Editing Alkali metal (section)

== Properties ==

=== Physical and chemical ===
The physical and chemical properties of the alkali metals can be readily explained by their having an ns<sup>1</sup> valence [[electron configuration]], which results in weak [[metallic bonding]]. Hence, all the alkali metals are soft and have low [[densities]],<ref name=rsc /> [[melting point|melting]]<ref name=rsc /> and [[boiling point]]s,<ref name=rsc /> as well as [[heat of sublimation|heats of sublimation]], [[heat of vaporization|vaporisation]], and [[dissociation (chemistry)|dissociation]].<ref name="Greenwood&Earnshaw" />{{rp|74}} They all crystallise in the [[body-centered cubic]] crystal structure,<ref name="Greenwood&Earnshaw" />{{rp|73}} and have distinctive [[flame test|flame colours]] because their outer s electron is very easily excited.<ref name="Greenwood&Earnshaw" />{{rp|75}} Indeed, these flame test colours are the most common way of identifying them since all their salts with common ions are soluble.<ref name="Greenwood&Earnshaw" />{{rp|75}} The ns<sup>1</sup> configuration also results in the alkali metals having very large [[atomic radius|atomic]] and [[ionic radii]], as well as very high [[thermal conductivity|thermal]] and [[electrical conductivity]].<ref name="Greenwood&Earnshaw" />{{rp|75}} Their chemistry is dominated by the loss of their lone valence electron in the outermost s-orbital to form the +1 oxidation state, due to the ease of ionising this electron and the very high second ionisation energy.<ref name="Greenwood&Earnshaw" />{{rp|76}} Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not well established due to its extreme [[radioactivity]];<ref name=rsc /> thus, the presentation of its properties here is limited. What little is known about francium shows that it is very close in behaviour to caesium, as expected. The physical properties of francium are even sketchier because the bulk element has never been observed; hence any data that may be found in the literature are certainly speculative extrapolations.<ref name=RubberBible84th />
{| class="wikitable"
|+ Properties of the alkali metals<ref name="Greenwood&Earnshaw" />{{rp|75}}<ref name=generalchemistry />
! Name
! [[Lithium]]
! [[Sodium]]
! [[Potassium]]
! [[Rubidium]]
! [[Caesium]]
! [[Francium]]
|-
| style="background:lightgrey; text-align:left;"|[[Atomic number]]
| 3 || 11 || 19 || 37 || 55 || 87
|-
| style="background:lightgrey; text-align:left;"|[[Standard atomic weight]]{{refn|The number given in [[bracket|parentheses]] refers to the [[standard uncertainty|measurement uncertainty]]. This uncertainty applies to the [[significant figure|least significant figure]](s) of the number prior to the parenthesised value (ie. counting from rightmost digit to left). For instance, {{val|1.00794|(7)}} stands for {{val|1.00794|0.00007}}, while {{val|1.00794|(72)}} stands for {{val|1.00794|0.00072}}.<ref>{{cite web|url=http://physics.nist.gov/cgi-bin/cuu/Info/Constants/definitions.html|title=Standard Uncertainty and Relative Standard Uncertainty|work=[[CODATA]] reference|publisher=[[National Institute of Standards and Technology]]|access-date=26 September 2011|archive-date=16 October 2011|archive-url=https://web.archive.org/web/20111016021440/http://physics.nist.gov/cgi-bin/cuu/Info/Constants/definitions.html|url-status=live}}</ref>|group=note}}<ref name="atomicweights2007">{{cite journal |last1=Wieser |first1=Michael E. |last2=Berglund |first2=Michael |year=2009 |title=Atomic weights of the elements 2007 (IUPAC Technical Report) |journal=[[Pure Appl. Chem.]] |volume=81 |issue=11 |pages= 2131–2156 |publisher=[[IUPAC]] |doi=10.1351/PAC-REP-09-08-03 |s2cid=98084907 |url=http://iupac.org/publications/pac/pdf/2009/pdf/8111x2131.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://iupac.org/publications/pac/pdf/2009/pdf/8111x2131.pdf |archive-date=2022-10-09 |url-status=live |access-date=7 February 2012}}</ref><ref name="atomicweights2009">{{cite journal |last1=Wieser |first1=Michael E. |last2=Coplen |first2=Tyler B. |year=2011 |title=Atomic weights of the elements 2009 (IUPAC Technical Report) |journal=[[Pure Appl. Chem.]] |volume=83 |issue=2 |pages=359–396 |publisher=[[IUPAC]] |doi=10.1351/PAC-REP-10-09-14 |s2cid=95898322 |url=http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf |archive-date=2022-10-09 |url-status=live |access-date=11 February 2012}}</ref>
