Editing Period 2 element

{{short description|Any of the chemical elements in the second row of the periodic table}}
{{Periodic table (micro)| title=Period&nbsp;2 in the [[periodic table]] | mark=Li,Be,B,C,N,O,F,Ne}}
{{Sidebar periodic table|expanded=structure }}
A '''period&nbsp;2 element''' is one of the [[chemical element]]s in the second row (or [[Periodic table period|period]]) of the [[periodic table|periodic table of the chemical elements]]. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behavior of the elements as their [[atomic number]] increases; a new row is started when chemical behavior begins to repeat, creating [[Group (periodic table)|columns]] of elements with similar properties.

The second period contains the elements [[lithium]], [[beryllium]], [[boron]], [[carbon]], [[nitrogen]], [[oxygen]], [[fluorine]], and [[neon]]. In a [[quantum mechanics|quantum mechanical]] description of [[atomic structure]], this period corresponds to the filling of the [[electron shell|second ({{math|1=''n'' = 2}}) shell]], more specifically its [[s-block|2s]] and [[p-block|2p]] subshells. Period&nbsp;2 elements (carbon, nitrogen, oxygen, fluorine and neon) obey the [[octet rule]] in that they need eight electrons to complete their [[valence shell]] (lithium and beryllium obey [[duet rule]], boron is [[Electron deficiency|electron deficient]].), where at most eight electrons can be accommodated: two in the 2s orbital and six in the 2p subshell.

==Periodic trends==
[[File:Period 2 Calculated Radii.png|thumb|Calculated atomic radii of period&nbsp;2 elements in picometers.]]
Period&nbsp;2 is the first period in the periodic table from which [[periodic trends]] can be drawn. [[Period 1 element|Period&nbsp;1]], which only contains two elements ([[hydrogen]] and [[helium]]), is too small to draw any conclusive trends from it, especially because the two elements behave nothing like other s-block elements.<ref>{{cite journal|title=Where to Put Hydrogen in a Periodic Table?|journal=Foundations of Chemistry|year=2006|author=Michael Laing|doi=10.1007/s10698-006-9027-5|volume=9|issue=2|pages=127–137|s2cid=93781427}}</ref><ref>{{cite web|url=http://old.iupac.org/reports/periodic_table/ |title=International Union of Pure and Applied Chemistry > Periodic Table of the Elements |publisher=IUPAC |access-date=2011-05-01}}</ref> Period&nbsp;2 has much more conclusive trends. For all elements in period&nbsp;2, as the atomic number increases, the [[atomic radius]] of the elements decreases, the [[electronegativity]] increases, and the [[ionization energy]] increases.<ref>{{cite book |title=Chemistry: Principles and reactions |url=https://archive.org/details/chemistryprincip00mast_907 |url-access=limited |last1=Masterson |first1=William |last2=Hurley |first2=Cecile |year=2009 |publisher=Brooks/Cole Cengage Learning |location=Belmont, CA |isbn=978-0-495-12671-3 |pages=[https://archive.org/details/chemistryprincip00mast_907/page/n48 24]–42|edition=sixth}}</ref>

Period&nbsp;2 only has two [[metal]]s (lithium and beryllium) of eight elements, less than for any subsequent period both by number and by proportion. It also has the most number of nonmetals, namely five, among all periods. The elements in period&nbsp;2 often have the most extreme properties in their respective groups; for example, fluorine is the most reactive [[halogen]], neon is the most inert [[noble gas]],<ref>{{cite journal |last1=Grochala |first1=Wojciech |date=1 November 2017 |title=On the position of helium and neon in the Periodic Table of Elements |journal=Foundations of Chemistry |volume=20 |issue=3 |pages=191–207 |doi=10.1007/s10698-017-9302-7 |doi-access=free }}</ref> and lithium is the least reactive [[alkali metal]].<ref name="Gray" >{{cite book |last=Gray |first=Theodore |title=The Elements: A Visual Exploration of Every Known Atom in the Universe |year=2009 |publisher=Black Dog & Leventhal Publishers |location=New York |isbn=978-1-57912-814-2 |url-access=registration |url=https://archive.org/details/elementsvisualex0000gray }}</ref>

All period&nbsp;2 elements completely obey the [[Madelung rule]]; in period&nbsp;2, lithium and beryllium [[s-block|fill the 2s subshell]], and boron, carbon, nitrogen, oxygen, fluorine, and neon [[p-block|fill the 2p subshell]]. The period shares this trait with periods 1 and [[Period 3 element|3]], none of which contain [[transition element]]s or [[inner transition element]]s, which often vary from the rule.<ref name="Gray"/>

