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=                            Alkali_metal                            =
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                            Introduction
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↓ Period
2
3
4
5
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colspan="2"|'Legend'    {
|Primordial}}; background:;" | primordial
|from decay}};  background:;padding:0 2px;" | element by radioactive
decay
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The alkali metals consist of the chemical elements lithium (Li),
sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium
(Fr). Together with hydrogen they constitute group 1, which lies in
the s-block of the periodic table. All alkali metals have their
outermost electron in an s-orbital: this shared electron configuration
results in their having very similar characteristic properties.
Indeed, the alkali metals provide the best example of group trends in
properties in the periodic table, with elements exhibiting
well-characterised homologous behaviour. This family of elements is
also known as the lithium family after its leading element.

The alkali metals are all shiny, soft, highly reactive metals at
standard temperature and pressure and readily lose their outermost
electron to form cations with charge +1. They can all be cut easily
with a knife due to their softness, exposing a shiny surface that
tarnishes rapidly in air due to oxidation by atmospheric moisture and
oxygen (and in the case of lithium, nitrogen). Because of their high
reactivity, they must be stored under oil to prevent reaction with
air, and are found naturally only in salts and never as the free
elements. Caesium, the fifth alkali metal, is the most reactive of all
the metals. All the alkali metals react with water, with the heavier
alkali metals reacting more vigorously than the lighter ones.

All of the discovered alkali metals occur in nature as their
compounds: in order of abundance, sodium is the most abundant,
followed by potassium, lithium, rubidium, caesium, and finally
francium, which is very rare due to its extremely high radioactivity;
francium occurs only in minute traces in nature as an intermediate
step in some obscure side branches of the natural decay chains.
Experiments have been conducted to attempt the synthesis of element
119, which is likely to be the next member of the group; none were
successful. However, ununennium may not be an alkali metal due to
relativistic effects, which are predicted to have a large influence on
the chemical properties of superheavy elements; even if it does turn
out to be an alkali metal, it is predicted to have some differences in
physical and chemical properties from its lighter homologues.

Most alkali metals have many different applications. One of the
best-known applications of the pure elements is the use of rubidium
and caesium in atomic clocks, of which caesium atomic clocks form the
basis of the second. A common application of the compounds of sodium
is the sodium-vapour lamp, which emits light very efficiently. Table
salt, or sodium chloride, has been used since antiquity. Lithium finds
use as a psychiatric medication and as an anode in lithium batteries.
Sodium, potassium and possibly lithium are essential elements, having
major biological roles as electrolytes, and although the other alkali
metals are not essential, they also have various effects on the body,
both beneficial and harmful.
__TOC__


                              History
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Sodium compounds have been known since ancient times; salt (sodium
chloride) has been an important commodity in human activities. While
potash has been used since ancient times, it was not understood for
most of its history to be a fundamentally different substance from
sodium mineral salts. Georg Ernst Stahl obtained experimental evidence
which led him to suggest the fundamental difference of sodium and
potassium salts in 1702, and Henri-Louis Duhamel du Monceau was able
to prove this difference in 1736. The exact chemical composition of
potassium and sodium compounds, and the status as chemical element of
potassium and sodium, was not known then, and thus Antoine Lavoisier
did not include either alkali in his list of chemical elements in
1789.

Pure potassium was first isolated in 1807 in England by Humphry Davy,
who derived it from caustic potash (KOH, potassium hydroxide) by the
use of electrolysis of the molten salt with the newly invented voltaic
pile. Previous attempts at electrolysis of the aqueous salt were
unsuccessful due to potassium's extreme reactivity. Potassium was the
first metal that was isolated by electrolysis. Later that same year,
Davy reported extraction of sodium from the similar substance caustic
soda (NaOH, lye) by a similar technique, demonstrating the elements,
and thus the salts, to be different.

Petalite () was discovered in 1800 by the Brazilian chemist José
Bonifácio de Andrada in a mine on the island of Utö, Sweden. However,
it was not until 1817 that Johan August Arfwedson, then working in the
laboratory of the chemist Jöns Jacob Berzelius, detected the presence
of a new element while analysing petalite ore. This new element was
noted by him to form compounds similar to those of sodium and
potassium, though its carbonate and hydroxide were less soluble in
water and more alkaline than the other alkali metals. Berzelius gave
the unknown material the name 'lithion'/'lithina', from the Greek word
'λιθoς' (transliterated as 'lithos', meaning "stone"), to reflect its
discovery in a solid mineral, as opposed to potassium, which had been
discovered in plant ashes, and sodium, which was known partly for its
high abundance in animal blood. He named the metal inside the material
'lithium'. Lithium, sodium, and potassium were part of the discovery
of periodicity, as they are among a series of triads of elements in
the same group that were noted by Johann Wolfgang Döbereiner in 1850
as having similar properties.