| 6.94(1){{refn|The value listed is the conventional value suitable for trade and commerce; the actual value may range from 6.938 to 6.997 depending on the isotopic composition of the sample.<ref name="atomicweights2009" />|group=note}} || 22.98976928(2) || 39.0983(1) || 85.4678(3) || 132.9054519(2) || [223]{{refn|The element does not have any stable [[nuclide]]s, and a value in brackets indicates the [[mass number]] of the longest-lived [[isotope]] of the element.<ref name="atomicweights2007" /><ref name="atomicweights2009" />|group=note}}
|-
| style="background:lightgrey; text-align:left;"|[[Electron configuration]]
| &#91;[[Helium|He]]&#93; 2s<sup>1</sup> || &#91;[[Neon|Ne]]&#93; 3s<sup>1</sup> || &#91;[[Argon|Ar]]&#93; 4s<sup>1</sup> || &#91;[[Krypton|Kr]]&#93; 5s<sup>1</sup> || &#91;[[Xenon|Xe]]&#93; 6s<sup>1</sup> || &#91;[[Radon|Rn]]&#93; 7s<sup>1</sup>
|-
| style="background:lightgrey; text-align:left;"|[[Melting point]] (°C)
| 180.54 || 97.72|| 63.38 || 39.31 || 28.44 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Boiling point]] (°C)
| 1342 || 883 || 759 || 688 || 671 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Density]]&nbsp;(g·cm<sup>−3</sup>)
| 0.534 || 0.968 || 0.89 || 1.532 || 1.93 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Heat of fusion]]&nbsp;(kJ·mol<sup>−1</sup>)
| 3.00 || 2.60 || 2.321 || 2.19 || 2.09 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Heat of vaporisation]]&nbsp;(kJ·mol<sup>−1</sup>)
| 136 || 97.42 || 79.1 || 69 || 66.1 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Heat of formation]]&nbsp;of&nbsp;monatomic&nbsp;gas&nbsp;(kJ·mol<sup>−1</sup>)
| 162 || 108 || 89.6 || 82.0 || 78.2 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Electrical resistivity]]&nbsp;at&nbsp;25&nbsp;°C&nbsp;(n[[ohm|Ω]]·cm)
| 94.7 || 48.8 || 73.9 || 131 || 208 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Atomic radius]]&nbsp;([[picometer|pm]])
| 152 || 186 || 227 || 248 || 265 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Ionic radius]] of hexacoordinate M<sup>+</sup> ion&nbsp;(pm)
| 76 || 102 || 138 || 152 || 167 ||?
|-
| style="background:lightgrey; text-align:left;"|First [[ionisation energy]]&nbsp;([[kilojoule per mole|kJ·mol<sup>−1</sup>]])
| 520.2 || 495.8 || 418.8 || 403.0 || 375.7 || 392.8<ref name="andreev">{{cite journal |last1=Andreev|first1=S.V. |last2=Letokhov|first2=V.S. |last3=Mishin|first3=V.I. |title= Laser resonance photoionization spectroscopy of Rydberg levels in Fr |journal= [[Phys. Rev. Lett.]] |year= 1987 |volume= 59 |pages= 1274–76 |doi= 10.1103/PhysRevLett.59.1274 |pmid=10035190 |bibcode=1987PhRvL..59.1274A |issue= 12}}</ref>
|-
| style="background:lightgrey; text-align:left;"|[[Electron affinity]]&nbsp;(kJ·mol<sup>−1</sup>)
| 59.62 || 52.87 || 48.38 || 46.89 || 45.51 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Enthalpy of dissociation]]&nbsp;of&nbsp;M<sub>2</sub>&nbsp;(kJ·mol<sup>−1</sup>)
| 106.5 || 73.6 || 57.3 || 45.6 || 44.77 ||?
|-
| style="background:lightgrey; text-align:left;"|Pauling [[electronegativity]]
| 0.98 || 0.93 || 0.82 || 0.82 || 0.79 ||?{{refn|[[Linus Pauling]] estimated the electronegativity of francium at 0.7 on the [[Pauling scale]], the same as caesium;<ref>{{cite book |last= Pauling |first= Linus |title= The Nature of the Chemical Bond|url= https://archive.org/details/natureofchemical00paul |url-access= registration |edition= Third |author-link= Linus Pauling |publisher= Cornell University Press |year= 1960 |isbn= 978-0-8014-0333-0 |page= [https://archive.org/details/natureofchemical00paul/page/93 93]}}</ref> the value for caesium has since been refined to 0.79, although there are no experimental data to allow a refinement of the value for francium.<ref>{{cite journal |last=Allred|first=A. L. |year= 1961 |journal= J. Inorg. Nucl. Chem.|volume= 17 |issue= 3–4 |pages= 215–221 |title= Electronegativity values from thermochemical data |doi= 10.1016/0022-1902(61)80142-5}}</ref> Francium has a slightly higher ionisation energy than caesium,<ref name="andreev" /> 392.811(4)&nbsp;kJ/mol as opposed to 375.7041(2)&nbsp;kJ/mol for caesium, as would be expected from [[relativistic effects]], and this would imply that caesium is the less electronegative of the two.|name=Fr-electronegativity|group=note}}
|-
| style="background:lightgrey; text-align:left;"|Allen [[electronegativity]]
|0.91
|0.87
|0.73
|0.71
|0.66
|0.67
|-
| style="background:lightgrey; text-align:left;"|[[Standard electrode potential]] (''E''°(M<sup>+</sup>→M<sup>0</sup>); [[volt|V]])<ref name=van92>Vanýsek, Petr (2011). [http://www.hbcpnetbase.com/articles/05_22_92.pdf “Electrochemical Series”], in [http://www.hbcpnetbase.com/ ''Handbook of Chemistry and Physics: 92nd Edition''] {{Webarchive|url=https://web.archive.org/web/20170724011402/http://www.hbcpnetbase.com/ |date=24 July 2017 }} (Chemical Rubber Company).</ref>
| −3.04 || −2.71 || −2.93 || −2.98 || −3.03 ||?