:{|
| colspan="3" | '''[[Chemical element]]''' || '''[[Block (periodic table)|Block]]''' || '''[[Electron configuration]]'''
|-bgcolor="{{element color|s-block}}"
|| 3 || '''Li''' || [[Lithium]] || [[s-block]] || [He] 2s<sup>1</sup>
|-bgcolor="{{element color|s-block}}"
|| 4 || '''Be''' || [[Beryllium]] || [[s-block]] || [He] 2s<sup>2</sup>
|-bgcolor="{{element color|p-block}}"
|| 5 || '''B''' || [[Boron]] || [[p-block]] || [He] 2s<sup>2</sup> 2p<sup>1</sup>
|-bgcolor="{{element color|p-block}}"
|| 6 || '''C''' || [[Carbon]] || [[p-block]] || [He] 2s<sup>2</sup> 2p<sup>2</sup>
|-bgcolor="{{element color|p-block}}"
|| 7 || '''N''' || [[Nitrogen]] || [[p-block]] || [He] 2s<sup>2</sup> 2p<sup>3</sup>
|-bgcolor="{{element color|p-block}}"
|| 8 || '''O''' || [[Oxygen]] || [[p-block]] || [He] 2s<sup>2</sup> 2p<sup>4</sup>
|-bgcolor="{{element color|p-block}}"
|| 9 || '''F''' || [[Fluorine]] || [[p-block]] || [He] 2s<sup>2</sup> 2p<sup>5</sup>
|-bgcolor="{{element color|p-block}}"
|| 10 || '''Ne''' || [[Neon]] || [[p-block]] || [He] 2s<sup>2</sup> 2p<sup>6</sup>
|}

===Lithium===
{{main article|Lithium}}
[[File:Lithium paraffin.jpg|thumb|left|150px|Lithium metal floating on paraffin oil]]

Lithium (Li) is an [[alkali metal]] with atomic number 3, occurring naturally in [[Isotopes of lithium|two isotopes]]: <sup>6</sup>Li and <sup>7</sup>Li. The two make up all natural occurrence of lithium on Earth, although [[Isotopes of lithium|further isotopes]] have been synthesized. In [[ionic compound]]s, lithium loses an [[electron]] to become positively charged, forming the [[cation]] Li<sup>+</sup>. Lithium is the first alkali metal in the periodic table,<ref group="note">Hydrogen is occasionally referred to as an alkali metal, although this is rare.</ref> and the first metal of any kind in the periodic table.<ref group="note">See note 1.</ref>  At [[standard temperature and pressure]], lithium is a soft, silver-white, highly reactive [[metal]]. With a [[density]] of 0.564 g⋅cm<sup>−3</sup>, lithium is the lightest metal and the least dense solid element.<ref name=weli>[http://www.webelements.com/lithium/ Lithium] at WebElements.</ref>

Lithium is one of the few elements [[Big Bang nucleosynthesis|synthesized]] in the [[Big Bang]].
Lithium is the 31st most abundant element on earth,<ref name=krebs>{{cite book | last = Krebs | first = Robert E. | year = 2006 | title = The History and Use of Our Earth's Chemical Elements: A Reference Guide | url = https://archive.org/details/historyuseourear00kreb_356 | url-access = limited | publisher = Greenwood Press | location = Westport, Conn. | isbn = 0-313-33438-2 | pages = [https://archive.org/details/historyuseourear00kreb_356/page/n71 47]–50}}</ref> occurring in concentrations of between 20 and 70 ppm by weight,<ref name=kamienski/> but due to its high reactivity it is only found naturally in [[Chemical compound|compounds]].<ref name=kamienski>Kamienski et al. "Lithium and lithium compounds". ''Kirk-Othmer Encyclopedia of Chemical Technology''. John Wiley & Sons, Inc. Published online '''2004'''. {{doi|10.1002/0471238961.1209200811011309.a01.pub2}}</ref>

