Editing Nonmetal (section)

==General properties==
===Physical===
{{hatnote|See also {{slink||Physical properties by element type}}}}

{{multiple image|perrow=3|total_width=330|caption_align=center
| align = right
| image_style = border:none;
|image1=Boron R105.jpg
 |alt1=Several dozen small angular stone like shapes, grey with scattered silver flecks and highlights.
 |caption1= Boron in its β-[[rhombohedral]] phase
|image2=Graphite2.jpg
 |alt2=A shiny grey-black cuboid nugget with a rough surface.
 |caption2= Metallic appearance of [[carbon]] as [[graphite]]
|image3=Liquid oxygen in a beaker 4.jpg
 |alt3=A pale blue liquid in a clear beaker
 |caption3= Blue color of [[liquid oxygen]]
|image4=Liquid fluorine tighter crop.jpg
 |alt4=A glass tube, is inside a larger glass tube, has some clear yellow liquid in it
 |caption4= Pale yellow liquid fluorine in a [[cryogenics|cryogenic bath]]
|image5=Sulfur-sample.jpg
 |alt5=Yellow powdery chunks
 |caption5= [[Sulfur]] as yellow chunks
|image6=Bromine_in_a_vial.png
 |alt6=A small capped jar a quarter filled with a very dark liquid
 |caption6= Liquid [[bromine]] at room temperature
|image7=Iodinecrystals.JPG
 |caption7= Metallic appearance of [[iodine]] under white light
 |alt7=Shiny violet-black colored crystalline shards.
|image8=An acrylic cube specially prepared for element collectors containing an ampoule filled with liquefied xenon.JPG
 |alt8=A partly filled ampoule containing a colorless liquid
 |caption8= Liquefied xenon
| header = {{font|size=100%|font=Sans-serif|text=Variety in color and form<br>of some nonmetallic elements}}
}}

Nonmetals vary greatly in appearance, being colorless, colored or shiny.<!-- It would be nice to say "due to the structure of their electrons" or something like that, as is mentioned w/r/t shiny solids and colorless gases, but nothing is stated about the electronics of the colored ones -->
For the colorless nonmetals (hydrogen, nitrogen, oxygen, and the noble gases), no absorption of light happens in the visible part of the spectrum, and all visible light is transmitted.<ref>[[#Wibaut|Wibaut 1951, p. 33]]: "Many substances&nbsp;...are colourless and therefore show no selective absorption in the visible part of the spectrum."</ref>
The colored nonmetals (sulfur, fluorine, chlorine, bromine) absorb some colors (wavelengths) and transmit the complementary or opposite colors. For example, chlorine's "familiar yellow-green colour&nbsp;... is due to a broad region of absorption in the violet and blue regions of the spectrum".<ref>[[#Elliot|Elliot 1929, p. 629]]</ref>{{efn|The absorbed light may be converted to heat or re-emitted in all directions so that the emission spectrum is thousands of times weaker than the incident light radiation.<ref>[[#Fox|Fox 2010, p. 31]]</ref>}} The shininess of boron, graphite (carbon), silicon, black phosphorus, germanium, arsenic, selenium, antimony, tellurium, and iodine{{efn|Solid iodine has a silvery metallic appearance under white light at room temperature. At ordinary and higher temperatures it [[sublimation (phase transition)|sublimes]] from the solid phase directly into a violet-colored vapor.<ref>[[#Tidy|Tidy 1887, pp. 107–108]]; [[#Koenig|Koenig 1962, p. 108]]</ref>}} is a result of the electrons reflecting incoming visible light.<ref>[[#Wiberg|Wiberg 2001, p. 416]]; Wiberg is here referring to iodine.</ref>

About half of nonmetallic elements are gases under [[standard temperature and pressure]]; most of the rest are solids. Bromine, the only liquid, is usually topped by a layer of its reddish-brown fumes. The gaseous and liquid nonmetals have very low densities, [[melting point|melting]] and [[boiling point]]s, and are poor conductors of heat and electricity.<ref name="Kneen">[[#Kneen|Kneen, Rogers & Simpson 1972, pp. 261–264]]</ref> The solid nonmetals have low densities and low mechanical strength (being either hard and brittle, or soft and crumbly),<ref name="ReferenceA">[[#Johnson1966|Johnson 1966, p. 4]]</ref> and a wide range of electrical conductivity.{{efn|The solid nonmetals have electrical conductivity values ranging from 10<sup>−18</sup> S•cm<sup>−1</sup> for sulfur<ref name="A&W"/> to 3 × 10<sup>4</sup> in graphite<ref name="Jenkins">[[#Jenkins|Jenkins & Kawamura 1976, p. 88]]</ref> or 3.9 × 10<sup>4</sup> for [[arsenic]];<ref>[[#Carapella|Carapella 1968, p. 30]]</ref> cf. 0.69 × 10<sup>4</sup> for [[manganese]] to 63 × 10<sup>4</sup> for [[silver]], both metals.<ref name="A&W">[[#Aylward|Aylward & Findlay 2008, pp. 6–12]]</ref> The conductivity of graphite and arsenic (both semimetals) exceed that of manganese.}}

