch |first27=N. |last27=Trautmann |first28=M. |last28=Wegrzecki |first29=P. |last29=Wieczorek |date=23 September 2024 |title=Manifestation of relativistic effects in the chemical properties of nihonium and moscovium revealed by gas chromatography studies |journal=Frontiers in Chemistry |volume=12 |issue= |pages= |doi=10.3389/fchem.2024.1474820 |doi-access=free |pmid=39391836 |pmc=11464923 |bibcode=2024FrCh...1274820Y }}</ref>

The trend that atomic radii decrease from left to right is also present in [[ionic radius|ionic radii]], though it is more difficult to examine because the most common ions of consecutive elements normally differ in charge. Ions with the same electron configuration decrease in size as their atomic number rises, due to increased attraction from the more positively charged nucleus: thus for example ionic radii decrease in the series Se<sup>2−</sup>, Br<sup>−</sup>, Rb<sup>+</sup>, Sr<sup>2+</sup>, Y<sup>3+</sup>, Zr<sup>4+</sup>, Nb<sup>5+</sup>, Mo<sup>6+</sup>, Tc<sup>7+</sup>. Ions of the same element get smaller as more electrons are removed, because the attraction from the nucleus begins to outweigh the repulsion between electrons that causes electron clouds to expand: thus for example ionic radii decrease in the series V<sup>2+</sup>, V<sup>3+</sup>, V<sup>4+</sup>, V<sup>5+</sup>.<ref>Wulfsberg, pp. 33–34</ref>

=== Ionisation energy ===
[[File:First Ionization Energy blocks.svg|thumb|right|512px|Graph of first ionisation energies of the elements in electronvolts (predictions used for elements 109–118)]]
The first [[ionisation energy]] of an atom is the energy required to remove an electron from it. This varies with the atomic radius: ionisation energy increases left to right and down to up, because electrons that are closer to the nucleus are held more tightly and are more difficult to remove. Ionisation energy thus is minimized at the first element of each period – hydrogen and the [[alkali metal]]s – and then generally rises until it reaches the [[noble gas]] at the right edge of the period.<ref name="cartoon" /> There are some exceptions to this trend, such as oxygen, where the electron being removed is paired and thus interelectronic repulsion makes it easier to remove than expected.<ref name="Greenwood294">Greenwood and Earnshaw, pp. 24–5</ref>

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.<ref name="chemguideIE">{{cite web |url=https://www.chemguide.co.uk/atoms/properties/ies.html |title=Ionisation Energy |last=Clark |first=Jim |date=2016 |website=Chemguide |access-date=30 March 2021 |archive-date=22 April 2021 |archive-url=https://web.archive.org/web/20210422032340/https://www.chemguide.co.uk/atoms/properties/ies.html |url-status=live }}</ref>

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.<ref name="chemguideIE" />

=== Electron affinity ===
[[File:Electron affinity of the elements.svg|thumb|384px|right|Trend in electron affinities]]
The opposite property to ionisation energy is the [[electron affinity]], which is the energy released when adding an electron to the atom.<ref name="chemguideea" /> A passing electron will be more readily attracted to an atom if it feels the pull of the nucleus more strongly, and especially if there is an available partially filled outer orbital that can accommodate it. Therefore, electron affinity tends to increase down to up and left to right. The exception is the last column, the noble gases, which have a full shell and have no room for another electron. This gives the [[halogen]]s in the next-to-last column the highest electron affinities.<ref name="cartoon" />

Some atoms, like the noble gases, have no electron affinity: they cannot form stable gas-phase anions.<ref>{{cite journal |last1=Cárdenas |first1=Carlos |last2=Ayers |first2=Paul |first3=Frank |last3=De Proft |first4=David J. |last4=Tozer |first5=Paul |last5=Geerlings |date=2010 |title=Should negative electron affinities be used for evaluating the chemical hardness? |journal=Physical Chemistry Chemical Physics |volume=13 |issue=6 |pages=2285–2293 |doi=10.1039/C0CP01785J|pmid=21113528 }}</ref> (They can form metastable [[Resonance (particle physics)|resonances]] if the incoming electron arrives with enough kinetic energy, but these inevitably and rapidly [[Autoionization|autodetach]]: for example, the lifetime of the most long-lived He<sup>−</sup> level is about 359&nbsp;microseconds.)<ref>{{cite journal |last1=Schmidt |first1=H. T. |last2=Reinhed |first2=P. |first3=A. |last3=Orbán |first4=S. |last4=Rosén |first5=R. D. |last5=Thomas |first6=H. A. B. |last6=Johansson |first7=J. |last7=Werner |first8=D. |last8=Misra |first9=M. |last9=Björkhage |first10=L. |last10=Brännholm |first11=P. |last11=Löfgren |first12=L. |last12=Liljeby |first13=H. |last13=Cederquist |date=2012 |title=The lifetime of the helium anion |journal=Journal of Physics: Conference Series |volume=388 |issue= 1|pages=012006 |doi=10.1088/1742-6596/388/1/012006 |doi-access=free |bibcode=2012JPhCS.388a2006S }}</ref> The noble gases, having high ionisation energies and no electron affinity, have little inclination towards gaining or losing electrons and are generally unreactive.<ref name="cartoon" />

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 (because of extreme interelectronic repulsion for the very small fluorine atom), but is more reactive.<ref name="chemguideea">{{cite web |url=https://www.chemguide.co.uk/atoms/properties/eas.html |title=Electron Affinity |last=Clark |first=Jim |date=2012 |website=Chemguide |access-date=30 March 2021 |archive-date=23 April 2021 |archive-url=https://web.archive.org/web/20210423195854/https://www.chemguide.co.uk/atoms/properties/eas.html |url-status=live }}</ref>

===Valence and oxidation states===
{{Multiple image|total_width = 256
<!-- Layout parameters -->
| align             = right
| direction         = horizontal
| width             =