Rubidium and caesium were the first elements to be discovered using
the spectroscope, invented in 1859 by Robert Bunsen and Gustav
Kirchhoff. The next year, they discovered caesium in the mineral water
from Bad Dürkheim, Germany. Their discovery of rubidium came the
following year in Heidelberg, Germany, finding it in the mineral
lepidolite. The names of rubidium and caesium come from the most
prominent lines in their emission spectra: a bright red line for
rubidium (from the Latin word 'rubidus', meaning dark red or bright
red), and a sky-blue line for caesium (derived from the Latin word
'caesius', meaning sky-blue).

Around 1865 John Newlands produced a series of papers where he listed
the elements in order of increasing atomic weight and similar physical
and chemical properties that recurred at intervals of eight; he
likened such periodicity to the octaves of music, where notes an
octave apart have similar musical functions. His version put all the
alkali metals then known (lithium to caesium), as well as copper,
silver, and thallium (which show the +1 oxidation state characteristic
of the alkali metals), together into a group. His table placed
hydrogen with the halogens.

After 1869, Dmitri Mendeleev proposed his periodic table placing
lithium at the top of a group with sodium, potassium, rubidium,
caesium, and thallium. Two years later, Mendeleev revised his table,
placing hydrogen in group 1 above lithium, and also moving thallium to
the boron group. In this 1871 version, copper, silver, and gold were
placed twice, once as part of group IB, and once as part of a "group
VIII" encompassing today's groups 8 to 11. After the introduction of
the 18-column table, the group IB elements were moved to their current
position in the d-block, while alkali metals were left in 'group IA'.
Later the group's name was changed to 'group 1' in 1988. The trivial
name "alkali metals" comes from the fact that the hydroxides of the
group 1 elements are all strong alkalis when dissolved in water.

There were at least four erroneous and incomplete discoveries before
Marguerite Perey of the Curie Institute in Paris, France discovered
francium in 1939 by purifying a sample of actinium-227, which had been
reported to have a decay energy of 220 keV. However, Perey noticed
decay particles with an energy level below 80 keV. Perey thought this
decay activity might have been caused by a previously unidentified
decay product, one that was separated during purification, but emerged
again out of the pure actinium-227. Various tests eliminated the
possibility of the unknown element being thorium, radium, lead,
bismuth, or thallium. The new product exhibited chemical properties of
an alkali metal (such as coprecipitating with caesium salts), which
led Perey to believe that it was element 87, caused by the alpha decay
of actinium-227. Perey then attempted to determine the proportion of
beta decay to alpha decay in actinium-227. Her first test put the
alpha branching at 0.6%, a figure that she later revised to 1%.
:

The next element below francium (eka-francium) in the periodic table
would be ununennium (Uue), element 119. The synthesis of ununennium
was first attempted in 1985 by bombarding a target of einsteinium-254
with calcium-48 ions at the superHILAC accelerator at the Lawrence
Berkeley National Laboratory in Berkeley, California. No atoms were
identified, leading to a limiting yield of 300 nb.

: +  → * → 'no atoms'

It is highly unlikely that this reaction will be able to create any
atoms of ununennium in the near future, given the extremely difficult
task of making sufficient amounts of einsteinium-254, which is
favoured for production of ultraheavy elements because of its large
mass, relatively long half-life of 270 days, and availability in
significant amounts of several micrograms, to make a large enough
target to increase the sensitivity of the experiment to the required
level; einsteinium has not been found in nature and has only been
produced in laboratories, and in quantities smaller than those needed
for effective synthesis of superheavy elements. However, given that
ununennium is only the first period 8 element on the extended periodic
table, it may well be discovered in the near future through other
reactions, and indeed an attempt to synthesise it is currently ongoing
in Japan. Currently, none of the period 8 elements has been discovered
yet, and it is also possible, due to drip instabilities, that only the
lower period 8 elements, up to around element 128, are physically
possible. No attempts at synthesis have been made for any heavier
alkali metals: due to their extremely high atomic number, they would
require new, more powerful methods and technology to make.