This diversity stems from variability in crystallographic structures and bonding arrangements. Covalent nonmetals existing as discrete atoms like xenon, or as small molecules, such as oxygen, sulfur, and bromine, have low melting and boiling points; many are gases at room temperature, as they are held together by weak [[London dispersion force]]s acting between their atoms or molecules, although the molecules themselves have strong covalent bonds.<ref>[[#ZumDeC|Zumdahl & DeCoste 2010, pp. 455, 456, 469, A40]]; [[#Earl&W|Earl & Wilford 2021, p. 3-24]]</ref> In contrast, nonmetals that form extended structures, such as long chains of selenium atoms,<ref>{{Cite journal |last1=Corb |first1=B.W. |last2=Wei |first2=W.D. |last3=Averbach |first3=B.L. |date=1982 |title=Atomic models of amorphous selenium |url=https://linkinghub.elsevier.com/retrieve/pii/0022309382900163 |journal=Journal of Non-Crystalline Solids |language=en |volume=53 |issue=1–2 |pages=29–42 |doi=10.1016/0022-3093(82)90016-3|bibcode=1982JNCS...53...29C }}</ref> sheets of carbon atoms in graphite,<ref>[[#Wiberg|Wiberg 2001, pp. 780]]</ref> or three-dimensional lattices of silicon atoms<ref>[[#Wiberg|Wiberg 2001, pp. 824, 785]]</ref> have higher melting and boiling points, and are all solids. Nonmetals closer to the left or bottom of the periodic table (and so closer to the metals) often have [[Metallic bonding|metallic interactions]] between their molecules, chains, or layers; this occurs in boron,<ref>[[#Siekierski|Siekierski & Burgess 2002, p. 86]]</ref> carbon,<ref>[[#Charlier|Charlier, Gonze & Michenaud 1994]]</ref> phosphorus,<ref>[[#Taniguchi|Taniguchi et al. 1984, p. 867]]: "...&nbsp;black phosphorus&nbsp;... [is] characterized by the wide valence bands with rather delocalized nature."; [[#Carmalt|Carmalt & Norman 1998, p. 7]]: "Phosphorus&nbsp;... should therefore be expected to have some metalloid properties."; [[#Du|Du et al. 2010]]: Interlayer interactions in black phosphorus, which are attributed to van der Waals-Keesom forces, are thought to contribute to the smaller band gap of the bulk material (calculated 0.19&nbsp;eV; observed 0.3&nbsp;eV) as opposed to the larger band gap of a single layer (calculated ~0.75&nbsp;eV).</ref> arsenic,<ref>[[#Wiberg|Wiberg 2001, pp. 742]]</ref> selenium,<ref>[[#Evans|Evans 1966, pp. 124–25]]</ref> antimony,<ref>[[#Wiberg|Wiberg 2001, pp. 758]]</ref> tellurium<ref>[[#Stuke|Stuke 1974, p. 178]]; [[#Donohue|Donohue 1982, pp. 386–87]]; [[#Cotton|Cotton et al. 1999, p. 501]]</ref> and iodine.<ref>[[#Steudel|Steudel 2020, p. 601]]: "...&nbsp;Considerable orbital overlap can be expected. Apparently, intermolecular multicenter bonds exist in crystalline iodine that extend throughout the layer and lead to the delocalization of electrons akin to that in metals. This explains certain physical properties of iodine: the dark color, the luster and a weak electric conductivity, which is 3400 times stronger within the layers then perpendicular to them. Crystalline iodine is thus a two-dimensional semiconductor."; [[#Segal|Segal 1989, p. 481]]: "Iodine exhibits some metallic properties&nbsp;..."</ref>

{|class="wikitable floatright" style="line-height: 1.3; font-size: 95%; margin-left:20px; margin-bottom:1.2em"
|+ Some general physical differences<br />between elemental metals and nonmetals<ref name="Kneen"/>
|-
! Aspect !! Metals !! Nonmetals
|-
|Appearance<br>and form
|Shiny if freshly prepared<br>or fractured; few colored;<ref>[[#Taylor|Taylor 1960, p. 207]]; [[#Brannt|Brannt 1919, p. 34]]</ref><br>all but one solid<ref name="Green">[[#Green|Green 2012, p. 14]]</ref>
|Shiny, colored or<br>transparent;<ref>[[#Spencer|Spencer, Bodner & Rickard 2012, p. 178]]</ref> all but<br>one solid or gaseous<ref name="Green"/>
|-
|[[Density]]
| Often higher
| Often lower
|-
| [[Plasticity (physics)|Plasticity]]
| Mostly malleable<br />and ductile
| Often brittle solids
|-
| [[Electrical conductivity|Electrical]]<br>[[Electrical conductivity|conductivity]]<ref>[[#Redmer|Redmer, Hensel & Holst 2010, preface]]</ref>
| Good
| Poor to good
|-
| [[Electronic band structure|Electronic]]<br>[[Electronic band structure|structure]]<ref name="K&W"/>
| Metal or [[semimetal]]ic
| Semimetal,<br>[[semiconductor]],<br>or [[insulator (electricity)|insulator]]
|}
Covalently bonded nonmetals often share only the electrons required to achieve a noble gas electron configuration.<ref>[[#DeKock|DeKock & Gray 1989, pp. 423, 426—427]]</ref> For example, nitrogen forms diatomic molecules featuring a triple bonds between each atom, both of which thereby attain the configuration of the noble gas neon. In contrast antimony has buckled layers in which each antimony atom is singly bonded with three other nearby atoms.<ref>[[#Boreskov|Boreskov 2003, p. 45]]</ref>