<!--image 1-->
| image1            = Oxid olovnatý.JPG
| width1            =
| alt1              = 
| link1             = 
| thumbtime1        =
| caption1          = 
<!--image 2-->
| image2            = Lead dioxide.jpg
| width2            = <!-- displayed width of image; overridden by "width" above -->
| alt2              = 
| link2             = 
| thumbtime2        =
| caption2          = 
<!-- and so on, to a maximum of 10 images (image10) -->

<!-- Footer -->
| footer_background = <!-- footer background as a 'hex triplet' web color prefixed by # e.g. #33CC00 -->
| footer_align      = <!-- left (default), center, right -->
| footer            = [[Lead(II) oxide]] (PbO, left) and [[lead(IV) oxide]] (PbO<sub>2</sub>, right), the two stable oxides of [[lead]]
}}
The [[valence (chemistry)|valence]] of an element can be defined either as the number of hydrogen atoms that can combine with it to form a simple binary hydride, or as twice the number of oxygen atoms that can combine with it to form a simple binary oxide (that is, not a [[peroxide]] or a [[superoxide]]).<ref name=johnson/> The valences of the main-group elements are directly related to the group number: the hydrides in the main groups 1–2 and 13–17 follow the formulae MH, MH<sub>2</sub>, MH<sub>3</sub>, MH<sub>4</sub>, MH<sub>3</sub>, MH<sub>2</sub>, and finally MH. The highest oxides instead increase in valence, following the formulae M<sub>2</sub>O, MO, M<sub>2</sub>O<sub>3</sub>, MO<sub>2</sub>, M<sub>2</sub>O<sub>5</sub>, MO<sub>3</sub>, M<sub>2</sub>O<sub>7</sub>.{{efn|There are many lower oxides as well: for example, [[phosphorus]] in group 15 forms two oxides, [[phosphorus trioxide|P<sub>2</sub>O<sub>3</sub>]] and [[phosphorus pentoxide|P<sub>2</sub>O<sub>5</sub>]].<ref name="Greenwood27">Greenwood and Earnshaw, pp. 27–9</ref>}} 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.<ref name="Greenwood27" />