In the Solar System
=====================
The Oddo-Harkins rule holds that elements with even atomic numbers are
more common that those with odd atomic numbers, with the exception of
hydrogen. This rule argues that elements with odd atomic numbers have
one unpaired proton and are more likely to capture another, thus
increasing their atomic number. In elements with even atomic numbers,
protons are paired, with each member of the pair offsetting the spin
of the other, enhancing stability. All the alkali metals have odd
atomic numbers and they are not as common as the elements with even
atomic numbers adjacent to them (the noble gases and the alkaline
earth metals) in the Solar System. The heavier alkali metals are also
less abundant than the lighter ones as the alkali metals from rubidium
onward can only be synthesised in supernovae and not in stellar
nucleosynthesis. Lithium is also much less abundant than sodium and
potassium as it is poorly synthesised in both Big Bang nucleosynthesis
and in stars: the Big Bang could only produce trace quantities of
lithium, beryllium and boron due to the absence of a stable nucleus
with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this
bottleneck by the triple-alpha process, fusing three helium nuclei to
form carbon, and skipping over those three elements.


On Earth
==========
The Earth formed from the same cloud of matter that formed the Sun,
but the planets acquired different compositions during the formation
and evolution of the Solar System. In turn, the natural history of the
Earth caused parts of this planet to have differing concentrations of
the elements. The mass of the Earth is approximately 5.98 kg. It is
composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%),
magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and
aluminium (1.4%); with the remaining 1.2% consisting of trace amounts
of other elements. Due to planetary differentiation, the core region
is believed to be primarily composed of iron (88.8%), with smaller
amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace
elements.

The alkali metals, due to their high reactivity, do not occur
naturally in pure form in nature. They are lithophiles and therefore
remain close to the Earth's surface because they combine readily with
oxygen and so associate strongly with silica, forming relatively
low-density minerals that do not sink down into the Earth's core.
Potassium, rubidium and caesium are also incompatible elements due to
their large ionic radii.

Sodium and potassium are very abundant on Earth, both being among the
ten most common elements in Earth's crust; sodium makes up
approximately 2.6% of the Earth's crust measured by weight, making it
the sixth most abundant element overall and the most abundant alkali
metal. Potassium makes up approximately 1.5% of the Earth's crust and
is the seventh most abundant element. Sodium is found in many
different minerals, of which the most common is ordinary salt (sodium
chloride), which occurs in vast quantities dissolved in seawater.
Other solid deposits include halite, amphibole, cryolite, nitratine,
and zeolite. Many of these solid deposits occur as a result of ancient
seas evaporating, which still occurs now in places such as Utah's
Great Salt Lake and the Dead Sea. Despite their near-equal abundance
in Earth's crust, sodium is far more common than potassium in the
ocean, both because potassium's larger size makes its salts less
soluble, and because potassium is bound by silicates in soil and what
potassium leaches is absorbed far more readily by plant life than
sodium.

Despite its chemical similarity, lithium typically does not occur
together with sodium or potassium due to its smaller size. Due to its
relatively low reactivity, it can be found in seawater in large
amounts; it is estimated that lithium concentration in seawater is
approximately 0.14 to 0.25 parts per million (ppm) or 25 micromolar.
Its diagonal relationship with magnesium often allows it to replace
magnesium in ferromagnesium minerals, where its crustal concentration
is about 18 ppm, comparable to that of gallium and niobium.
Commercially, the most important lithium mineral is spodumene, which
occurs in large deposits worldwide.

Rubidium is approximately as abundant as zinc and more abundant than
copper. It occurs naturally in the minerals leucite, pollucite,
carnallite, zinnwaldite, and lepidolite, although none of these
contain only rubidium and no other alkali metals. Caesium is more
abundant than some commonly known elements, such as antimony, cadmium,
tin, and tungsten, but is much less abundant than rubidium.

Francium-223, the only naturally occurring isotope of francium, is the
product of the alpha decay of actinium-227 and can be found in trace
amounts in uranium minerals. In a given sample of uranium, there is
estimated to be only one francium atom for every 1018 uranium atoms.
It has been calculated that there are at most 30 grams of francium in
the earth's crust at any time, due to its extremely short half-life of
22 minutes.