Good electrical conductivity occurs when there is [[metallic bond]]ing,<ref name="Ashcroft and Mermin">[[Ashcroft and Mermin]]</ref> however the electrons in some nonmetals are not metallic.<ref name="Ashcroft and Mermin"/> Good electrical and thermal conductivity associated with metallic electrons is seen in carbon (as graphite, along its planes), arsenic, and antimony.{{efn|Thermal conductivity values for metals range from 6.3 W m<sup>−1</sup> K<sup>−1</sup> for [[neptunium]] to 429 for [[silver]]; cf. antimony 24.3, arsenic 50, and carbon 2000.<ref name="A&W"/> Electrical conductivity values of metals range from 0.69 S•cm<sup>−1</sup> × 10<sup>4</sup> for [[manganese]] to 63 × 10<sup>4</sup> for [[silver]]; cf. carbon 3 × 10<sup>4</sup>,<ref name="Jenkins"/> arsenic 3.9 × 10<sup>4</sup> and antimony 2.3 × 10<sup>4</sup>.<ref name="A&W"/>}} Good thermal conductivity occurs in boron, silicon, phosphorus, and germanium;<ref name="A&W"/> such conductivity is transmitted though vibrations of the crystalline lattices ([[phonons]] of these elements.<ref>[[#Yang|Yang 2004, p. 9]]</ref> Moderate electrical conductivity is observed in the semiconductors<ref>[[#Wiberg|Wiberg 2001, pp. 416, 574, 681, 824, 895, 930]]; [[#Siekierski|Siekierski & Burgess 2002, p. 129]]</ref> boron, silicon, phosphorus, germanium, selenium, tellurium, and iodine.

Many of the nonmetallic elements are hard and brittle,<ref name="ReferenceA"/> where [[dislocations]] cannot readily move so they tend to undergo [[brittle fracture]] rather than deforming.<ref>{{Cite book |last1=Weertman |first1=Johannes |title=Elementary dislocation theory |last2=Weertman |first2=Julia R. |date=1992 |publisher=Oxford University Press |isbn=978-0-19-506900-6 |location=New York}}</ref> Some do deform such as [[allotropes of phosphorus#White phosphorus|white phosphorus]] (soft as wax, pliable and can be cut with a knife at room temperature),<ref name="Holderness 1979, p. 255">[[#Faraday|Faraday 1853, p. 42]]; [[#Holderness|Holderness & Berry 1979, p. 255]]</ref> [[allotropes of sulfur#Amorphous sulfur|plastic sulfur]],<ref name="ReferenceE">[[#Partington1944|Partington 1944, p. 405]]</ref> and selenium which can be drawn into wires from its molten state.<ref name="ReferenceF">[[#Regnault|Regnault 1853, p. 208]]</ref> Graphite is a standard [[solid lubricant]] where dislocations move very easily in the basal planes.<ref name=":1">{{Cite journal |last1=Scharf |first1=T. W. |last2=Prasad |first2=S. V. |date=January 2013 |title=Solid lubricants: a review |url=http://link.springer.com/10.1007/s10853-012-7038-2 |journal=Journal of Materials Science |language=en |volume=48 |issue=2 |pages=511–531 |doi=10.1007/s10853-012-7038-2 |bibcode=2013JMatS..48..511S |issn=0022-2461}}</ref>

====Allotropes====
{{multiple image
| perrow            = 3
| total_width       = 350
| align             = right
| image_style       = border:none;
| image1            = Diamond-dimd15a.jpg
| alt1              = A clear triangular crystal with a flat face and slightly rough edges
| caption1          = A transparent electrical{{nbsp}}insulator
| image2            = C60-Fulleren-kristallin (cropped).JPG
| alt2              = a haphazard aggregate of brownish crystals
| caption2          = A brownish semiconductor
| image3            = Graphite-233436.jpg
| alt3              = A black multi-layered lozenge-shaped rock
| caption3          = A blackish semimetal
| header            = Three allotropes of carbon
| footer            = From left to right, [[diamond]], [[buckminsterfullerene]], and [[graphite]]
| caption_align     = center
| footer_align      = center
}}
{{Main list|Allotropy#Non-metals|Single-layer materials}}
Over half of the nonmetallic elements exhibit a range of less stable allotropic forms, each with distinct physical properties.<ref>[[#Barton|Barton 2021, p. 200]]</ref> For example, carbon, the most stable form of which is [[graphite]], can manifest as [[diamond]], [[buckminsterfullerene]],<ref>[[#Wiberg|Wiberg 2001, p. 796]]</ref> [[amorphous]]<ref>[[#Shang|Shang et al. 2021]]</ref> and [[Paracrystallinity|paracrystalline]]<ref>[[#Tang|Tang et al. 2021]]</ref> variations. Allotropes also occur for nitrogen, oxygen, phosphorus, sulfur, selenium and iodine.<ref>[[#Steudel|Steudel 2020, ''passim'']]; [[#Carrasco|Carrasco et al. 2023]]; [[#Shanabrook|Shanabrook, Lannin & Hisatsune 1981, pp. 130–133]]</ref>