The electron configuration suggests a ready explanation from the number of electrons available for bonding;<ref name="Greenwood27" /> indeed, the number of valence electrons starts at 1 in group 1, and then increases towards the right side of the periodic table, only resetting at 3 whenever each new block starts. Thus in period 6, Cs–Ba have 1–2 valence electrons; La–Yb have 3–16; Lu–Hg have 3–12; and Tl–Rn have 3–8.<ref name=wulfsberg26>Wulfsberg, p. 26</ref> 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;<ref>Wulfsberg, p. 28</ref> the same situation affects oxygen, fluorine, and the light noble gases up to krypton.<ref>Wulfsberg, p. 274</ref>
{| class="wikitable" style="margin:auto;text-align:center;"
|+ Number of valence electrons
!
! [[Alkali metal|1]]
! [[Alkaline earth metal|2]]
! colspan=14 |
! [[Group 3 element|3]]
! [[Group 4 element|4]]
! [[Group 5 element|5]]
! [[Group 6 element|6]]
! [[Group 7 element|7]]
! [[Group 8 element|8]]
! [[Group 9 element|9]]
! [[Group 10 element|10]]
! [[Group 11 element|11]]
! [[Group 12 element|12]]
! [[Boron group|13]]
! [[Carbon group|14]]
! [[Pnictogen|15]]
! [[Chalcogen|16]]
! [[Halogen|17]]
! [[Noble gas|18]]
|-
! [[Period 1 element|1]]
| bgcolor="{{element color|s-block}}" | H<br />1
| colspan=30 style="border-width:0" |
| bgcolor="{{element color|s-block}}" | He<br />2
|-
! [[Period 2 element|2]]
| bgcolor="{{element color|s-block}}" | Li<br />1
| bgcolor="{{element color|s-block}}" | Be<br />2
| colspan=24 style="border-width:0" |
| bgcolor="{{element color|p-block}}" | B<br />3
| bgcolor="{{element color|p-block}}" | C<br />4
| bgcolor="{{element color|p-block}}" | N<br />5
| bgcolor="{{element color|p-block}}" | O<br />6
| bgcolor="{{element color|p-block}}" | F<br />7
| bgcolor="{{element color|p-block}}" | Ne<br />8
|-
! [[Period 3 element|3]]
| bgcolor="{{element color|s-block}}" | Na<br />1
| bgcolor="{{element color|s-block}}" | Mg<br />2
| colspan=24 style="border-width:0" |
| bgcolor="{{element color|p-block}}" | Al<br />3
| bgcolor="{{element color|p-block}}" | Si<br />4
| bgcolor="{{element color|p-block}}" | P<br />5
| bgcolor="{{element color|p-block}}" | S<br />6
| bgcolor="{{element color|p-block}}" | Cl<br />7
| bgcolor="{{element color|p-block}}" | Ar<br />8
|-
! [[Period 4 element|4]]
| bgcolor="{{element color|s-block}}" | K<br />1
| bgcolor="{{element color|s-block}}" | Ca<br />2
| colspan=14 style="border-width:0" |
| bgcolor="{{element color|d-block}}" | Sc<br />3
| bgcolor="{{element color|d-block}}" | Ti<br />4
| bgcolor="{{element color|d-block}}" | V<br />5
| bgcolor="{{element color|d-block}}" | Cr<br />6
| bgcolor="{{element color|d-block}}" | Mn<br />7
| bgcolor="{{element color|d-block}}" | Fe<br />8
| bgcolor="{{element color|d-block}}" | Co<br />9
| bgcolor="{{element color|d-block}}" | Ni<br />10
| bgcolor="{{element color|d-block}}" | Cu<br />11
| bgcolor="{{element color|d-block}}" | Zn<br />12
| bgcolor="{{element color|p-block}}" | Ga<br />3
| bgcolor="{{element color|p-block}}" | Ge<br />4
| bgcolor="{{element color|p-block}}" | As<br />5
| bgcolor="{{element color|p-block}}" | Se<br />6
| bgcolor="{{element color|p-block}}" | Br<br />7
| bgcolor="{{element color|p-block}}" | Kr<br />8
|-
! [[Period 5 element|5]]
| bgcolor="{{element color|s-block}}" | Rb<br />1
| bgcolor="{{element color|s-block}}" | Sr<br />2
| colspan=14 style="border-width:0" |
| bgcolor="{{element color|d-block}}" | Y<br />3
| bgcolor="{{element color|d-block}}" | Zr<br />4
| bgcolor="{{element color|d-block}}" | Nb<br />5
| bgcolor="{{element color|d-block}}" | Mo<br />6
| bgcolor="{{element color|d-block}}" | Tc<br />7
| bgcolor="{{element color|d-block}}" | Ru<br />8
| bgcolor="{{element color|d-block}}" | Rh<br />9
| bgcolor="{{element color|d-block}}" | Pd<br />10
| bgcolor="{{element color|d-block}}" | Ag<br />11
| bgcolor="{{element color|d-block}}" | Cd<br />12
| bgcolor="{{element color|p-block}}" | In<br />3
| bgcolor="{{element color|p-block}}" | Sn<br />4
| bgcolor="{{element color|p-block}}" | Sb<br />5
| bgcolor="{{element color|p-block}}" | Te<br />6
| bgcolor="{{element color|p-block}}" | I<br />7
| bgcolor="{{element color|p-block}}" | Xe<br />8
|-
! [[Period 6 element|6]]
| bgcolor="{{element color|s-block}}" | Cs<br />1
| bgcolor="{{element color|s-block}}" | Ba<br />2
| bgcolor="{{element color|f-block}}" | La<br />3
| bgcolor="{{element color|f-block}}" | Ce<br />4
| bgcolor="{{element color|f-block}}" | Pr<br />5
| bgcolor="{{element color|f-block}}" | Nd<br />6
| bgcolor="{{element color|f-block}}" | Pm<br />7
| bgcolor="{{element color|f-block}}" | Sm<br />8
| bgcolor="{{element color|f-block}}" | Eu<br />9
| bgcolor="{{element color|f-block}}" | Gd<br />10
| bgcolor="{{element color|f-block}}" | Tb<br />11
| bgcolor="{{element color|f-block}}" | Dy<br />12
| bgcolor="{{element color|f-block}}" | Ho<br />13
| bgcolor="{{element color|f-block}}" | Er<br />14
| bgcolor="{{element color|f-block}}" | Tm<br />15
| bgcolor="{{element color|f-block}}" | Yb<br />16
| bgcolor="{{element color|d-block}}" | Lu<br />3
| bgcolor="{{element color|d-block}}" | Hf<br />4
| bgcolor="{{element color|d-block}}" | Ta<br />5
| bgcolor="{{element color|d-block}}" | W<br />6
| bgcolor="{{element color|d-block}}" | Re<br />7
| bgcolor="{{element color|d-block}}" | Os<br />8
| bgcolor="{{element color|d-block}}" | Ir<br />9
| bgcolor="{{element color|d-block}}" | Pt<br />10
| bgcolor="{{element color|d-block}}" | Au<br />11
| bgcolor="{{element color|d-block}}" | Hg<br />12
| bgcolor="{{element color|p-block}}" | Tl<br />3
| bgcolor="{{element color|p-block}}" | Pb<br />4
| bgcolor="{{element color|p-block}}" | Bi<br />5
| bgcolor="{{element color|p-block}}" | Po<br />6
| bgcolor="{{element color|p-block}}" | At<br />7
| bgcolor="{{element color|p-block}}" | Rn<br />8
|-
! [[Period 7 element|7]]
| bgcolor="{{element color|s-block}}" | Fr<br />1
| bgcolor="{{element color|s-block}}" | Ra<br />2
| bgcolor="{{element color|f-block}}" | Ac<br />3
| bgcolor="{{element color|f-block}}" | Th<br />4
| bgcolor="{{element color|f-block}}" | Pa<br />5
| bgcolor="{{element color|f-block}}" | U<br />6
| bgcolor="{{element color|f-block}}" | Np<br />7
| bgcolor="{{element color|f-block}}" | Pu<br />8
| bgcolor="{{element color|f-block}}" | Am<br />9
| bgcolor="{{element color|f-block}}" | Cm<br />10
| bgcolor="{{element color|f-block}}" | Bk<br />11
| bgcolor="{{element color|f-block}}" | Cf<br />12
| bgcolor="{{element color|f-block}}" | Es<br />13
| bgcolor="{{element color|f-block}}" | Fm<br />14
| bgcolor="{{element color|f-block}}" | Md<br />15
| bgcolor="{{element color|f-block}}" | No<br />16
| bgcolor="{{element color|d-block}}" | Lr<br />3
| bgcolor="{{element color|d-block}}" | Rf<br />4
| bgcolor="{{element color|d-block}}" | Db<br />5
| bgcolor="{{element color|d-block}}" | Sg<br />6
| bgcolor="{{element color|d-block}}" | Bh<br />7
| bgcolor="{{element color|d-block}}" | Hs<br />8
| bgcolor="{{element color|d-block}}" | Mt<br />9
| bgcolor="{{element color|d-block}}" | Ds<br />10
| bgcolor="{{element color|d-block}}" | Rg<br />11
| bgcolor="{{element color|d-block}}" | Cn<br />12
| bgcolor="{{element color|p-block}}" | Nh<br />3
| bgcolor="{{element color|p-block}}" | Fl<br />4
| bgcolor="{{element color|p-block}}" | Mc<br />5
| bgcolor="{{element color|p-block}}" | Lv<br />6
| bgcolor="{{element color|p-block}}" | Ts<br />7
| bgcolor="{{element color|p-block}}" | Og<br />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.<ref name="Greenwood113">Greenwood and Earnshaw, p. 113</ref> For example, magnesium forms Mg<sup>2+</sup> rather than Mg<sup>+</sup> cations when dissolved in water, because the latter would spontaneously [[disproportionation|disproportionate]] into Mg<sup>0</sup> and Mg<sup>2+</sup> 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<sup>+</sup>, 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 Mg<sup>2+</sup> 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<sup>+</sup> uncovers 6s, and Tl<sup>3+</sup> 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.<ref name=sb45>Siekierski and Burgess, pp. 45–54</ref>{{efn|The normally "forbidden" intermediate oxidation states may be stabilized by forming [[Dimer (chemistry)|dimers]], as in [Cl<sub>3</sub>Ga–GaCl<sub>3</sub>]<sup>2−</sup> (gallium in the +2 oxidation state) or [[disulfur decafluoride|S<sub>2</sub>F<sub>10</sub>]] (sulfur in the +5 oxidation state).<ref name=sb45/> Some compounds that appear to be in such intermediate oxidation states are actually mixed-valence compounds, such as [[antimony tetroxide|Sb<sub>2</sub>O<sub>4</sub>]], which contains both Sb(III) and Sb(V).<ref name="Amador">{{cite journal | last1 = Amador | first1 = J. | last2 = Puebla | first2 = E. Gutierrez | last3 = Monge | first3 = M. A. | last4 = Rasines | first4 = I. | last5 = Valero | first5 = C. Ruiz | year = 1988 | title = Diantimony Tetraoxides Revisited | journal = Inorganic Chemistry | volume = 27 | issue = 8 | pages = 1367–1370 | doi = 10.1021/ic00281a011 }}</ref>}}