Lithium [[Salt (chemistry)|salts]] are used in the pharmacology industry as [[Mood stabilizer|mood stabilising]] [[Medication|drugs]].<ref>{{cite journal
| title = Lithium salts in the treatment of psychotic excitement
| author = Cade J. F. J.
| journal = Medical Journal of Australia
| year = 1949
| volume = 2 |pmid=18142718
| pmc = 2560740
| issue = 10
| pages = 349–52
| url =https://www.who.int/docstore/bulletin/pdf/2000/issue4/classics.pdf | doi = 10.1080/j.1440-1614.1999.06241.x
}}</ref><ref>{{cite journal
| title =Lithium treatment for bipolar disorder
|author1=P. B. Mitchell |author2=D. Hadzi-Pavlovic | journal = Bulletin of the World Health Organization
| year = 2000
| volume = 78
| issue =4
| pages = 515–7 |pmid=10885179
| url =https://www.who.int/docstore/bulletin/pdf/2000/issue4/classics.pdf | pmc =2560742}}</ref> They are used in the treatment of [[bipolar disorder]], where they have a role in treating [[depression (mood)|depression]] and [[mania]] and may reduce the chances of [[suicide]].<ref>{{cite journal |vauthors=Baldessarini RJ, Tondo L, Davis P, Pompili M, Goodwin FK, Hennen J |date=October 2006 |title=Decreased risk of suicides and attempts during long-term lithium treatment: a meta-analytic review |journal=Bipolar Disorders |volume=8 |issue=5 Pt 2 |pages=625–39 |pmid=17042835 |doi=10.1111/j.1399-5618.2006.00344.x|doi-access=free }}</ref> The most common compounds used are [[lithium carbonate]], Li<sub>2</sub>CO<sub>3</sub>, [[lithium citrate]], Li<sub>3</sub>C<sub>6</sub>H<sub>5</sub>O<sub>7</sub>, [[lithium sulphate]], Li<sub>2</sub>SO<sub>4</sub>, and [[lithium orotate]], LiC<sub>5</sub>H<sub>3</sub>N<sub>2</sub>O<sub>4</sub>·H<sub>2</sub>O. Lithium is also used in [[lithium battery|batteries]] as an [[anode]] and its [[alloy]]s with [[aluminium]], [[cadmium]], [[copper]] and [[manganese]] are used to make high performance parts for [[aircraft]], most notably the [[Space Shuttle external tank|external tank]] of the [[Space Shuttle]].<ref name=weli/>

===Beryllium===
{{main article|Beryllium}}
[[File:Be-140g.jpg|thumb|left|150px|Large piece of beryllium]]

Beryllium (Be) is the chemical element with atomic number 4, occurring in the form of <sup>9</sup>Be. At standard temperature and pressure, beryllium is a strong, steel-grey, light-weight, [[brittle]], [[Bivalent (chemistry)|bivalent]] [[alkaline earth metal]], with a density of 1.85 g⋅cm<sup>−3</sup>.<ref name=webe>[http://www.webelements.com/beryllium/ Beryllium] at WebElements.</ref> It also has one of the highest [[melting point]]s of all the [[light metal]]s. Beryllium's most common [[isotope]] is <sup>9</sup>Be, which contains 4 protons and 5 neutrons. It makes up almost 100% of all naturally occurring beryllium and is its only stable isotope; however [[Isotopes of beryllium|other isotopes]] have been synthesised. In ionic compounds, beryllium loses its two [[valence electron]]s to form the cation, Be<sup>2+</sup>.

Small amounts of beryllium were [[Big Bang nucleosynthesis|synthesised]] during the [[Big Bang]], although most of it [[Nuclear decay|decayed]] or reacted further to create larger nuclei, like carbon, nitrogen or oxygen. Beryllium is a component of 100 out of 4000 known [[mineral]]s, such as [[bertrandite]], Be<sub>4</sub>Si<sub>2</sub>O<sub>7</sub>(OH)<sub>2</sub>, [[beryl]], Al<sub>2</sub>Be<sub>3</sub>Si<sub>6</sub>O<sub>18</sub>, [[chrysoberyl]], Al<sub>2</sub>BeO<sub>4</sub>, and [[phenakite]], Be<sub>2</sub>SiO<sub>4</sub>. Precious forms of beryl are [[Aquamarine (gemstone)|aquamarine]], [[red beryl]] and [[emerald]]. The most common sources of beryllium used commercially are beryl and bertrandite and production of it involves the [[redox|reduction]] of [[beryllium fluoride]] with [[magnesium]] metal or the [[electrolysis]] of molten [[beryllium chloride]], containing some [[sodium chloride]] as beryllium chloride is a poor [[Electrical conductor|conductor of electricity]].<ref name=webe/>

Due to its stiffness, light weight, and dimensional stability over a wide temperature range, beryllium metal is used in as a structural material in aircraft, missiles and [[communication satellite]]s.<ref name=webe/> It is used as an alloying agent in [[beryllium copper]], which is used to make electrical components due to its high electrical and heat conductivity.<ref>[http://www.copper.org/resources/properties/microstructure/be_cu.html Standards and properties] of beryllium copper.</ref> Sheets of beryllium are used in [[X-ray]] detectors to filter out [[visible light]] and let only X-rays through.<ref name=webe/> It is used as a [[neutron moderator]] in [[nuclear reactor]]s because light nuclei are more effective at slowing down neutrons than heavy nuclei.<ref name=webe/> Beryllium's low weight and high rigidity also make it useful in the construction of [[tweeter]]s in [[loudspeaker]]s.<ref>[http://www.hometheaterhifi.com/volume_14_3/feature-article-beryllium-9-2007.html Information] about beryllium tweeters.</ref>