===Chemical===
{{hatnote|See also {{slink||Chemical properties by element type}}}}
{|class="wikitable floatright" style="line-height: 1.3; font-size: 95%; margin-left:20px"
|+ Some general chemistry-based<br/>differences between metals and nonmetals<ref name="Kneen"/>
|-
! colspan=2 | Aspect !! Metals !! Nonmetals
|-
| colspan=2 |Reactivity<ref>[[#Weller|Weller et al. 2018, preface]]</ref>
| colspan=2 style="text-align: center"| Wide range: very reactive to noble
|-
| rowspan =2 | [[Oxide]]s || lower
| [[base (chemistry)|Basic]]
| rowspan =2 | [[Acid]]ic; never basic<ref name="ReferenceC">[[#Abbott|Abbott 1966, p. 18]]</ref>
|-
| higher || Increasingly acidic
|-
| colspan=2 |Compounds<br>with metals<ref>[[#Ganguly|Ganguly 2012, p. 1-1]]</ref>
| [[Alloy]]s
| [[Covalent bond|Covalent]] or [[Ionic compounds|Ionic]]
|-
| colspan=2 | [[Ionization energy]]<ref name="AylwardIE"/>
| Low to high
| Moderate to very high
|-
| colspan=2 | [[Electronegativity]]<ref name="AylwardEN"/>
| Low to high
| Moderate to very high
|}
Nonmetals have relatively high values of electronegativity, and their oxides are usually acidic. Exceptions may occur if a nonmetal is not very electronegative, or if its [[oxidation state]] is low, or both. These non-acidic oxides of nonmetals may be [[amphoteric]] (like water, H<sub>2</sub>O<ref>[[#Eagleson1994|Eagleson 1994, 1169]]</ref>) or neutral (like [[nitrous oxide]], N<sub>2</sub>O<ref>[[#Moody|Moody 1991, p. 365]]</ref>{{efn|While [[carbon monoxide|CO]] and [[nitrogen monoxide|NO]] are commonly referred to as being neutral, CO is a slightly acidic oxide, reacting with bases to produce formates (CO + OH<sup>−</sup> → HCOO<sup>−</sup>);<ref>[[#House2013|House 2013, p. 427]]</ref> and in water, NO reacts with oxygen to form nitrous acid HNO<sub>2</sub> (4NO + O<sub>2</sub> + 2H<sub>2</sub>O → 4HNO<sub>2</sub>).<ref>[[#LewisRS|Lewis & Deen 1994, p. 568]]</ref>}}), but never basic.

They tend to gain electrons during chemical reactions, in contrast to metallic elements which tend to donate electrons. This behavior is related to the stability of [[electron configuration]]s in the noble gases, which have complete outer [[electron shell|shells]], empirically described by the [[Octet rule#Other rules|duet]] and [[octet rule]]s of thumb, more correctly explained in terms of [[valence bond theory]].<ref>[[#SmithDW|Smith 1990, pp. 177–189]]</ref>

The chemical differences between metals and nonmetals stem from variations in how strongly atoms attract and retain electrons. Across a period of the periodic table, the nuclear charge increases as more protons are added to the nucleus.<ref>[[#Young2018|Young et al. 2018, p. 753]]</ref> However, because the number of inner electron shells remains constant, the [[effective nuclear charge]] experienced by the outermost electrons also increases, pulling them closer to the nucleus. This leads to a corresponding reduction in atomic radius,<ref>[[#Brown et al.|Brown et al. 2014, p. 227]]</ref> and a greater tendency of these elements to gain electrons during chemical reactions, forming negatively charged ions.<ref>[[#Moore|Moore 2016]]; [[#Burford|Burford, Passmore & Sanders 1989, p. 54]]</ref> Nonmetals, which occupy the right-hand side of the periodic table, exemplify this behavior.

Nonmetals typically exhibit higher [[ionization energy|ionization energies]], [[electron affinity|electron affinities]], and [[standard electrode potential]]s than metals. The higher these values are (including electronegativity) the more nonmetallic the element tends to be.<ref>[[#Yoder|Yoder, Suydam & Snavely 1975, p. 58]]</ref> For example, the chemically very active nonmetals fluorine, chlorine, bromine, and iodine have an average electronegativity of 3.19—a figure{{efn|Electronegativity values of fluorine to iodine are: 3.98 + 3.16 + 2.96 + 2.66 &#61; 12.76/4 3.19.}} higher than that of any metallic element.