[[File:Transition metal oxidation states.svg|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 d<sup>x+1</sup>s<sup>1</sup> or d<sup>x+2</sup>s<sup>0</sup> 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.<ref name=sb134>Siekierski and Burgess, pp. 134–137</ref> Because ionizing the transition metals further does not uncover any new inner subshells, their oxidation states tend to vary by steps of 1 instead.<ref name=sb45/> 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&nbsp;orbitals, that are similar in energy to ns.<ref name=sb178/> 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 ionize 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|OsO<sub>4</sub>]]), and then decrease to +2 at the end.<ref name=sb134/> The lanthanides and late actinides usually have high fourth ionisation energies and hence rarely surpass the +3 oxidation state, whereas early actinides have low fourth ionisation energies and so for example neptunium and plutonium can reach +7.<ref name=johnson/><ref name=sb134/><ref name=sb178>Siekierski and Burgess, pp. 178–180</ref> 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.<ref name=rareearths/>

As elements in the same group share the same valence configurations, they usually exhibit similar chemical behaviour. For example, the [[alkali metal]]s 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 heterogeneous. 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|group 14]] comprise a nonmetal ([[carbon]]), two semiconductors ([[silicon]] and [[germanium]]), and two metals ([[tin]] and [[lead]]); they are nonetheless united by having four valence electrons.<ref name="Scerri14">Scerri, pp. 14–15</ref> This often leads to similarities in maximum and minimum oxidation states (e.g. [[sulfur]] and [[selenium]] in [[chalcogen|group 16]] both have maximum oxidation state +6, as in [[sulfur trioxide|SO<sub>3</sub>]] and [[selenium trioxide|SeO<sub>3</sub>]], and minimum oxidation state −2, as in [[sulfide]]s and [[selenide]]s); but not always (e.g. [[oxygen]] is not known to form oxidation state +6, despite being in the same group as sulfur and selenium).<ref name=jensenlaw/>

=== Electronegativity ===
[[Image:Electrostatic Potential.jpg|thumb|alt=A water molecule is put into a see-through egg shape, which is colour-coded by electrostatic potential. A concentration of red is near the top of the shape, where the oxygen atom is, and gradually shifts through yellow, green, and then to blue near the lower-right and lower-left corners of the shape where the hydrogen atoms are.|upright=1.5|right|Electrostatic potential map of a water molecule, where the oxygen atom has a more negative charge (red) than the positive (blue) hydrogen atoms]]
Another important property of elements is their [[electronegativity]]. Atoms can form [[covalent bond]]s to each other by sharing electrons in pairs, creating an overlap of valence orbitals. The degree to which each atom attracts the shared electron pair depends on the atom's electronegativity<ref name="Greenwood25" /> – the tendency of an atom towards gaining or losing electrons.<ref name="cartoon" /> The more electronegative atom will tend to attract the electron pair more, and the less electronegative (or more electropositive) one will attract it less. In extreme cases, the electron can be thought of as having been passed completely from the more electropositive atom to the more electronegative one, though this is a simplification. The bond then binds two ions, one positive (having given up the electron) and one negative (having accepted it), and is termed an [[ionic bond]].<ref name="cartoon" />

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.<ref name="Greenwood25" />

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.<ref name="cartoon" /> 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.<ref>{{cite journal |doi=10.1021/ja00207a003 |title=Electronegativity is the average one-electron energy of the valence-shell electrons in ground-state free atoms|year=1989|author=Allen, Leland C.|journal=Journal of the American Chemical Society |volume=111|pages=9003–9014 |issue=25|bibcode=1989JAChS.111.9003A }}</ref>

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.<ref name="Greenwood25">Greenwood and Earnshaw, pp. 25–6</ref> 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.<ref>{{cite book |last1=Dieter |first1=R. K. |last2=Watson |first2=R. T. |chapter=Transmetalation reactions producing organocopper compounds |pages=443–526 |editor-last1=Rappoport |editor-first1=Z. |editor-last2=Marek |editor-first2=I. |title=The Chemistry of Organocopper Compounds |volume=1 |year=2009 |publisher=John Wiley & Sons |isbn=978-0-470-77296-6 |chapter-url=https://books.google.com/books?id=263AXB0Q6tAC |access-date=6 April 2022 |archive-date=17 October 2022 |archive-url=https://web.archive.org/web/20221017193845/https://books.google.com/books?id=263AXB0Q6tAC |url-status=live }}<!--specifically page 509--></ref>