|-
| style="background:lightgrey; text-align:left;"|[[Flame test]] colour<br />Principal emission/absorption wavelength&nbsp;([[nanometer|nm]])
| Crimson<br />670.8 || Yellow<br />589.2 || Violet<br />766.5 || Red-violet<br />780.0 || Blue<br />455.5 ||?
|}
The alkali metals are more similar to each other than the elements in any other [[group (periodic table)|group]] are to each other.<ref name=rsc /> Indeed, the similarity is so great that it is quite difficult to separate potassium, rubidium, and caesium, due to their similar [[ionic radii]]; lithium and sodium are more distinct. For instance, when moving down the table, all known alkali metals show increasing [[atomic radius]],<ref name=chemguide /> decreasing [[electronegativity]],<ref name=chemguide>{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/properties.html |title=Atomic and Physical Properties of the Group 1 Elements |last=Clark |first=Jim |year=2005 |work=chemguide |access-date=30 January 2012 |archive-date=9 October 2014 |archive-url=https://web.archive.org/web/20141009183758/http://www.chemguide.co.uk/inorganic/group1/properties.html |url-status=live }}</ref> increasing [[Reactivity (chemistry)|reactivity]],<ref name=rsc /> and decreasing melting and boiling points<ref name=chemguide /> as well as heats of fusion and vaporisation.<ref name="Greenwood&Earnshaw" />{{rp|75}} In general, their [[densities]] increase when moving down the table, with the exception that potassium is less dense than sodium.<ref name=chemguide /> One of the very few properties of the alkali metals that does not display a very smooth trend is their [[reduction potential]]s: lithium's value is anomalous, being more negative than the others.<ref name="Greenwood&Earnshaw" />{{rp|75}} This is because the Li<sup>+</sup> ion has a very high [[hydration energy]] in the gas phase: though the lithium ion disrupts the structure of water significantly, causing a higher change in entropy, this high hydration energy is enough to make the reduction potentials indicate it as being the most electropositive alkali metal, despite the difficulty of ionising it in the gas phase.<ref name="Greenwood&Earnshaw" />{{rp|75}}

The stable alkali metals are all silver-coloured metals except for caesium, which has a pale golden tint:<ref name="theodoregray-caesium">{{cite web|url=http://www.theodoregray.com/periodictable/Elements/055/index.s7.html|title=Facts, pictures, stories about the element Cesium in the Periodic Table|last=Gray|first=Theodore|author-link=Theodore Gray|work=The Wooden Periodic Table Table|access-date=13 January 2012|archive-date=28 January 2014|archive-url=https://web.archive.org/web/20140128232712/http://www.theodoregray.com/periodictable/Elements/055/index.s7.html|url-status=live}}</ref> it is one of only three metals that are clearly coloured (the other two being copper and gold).<ref name="Greenwood&Earnshaw" />{{rp|74}} Additionally, the heavy [[alkaline earth metal]]s [[calcium]], [[strontium]], and [[barium]], as well as the divalent [[lanthanide]]s [[europium]] and [[ytterbium]], are pale yellow, though the colour is much less prominent than it is for caesium.<ref name="Greenwood&Earnshaw" />{{rp|74}} Their lustre tarnishes rapidly in air due to oxidation.<ref name=rsc />

[[File:Potassium water 20.theora.ogv|thumb|right|Potassium reacts violently with water at room temperature]]
[[File:Cesium water.theora.ogv|thumb|right|Caesium reacts explosively with water even at low temperatures]]
All the alkali metals are highly reactive and are never found in elemental forms in nature.<ref name="krebs" /> Because of this, they are usually stored in [[mineral oil]] or [[kerosene]] (paraffin oil).<ref name="OU">{{cite web |url=http://www.open.edu/openlearn/science-maths-technology/science/chemistry/alkali-metals |title=Alkali metals |author=The OpenLearn team |year=2012 |work=OpenLearn |publisher=The Open University |access-date=9 July 2012 |archive-date=29 November 2014 |archive-url=https://web.archive.org/web/20141129052111/http://www.open.edu/openlearn/science-maths-technology/science/chemistry/alkali-metals |url-status=live }}</ref> They react aggressively with the [[halogen]]s to form the [[alkali metal halide]]s, which are white [[ionic crystal]]line compounds that are all [[soluble]] in water except [[lithium fluoride]] (LiF).<ref name=rsc /> The alkali metals also react with water to form strongly [[alkali]]ne [[hydroxide]]s and thus should be handled with great care. The heavier alkali metals react more vigorously than the lighter ones; for example, when dropped into water, caesium produces a larger explosion than potassium if the same number of moles of each metal is used.