Beryllium and beryllium compounds are classified by the [[International Agency for Research on Cancer]] as [[List of IARC Group 1 carcinogens|Group 1 carcinogens]]; they are carcinogenic to both animals and humans.<ref>{{cite web
 | url = http://www.inchem.org/documents/iarc/vol58/mono58-1.html
 | publisher = International Agency for Research on Cancer
 | title = IARC Monograph, Volume 58
 | year = 1993
 | access-date = 2008-09-18}}</ref> Chronic [[berylliosis]] is a [[pulmonary]] and [[systemic circulation|systemic]] [[granulomatous]] disease caused by exposure to beryllium. Between 1% – 15% of people are sensitive to beryllium and may develop an inflammatory reaction in their [[respiratory system]] and [[skin]], called chronic beryllium disease or [[berylliosis]]. The body's [[immune system]] recognises the beryllium as foreign particles and mounts an attack against them, usually in the lungs where they are breathed in. This can cause fever, fatigue, weakness, night sweats and difficulty in breathing.<ref>[https://web.archive.org/web/20010331191955/http://www.chronicberylliumdisease.com/medical/med_bediseases.htm#cbd Information] about chronic beryllium disease.</ref>

===Boron===
{{main article|Boron}}
[[File:Bor 1.jpg|thumb|left|150px|Boron chunks]]

Boron (B) is the chemical element with atomic number 5, occurring as <sup>10</sup>B and <sup>11</sup>B. At standard temperature and pressure, boron is a [[trivalent]] [[metalloid]] that has several different [[allotropy|allotropes]]. [[Amorphous]] boron is a brown powder formed as a product of many chemical reactions. [[Crystalline]] boron is a very hard, black material with a high melting point and exists in many [[Polymorphism (materials science)|polymorphs]]: Two [[rhombohedral]] forms, α-boron and β-boron containing 12 and 106.7 atoms in the rhombohedral unit cell respectively, and 50-atom [[tetragonal]] boron are the most common. Boron has a density of 2.34<sup>−3</sup>.<ref name=web>[http://www.webelements.com/boron/  Boron] at WebElements.</ref> Boron's most common [[isotope]] is <sup>11</sup>B at 80.22%, which contains 5 protons and 6 neutrons. The other common isotope is <sup>10</sup>B at 19.78%, which contains 5 protons and 5 neutrons.<ref name=rem>[http://www.rareearth.org/boron_properties.htm Properties] of boron.</ref> These are the only stable isotopes of boron; however [[Isotopes of boron|other isotopes]] have been synthesised. Boron forms covalent bonds with other [[Nonmetal (chemistry)|nonmetal]]s and has [[oxidation state]]s of 1, 2, 3 and 4.<ref>{{cite web |url=http://bernath.uwaterloo.ca/media/78.html |format=PDF |title=Fourier Transform Spectroscopy: B<sup>4</sup>Σ<sup>−</sup>−X<sup>4</sup>Σ<sup>−</sup> |author1=W.T.M.L. Fernando |author2=L.C. O'Brien |author3=P.F. Bernath |publisher=University of Arizona, Tucson |access-date=2007-12-10 }}{{Dead link|date=August 2018 |bot=InternetArchiveBot |fix-attempted=yes }}</ref><ref>{{cite web |url=http://bernath.uwaterloo.ca/media/125.html |format=PDF |title=Infrared Emission Spectroscopy of BF and AIF |author=K.Q. Zhang, B.Guo, V. Braun, M. Dulick, P.F. Bernath |access-date=2007-12-10 }}{{Dead link|date=August 2018 |bot=InternetArchiveBot |fix-attempted=yes }}</ref><ref>{{cite web |url=http://lb.chemie.uni-hamburg.de/search/index.php?content=166/dGp23678 |title=Compound Descriptions: B<sub>2</sub>F<sub>4</sub> |access-date=2007-12-10 |publisher=Landol Börnstein Substance/Property Index}}</ref>
Boron does not occur naturally as a free element, but in compounds such as [[borate]]s. The most common sources of boron are [[tourmaline]], [[borax]], Na<sub>2</sub>B<sub>4</sub>O<sub>5</sub>(OH)<sub>4</sub>·8H<sub>2</sub>O, and [[kernite]], Na<sub>2</sub>B<sub>4</sub>O<sub>5</sub>(OH)<sub>4</sub>·2H<sub>2</sub>O.<ref name=web/> it is difficult to obtain pure boron. It can be made through the [[magnesium]] [[redox|reduction]] of [[boron trioxide]], B<sub>2</sub>O<sub>3</sub>. This oxide is made by melting [[boric acid]], B(OH)<sub>3</sub>, which in turn is obtained from borax. Small amounts of pure boron can be made by the [[thermal decomposition]] of boron bromide, BBr<sub>3</sub>, in hydrogen gas over hot [[tantalum]] wire, which acts as a [[catalyst]].<ref name=web/> The most commercially important sources of boron are: [[sodium tetraborate]] pentahydrate, Na<sub>2</sub>B<sub>4</sub>O<sub>7</sub> · 5H<sub>2</sub>O, which is used in large amounts in making insulating [[fiberglass]] and [[sodium perborate]] [[Bleach (chemical)|bleach]]; [[boron carbide]], a [[ceramic]] material, is used to make armour materials, especially in [[bulletproof vest]]s for soldiers and police officers; [[orthoboric acid]], H<sub>3</sub>BO<sub>3</sub> or boric acid, used in the production of textile [[fiberglass]] and [[flat panel display]]s; sodium tetraborate decahydrate, Na<sub>2</sub>B<sub>4</sub>O<sub>7</sub> · 10H<sub>2</sub>O or borax, used in the production of adhesives; and the isotope boron-10 is used as a control for nuclear reactors, as a shield for nuclear radiation, and in instruments used for detecting neutrons.<ref name=rem/>