The number of compounds formed by nonmetals is vast.<ref>[[#Brady|Brady & Senese 2009, p. 69]]</ref> The first 10 places in a "top 20" table of elements most frequently encountered in 895,501,834 compounds, as listed in the [[Chemical Abstracts Service]] register for November 2, 2021, were occupied by nonmetals. Hydrogen, carbon, oxygen, and nitrogen collectively appeared in most (80%) of compounds. Silicon, a metalloid, ranked 11th. The highest-rated metal, with an occurrence frequency of 0.14%, was iron, in 12th place.<ref>[[#CAS|Chemical Abstracts Service 2021]]</ref>

===Complications===

Adding complexity to the chemistry of the nonmetals are anomalies occurring in the first row of each [[periodic table block]]; non-uniform periodic trends; higher oxidation states; multiple bond formation; and property overlaps with metals.

====First row anomaly====
<div style="line-height:1px;">[[Image:1x1.png|link=|alt=A table with seven rows and ten columns. Rows are labeled on the left with a period number from 1 through 7. Columns are labeled on the bottom with a group number. Most cells represent a single chemical element and have two lines of information: the element's symbol on the top and its atomic number on the bottom. The table as a whole is divided into four rectangular areas separated from each other by narrow gaps. The first rectangle fills all seven rows of the first two columns. The rectangle is labeled "s-block" at the top and its two columns are labeled with group numbers "(1)" and "(2)" on the bottom. The cells in the first row - hydrogen and helium, with symbols H and He and atomic numbers 1 and 2 respectively - are both shaded red. The second rectangle fills the bottom two rows (periods 6 and 7) of the third column. Just above these cells is the label "f-block"; there is no group label on the bottom. The topmost cell - labeled "La-Yb" for elements 57-70 - is shaded green. The third rectangle fills the bottom four rows (periods 4 through 7) of the fourth column. Just above these cells is the label "d-block"; at the bottom is the label "(3-12)" for the group numbers of these elements. The topmost cell - labeled "Sc-Zn" for elements 21-30 - is shaded blue. The fourth and last rectangle fills the bottom six rows (periods 2 through 7) of the last six columns. Just above these cells is the label "p-block"; at the bottom are labels "(13)" through "(18) for the group numbers of these elements. The cells in the topmost row - for the elements boron (B,5), carbon (C,6), nitrogen (N,7), oxygen (O,8), fluorine (Fl,9), and neon (Ne,10) - are shaded yellow. Bold lines encircle the cells of the nonmetals - the top two cells on the left and 21 cells in the upper right of the table.]]</div>
{| class=" floatright" style="border-collapse:collapse; text-align:center;font-size:80%;line-height:1.1;margin-top:1.2em;"
| colspan=14 style="padding-bottom:3px;border:none;text-align:center;font-size:105%" | '''Condensed periodic table highlighting<br>the first row of each block: {{color box|{{element color|s-block}}|s}} {{color box|{{element color|p-block}}|p}} {{color box|{{element color|d-block}}|d}} and {{color box|{{element color|f-block}}|f}}'''
|-
| colspan=1 | '''Period'''
| colspan=2 | '''{{nowrap|s-block}}'''
| rowspan=9 style="padding:1px;" |
| colspan=1 |
| rowspan=9 style="padding:1px;" |
| colspan=1 |
| rowspan=9 style="padding:1px;" |
| colspan=6 |
|-
| '''1'''
| style="border:solid black;border-width:2px 1px 2px 2px;background-color:{{element color|s-block}};" | H <br>1
| style="border:solid black;border-width:2px 2px 2px 1px;background-color:{{element color|s-block}};" | He<br>2
| 
| 
| colspan=6 | <br>'''p-block'''
|-
| '''2'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Li<br>3
| style="border:solid black;border-width:1px 1px 1px 1px;" | Be<br>4
|
|
| style="border:solid black;border-width:2px 1px 2px 2px;background-color:{{element color|p-block}};" | B <br>5
| style="border:solid black;border-width:2px 1px 1px 1px;background-color:{{element color|p-block}};" | C <br>6
| style="border:solid black;border-width:2px 1px 1px 1px;background-color:{{element color|p-block}};" | N <br>7
| style="border:solid black;border-width:2px 1px 1px 1px;background-color:{{element color|p-block}};" | O <br>8
| style="border:solid black;border-width:2px 1px 1px 1px;background-color:{{element color|p-block}};" | F <br>9
| style="border:solid black;border-width:2px 2px 1px 1px;background-color:{{element color|p-block}};" | Ne<br>10
|-
| '''3'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Na<br>11
| style="border:solid black;border-width:1px 1px 1px 1px;" | Mg<br>12
|
| <br>'''{{nowrap|d-block}}'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Al<br>13
