Editing Barium (section)

==Characteristics==

===Physical properties===
[[File:Barium 1.jpg|thumb|left|Oxidized barium]]
Barium is a soft, silvery-white metal, with a slight golden shade when ultrapure.<ref name="Ullman2005">{{Ullmann |last1=Kresse|first1=Robert|last2=Baudis|first2=Ulrich|last3=Jäger|first3=Paul|last4=Riechers|first4=H. Hermann|last5=Wagner|first5=Heinz|last6=Winkler|first6=Jochen|last7=Wolf|first7=Hans Uwe|title=Barium and Barium Compounds|date=2007 |doi=10.1002/14356007.a03_325.pub2|isbn=9783527306732}}</ref>{{rp|2}} The silvery-white color of barium metal rapidly vanishes upon [[redox|oxidation]] in air yielding a dark gray layer containing the [[Barium oxide|oxide]]. Barium has a medium [[specific weight]] and high electrical conductivity. Because barium is difficult to purify, many of its properties have not been accurately determined.<ref name="Ullman2005" />{{rp|2}}

At room temperature and pressure, barium metal adopts a [[body-centered cubic]] structure, with a barium–barium distance of 503 [[picometer]]s, expanding with heating at a rate of approximately 1.8{{e|-5}}/°C.<ref name="Ullman2005" />{{rp|2}} It is a soft metal with a [[Mohs hardness]] of 1.25.<ref name="Ullman2005" />{{rp|2}} Its melting temperature of {{convert|1000|K|C F}}<ref name="Lide2004">{{cite book|last = Lide |first= D. R. |title = CRC Handbook of Chemistry and Physics |url = https://archive.org/details/crchandbookofche81lide |url-access = registration |edition = 84th |location = Boca Raton (FL) |publisher = CRC Press |date = 2004 |isbn = 978-0-8493-0484-2}}</ref>{{rp|4–43}} is intermediate between those of the lighter strontium ({{convert|1050|K|C F|disp=or}})<ref name="Lide2004" />{{rp|4–86}} and heavier radium ({{convert|973|K|C F|disp=or}});<ref name="Lide2004" />{{rp|4–78}} however, its boiling point of {{convert|2170|K|C F}} exceeds that of strontium ({{convert|1655|K|C F|disp=or}}).<ref name="Lide2004" />{{rp|4–86}} The density (3.62&nbsp;g/cm<sup>3</sup>)<ref name="Lide2004" />{{rp|4–43}} is again intermediate between those of strontium (2.36&nbsp;g/cm<sup>3</sup>)<ref name="Lide2004" />{{rp|4–86}} and radium (≈5&nbsp;g/cm<sup>3</sup>).<ref name="Lide2004" />{{rp|4–78}}

===Chemical reactivity===
Barium is chemically similar to magnesium, calcium, and strontium, but more reactive. Its compounds are almost invariably found in the +2 oxidation state.  As expected for a highly electropositive metal, barium's reaction with [[chalcogen]]s is highly [[exothermic reaction|exothermic]] (release energy). Barium reacts with atmospheric oxygen in air at room temperature. For this reason, metallic barium is often stored under oil or in an inert atmosphere.<ref name="Ullman2005" />{{rp|2}} Reactions with other [[Nonmetal (chemistry)|nonmetal]]s, such as carbon, nitrogen, phosphorus, silicon, and hydrogen, proceed upon heating.<ref name="Ullman2005" />{{rp|2–3}} Reactions with water and alcohols are also exothermic and release hydrogen gas:<ref name="Ullman2005" />{{rp|3}}
: Ba + 2 ROH → Ba(OR)<sub>2</sub> + H<sub>2</sub>↑ (R is an alkyl group or a hydrogen atom)

Barium reacts with [[ammonia]] to form the electride [Ba(NH<sub>3</sub>)<sub>6</sub>](e<sup>−</sup>)<sub>2</sub>, which near room temperature gives the amide Ba(NH<sub>2</sub>)<sub>2</sub>.<ref>{{Greenwood&Earnshaw2nd|page=113}}</ref>

The metal is readily attacked by acids. [[Sulfuric acid]] is a notable exception because [[Passivation (chemistry)|passivation]] stops the reaction by forming the insoluble [[barium sulfate]] on the surface.<ref>{{Ullmann |last=Müller |first=Hermann |date= 2000 |title=Sulfuric Acid and Sulfur Trioxide |doi=10.1002/14356007.a25_635 |isbn=9783527306732}}</ref> Barium combines with several other metals, including [[aluminium]], [[zinc]], [[lead]], and [[tin]], forming [[intermetallics|intermetallic phases]] and alloys.<ref>{{cite book|author= Ferro, Riccardo|author2= Saccone, Adriana|name-list-style= amp|page=355|title=Intermetallic Chemistry|publisher=Elsevier|date=2008|isbn=978-0-08-044099-6}}</ref>