Boron is an essential plant [[micronutrient]], required for cell wall strength and development, cell division, seed and fruit development, sugar transport and hormone development.<ref>{{cite journal
 | title = Functions of Boron in Plant Nutrition
 | first = Dale G.
 | last = Blevins
 |author2=Lukaszewski, Krystyna M.
  | journal = Annual Review of Plant Physiology and Plant Molecular Biology
 | volume = 49
 | pages = 481–500
 | year = 1998
 | doi = 10.1146/annurev.arplant.49.1.481
 | pmid = 15012243 }}</ref> However, high soil concentrations of over 1.0 [[parts per million|ppm]] can cause necrosis in leaves and poor growth. Levels as low as 0.8 ppm can cause these symptoms to appear in plants particularly boron-sensitive. Most plants, even those tolerant of boron in the soil, will show symptoms of boron toxicity when boron levels are higher than 1.8 ppm.<ref name=rem/> In animals, boron is an [[ultratrace element]]; in human diets, daily intake ranges from 2.1 to 4.3&nbsp;mg boron/kg body weight (bw)/day.<ref>{{cite journal | title = 850-5 |vauthors=Zook EG, Lehman J| journal = J. Assoc. Off Agric. Chem | volume = 48 | year = 1965}}</ref> It is also used as a supplement for the prevention and treatment of osteoporosis and arthritis.<ref>{{cite web | url = http://www.pdrhealth.com/drug_info/nmdrugprofiles/nutsupdrugs/bor_0040.shtml
| title = Boron | access-date = 2008-09-18 | publisher = PDRhealth |archive-url = https://web.archive.org/web/20071011101928/http://pdrhealth.com/drug_info/nmdrugprofiles/nutsupdrugs/bor_0040.shtml |archive-date=October 11, 2007 }}</ref>

===Carbon===
{{main article|Carbon}}
[[File:Diamond-and-graphite-with-scale.jpg|thumb|left|150px|Diamond and graphite, two different [[allotrope]]s of carbon]]