<ref name=rsc /><ref name="alkalibangs">{{cite web|last=Gray|first=Theodore|title=Alkali Metal Bangs|url=http://www.theodoregray.com/periodictable/AlkaliBangs/index.html|publisher=[[Theodore Gray]]|access-date=13 May 2012|archive-date=31 October 2014|archive-url=https://web.archive.org/web/20141031113655/http://www.theodoregray.com/PeriodicTable/AlkaliBangs/index.html|url-status=live}}</ref><ref name="pubs.usgs" /> The alkali metals have the lowest first [[ionisation energies]] in their respective periods of the [[periodic table]]<ref name="RubberBible84th">{{cite book |editor= Lide, D. R. |title= CRC Handbook of Chemistry and Physics |edition= 84th |location= Boca Raton, FL |publisher= CRC Press |year= 2003}}</ref> because of their low [[effective nuclear charge]]<ref name=rsc /> and the ability to attain a [[noble gas]] configuration by losing just one [[electron]].<ref name=rsc /> Not only do the alkali metals react with water, but also with proton donors like [[Alcohol (chemistry)|alcohol]]s and [[phenols]], gaseous [[ammonia]], and [[alkyne]]s, the last demonstrating the phenomenal degree of their reactivity. Their great power as reducing agents makes them very useful in liberating other metals from their oxides or halides.<ref name="Greenwood&Earnshaw" />{{rp|76}}

The second ionisation energy of all of the alkali metals is very high<ref name=rsc /><ref name="RubberBible84th" /><!--the second ionisation energy for francium is not given in [[ionization energies of the elements (data page)]]--> as it is in a full shell that is also closer to the nucleus;<ref name=rsc /> thus, they almost always lose a single electron, forming cations.<ref name="Greenwood&Earnshaw" />{{rp|28}} The [[alkalide]]s are an exception: they are unstable compounds which contain alkali metals in a −1 oxidation state, which is very unusual as before the discovery of the alkalides, the alkali metals were not expected to be able to form [[anion]]s and were thought to be able to appear in [[salts]] only as cations. The alkalide anions have filled [[s-orbital|s-subshells]], which gives them enough stability to exist. All the stable alkali metals except lithium are known to be able to form alkalides,<ref>{{cite journal |journal= [[J. Am. Chem. Soc.]] |title= Crystalline salt of the sodium anion (Na<sup>−</sup>) |year= 1974 |volume= 96 |issue= 2 |pages= 608–609 |doi= 10.1021/ja00809a060|last1= Dye |first1= James L. |last2= Ceraso |first2= Joseph M. |last3= Lok |first3= Mei |last4= Barnett |first4= B. L. |last5= Tehan |first5= Frederick J. |bibcode= 1974JAChS..96..608D }}</ref><ref>{{cite journal |title= Alkali anions. Preparation and crystal structure of a compound which contains the cryptated sodium cation and the sodium anion |journal= [[J. Am. Chem. Soc.]] |year= 1974 |volume= 96 |issue= 23 |pages= 7203–7208 |doi= 10.1021/ja00830a005|last1= Tehan |first1= Frederick J. |last2= Barnett |first2= B. L. |last3= Dye |first3= James L. |bibcode= 1974JAChS..96.7203T }}</ref><ref>{{cite journal |journal= [[Angew. Chem. Int. Ed. Engl.]] |year= 1979 |last=Dye|first=J. L. |title= Compounds of Alkali Metal Anions |volume= 18 |issue= 8 |pages= 587–598 |doi= 10.1002/anie.197905871}}</ref> and the alkalides have much theoretical interest due to their unusual [[stoichiometry]] and low [[ionization potential|ionisation potentials]]. Alkalides are chemically similar to the [[electride]]s, which are salts with trapped [[electron]]s acting as anions.<ref name="Redko">{{cite journal |year= 2003 |title= Barium azacryptand sodide, the first alkalide with an alkaline Earth cation, also contains a novel dimer, (Na<sub>2</sub>)<sup>2−</sup> |journal= [[J. Am. Chem. Soc.]] |volume= 125 |issue= 8 |pages= 2259–2263 |doi= 10.1021/ja027241m |pmid= 12590555 |url= https://www.researchgate.net/publication/10896204 |last1= Redko |first1= M. Y. |last2= Huang |first2= R. H. |last3= Jackson |first3= J. E. |last4= Harrison |first4= J. F. |last5= Dye |first5= J. L. |bibcode= 2003JAChS.125.2259R |archive-date= 25 April 2018 |access-date= 16 November 2016 |archive-url= https://web.archive.org/web/20180425183546/https://www.researchgate.net/publication/10896204 |url-status= live }}</ref> A particularly striking example of an alkalide is "inverse [[sodium hydride]]", H<sup>+</sup>Na<sup>−</sup> (both ions being [[coordination complex|complexed]]), as opposed to the usual sodium hydride, Na<sup>+</sup>H<sup>−</sup>:<ref name="HNa">{{cite journal |year= 2002 |title="Inverse sodium hydride": a crystalline salt that contains H<sup>+</sup> and Na<sup>−</sup> |journal= [[J. Am. Chem. Soc.]] |volume= 124 |issue= 21 |pages= 5928–5929 |doi= 10.1021/ja025655+|pmid=12022811 |last1=Redko |first1=M. Y. |last2=Vlassa |first2=M. |last3=Jackson |first3=J. E. |last4=Misiolek |first4=A. W. |last5=Huang |first5=R. H. |last6=Dye |first6=J. L. }}</ref> it is unstable in isolation, due to its high energy resulting from the displacement of two electrons from hydrogen to sodium, although several derivatives are predicted to be [[metastable]] or stable.