| style="border:solid black;border-width:1px 1px 1px 2px;" | Si<br>14
| style="border:solid black;border-width:1px 1px 1px 1px;" | P <br>15
| style="border:solid black;border-width:1px 1px 1px 1px;" | S <br>16
| style="border:solid black;border-width:1px 1px 1px 1px;" | Cl<br>17
| style="border:solid black;border-width:1px 2px 1px 1px;" | Ar<br>18
|-
| '''4'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | K <br>19
| style="border:solid black;border-width:1px 1px 1px 1px;" | Ca<br>20
|
| style="border:solid black;border-width:1px 1px 1px 1px;background-color:{{element color|d-block}};" | Sc-Zn<br>21-30
| style="border:solid black;border-width:1px 1px 1px 1px;" | Ga<br>31
| style="border:solid black;border-width:1px 1px 2px 2px;" | Ge<br>32
| style="border:solid black;border-width:1px 1px 1px 1px;" | As<br>33
| style="border:solid black;border-width:1px 1px 1px 1px;" | Se<br>34
| style="border:solid black;border-width:1px 1px 1px 1px;" | Br<br>35
| style="border:solid black;border-width:1px 2px 1px 1px;" | Kr<br>36
|-
| '''5'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Rb<br>37
| style="border:solid black;border-width:1px 1px 1px 1px;" | Sr<br>38
| <br>'''{{nowrap|f-block}}'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Y-Cd<br>39-48
| style="border:solid black;border-width:1px 1px 1px 1px;" | In<br>49
| style="border:solid black;border-width:1px 1px 1px 1px;" | Sn<br>50
| style="border:solid black;border-width:1px 1px 2px 2px;" | Sb<br>51
| style="border:solid black;border-width:1px 1px 2px 1px;" | Te<br>52
| style="border:solid black;border-width:1px 1px 2px 1px;" | I <br>53
| style="border:solid black;border-width:1px 2px 1px 1px;" | Xe<br>54
|-
| '''6'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Cs<br>55
| style="border:solid black;border-width:1px 1px 1px 1px;" | Ba<br>56
| style="border:solid black;border-width:1px 1px 1px 1px;background-color:{{element color|f-block}};" | La-Yb<br>57-70
| style="border:solid black;border-width:1px 1px 1px 1px;" | ''Lu-Hg<br>71-80''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Tl<br>81
| style="border:solid black;border-width:1px 1px 1px 1px;" | Pb<br>82
| style="border:solid black;border-width:1px 1px 1px 1px;" | Bi<br>83
| style="border:solid black;border-width:1px 1px 1px 1px;" | Po<br>84
| style="border:solid black;border-width:1px 1px 1px 1px;" | At<br>85
| style="border:solid black;border-width:1px 2px 2px 2px;" | Rn<br>86
|-
| '''7'''
| style="border:solid black;border-width:1px 1px 1px 1px;" | Fr<br>87
| style="border:solid black;border-width:1px 1px 1px 1px;" | Ra<br>88
| style="border:solid black;border-width:1px 1px 1px 1px;" | Ac-No<br>89-102
| style="border:solid black;border-width:1px 1px 1px 1px;" | Lr-Cn<br>103-112
| style="border:solid black;border-width:1px 1px 1px 1px;" | Nh<br>113
| style="border:solid black;border-width:1px 1px 1px 1px;" | Fl<br>114
| style="border:solid black;border-width:1px 1px 1px 1px;" | Mc<br>115
| style="border:solid black;border-width:1px 1px 1px 1px;" | Lv<br>116
| style="border:solid black;border-width:1px 1px 1px 1px;" | Ts<br>117
| style="border:solid black;border-width:1px 1px 1px 1px;" | Og<br>118
|-
| ''Group''
| ''(1)''
| ''(2)''
|
| ''(3-12)''
| ''(13)''
| ''(14)''
| ''(15)''
| ''(16)''
| ''(17)''
| ''(18)''
|-
| colspan=14 style="border:none;"|
|-
| colspan=14 style="border:none; text-align:Center;font-size:105%;"| The [[Kainosymmetry|first-row anomaly]] strength by block is '''s''' >> '''p''' > '''d''' > '''f'''.<ref>[[#Jensen|Jensen 1986, p. 506]]</ref>{{efn|Helium is shown above beryllium for electron configuration consistency purposes; as a noble gas it is usually placed above neon, in group 18.}}
|}
Starting with hydrogen, the [[Kainosymmetry|first row anomaly]] primarily arises from the electron configurations of the elements concerned. Hydrogen is notable for its diverse bonding behaviors. It most commonly forms covalent bonds, but it can also lose its single electron in an [[aqueous solution]], leaving behind a bare proton with high polarizing power.<ref>[[#Lee|Lee 1996, p. 240]]</ref> Consequently, this proton can attach itself to the lone electron pair of an oxygen atom in a water molecule, laying the foundation for [[acid-base chemistry]].<ref>[[#Greenwood|Greenwood & Earnshaw 2002, p. 43]]</ref> Moreover, a hydrogen atom in a molecule can form a [[hydrogen bonding|second, albeit weaker, bond]] with an atom or group of atoms in another molecule. Such bonding, "helps give [[snowflake]]s their hexagonal symmetry, binds [[DNA]] into a [[double helix]]; shapes the three-dimensional forms of [[protein]]s; and even raises water's boiling point high enough to make a decent cup of tea."<ref>[[#Cressey|Cressey 2010]]</ref>