Physical and chemical
=======================
The physical and chemical properties of the alkali metals can be
readily explained by their having an ns1 valence electron
configuration, which results in weak metallic bonding. Hence, all the
alkali metals are soft and have low densities, melting and boiling
points, as well as heats of sublimation, vaporisation, and
dissociation. They all crystallise in the body-centered cubic crystal
structure, and have distinctive flame colours because their outer s
electron is very easily excited. Indeed, these flame test colours are
the most common way of identifying them since all their salts with
common ions are soluble. The ns1 configuration also results in the
alkali metals having very large atomic and ionic radii, as well as
very high thermal and electrical conductivity. 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. 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; 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.
Properties of the alkali metals
Name    Lithium Sodium  Potassium       Rubidium        Caesium Francium
|Atomic number  3        11      19      37      55      87
|Standard atomic weight 6.94(1)          22.98976928(2)          39.0983(1)
85.4678(3)       132.9054519(2)          [223]
|Electron configuration [He] 2s1         [Ne] 3s1        [Ar] 4s1        [Kr] 5s1
[Xe] 6s1         [Rn] 7s1
|Melting point (°C)    180.54   97.72   63.38   39.31   28.44  ?
|Boiling point (°C)    1342     883     759     688     671    ?
|Density (g·cm−3)    0.534    0.968   0.89    1.532   1.93   ?
|Heat of fusion (kJ·mol−1)   3.00     2.60    2.321   2.19    2.09   ?
|Heat of vaporisation (kJ·mol−1)     136      97.42   79.1    69      66.1   ?
|Heat of formation of monatomic gas (kJ·mol−1)       162      108     89.6
82.0     78.2   ?
|Electrical resistivity at 25 °C (nΩ·cm)     94.7     48.8    73.9    131
208     ?
|Atomic radius (pm)     152      186     227     248     265    ?
|Ionic radius of hexacoordinate M+ ion (pm)     76       102     138     152
167     ?
|First ionisation energy (kJ·mol−1)  520.2    495.8   418.8   403.0
375.7    392.8
|Electron affinity (kJ·mol−1)        59.62    52.87   48.38   46.89   45.51
?
|Enthalpy of dissociation of M2 (kJ·mol−1)   106.5    73.6    57.3    45.6
44.77   ?
|Pauling electronegativity      0.98     0.93    0.82    0.82    0.79   ?
|Allen electronegativity        |0.91   |0.87   |0.73   |0.71   |0.66   |0.67
|Standard electrode potential ('E'°(M+→M0); V)       −3.04          −2.71         −2.93
−2.98          −3.03        ?
|Flame test colour Principal emission/absorption wavelength (nm)
Crimson 670.8    Yellow 589.2    Violet 766.5    Red-violet 780.0        Blue
455.5   ?
The alkali metals are more similar to each other than the elements in
any other group are to each other. 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, decreasing electronegativity,
increasing reactivity, and decreasing melting and boiling points as
well as heats of fusion and vaporisation. In general, their densities
increase when moving down the table, with the exception that potassium
is less dense than sodium. One of the very few properties of the
alkali metals that does not display a very smooth trend is their
reduction potentials: lithium's value is anomalous, being more
negative than the others. This is because the Li+ 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.

The stable alkali metals are all silver-coloured metals except for
caesium, which has a pale golden tint: it is one of only three metals
that are clearly coloured (the other two being copper and gold).
Additionally, the heavy alkaline earth metals calcium, strontium, and
barium, as well as the divalent lanthanides europium and ytterbium,
are pale yellow, though the colour is much less prominent than it is
for caesium. Their lustre tarnishes rapidly in air due to oxidation.

All the alkali metals are highly reactive and are never found in
elemental forms in nature. Because of this, they are usually stored in
mineral oil or kerosene (paraffin oil). They react aggressively with
the halogens to form the alkali metal halides, which are white ionic
crystalline compounds that are all soluble in water except lithium
fluoride (LiF). The alkali metals also react with water to form
strongly alkaline hydroxides 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. The alkali metals have the lowest first ionisation energies in
their respective periods of the periodic table because of their low
effective nuclear charge and the ability to attain a noble gas
configuration by losing just one electron. Not only do the alkali
metals react with water, but also with proton donors like alcohols and
phenols, gaseous ammonia, and alkynes, 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.

The second ionisation energy of all of the alkali metals is very high
as it is in a full shell that is also closer to the nucleus; thus,
they almost always lose a single electron, forming cations. The
alkalides 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 anions and were thought to be able to appear in salts
only as cations. The alkalide anions have filled s-subshells, which
gives them enough stability to exist. All the stable alkali metals
except lithium are known to be able to form alkalides, and the
alkalides have much theoretical interest due to their unusual
stoichiometry and low ionisation potentials. Alkalides are chemically
similar to the electrides, which are salts with trapped electrons
acting as anions. A particularly striking example of an alkalide is
"inverse sodium hydride", H+Na− (both ions being complexed), as
opposed to the usual sodium hydride, Na+H−: 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.

In aqueous solution, the alkali metal ions form aqua ions of the
formula [M(H2O)'n']+, where 'n' is the solvation number. Their
coordination numbers 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 bonds 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 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. The solvation number for Li+ has been experimentally
determined to be 4, forming the tetrahedral [Li(H2O)4]+: 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 pairs, and the higher solvation numbers may be interpreted
in terms of water molecules that approach [Li(H2O)4]+ 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(H2O)6]+ ion. 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(H2O)8]+ and [Rb(H2O)8]+ ions, which have the square
antiprismatic structure, and that caesium forms the 12-coordinate
[Cs(H2O)12]+ ion.