<ref name="HNa" /><ref name="HNa-theory">{{cite journal |url=http://simons.hec.utah.edu/papers/266.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://simons.hec.utah.edu/papers/266.pdf |archive-date=2022-10-09 |url-status=live|title=Inverse Sodium Hydride: A Theoretical Study|journal=J. Am. Chem. Soc.|year=2003|volume=125|pages=3954–3958|doi=10.1021/ja021136v|pmid=12656631|issue=13|last1=Sawicka|first1=A.|last2=Skurski|first2=P.|last3=Simons|first3=J.|bibcode=2003JAChS.125.3954S }}</ref>

In aqueous solution, the alkali metal ions form [[metal ions in aqueous solution|aqua ions]] of the formula [M(H<sub>2</sub>O)<sub>''n''</sub>]<sup>+</sup>, where ''n'' is the solvation number. Their [[coordination number]]s and shapes agree well with those expected from their ionic radii. In aqueous solution the water molecules directly attached to the metal ion are said to belong to the [[first coordination sphere]], also known as the first, or primary, solvation shell. The bond between a water molecule and the metal ion is a [[dative covalent bond]], with the oxygen atom donating both electrons to the bond. Each coordinated water molecule may be attached by [[hydrogen bond]]s to other water molecules. The latter are said to reside in the second coordination sphere. However, for the alkali metal cations, the second coordination sphere is not well-defined as the +1 charge on the cation is not high enough to [[Polarizability|polarise]] the water molecules in the primary solvation shell enough for them to form strong hydrogen bonds with those in the second coordination sphere, producing a more stable entity.<ref>{{cite book |last=Burgess |first=John |title=Metal Ions in Solution |year=1978 |publisher=Ellis Horwood |location=Chichester |page=20 |isbn=978-0-85312-027-8}}</ref><ref name=Richens />{{rp|25}} The solvation number for Li<sup>+</sup> has been experimentally determined to be 4, forming the [[tetrahedral]] [Li(H<sub>2</sub>O)<sub>4</sub>]<sup>+</sup>: while solvation numbers of 3 to 6 have been found for lithium aqua ions, solvation numbers less than 4 may be the result of the formation of contact [[ion pair]]s, and the higher solvation numbers may be interpreted in terms of water molecules that approach [Li(H<sub>2</sub>O)<sub>4</sub>]<sup>+</sup> through a face of the tetrahedron, though molecular dynamic simulations may indicate the existence of an [[octahedral]] hexaaqua ion. There are also probably six water molecules in the primary solvation sphere of the sodium ion, forming the octahedral [Na(H<sub>2</sub>O)<sub>6</sub>]<sup>+</sup> ion.<ref name=generalchemistry /><ref name=Richens>{{cite book |last=Richens |first=David. T. |title=The Chemistry of Aqua Ions |year=1997 |publisher=Wiley |isbn=978-0-471-97058-3}}</ref>{{rp|126–127}} While it was previously thought that the heavier alkali metals also formed octahedral hexaaqua ions, it has since been found that potassium and rubidium probably form the [K(H<sub>2</sub>O)<sub>8</sub>]<sup>+</sup> and [Rb(H<sub>2</sub>O)<sub>8</sub>]<sup>+</sup> ions, which have the [[square antiprism]]atic structure, and that caesium forms the 12-coordinate [Cs(H<sub>2</sub>O)<sub>12</sub>]<sup>+</sup> ion.<ref>{{cite journal |last=Persson |first=Ingmar |date=2010 |title=Hydrated metal ions in aqueous solution: How regular are their structures? |url=http://pac.iupac.org/publications/pac/pdf/2010/pdf/8210x1901.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://pac.iupac.org/publications/pac/pdf/2010/pdf/8210x1901.pdf |archive-date=2022-10-09 |url-status=live |journal=Pure Appl. Chem. |volume=82 |issue=10 |pages=1901–1917 |doi=10.1351/PAC-CON-09-10-22 |s2cid=98411500 |access-date=23 August 2014}}</ref>
{{clear left}}

==== Lithium ====
The chemistry of lithium shows several differences from that of the rest of the group as the small Li<sup>+</sup> cation [[chemical polarity|polarises]] [[anion]]s and gives its compounds a more [[covalent]] character.<ref name=rsc /> Lithium and [[magnesium]] have a [[diagonal relationship]] due to their similar atomic radii,<ref name=rsc /> so that they show some similarities. For example, lithium forms a stable [[nitride]], a property common among all the [[alkaline earth metal]]s (magnesium's group) but unique among the alkali metals.<ref name="alkalireact" /> In addition, among their respective groups, only lithium and magnesium form [[organometallic compound]]s with significant covalent character (e.g. Li[[methyl group|Me]] and MgMe<sub>2</sub>).<ref name="Shriver&Atkins">{{cite book |title=Inorganic Chemistry |first1=Duward |last1=Shriver |first2=Peter |last2=Atkins |publisher=W. H. Freeman |year=2006 |isbn=978-0-7167-4878-6 |page=259 |access-date=10 November 2012 |url=https://books.google.com/books?id=NwOTQAAACAAJ }}{{Dead link|date=August 2023 |bot=InternetArchiveBot |fix-attempted=yes }}</ref>