Carbon is the chemical element with atomic number 6, occurring as <sup>12</sup>C, <sup>13</sup>C and <sup>14</sup>C.<ref name=wec>[http://www.webelements.com/carbon/ Carbon] at WebElements.</ref> At standard temperature and pressure, carbon is a solid, occurring in [[Allotropes of carbon|many different allotropes]], the most common of which are [[graphite]], [[diamond]], the [[fullerenes]] and [[amorphous carbon]].<ref name=wec/> Graphite is a soft, [[hexagonal crystal system|hexagonal crystalline]], opaque black [[semimetal]] with very good [[electrical conductor|conductive]] and [[thermodynamic equilibrium|thermodynamically stable]] properties. Diamond however is a highly [[transparency (optics)|transparent]] [[Transparency and translucency|colourless]] [[cubic crystal system|cubic crystal]] with poor conductive properties, is the [[Mohs scale of mineral hardness|hardest known naturally occurring mineral]] and has the highest [[refractive index]] of all [[gemstones]]. In contrast to the [[crystal lattice]] structure of diamond and graphite, the [[fullerenes]] are [[molecules]], named after [[Richard Buckminster Fuller]] whose architecture the molecules resemble. There are several different fullerenes, the most widely known being the "buckeyball" C<sub>60</sub>. Little is known about the fullerenes and they are a current subject of research.<ref name=wec/> There is also amorphous carbon, which is carbon without any crystalline structure.<ref>{{cite book |chapter=Amorphous carbon |chapter-url=http://iupac.org/goldbook/A00294.html |title=IUPAC Compendium of Chemical Terminology |publisher=International Union of Pure and Applied Chemistry |year=1997|edition=2nd |access-date=2008-09-24}}</ref> In [[mineralogy]], the term is used to refer to [[soot]] and [[coal]], although these are not truly amorphous as they contain small amounts of graphite or diamond.<ref>{{cite journal |url=http://gltrs.grc.nasa.gov/reports/1996/CR-198469.html |format=PDF |title=Soot Precursor Material: Spatial Location via Simultaneous LIF-LII Imaging and Characterization via TEM |journal=NASA Contractor Report |last=Vander Wal |first=R. |issue=198469 |date=May 1996 |access-date=2008-09-24 }}{{dead link|date=June 2021|bot=medic}}{{cbignore|bot=medic}}</ref><ref>{{cite book |chapter=diamond-like carbon films |chapter-url=https://goldbook.iupac.org/terms/view/D01673 |title=IUPAC Compendium of Chemical Terminology |publisher=International Union of Pure and Applied Chemistry |year=1997|edition=2nd |doi=10.1351/goldbook.D01673 |access-date=2008-09-24}}</ref> Carbon's most common isotope at 98.9% is <sup>12</sup>C, with six protons and six neutrons.<ref name=slide>[http://www.scienceschool.usyd.edu.au/media/17-dasgupta-slides.pdf Presentation about isotopes] {{webarchive|url=https://web.archive.org/web/20080719061754/http://www.scienceschool.usyd.edu.au/media/17-dasgupta-slides.pdf |date=2008-07-19 }} by Mahananda Dasgupta of the Department of Nuclear Physics at Australian National University.</ref> <sup>13</sup>C is also stable, with six protons and seven neutrons, at 1.1%.<ref name=slide/> Trace amounts of <sup>14</sup>C also occur naturally but this [[radioisotope|isotope is radioactive]] and decays with a half life of 5730 years; it is used for [[radiocarbon dating]].<ref>{{cite journal |last=Plastino |first=W. |author2=Kaihola, L. |author3=Bartolomei, P. |author4=Bella, F. |year=2001 |title=Cosmic Background Reduction In The Radiocarbon Measurement By Scintillation Spectrometry At The Underground Laboratory Of Gran Sasso |journal=Radiocarbon |volume=43 |issue=2A |pages=157–161 |doi=10.1017/S0033822200037954 |doi-access=free }}</ref> Other [[isotopes of carbon]] have also been synthesised. Carbon forms covalent bonds with other non-metals with an oxidation state of −4, −2, +2 or +4.<ref name=wec/>

Carbon is the fourth most abundant element in the universe by mass after [[hydrogen]], [[helium]] and oxygen<ref>[http://plymouthlibrary.org/faqelements.htm Ten most abundant elements in the universe, taken from ''The Top 10 of Everything'', 2006, Russell Ash, page 10. Retrieved October 15, 2008.] {{webarchive|url=https://web.archive.org/web/20100210170504/http://plymouthlibrary.org/faqelements.htm |date=February 10, 2010 }}</ref> and is the second [[Chemical makeup of the human body|most abundant element in the human body]] by mass after oxygen,<ref>{{cite book
 | last = Chang
 | first = Raymond
 | title = Chemistry, Ninth Edition
 | publisher = McGraw-Hill
 | year = 2007
 | pages = 52
 | isbn = 978-0-07-110595-8 }}</ref> the third most abundant by number of atoms.<ref name="Freitas">{{cite book |first=Robert A. Jr. |last=Freitas |title=Nanomedicine |url=http://www.foresight.org/Nanomedicine/Ch03_1.html |publisher=Landes Bioscience |year=1999 |page=Tables 3–1 & 3–2 |isbn=1-57059-680-8 |no-pp=true }}</ref> There are an almost infinite number of compounds that contain carbon due to carbon's ability to form long stable chains of C&nbsp;— C bonds.<ref name="hydrocarbon"/><ref name=acell>{{cite book |last = Alberts |first = Bruce|author2=Alexander Johnson |author3=Julian Lewis |author4=Martin Raff |author5=Keith Roberts |author6=Peter Walter  |title = Molecular Biology of the Cell | chapter=The Chemical Components of a Cell |publisher = Garland Science |url = https://www.ncbi.nlm.nih.gov/books/bv.fcgi?highlight=carbon&rid=mboc4.section.165|year = 2002}}</ref> The simplest carbon-containing molecules are the [[hydrocarbon]]s, which contain carbon and hydrogen,<ref name="hydrocarbon"/> although they sometimes contain other elements in [[functional group]]s. Hydrocarbons are used as [[fossil fuels]] and to manufacture [[plastics]] and [[petrochemicals]]. All [[organic compound]]s, those essential for life, contain at least one atom of carbon.<ref name="hydrocarbon">{{cite web| title = Structure and Nomenclature of Hydrocarbons | publisher = Purdue University| url = http://chemed.chem.purdue.edu/genchem/topicreview/bp/1organic/organic.html| access-date = 2008-03-23}}</ref><ref name=acell/> When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds<ref name=acell/> including [[sugar]]s, [[lignan]]s, [[chitin]]s, [[Alcohol (chemistry)|alcohol]]s, [[fat]]s, and aromatic [[ester]]s, [[carotenoids]] and [[terpenes]]. With [[nitrogen]] it forms [[alkaloid]]s, and with the addition of sulfur also it forms [[antibiotic]]s, [[amino acid]]s, and [[rubber]] products. With the addition of phosphorus to these other elements, it forms [[DNA]] and [[RNA]], the chemical-code carriers of life, and [[adenosine triphosphate]] (ATP), the most important energy-transfer molecule in all living cells.<ref name=acell/>