===Compounds===
{| class="wikitable" style="float:left; margin-top:0; margin-right:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ style="margin-bottom: 5px;"|Selected alkaline earth and zinc salts densities,&nbsp;g/cm<sup>3</sup>
|-
!
! [[oxide|{{chem|O|2-}}]]
! [[sulfide|{{chem|S|2-}}]]
! [[fluoride|{{chem|F|-}}]]
! [[chloride|{{chem|Cl|-}}]]
! [[sulfate|{{chem|SO|4|2-}}]]
! [[carbonate|{{chem|CO|3|2-}}]]
! [[peroxide|{{chem|O|2|2-}}]]
! [[hydride|{{chem|H|-}}]]
|-
! scope="row"|[[calcium|{{chem|Ca|2+}}]]<ref name="Lide2004" />{{rp|4–48–50}}
|3.34
|2.59
|3.18
|2.15
|2.96
|2.83
|2.9
|1.7
|-
! scope="row"|[[strontium|{{chem|Sr|2+}}]]<ref name="Lide2004" />{{rp|4–86–88}}
|5.1
|3.7
|4.24
|3.05
|3.96
|3.5
|4.78
|3.26
|-
! scope="row" style="background:#ff9;"| '''''{{chem|Ba|2+}}'''''<ref name="Lide2004" />{{rp|4–43–45}}
| style="background:#ff9;"| ''5.72''
| style="background:#ff9;"| ''4.3''
| style="background:#ff9;"| ''4.89''
| style="background:#ff9;"| ''3.89''
| style="background:#ff9;"| ''4.49''
| style="background:#ff9;"| ''4.29''
| style="background:#ff9;"| ''4.96''
| style="background:#ff9;"| ''4.16''
|-
! scope="row"|[[zinc|{{chem|Zn|2+}}]]<ref name="Lide2004" />{{rp|4–95–96}}
|5.6
|4.09
|4.95
|2.09
|3.54
|4.4
|1.57
|—
|}

Barium salts are typically white when solid and colorless when dissolved.<ref>{{cite book|page=87|title=Qualitative analysis and the properties of ions in aqueous solution|author=Slowinski, Emil J.|author2=Masterton, William L.|publisher=Saunders|date=1990|edition=2nd|isbn=978-0-03-031234-2}}</ref> They are denser than the [[strontium]] or [[calcium]] analogs, except for the [[halide]]s (see table; [[zinc]] is given for comparison).

[[Barium hydroxide]] ("baryta") was known to alchemists, who produced it by heating barium carbonate. Unlike calcium hydroxide, it absorbs very little CO<sub>2</sub> in aqueous solutions and is therefore insensitive to atmospheric fluctuations. This property is used in calibrating pH equipment.

Barium compounds burn with a green to pale green [[flame test|flame]], which is an efficient test to detect a barium compound. The color results from [[spectral line]]s at 455.4, 493.4, 553.6, and 611.1&nbsp;nm.<ref name="Ullman2005" />{{rp|3}} <!---BaO forms a peroxide when heated in air.<ref name=O2/>--->

[[Group 2 organometallic chemistry#Organobarium|Organobarium compounds]] are a growing field of knowledge: recently discovered are dialkylbariums and alkylhalobariums.<ref name="Ullman2005" />{{rp|3}}

===Isotopes===
{{Main|Isotopes of barium}}

Barium found in the Earth's crust is a mixture of seven [[primordial nuclides]], barium-130, 132, and 134 through 138.<ref name="iso" /> Barium-130 undergoes very slow [[radioactive decay]] to [[xenon]]-130 by double [[Beta decay|beta plus decay]], with a half-life of (0.5–2.7)×10<sup>21</sup> years (about 10<sup>11</sup> times the age of the universe). Its abundance is ≈0.1% that of natural barium.<ref name="iso">{{CIAAW2003}}</ref>  Theoretically, barium-132 can similarly undergo double beta decay to xenon-132; this decay has not been detected.{{NUBASE2016|ref}} The radioactivity of these isotopes is so weak that they pose no danger to life.

Of the stable isotopes, barium-138 composes 71.7% of all barium; other isotopes have decreasing abundance with decreasing [[mass number]].<ref name="iso" />

In total, barium has 40 known isotopes, ranging in mass between 114 and 153. The most stable [[synthetic radioisotope|artificial radioisotope]] is barium-133 with a half-life of approximately 10.51&nbsp;years. Five other isotopes have half-lives longer than a day.{{NUBASE2016|ref}} Barium also has 10 [[meta state]]s, of which barium-133m1 is the most stable with a half-life of about 39 hours.{{NUBASE2016|ref}}<!---<sup>133</sup>Ba is a standard calibrant for [[gamma-ray]] detectors in nuclear physics studies.--->