Hydrogen and helium, as well as boron through neon, have small atomic radii. The ionization energies and electronegativities among these elements are higher than the [[periodic trends]] would otherwise suggest.

While it would normally be expected, on electron configuration consistency grounds, that hydrogen and helium would be placed atop the s-block elements, the significant first row anomaly shown by these two elements justifies alternative placements. Hydrogen is occasionally positioned above fluorine, in group 17, rather than above lithium in group 1. Helium is almost always placed above neon, in group 18, rather than above beryllium in group 2.<ref>[[#Petruševski|Petruševski & Cvetković 2018]]; [[#Grochala|Grochala 2018]]</ref>

====Secondary periodicity====
[[File:EN values of chalcogens.png|thumb|upright=0.8|Electronegativity values of the group 16 [[chalcogen]] elements showing a W-shaped alternation or secondary periodicity going down the group|alt=A graph with a vertical electronegativity axis and a horizontal atomic number axis. The five elements plotted are {{abbr|O|oxygen}}, {{abbr|S|sulfur}}, {{abbr|Se|selenium}}, {{abbr|Te|tellurium}} and {{abbr|Po|polonium}}. The electronegativity of {{abbr|Se|selenium}} looks too high, and causes a bump in what otherwise be a smooth curve.]]
An alternation in certain periodic trends, sometimes referred to as [[Periodic table#Atomic radius|secondary periodicity]], becomes evident when descending groups 13 to 15, and to a lesser extent, groups 16 and 17.<ref>[[#Kneen|Kneen, Rogers & Simpson 1972, pp. 226, 360]]; [[#Siekierski|Siekierski & Burgess 2002, pp. 52, 101, 111, 124, 194]]</ref>{{efn|The net result is an even-odd difference between periods (except in the [[s-block]]): 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. Many properties in the p-block then 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>[[#Scerri2020|Scerri 2020, pp. 407–420]]</ref>}} Immediately after the first row of [[Block (periodic table)#d-block|d-block]] metals, from scandium to zinc, the 3d electrons in the [[Block (periodic table)#p-block|p-block]] elements—specifically, gallium (a metal), germanium, arsenic, selenium, and bromine—prove less effective at [[shielding effect|shielding]] the increasing positive nuclear charge.

The Soviet chemist {{Interlanguage link|Shchukarev|2=ru|3=Щукарев, Сергей Александрович|preserve=1}} gives two more tangible examples:<ref>[[#Shchukarev|Shchukarev 1977, p. 229]]</ref>
:<span style="font-size:95%">"The toxicity of some arsenic compounds, and the absence of this property in analogous compounds of phosphorus [P] and antimony [Sb]; and the ability of [[selenic acid]] [{{chem2|H2SeO4}}] to bring metallic gold [Au] into solution, and the absence of this property in sulfuric [[sulfuric acid|[{{chem2|H2SO4}}]]] and [[telluric acid|[{{chem2|H2TeO4}}]]] acids."</span>

====Higher oxidation states====
:''Roman numerals such as III, V and VIII denote oxidation states''

Some nonmetallic elements exhibit [[oxidation state]]s that deviate from those predicted by the octet rule, which typically results in an oxidation state of –3 in group 15, –2 in group 16, –1 in group 17, and 0 in group 18. Examples include [[ammonia]] NH<sub>3</sub>, [[hydrogen sulfide]] H<sub>2</sub>S, [[hydrogen fluoride]] HF, and elemental xenon Xe. Meanwhile, the maximum possible oxidation state increases from +5 in [[pnictogen|group 15]], to +8 in [[noble gas|group 18]]. The +5 oxidation state is observable from period 2 onward, in compounds such as [[nitric acid]] HN(V)O<sub>3</sub> and [[phosphorus pentafluoride]] PCl<sub>5</sub>.{{efn|Oxidation states do not reflect the actual net charge of atoms in molecules or ions, they represents the valence which refers more to how many bonds there are. For instance carbon typically has a valence of +4, but that only means that it forms three bonds. Electronegative elements such as fluorine are conventionally associated with negative valence, while electropositive ones have positive valence.}} [[Oxidation state#List of oxidation states of the elements|Higher oxidation states]] in later groups emerge from period 3 onwards, as seen in [[sulfur hexafluoride]] SF<sub>6</sub>, [[iodine heptafluoride]] IF<sub>7</sub>, and [[xenon tetroxide|xenon(VIII) tetroxide]] XeO<sub>4</sub>. For heavier nonmetals, their larger atomic radii and lower electronegativity values enable the formation of compounds with higher oxidation numbers, supporting higher bulk [[coordination number]]s.<ref name="Cox" />

====Multiple bond formation====
[[File:Pentazenium.png|thumb|right|alt=A chain of five N's in a wing shape|Molecular structure of [[pentazenium]], a homopolyatomic cation of nitrogen with the formula N<sub>5</sub><sup>+</sup> and structure N−N−N−N−N.<ref>[[#Vij|Vij et al. 2001]]</ref>]]Period 2 nonmetals, particularly carbon, nitrogen, and oxygen, show a propensity to form multiple bonds. The compounds formed by these elements often exhibit unique [[stoichiometries]] and structures, as seen in the various nitrogen oxides,<ref name="Cox">[[#Cox2004|Cox 2004, p. 146]]</ref> which are not commonly found in elements from later periods.