===Nitrogen===
{{main article|Nitrogen}}
[[File:Liquidnitrogen.jpg|thumb|left|150px|Liquid nitrogen being poured]]

Nitrogen is the chemical element with atomic number 7, the symbol '''N''' and [[atomic mass]] 14.00674&nbsp;u. Elemental nitrogen is a colorless, odorless, tasteless and mostly [[Inert gas|inert]] [[diatomic]] gas at [[standard conditions]], constituting 78.08% by volume of [[Earth's atmosphere]]. The element nitrogen was discovered as a separable component of air, by Scottish physician [[Daniel Rutherford]], in 1772.<ref>{{Cite book |url=https://archive.org/details/elementsofchemis0000lavo |url-access=registration |page=[https://archive.org/details/elementsofchemis0000lavo/page/15 15] |title=Elements of chemistry, in a new systematic order: containing all the modern discoveries |author=Lavoisier, Antoine Laurent |author-link=Antoine Lavoisier |publisher=Courier Dover Publications |year=1965 |isbn=0-486-64624-6}}</ref> It occurs naturally in form of two isotopes: nitrogen-14 and nitrogen-15.<ref name=wen>[http://www.webelements.com/nitrogen/ Nitrogen] at WebElements.</ref>

Many industrially important compounds, such as [[ammonia]], [[nitric acid]], organic nitrates ([[propellant]]s and [[explosive]]s), and [[cyanide]]s, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the {{chem|N|2}} molecule into useful [[Chemical compound|compounds]], but at the same time causing release of large amounts of often useful energy when the compounds burn, explode, or decay back into nitrogen gas.

Nitrogen occurs in all living organisms, and the [[nitrogen cycle]] describes movement of the element from air into the [[biosphere]] and organic compounds, then back into the atmosphere. Synthetically produced [[nitrate]]s are key ingredients of industrial [[fertilizer]]s, and also key pollutants in causing the [[eutrophication]] of water systems. Nitrogen is a constituent element of [[amino acid]]s and thus of [[protein]]s, and of [[nucleic acid]]s ([[DNA]] and [[RNA]]). It resides in the [[chemical structure]] of almost all [[neurotransmitter]]s, and is a defining component of [[alkaloid]]s, biological molecules produced by many organisms.<ref name="Lightning">{{Cite book |title = Lightning: Physics and Effects |first = Vladimir A. |last = Rakov |author2=Uman, Martin A.  |publisher = Cambridge University Press |year = 2007 |isbn = 978-0-521-03541-5 |page = 508 |url=https://books.google.com/books?id=TuMa5lAa3RAC&pg=PA508}}</ref>

===Oxygen===
{{main article|Oxygen}}Oxygen is the chemical element with atomic number 8, occurring mostly as <sup>16</sup>O, but also <sup>17</sup>O and <sup>18</sup>O.

Oxygen is the third-most common element by mass in the universe (although there are more carbon atoms, each carbon atom is lighter). It is highly electronegative and non-metallic, usually diatomic, gas down to very low temperatures. Only fluorine is more reactive among non-metallic elements. It is two electrons short of a full octet and readily takes electrons from other elements. It reacts violently with [[alkali metals]] and [[white phosphorus]] at room temperature and less violently with alkali earth metals heavier than magnesium. At higher temperatures it burns most other metals and many non-metals (including hydrogen, carbon, and sulfur). Many oxides are extremely stable substances difficult to decompose—like [[water]], [[carbon dioxide]], [[alumina]], [[silica]], and iron oxides (the latter often appearing as [[rust]]). Oxygen is part of substances best described as some salts of metals and oxygen-containing acids (thus nitrates, sulfates, phosphates, silicates, and carbonates.

Oxygen is essential to all life. Plants and [[phytoplankton]] photosynthesize water and carbon dioxide and water, both oxides, in the presence of sunlight to form [[sugar]]s with the release of  oxygen. The sugars are then turned into such substances as cellulose and (with nitrogen and often sulfur) proteins and other essential substances of life. Animals especially but also fungi and bacteria ultimately depend upon photosynthesizing plants and phytoplankton for food and oxygen.