=== Metallicity ===
[[File:Diamond cubic animation.gif|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.<br />(In grey tin, the band gap vanishes and metallization occurs.<ref>{{cite journal |last1=Carrasco |first1=Rigo A. |last2=Zamarripa |first2=Cesy M. |first3=Stefan |last3=Zollner |first4=José |last4=Menéndez |first5=Stephanie A. |last5=Chastang |first6=Jinsong |last6=Duan |first7=Gordon J. |last7=Grzybowski |first8=Bruce B. |last8=Claflin |first9=Arnold M. |last9=Kiefer |date=2018 |title=The direct bandgap of gray α-tin investigated by infrared ellipsometry |url=https://pubs.aip.org/aip/apl/article/113/23/232104/36404/The-direct-bandgap-of-gray-tin-investigated-by |journal=Applied Physics Letters |volume=113 |issue=23 |pages=232104 |doi=10.1063/1.5053884 |bibcode=2018ApPhL.113w2104C |s2cid=125130534 |access-date=|url-access=subscription }}</ref> 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, whose atoms bond in pairs) or giant structures stretching indefinitely (like carbon or silicon). The noble gases simply stay as single atoms, as they already have a full shell.<ref name="cartoon" /> Substances composed of discrete molecules or single atoms are held together by weaker attractive forces between the molecules, such as the [[London dispersion force]]: as electrons move within the molecules, they create momentary imbalances of electrical charge, which induce similar imbalances on nearby molecules and create synchronized movements of electrons across many neighbouring molecules.<ref>{{cite web|url=https://www.chemguide.co.uk/atoms/bonding/vdw.html|title=Intermolecular bonding – van der Waals forces|access-date=17 November 2021|archive-date=22 January 2022|archive-url=https://web.archive.org/web/20220122154740/https://www.chemguide.co.uk/atoms/bonding/vdw.html|url-status=live}}</ref>

[[File:Graphite-and-diamond-with-scale.jpg|thumb|right|Graphite and diamond, two allotropes of carbon]]
The more electropositive atoms, however, tend to instead lose electrons, creating a "sea" of electrons engulfing cations.<ref name="cartoon" /> 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.<ref name="chemguidemetal">{{cite web |url=https://www.chemguide.co.uk/atoms/bonding/metallic.html |title=Metallic Bonding |last=Clark |first=Jim |date=2019 |website=Chemguide |access-date=30 March 2021 |archive-date=21 April 2021 |archive-url=https://web.archive.org/web/20210421105423/https://www.chemguide.co.uk/atoms/bonding/metallic.html |url-status=live }}</ref> This negatively charged "sea" pulls on all the ions and keeps them together in a [[metallic bond]]. Elements forming such bonds are often called [[metal]]s; those which do not are often called [[Nonmetal (chemistry)|nonmetal]]s.<ref name="cartoon" /> Some elements can form multiple simple substances with different structures: these are called [[allotrope]]s. For example, [[diamond]] and [[graphite]] are two allotropes of carbon.<ref name="Scerri14" />{{efn|The boundary between dispersion forces and metallic bonding is gradual, like that between ionic and covalent bonding. Characteristic metallic properties do not appear in small mercury clusters, but do appear in large ones.<ref>{{cite journal |last1=Pastor |first1=G. M. |last2=Stampfli |first2=P. |last3=Bennemann |first3=K. |date=1988 |title=On the transition from Van der Waals- to metallic bonding in Hg-clusters as a function of cluster size |url= |journal=Physica Scripta |volume=38 |issue=4 |pages=623–626 |doi=10.1088/0031-8949/38/4/022 |bibcode=1988PhyS...38..623P |s2cid=250842014 }}</ref>}}

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 orbital]]s 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 H<sub>2</sub> molecule, and boron forms a giant covalent structure based on icosahedral B<sub>12</sub> 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.<ref name=Siekierski>Siekierski and Burgess, pp. 60–66</ref>

[[File:Solid state electronic band structure.svg|thumb|upright=2.0|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&nbsp;[[electronvolt|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 gap]]s and are [[semiconductor]]s 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.<ref name=Siekierski/> As the temperature rises, all nonmetals develop some semiconducting properties, to a greater or lesser extent depending on the size of the band gap. Thus metals and nonmetals may be distinguished by the temperature dependence of their electrical conductivity: a metal's conductivity lowers as temperature rises (because thermal motion makes it more difficult for the electrons to flow freely), whereas a nonmetal's conductivity rises (as more electrons may be excited to cross the gap).<ref name=steudel/>

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 (N<sub>2</sub>), oxygen (O<sub>2</sub>), white phosphorus and yellow arsenic (P<sub>4</sub> and As<sub>4</sub>), sulfur and red selenium (S<sub>8</sub> and Se<sub>8</sub>), and the stable halogens (F<sub>2</sub>, Cl<sub>2</sub>, Br<sub>2</sub>, and I<sub>2</sub>) 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, antimony, bismuth) that only extend in one or two rather than three dimensions. Both kinds of structures can be found as allotropes of phosphorus, arsenic, and selenium, although the long-chained allotropes are more stable in all three. 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, antimony, and bismuth are [[semimetal]]s (exhibiting quasi-metallic conduction, with a very small band overlap); and polonium and probably astatine are true metals.<ref name=Siekierski/> Finally, the natural group 18 elements all stay as individual atoms.<ref name=Siekierski/>{{efn|All this describes the situation at standard pressure. Under sufficiently high pressure, the band gaps of any solid drop to zero and metallization occurs. Thus for example at about 170&nbsp;[[bar (unit)|kbar]] iodine becomes a metal,<ref name=Siekierski/> and [[metallic hydrogen]] should form at pressures of about four million atmospheres.<ref>{{cite journal |last1=McMinis |first1=J. |last2=Clay |first2=R.C. |last3=Lee |first3=D. |last4=Morales |first4=M.A. |year=2015 |title=Molecular to Atomic Phase Transition in Hydrogen under High Pressure |journal=[[Physical Review Letters|Phys. Rev. Lett.]] |volume=114 |issue=10 |page=105305 |doi=10.1103/PhysRevLett.114.105305 |pmid=25815944 |bibcode=2015PhRvL.114j5305M|doi-access=free }}</ref> 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 [[metalloid]]s.<ref name="cartoon" /> 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.<ref>{{cite journal |last1=Hawkes |first1=Stephen J. |date=2001 |title=Semimetallicity? |journal=Journal of Chemical Education |volume=78 |issue=12 |page=1686 |doi=10.1021/ed078p1686|bibcode=2001JChEd..78.1686H }}</ref>