==== Property overlaps ====
While certain elements have traditionally been classified as nonmetals and others as metals, some overlapping of properties occurs. Writing early in the twentieth century, by which time the era of modern chemistry had been well-established<ref>[[#Dorsey|Dorsey 2023, pp. 12–13]]</ref> (although not as yet more precise [[quantum chemistry]]) Humphrey<ref>[[#Humphrey|Humphrey 1908]]</ref> observed that:
:<span style="font-size:95%">... these two groups, however, are not marked off perfectly sharply from each other; some nonmetals resemble metals in certain of their properties, and some metals approximate in some ways to the non-metals.</span>

[[Image:brown-boron.jpg|thumb|right|alt=An open glass jar with a brown powder in it|Boron (here in its less stable amorphous form) shares some similarities with metals{{efn|Greenwood<ref>[[#Greenwood2001|Greenwood 2001, p.&nbsp;2057]]</ref> commented that: "The extent to which metallic elements mimic boron (in having fewer electrons than orbitals available for bonding) has been a fruitful cohering concept in the development of metalloborane chemistry&nbsp;... Indeed, metals have been referred to as "honorary boron atoms" or even as "flexiboron atoms". The converse of this relationship is clearly also valid."}}]]
Examples of metal-like properties occurring in nonmetallic elements include:
* Silicon has an electronegativity (1.9) comparable with metals such as cobalt (1.88), copper (1.9), nickel (1.91) and silver (1.93);<ref name="AylwardEN"/>
* The electrical conductivity of graphite exceeds that of some metals;{{efn|For example, the conductivity of graphite is 3 × 10<sup>4</sup> S•cm<sup>−1.</sup><ref name=Bog/> whereas that of [[manganese]] is 6.9 × 10<sup>3</sup> S•cm<sup>−1</sup>.<ref>[[#Desai|Desai, James & Ho 1984, p.&nbsp;1160]]</ref>}}
* Selenium can be drawn into a wire;<ref name="ReferenceF"/>
* Radon is the most metallic of the noble gases and begins to show some [[cation]]ic behavior, which is unusual for a nonmetal;<ref>[[#Stein1983|Stein 1983, p. 165]]</ref> and
* In extreme conditions, just over half of nonmetallic elements can form homopolyatomic cations.{{efn|A homopolyatomic cation consists of two or more atoms of the same element bonded together and carrying a positive charge, for example, N<sub>5</sub><sup>+</sup>, O<sub>2</sub><sup>+</sup> and Cl<sub>4</sub><sup>+</sup>. This is unusual behavior for nonmetals since cation formation is normally associated with metals, and nonmetals are normally associated with anion formation. Homopolyatomic cations are further known for carbon, phosphorus, antimony, sulfur, selenium, tellurium, bromine, iodine and xenon.<ref>[[#Engesser|Engesser & Krossing 2013, p. 947]]</ref>}}

Examples of nonmetal-like properties occurring in metals are:
*[[Tungsten]] displays some nonmetallic properties, sometimes being brittle, having a high electronegativity, and forming only anions in aqueous solution,<ref>[[#S&P|Schweitzer & Pesterfield 2010, p. 305]]</ref> and predominately acidic oxides.<ref name="Porterfield"/><ref>[[#Rieck|Rieck 1967, p. 97]]: Tungsten trioxide dissolves in [[hydrofluoric acid]] to give an [[oxyfluoride]] [[Coordination complex|complex]].</ref>
*[[Gold]], the "king of metals" has the highest [[electrode potential]] among metals, suggesting a preference for gaining rather than losing electrons. Gold's ionization energy is one of the highest among metals, and its electron affinity and electronegativity are high, with the latter exceeding that of some nonmetals. It forms the Au<sup>–</sup> auride anion and exhibits a tendency to bond to itself, behaviors which are unexpected for metals. In aurides (MAu, where M = Li–Cs), gold's behavior is similar to that of a halogen.<ref>[[#Wiberg|Wiberg 2001, p. 1279]]</ref> The reason for this is that gold has a large enough nuclear potential that the electrons have to be considered with [[Relativistic quantum mechanics|relativistic]] effects included, which changes some of the properties.<ref>{{Cite journal |last=Pyper |first=N. C. |date=2020-09-18 |title=Relativity and the periodic table |url=https://royalsocietypublishing.org/doi/10.1098/rsta.2019.0305 |journal=Philosophical Transactions of the Royal Society A: Mathematical, Physical and Engineering Sciences |language=en |volume=378 |issue=2180 |pages=20190305 |doi=10.1098/rsta.2019.0305 |pmid=32811360 |bibcode=2020RSPTA.37890305P |issn=1364-503X}}</ref>

A relatively recent development involves certain compounds of heavier p-block elements, such as silicon, phosphorus, germanium, arsenic and antimony, exhibiting behaviors typically associated with [[coordination compound|transition metal complexes]]. This is linked to a small energy gap between their [[HOMO and LUMO|filled and empty]] [[molecular orbitals]], which are the regions in a molecule where electrons reside and where they can be available for chemical reactions. In such compounds, this allows for unusual reactivity with small molecules like hydrogen (H<sub>2</sub>), [[ammonia]] (NH<sub>3</sub>), and [[ethylene]] (C<sub>2</sub>H<sub>4</sub>), a characteristic previously observed primarily in transition metal compounds. These reactions may open new avenues in [[catalyst|catalytic]] applications.<ref>[[#Power|Power 2010]]; [[#Crow|Crow 2013]]; [[#Weetman|Weetman & Inoue 2018]]</ref>