Lithium fluoride is the only alkali metal halide that is poorly soluble in water,<ref name=rsc /> and [[lithium hydroxide]] is the only alkali metal hydroxide that is not [[deliquescent]].<ref name=rsc /> Conversely, [[lithium perchlorate]] and other lithium salts with large anions that cannot be polarised are much more stable than the analogous compounds of the other alkali metals, probably because Li<sup>+</sup> has a high [[solvation energy]].<ref name="Greenwood&Earnshaw" />{{rp|76}} This effect also means that most simple lithium salts are commonly encountered in hydrated form, because the anhydrous forms are extremely [[hygroscopic]]: this allows salts like [[lithium chloride]] and [[lithium bromide]] to be used in [[dehumidifier]]s and [[air-conditioner]]s.<ref name="Greenwood&Earnshaw" />{{rp|76}}

==== Francium ====
Francium is also predicted to show some differences due to its high [[atomic weight]], causing its electrons to travel at considerable fractions of the speed of light and thus making [[relativistic effects]] more prominent. In contrast to the trend of decreasing [[electronegativities]] and [[ionisation energies]] of the alkali metals, francium's electronegativity and ionisation energy are predicted to be higher than caesium's due to the relativistic stabilisation of the 7s electrons; also, its [[atomic radius]] is expected to be abnormally low.<!--Haire says this happens for Uue because of the analogous effect for 8s – seems likely for Fr too--> Thus, contrary to expectation, caesium is the most reactive of the alkali metals, not francium.<ref name="andreev" /><ref name="Uue">{{cite book |title= The Chemistry of the Actinide and Transactinide Elements |editor1-last= Morss|editor2-first= Norman M. |editor2-last= Edelstein |editor3-last= Fuger|editor3-first= Jean |last1= Hoffman|first1= Darleane C. |last2=Lee|first2=Diana M. |last3=Pershina|first3=Valeria |chapter= Transactinides and the future elements |publisher= Springer |year= 2006 |isbn= 978-1-4020-3555-5 |location= Dordrecht, The Netherlands |edition= 3rd }}</ref>{{rp|1729}}<ref name=Thayer /> All known physical properties of francium also deviate from the clear trends going from lithium to caesium, such as the first ionisation energy, electron affinity, and anion polarisability, though due to the paucity of known data about francium many sources give extrapolated values, ignoring that relativistic effects make the trend from lithium to caesium become inapplicable at francium.<ref name=Thayer /> Some of the few properties of francium that have been predicted taking relativity into account are the electron affinity (47.2 kJ/mol)<ref name=Landaualkalis>{{cite journal |last1= Landau |first1= A. |last2= Eliav |first2= E. |last3= Ishikawa |first3= Y. |last4= Kaldor |first4= U. |year= 2001 |title= Benchmark calculations of electron affinities of the alkali atoms sodium to eka-francium (element 119) |url= https://www.academia.edu/20466410 |journal= J. Chem. Phys. |volume= 115 |issue= 6 |page= 2389 |doi= 10.1063/1.1386413 |bibcode= 2001JChPh.115.2389L |archive-date= 31 July 2020 |access-date= 16 November 2016 |archive-url= https://web.archive.org/web/20200731131615/https://www.academia.edu/20466410/Benchmark_calculations_of_electron_affinities_of_the_alkali_atoms_sodium_to_eka-francium_element_119_ |url-status= live }}</ref> and the enthalpy of dissociation of the Fr<sub>2</sub> molecule (42.1 kJ/mol).<ref name=Liddle>{{cite book |last1=Jones |first1=Cameron |last2=Mountford |first2=Philip |last3=Stasch |first3=Andreas |last4=Blake |first4=Matthew P. |editor-last=Liddle |editor-first=Stephen T. |title=Molecular Metal-Metal Bonds: Compounds, Synthesis, Properties |publisher=John Wiley and Sons |date=22 June 2015 |pages=23–24 |chapter=s-block Metal-Metal Bonds |isbn=978-3-527-33541-1}}</ref> The CsFr molecule is polarised as Cs<sup>+</sup>Fr<sup>−</sup>, showing that the 7s subshell of francium is much more strongly affected by relativistic effects than the 6s subshell of caesium.<ref name=Thayer /> Additionally, francium superoxide (FrO<sub>2</sub>) is expected to have significant covalent character, unlike the other alkali metal superoxides, because of bonding contributions from the 6p electrons of francium.<ref name=Thayer />

=== Nuclear ===
<div style="float: right; margin: 5px;">