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 force]]s. 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,<ref name=Siekierski/> except perhaps for oganesson.<ref name="semiconductor">{{cite journal |last1=Mewes |first1=Jan-Michael |last2=Smits |first2=Odile Rosette |first3=Paul |last3=Jerabek |first4=Peter |last4=Schwerdtfeger |date=25 July 2019 |title=Oganesson is a Semiconductor: On the Relativistic Band-Gap Narrowing in the Heaviest Noble-Gas Solids |journal=Angewandte Chemie |volume=58 |issue=40 |pages=14260–14264|doi=10.1002/anie.201908327|pmid=31343819|pmc=6790653}}</ref>{{efn|Descriptions of the structures formed by the elements can be found throughout Greenwood and Earnshaw. There are two borderline cases. Arsenic's most stable form conducts electricity like a metal, but the bonding is significantly more localized to the nearest neighbours than it is for the similar structures of antimony and bismuth,<ref>{{cite book |last=Smith |first=J. D. |date=1973 |title=The Chemistry of Arsenic, Antimony and Bismuth |publisher=Pergamon Press |page=556 |isbn=}}</ref> and unlike normal metals it does not have a long liquid range, but rather sublimes instead. Hence its structure is better treated as network covalent.<ref>{{cite book |last1=Rayner-Canham |first1=Geoff |last2=Overton |first2=Tina |author-link= |date=2008 |title=Descriptive Inorganic Chemistry |edition=5th |url= |location=New York |publisher=W. H. Freeman and Company |page=194 |isbn=978-1-4292-2434-5}}</ref> Carbon as [[graphite]] shows metallic conduction parallel to its planes, but is a semiconductor perpendicular to them. Some computations predict copernicium and flerovium to be nonmetallic,<ref name=CRNL/><ref name=Florez/> but the most recent experiments on them suggest that they are metallic.<ref name=superheavy/><ref name=Ingo/><ref name=Yakushev/> Astatine is calculated to metallise at standard conditions,<ref name="Hermann">{{cite journal
|doi=10.1103/PhysRevLett.111.116404|title=Condensed Astatine: Monatomic and Metallic|year=2013|last1=Hermann|first1=A.|last2=Hoffmann|first2=R.|last3=Ashcroft|first3=N. W.|journal=Physical Review Letters|volume=111|issue=11|pages=116404-1–116404-5|bibcode=2013PhRvL.111k6404H|pmid=24074111}}</ref> so presumably tennessine should as well.<ref>{{cite news |last=Ball |first=Philip |date=13 September 2013 |title=
Metallic properties predicted for astatine |url=https://www.chemistryworld.com/news/metallic-properties-predicted-for-astatine/6582.article |work=Chemistry World |location= |access-date=7 April 2023}}</ref>}}