[[Fire]] uses oxygen to oxidize compounds typically of carbon and hydrogen to water and carbon dioxide (although other elements may be involved) whether in uncontrolled conflagrations that destroy buildings and forests or the controlled fire within engines or that supply electrical energy from turbines, heat for keeping buildings warm, or the motive force that drives vehicles.

Oxygen forms roughly 21% of the Earth's atmosphere; all of this oxygen is the result of photosynthesis. Pure oxygen has use in medical treatment of people who have respiratory difficulties. [[oxygen toxicity|Excess oxygen is toxic]].

Oxygen was originally associated with the formation of acids—until some acids were shown to not have oxygen in them. Oxygen is named for its formation of acids, especially with non-metals. Some oxides of some non-metals are extremely acidic, like [[sulfur trioxide]], which forms [[sulfuric acid]] on contact with water. Most oxides with metals are alkaline, some extremely so, like [[potash|potassium oxide]]. Some metallic oxides are amphoteric, like aluminum oxide, which means that they can react with both acids and bases.

Although oxygen is normally a diatomic gas, oxygen can form an allotrope known as [[ozone]]. Ozone is a triatomic gas even more reactive than oxygen. Unlike regular diatomic oxygen, ozone is a toxic material generally considered a pollutant. In the upper atmosphere, some oxygen forms ozone which has the property of absorbing dangerous ultraviolet rays within the [[ozone layer]]. Land life was impossible before the formation of an ozone layer.

===Fluorine===
{{main article|Fluorine}}
[[File:Liquid fluorine tighter crop.jpg|thumb|left|150px|Liquid fluorine in ampoule]]

Fluorine is the chemical element with atomic number 9. It occurs naturally in its only stable form <sup>19</sup>F.<ref>{{cite web |author=National Nuclear Data Center |title=NuDat 2.1 database – fluorine-19 |url=http://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=9&n=10 |publisher=[[Brookhaven National Laboratory]] |access-date=2011-05-01}}</ref>

Fluorine is a pale-yellow, diatomic gas under normal conditions and down to very low temperatures. Short one electron of the highly stable octet in each atom, fluorine molecules are unstable enough that they easily snap, with loose fluorine atoms tending to grab single electrons from just about any other element. Fluorine is the most reactive of all elements, and it even attacks many oxides to replace oxygen with fluorine. Fluorine even attacks silica, one of the favored materials for transporting strong acids, and burns asbestos. It attacks [[sodium chloride|common salt]], one of the most stable compounds, with the release of chlorine. It never appears uncombined in nature and almost never stays uncombined for long. It burns hydrogen simultaneously if either is liquid or gaseous—even at temperatures close to absolute zero.<ref>{{cite web|url=https://www.webelements.com/fluorine/|title=WebElements Periodic Table » Fluorine » the essentials|website=www.webelements.com}}</ref> It is extremely difficult to isolate from any compounds, let alone keep uncombined.

Fluorine gas is extremely dangerous because it attacks almost all organic material, including live flesh. Many of the binary compounds that it forms (called fluorides) are themselves highly toxic, including soluble fluorides and especially [[hydrogen fluoride]]. Fluorine forms very strong bonds with many elements. With sulfur it can form the extremely stable and chemically inert [[sulfur hexafluoride]]; with carbon it can form the remarkable material [[Teflon]] that is a stable and non-combustible solid with a high melting point and a very low coefficient of friction that makes it an excellent liner for cooking pans and raincoats. Fluorine-carbon compounds include some unique plastics.
it is also used as a reactant in the making of toothpaste.

===Neon===
{{main article|Neon}}
[[File:Neon discharge tube.jpg|thumb|left|150px|Neon [[discharge tube]]]]

{{Expand section |date=May 2011}}
Neon is the chemical element with atomic number 10, occurring as <sup>20</sup>Ne, <sup>21</sup>Ne and <sup>22</sup>Ne.<ref>{{cite web |url=http://nautilus.fis.uc.pt/st2.5/scenes-e/elem/e01093.html |title=Neon: Isotopes |access-date=2011-05-01 |publisher=Softciências |archive-url=https://web.archive.org/web/20121115190653/http://nautilus.fis.uc.pt/st2.5/scenes-e/elem/e01093.html |archive-date=2012-11-15 |url-status=dead }}</ref>

Neon is a monatomic gas. With a complete octet of outer electrons it is highly resistant to removal of any electron, and it cannot accept an electron from anything. Neon has no tendency to form any normal compounds under normal temperatures and pressures; it is effectively inert.  It is one of the so-called "noble gases".

Neon is a trace component of the atmosphere without any biological role.

==Notes==
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==References==
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==External links==
*{{Commonscatinline|Periodic table row 2}}

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