{| class="sortable wikitable" style="text-align:center;"
|+Primordial isotopes of the alkali metals
|-
! Z<br />
! Alkali metal<br />
! <small>[[stable isotope|Stable]]</small><br />
! <small>''[[primordial element|Decays]]''</small><br />
! class="unsortable" colspan="3"|<small>''unstable: italics''<div style="background:pink">odd–odd isotopes coloured pink</div></small>
|-
|  3 ||[[lithium]]      || [[isotopes of lithium|2]]    || — || {{SimpleNuclide|lithium|7}}|| style="background:pink;"|{{SimpleNuclide|lithium|6}}||&nbsp;
|-
| 11 ||[[sodium]]       || [[isotopes of sodium|1]]     || — ||{{SimpleNuclide|sodium|23}}||&nbsp;||&nbsp;
|-
| 19 ||[[potassium]]    || [[isotopes of potassium|2]]  || 1 ||{{SimpleNuclide|potassium|39}}||{{SimpleNuclide|potassium|41}}|| style="background:pink;"|''{{SimpleNuclide|potassium|40}}''
|-
| 37 ||[[rubidium]]     || [[isotopes of rubidium|1]]   || 1 ||{{SimpleNuclide|rubidium|85}}|||''{{SimpleNuclide|rubidium|87}}''||&nbsp;
|-
| 55 ||[[caesium]]      || [[isotopes of caesium|1]]    || — ||{{SimpleNuclide|caesium|133}}||&nbsp;||&nbsp;
|-
| 87 ||[[francium]]     || [[isotopes of francium|—]]   || — ||colspan="3"|''No primordial isotopes''<br />(''{{SimpleNuclide|francium|223}}'' is a [[radiogenic nuclide]])
|-
| colspan="7"|<small>Radioactive: {{nowrap|<sup>40</sup>K, [[half-life|t<sub>1/2</sub>]] 1.25 × 10<sup>9</sup> years;}} {{nowrap|<sup>87</sup>Rb, t<sub>1/2</sub> 4.9 × 10<sup>10</sup> years;}} {{nowrap|<sup>223</sup>Fr, t<sub>1/2</sub> 22.0 min.}}</small>
|}</div>
All the alkali metals have odd atomic numbers; hence, their isotopes must be either [[odd–odd nuclei|odd–odd]] (both proton and [[neutron number]] are odd) or [[odd–even nuclei|odd–even]] ([[proton number]] is odd, but neutron number is even). Odd–odd nuclei have even [[mass number]]s, whereas odd–even nuclei have odd mass numbers. Odd–odd [[primordial nuclide]]s are rare because most odd–odd nuclei are highly unstable with respect to [[beta decay]], because the decay products are even–even, and are therefore more strongly bound, due to [[Semi-empirical mass formula#Pairing term|nuclear pairing effects]].<ref name="Lide02">{{cite book |editor-last=Lide |editor-first=David R. |year=2002 |title=Handbook of Chemistry & Physics |edition=88th |publisher=CRC |url=http://www.hbcpnetbase.com/ |access-date=2008-05-23 |isbn=978-0-8493-0486-6 |oclc=179976746 |archive-date=24 July 2017 |archive-url=https://web.archive.org/web/20170724011402/http://www.hbcpnetbase.com/ |url-status=dead }}</ref>

Due to the great rarity of odd–odd nuclei, almost all the primordial isotopes of the alkali metals are odd–even (the exceptions being the light stable isotope lithium-6 and the long-lived [[radioisotope]] potassium-40). For a given odd mass number, there can be only a single [[beta-decay stable isobars|beta-stable nuclide]], since there is not a difference in binding energy between even–odd and odd–even comparable to that between even–even and odd–odd, leaving other nuclides of the same mass number ([[isobar (nuclide)|isobars]]) free to [[beta decay]] toward the lowest-mass nuclide. An effect of the instability of an odd number of either type of nucleons is that odd-numbered elements, such as the alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 [[monoisotopic element]]s that have only a single stable isotope, all but one have an odd atomic number and all but one also have an even number of neutrons. [[Beryllium]] is the single exception to both rules, due to its low atomic number.<ref name="Lide02" />

All of the alkali metals except lithium and caesium have at least one naturally occurring [[radioisotope]]: [[sodium-22]] and [[sodium-24]] are [[trace radioisotope]]s produced [[cosmogenic]]ally,<ref>{{cite web |url=http://www.nucleonica.net/unc.aspx |title=Universal Nuclide Chart |date=2007–2012 |work=Nucleonica |publisher=Institute for Transuranium Elements |access-date=2011-04-17 |archive-date=19 February 2017 |archive-url=https://web.archive.org/web/20170219043412/http://www.nucleonica.net/unc.aspx |url-status=dead }}</ref> potassium-40 and [[rubidium-87]] have very long [[half-lives]] and thus occur naturally,<ref name="nuclideschart" /> and all [[isotopes of francium]] are [[radioactive]].<ref name="nuclideschart" /> Caesium was also thought to be radioactive in the early 20th century,<ref name="Patt1926">{{cite journal |doi= 10.1021/cr60009a003 |title= The Radioactivity of the Alkali Metals |year= 1926 |last1= Patton |first1= I. Jocelyn |last2= Waldbauer |first2= L. J. |journal= Chemical Reviews |volume= 3 |pages= 81–93}}</ref><ref name="Kenn1908">{{cite journal |doi= 10.1080/14786440908636519 |title= On the radioactivity of potassium and other alkali metals |year= 1908 |last1= McLennan |first1= J. C. |last2= Kennedy |first2= W. T. |journal= [[Philosophical Magazine]] |series= 6 |volume= 16 |issue= 93 |pages= 377–395 |url= https://zenodo.org/record/1430860 |archive-date= 28 October 2020 |access-date= 28 June 2019 |archive-url= https://web.archive.org/web/20201028103710/https://zenodo.org/record/1430860 |url-status= live }}</ref> although it has no naturally occurring radioisotopes.