{{Periodic table (simple substance bonding)}}
<gallery mode="packed">
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.<ref name="cartoon" /> 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.<ref name="chemguidem">{{cite web |url=https://www.chemguide.co.uk/atoms/structures/metals.html |title=Metallic Structures |last=Clark |first=Jim |date=2012 |website=Chemguide |access-date=30 March 2021 |archive-date=24 April 2021 |archive-url=https://web.archive.org/web/20210424070514/https://www.chemguide.co.uk/atoms/structures/metals.html |url-status=live }}</ref> 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.<ref name="chemguidem" /> Metals tend to be reactive towards nonmetals.<ref name="cartoon" /> Some exceptions can be found to these generalizations: for example, beryllium, chromium,<ref name=raynercanham/> manganese,<ref name="Holl">{{cite book|publisher=Walter de Gruyter|date=1985|edition=91–100 |pages=1110–1117|isbn=978-3-11-007511-3|title=Lehrbuch der Anorganischen Chemie|first=Arnold F.|last=Holleman|author2=Wiberg, Egon|author3=Wiberg, Nils|language=de|chapter=Mangan}}</ref> antimony,<ref name="wiberg_holleman">{{cite book|title=Inorganic chemistry|author=Wiberg, Egon|author2=Wiberg, Nils|author3=Holleman, Arnold Frederick|name-list-style=amp|publisher=Academic Press|date=2001|isbn=978-0-12-352651-9|page=758}}</ref> bismuth,<ref name="CRC">{{cite book| first = C. R.| last = Hammond| pages = [https://archive.org/details/crchandbookofche81lide/page/4 4–1<!-- not a range -->]| title = The Elements, in Handbook of Chemistry and Physics| edition = 81st| location = Boca Raton (FL, US)| publisher = CRC press| isbn = 978-0-8493-0485-9| date = 2004| url-access = registration| url = https://archive.org/details/crchandbookofche81lide/page/4}}</ref> and uranium are brittle (not an exhaustive list);<ref name=raynercanham/> chromium is extremely hard;<ref name=r1>{{cite book|editor=G.V. Samsonov|chapter=Mechanical Properties of the Elements|doi=10.1007/978-1-4684-6066-7_7|isbn=978-1-4684-6066-7|url=https://ihtik.lib.ru/2011.08_ihtik_nauka-tehnika/2011.08_ihtik_nauka-tehnika_3560.rar|publisher=IFI-Plenum|place=New York, USA|year=1968|pages=387–446|archive-url=https://web.archive.org/web/20150402123344/https://ihtik.lib.ru/2011.08_ihtik_nauka-tehnika/2011.08_ihtik_nauka-tehnika_3560.rar|archive-date=2 April 2015 |title=Handbook of the Physicochemical Properties of the Elements }}</ref> gallium, rubidium, caesium, and mercury are liquid at or close to room temperature;{{efn|See [[melting points of the elements (data page)]]. The same is probably true of francium, but due to its extreme instability, this has never been experimentally confirmed. Copernicium and flerovium are expected to be liquids,<ref name="CRNL">{{cite journal |last1=Mewes |first1=J.-M. |last2=Smits |first2=O. R. |last3=Kresse |first3=G. |last4=Schwerdtfeger |first4=P. |title=Copernicium is a Relativistic Noble Liquid |journal=Angewandte Chemie International Edition |date=2019 |volume=58|issue=50|pages=17964–17968|doi=10.1002/anie.201906966 |pmid=31596013 |pmc=6916354 |url=}}</ref><ref name=Florez>{{cite journal |last1=Florez |first1=Edison |last2=Smits |first2=Odile R. |last3=Mewes |first3=Jan-Michael |last4=Jerabek |first4=Paul |last5=Schwerdtfeger |first5=Peter |date=2022 |title=From the gas phase to the solid state: The chemical bonding in the superheavy element flerovium |journal=The Journal of Chemical Physics |volume=157 |issue=6 |page=064304 |doi=10.1063/5.0097642|pmid=35963734 |bibcode=2022JChPh.157f4304F |s2cid=250539378 }}</ref> similar to mercury, and experimental evidence suggests that they are metals.<ref name="superheavy">
{{Cite web
|last1=Gäggeler
|first1=H. W.
|year=2007
|title=Gas Phase Chemistry of Superheavy Elements
|url=https://lch.web.psi.ch/files/lectures/TexasA&M/TexasA&M.pdf
|pages=26–28
|publisher=[[Paul Scherrer Institute]]
|archive-url=https://web.archive.org/web/20120220090755/https://lch.web.psi.ch/files/lectures/TexasA%26M/TexasA%26M.pdf
|archive-date=20 February 2012
}}</ref><ref name=Ingo>{{cite news |last=Ingo |first=Peter |date=15 September 2022 |title=Study shows flerovium is the most volatile metal in the periodic table |url=https://phys.org/news/2022-09-flerovium-volatile-metal-periodic-table.html |work=phys.org<!--but provided by GSI Helmholtz--> |location= |access-date=22 November 2022}}</ref><ref name=Yakushev>{{cite journal |last1=Yakushev |first1=A. |last2=Lens |first2=L. |first3=Ch. E. |last3=Düllmann |first4=J. |last4=Khuyagbaatar |first5=E. |last5=Jäger |first6=J. |last6=Krier |first7=J. |last7=Runke |first8=H. M. |last8=Albers |first9=M. |last9=Asai |first10=M. |last10=Block |first11=J. |last11=Despotopulos |first12=A. |last12=Di Nitto |first13=K. |last13=Eberhardt |first14=U. |last14=Forsberg |first15=P. |last15=Golubev |first16=M. |last16=Götz |first17=S. |last17=Götz |first18=H. |last18=Haba |first19=L. |last19=Harkness-Brennan |first20=R.-D. |last20=Herzberg |first21=F. P. |last21=Heßberger |first22=D. |last22=Hinde |first23=A. |last23=Hübner |first24=D. |last24=Judson |first25=B. |last25=Kindler |first26=Y. |last26=Komori |first27=J. |last27=Konki |first28=J. V. |last28=Kratz |first29=N. |last29=Kurz |first30=M. |last30=Laatiaoui |first31=S. |last31=Lahiri |first32=B. |last32=Lommel |first33=M. |last33=Maiti |first34=A. K. |last34=Mistry |first35=Ch. |last35=Mokry |first36=K. J. |last36=Moody |first37=Y. |last37=Nagame |first38=J. P. |last38=Omtvedt |first39=P. |last39=Papadakis |first40=V. |last40=Pershina |first41=D. |last41=Rudolph |first42=L. G. |last42=Samiento |first43=T. K. |last43=Sato |first44=M. |last44=Schädel |first45=P. |last45=Scharrer |first46=B. |last46=Schausten |first47=D. A. |last47=Shaughnessy |first48=J. |last48=Steiner |first49=P. |last49=Thörle-Pospiech |first50=A. |last50=Toyoshima |first51=N. |last51=Trautmann |first52=K. |last52=Tsukada |first53=J. |last53=Uusitalo |first54=K.-O. |last54=Voss |first55=A. |last55=Ward |first56=M. |last56=Wegrzecki |first57=N. |last57=Wiehl |first58=E. |last58=Williams |first59=V. |last59=Yakusheva |display-authors=3 |date=25 August 2022 |title=On the adsorption and reactivity of element 114, flerovium |journal=Frontiers in Chemistry |volume=10 |issue=976635 |page=976635 |doi=10.3389/fchem.2022.976635 |pmid=36092655 |pmc=9453156 |bibcode=2022FrCh...10.6635Y |doi-access=free }}</ref>}} and [[noble metal]]s such as gold are chemically very inert.<ref>{{cite journal |doi=10.1038/376238a0 |title=Why gold is the noblest of all the metals |date=1995 |last1=Hammer |first1=B. |last2=Norskov |first2=J. K. |journal=Nature |volume=376 |issue=6537 |pages=238–240 |bibcode=1995Natur.376..238H|s2cid=4334587 }}</ref><ref>{{cite journal |doi=10.1103/PhysRevB.6.4370 |title=Optical Constants of the Noble Metals |date=1972 |last1=Johnson |first1=P. B. |last2=Christy |first2=R. W. |journal=Physical Review B |volume=6 |issue=12 |pages=4370–4379 |bibcode=1972PhRvB...6.4370J}}</ref>