<ref name="nuclideschart">{{cite web|url=http://www.nndc.bnl.gov/chart/|title=Interactive Chart of Nuclides|publisher=Brookhaven National Laboratory|last=Sonzogni|first=Alejandro|location=National Nuclear Data Center|access-date=4 October 2012|archive-date=21 July 2011|archive-url=https://web.archive.org/web/20110721051025/http://www.nndc.bnl.gov/chart/|url-status=dead}}</ref> (Francium had not been discovered yet at that time.) The natural long-lived radioisotope of potassium, potassium-40, makes up about 0.012% of natural potassium,<ref>{{cite web |url=http://www.ead.anl.gov/pub/doc/potassium.pdf |archive-url=https://ghostarchive.org/archive/20221009/http://www.ead.anl.gov/pub/doc/potassium.pdf |archive-date=2022-10-09 |url-status=live |title=Potassium-40 |date=August 2005 |work=Human Health Fact Sheet |publisher=[[Argonne National Laboratory]], Environmental Science Division |access-date=7 February 2012}}</ref> and thus natural potassium is weakly radioactive. This natural radioactivity became a basis for a mistaken claim of the discovery for element 87 (the next alkali metal after caesium) in 1925.<ref name="fontani" /><ref name="vanderkrogt-Fr">{{cite web |last= Van der Krogt |first= Peter |title= Francium |work= Elementymology & Elements Multidict |date= 10 January 2006 |url= http://elements.vanderkrogt.net/element.php?sym=Fr |access-date= 8 April 2007 |archive-date= 23 January 2010 |archive-url= https://web.archive.org/web/20100123003337/http://elements.vanderkrogt.net/element.php?sym=Fr |url-status= live }}</ref> Natural rubidium is similarly slightly radioactive, with 27.83% being the long-lived radioisotope rubidium-87.<ref name="Greenwood&Earnshaw" />{{rp|74}}

[[Caesium-137]], with a half-life of 30.17&nbsp;years, is one of the two principal [[medium-lived fission product]]s, along with [[strontium-90]], which are responsible for most of the [[radioactivity]] of [[spent nuclear fuel]] after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left from the [[Chernobyl accident]]. Caesium-137 undergoes high-energy beta decay and eventually becomes stable [[barium-137]]. It is a strong emitter of gamma radiation. Caesium-137 has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay.<ref name="Cs-137">{{cite web|title=Radionuclide Half-Life Measurements|url=https://www.nist.gov/pml/data/halflife-html.cfm|author=National Institute of Standards and Technology|date=6 September 2009|access-date=2011-11-07|archive-date=12 August 2016|archive-url=https://web.archive.org/web/20160812133216/http://nist.gov/pml/data/halflife-html.cfm|url-status=dead}}</ref> Caesium-137 has been used as a [[Flow tracer|tracer]] in hydrologic studies, analogous to the use of [[tritium]].<ref>[https://web.archive.org/web/20160329120038/http://www.bt.cdc.gov/radiation/isotopes/cesium.asp Radioisotope Brief: Cesium-137 (Cs-137)]. U.S. National Center for Environmental Health</ref> Small amounts of [[caesium-134]] and caesium-137 were released into the environment during nearly all [[nuclear weapon test]]s and some [[nuclear accident]]s, most notably the [[Goiânia accident]] and the [[Chernobyl disaster]]. As of 2005, caesium-137 is the principal source of radiation in the [[zone of alienation]] around the [[Chernobyl nuclear power plant]].<ref name="IAEA">{{cite book |title=The Radiological Accident in Goiânia |publisher=[[IAEA]] |year=1988 |url=http://www-pub.iaea.org/MTCD/publications/PubDetAR.asp?pubId=3684 |archive-date=20 January 2011 |access-date=12 January 2013 |archive-url=https://web.archive.org/web/20110120085823/http://www-pub.iaea.org/MTCD/publications/PubDetAR.asp?pubId=3684 |url-status=live }}</ref> Its chemical properties as one of the alkali metals make it one of the most problematic of the short-to-medium-lifetime fission products because it easily moves and spreads in nature due to the high water solubility of its salts, and is taken up by the body, which mistakes it for its essential congeners sodium and potassium.<ref name="RPD">{{cite book |title=Radionuclide and Radiation Protection Data Handbook 2002 |isbn=978-1-870965-87-3 |year=2002 |first1=D. |last1=Delacroix |first2=J. P. |last2=Guerre |first3=P. |last3=Leblanc |first4=C. |last4=Hickman |publisher=Nuclear Technology Publishing |edition=2nd}}</ref>{{rp|114}}