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,<ref>{{cite web |url=https://www.chemguide.co.uk/inorganic/period3/elementsphys.html |title=Atomic and Physical Properties of the Period 3 Elements |last=Clark |first=Jim |date=2018 |website=Chemguide |access-date=30 March 2021 |archive-date=22 April 2021 |archive-url=https://web.archive.org/web/20210422142013/https://www.chemguide.co.uk/inorganic/period3/elementsphys.html |url-status=live }}</ref> and many are liquids or gases at room temperature.<ref name="cartoon" /> Nonmetals are often dull-looking. They tend to be reactive towards metals, except for the noble gases, which are inert towards most substances.<ref name="cartoon" /> They are brittle when solid as their atoms are held tightly in place. They are less dense and conduct electricity poorly,<ref name="cartoon" /> because there are no mobile electrons.<ref name="group4">{{cite web |url=https://www.chemguide.co.uk/inorganic/group4/properties.html |title=The Trend From Non-Metal to Metal In the Group 4 Elements |last=Clark |first=Jim |date=2015 |website=Chemguide |access-date=30 March 2021 |archive-date=27 April 2021 |archive-url=https://web.archive.org/web/20210427234147/https://www.chemguide.co.uk/inorganic/group4/properties.html |url-status=live }}</ref> Near the borderline, band gaps are small and thus many elements in that region are semiconductors, such as silicon, germanium,<ref name="group4" /> and tellurium.<ref name=Siekierski/> Selenium has both a semiconducting grey allotrope and an insulating red allotrope; arsenic has a metallic grey allotrope, a semiconducting black allotrope, and an insulating yellow allotrope (though the last is unstable at ambient conditions).<ref name=steudel/> Again there are exceptions; for example, diamond has the highest thermal conductivity of all known materials, greater than any metal.<ref name=PNU>{{cite journal |doi=10.1103/PhysRevLett.70.3764 |title=Thermal conductivity of isotopically modified single crystal diamond |year=1993 |last1=Wei |first1=Lanhua |last2=Kuo |first2=P. K. |last3=Thomas |first3=R. L. |last4=Anthony |first4=T. R. |last5=Banholzer |first5=W. F. |journal=Physical Review Letters |volume=70 |issue=24 |pages=3764–3767 |pmid=10053956 |bibcode=1993PhRvL..70.3764W}}</ref>

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.<ref name="cartoon" /> 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.{{efn|See [[lists of metalloids]]. For example, a periodic table used by the American Chemical Society includes polonium as a metalloid,<ref name="ACS" /> but one used by the Royal Society of Chemistry does not,<ref>{{cite web |url=https://www.rsc.org/periodic-table |title=Periodic Table |date=2021 |website=www.rsc.org |publisher=[[Royal Society of Chemistry]] |access-date=27 March 2021 |archive-date=21 March 2021 |archive-url=https://web.archive.org/web/20210321033913/https://www.rsc.org/periodic-table |url-status=live }}</ref> and that included in the ''[[Encyclopædia Britannica]]'' does not refer to metalloids or semi-metals at all.<ref name="EB" /> Classification can change even within a single work. For example, Sherwin and Weston's ''Chemistry of the Non-Metallic Elements'' (1966) has a periodic table on p. 7 classifying antimony as a nonmetal, but on p. 115 it is called a metal.<ref>{{cite book |last1=Sherwin |first1=E. |last2=Weston |first2=G. J. |editor=Spice, J. E. |date=1966 |title=Chemistry of the Non-Metallic Elements |publisher=Pergamon Press |isbn=978-1-4831-3905-0}}</ref>|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.<ref name=raynercanham/><ref name=hawkes>{{cite journal |last1=Hawkes |first1=Stephen J. |date=2001 |title=Semimetallicity? |url= |journal=Journal of Chemical Education |volume=78 |issue=12 |pages=1686–1687 |doi=10.1021/ed078p1686 |bibcode=2001JChEd..78.1686H |access-date=}}</ref><ref name=steudel>{{cite book |last1=Steudel |first1=Ralf |first2=David |last2=Scheschkewitz |author-link= |date=2020 |title=Chemistry of the Non-Metals |url= |location= |publisher=Walter de Gruyter |pages=154–155, 425, 436 |isbn=978-3-11-057805-8 |quote=In Group 15 of the Periodic Table, as in both neighboring groups, the metallic character increases when going down. More specifically, there is a transition from a purely non-metallic element (N) via elements with nonmetallic and metallic modifications to purely metallic elements (Sb, Bi). This chapter addresses the two elements besides nitrogen, which are clearly nonmetallic under standard conditions: phosphorus and arsenic. The chemistry of arsenic, however, is only briefly described as many of the arsenic compounds resemble the corresponding phosphorus species.}}</ref> 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<ref name=hawkes/> – though this situation also holds for phosphorus,<ref name=steudel/> which is a much rarer inclusion among the metalloids.{{efn|name=metalloids}}

=== Further manifestations of periodicity ===
There are some other relationships throughout the periodic table between elements that are not in the same group, such as the [[diagonal relationship]]s between elements that are diagonally adjacent (e.g. lithium and magnesium).<ref name="PTSS2">Scerri, pp. 407–420</ref> 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,<ref name="PTSS2" /> as all three have six valence electrons.<ref name="Jensen" /> Relationships between elements with the same number of valence electrons but different types of valence orbital have been called secondary or isodonor relationships: they usually have the same maximum oxidation states, but not the same minimum oxidation states. For example, chlorine and manganese both have +7 as their maximum oxidation state (e.g. [[dichlorine heptoxide|Cl<sub>2</sub>O<sub>7</sub>]] and [[manganese heptoxide|Mn<sub>2</sub>O<sub>7</sub>]]), but their respective minimum oxidation states are −1 (e.g. [[hydrogen chloride|HCl]]) and −3 (K<sub>2</sub>[Mn(CO)<sub>4</sub>]). Elements with the same number of valence vacancies but different numbers of valence electrons are related by a tertiary or isoacceptor relationship: they usually have similar minimum but not maximum oxidation states. For example, hydrogen and chlorine both have −1 as their minimum oxidation state (in [[hydride]]s and [[chloride]]s), but hydrogen's maximum oxidation state is +1 (e.g. [[water|H<sub>2</sub>O]]) while chlorine's is +7.<ref name=jensenlaw/>

Many other physical properties of the elements exhibit periodic variation in accordance with the periodic law, such as [[melting point]]s, [[boiling point]]s, [[heat of fusion|heats of fusion]], [[heat of vaporisation|heats of vaporization]], [[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.<ref name="Greenwood25" /> Chemical properties are more difficult to describe quantitatively, but likewise exhibit their own periodicities. Examples include the variation in the [[acid]]ic and [[base (chemistry)|basic]] properties of the elements and their compounds, the stabilities of compounds, and methods of isolating the elements.<ref name="Greenwood27" /> 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.<ref name="Greenwood29bis">Greenwood and Earnshaw, pp. 29–31</ref>