Lithium
=========
The chemistry of lithium shows several differences from that of the
rest of the group as the small Li+ cation polarises anions and gives
its compounds a more covalent character. Lithium and magnesium have a
diagonal relationship due to their similar atomic radii, so that they
show some similarities. For example, lithium forms a stable nitride, a
property common among all the alkaline earth metals (magnesium's
group) but unique among the alkali metals. In addition, among their
respective groups, only lithium and magnesium form organometallic
compounds with significant covalent character (e.g. LiMe and MgMe2).

Lithium fluoride is the only alkali metal halide that is poorly
soluble in water, and lithium hydroxide is the only alkali metal
hydroxide that is not deliquescent. 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+ has a high solvation energy. 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 dehumidifiers and air-conditioners.


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.
Thus, contrary to expectation, caesium is the most reactive of the
alkali metals, not francium. 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. Some of the few properties of francium that have been
predicted taking relativity into account are the electron affinity
(47.2 kJ/mol) and the enthalpy of dissociation of the Fr2 molecule
(42.1 kJ/mol). The CsFr molecule is polarised as Cs+Fr−, showing that
the 7s subshell of francium is much more strongly affected by
relativistic effects than the 6s subshell of caesium. Additionally,
francium superoxide (FrO2) is expected to have significant covalent
character, unlike the other alkali metal superoxides, because of
bonding contributions from the 6p electrons of francium.


Nuclear
=========
Primordial isotopes of the alkali metals
Z       Alkali metal    Stable  'Decays'
3       lithium          2       --
11      sodium           1       --
19      potassium        2       1                      '
37      rubidium         1       1              |'
55      caesium          1       --
87      francium         --      --     colspan="3"|'No primordial isotopes' ('
is a radiogenic nuclide)
colspan="7"|Radioactive:
All the alkali metals have odd atomic numbers; hence, their isotopes
must be either odd-odd (both proton and neutron number are odd) or
odd-even (proton number is odd, but neutron number is even). Odd-odd
nuclei have even mass numbers, whereas odd-even nuclei have odd mass
numbers. Odd-odd primordial nuclides 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 nuclear pairing effects.

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-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 (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
elements 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.

All of the alkali metals except lithium and caesium have at least one
naturally occurring radioisotope: sodium-22 and sodium-24 are trace
radioisotopes produced cosmogenically, potassium-40 and rubidium-87
have very long half-lives and thus occur naturally, and all isotopes
of francium are radioactive. Caesium was also thought to be
radioactive in the early 20th century, although it has no naturally
occurring radioisotopes. (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, 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. Natural rubidium is similarly slightly
radioactive, with 27.83% being the long-lived radioisotope
rubidium-87.

Caesium-137, with a half-life of 30.17 years, is one of the two
principal medium-lived fission products, 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. Caesium-137 has been used as a tracer in hydrologic studies,
analogous to the use of tritium. Small amounts of caesium-134 and
caesium-137 were released into the environment during nearly all
nuclear weapon tests and some nuclear accidents, 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. 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.


                          Periodic trends
======================================================================
The alkali metals are more similar to each other than the elements in
any other group are to each other. For instance, when moving down the
table, all known alkali metals show increasing atomic radius,
decreasing electronegativity, increasing reactivity, and decreasing
melting and boiling points as well as heats of fusion and
vaporisation. In general, their densities increase when moving down
the table, with the exception that potassium is less dense than
sodium.


Atomic and ionic radii
========================
The atomic radii of the alkali metals increase going down the group.
Because of the shielding effect, when an atom has more than one
electron shell, each electron feels electric repulsion from the other
electrons as well as electric attraction from the nucleus. In the
alkali metals, the outermost electron only feels a net charge of +1,
as some of the nuclear charge (which is equal to the atomic number) is
cancelled by the inner electrons; the number of inner electrons of an
alkali metal is always one less than the nuclear charge. Therefore,
the only factor which affects the atomic radius of the alkali metals
is the number of electron shells. Since this number increases down the
group, the atomic radius must also increase down the group.

The ionic radii of the alkali metals are much smaller than their
atomic radii. This is because the outermost electron of the alkali
metals is in a different electron shell than the inner electrons, and
thus when it is removed the resulting atom has one fewer electron
shell and is smaller. Additionally, the effective nuclear charge has
increased, and thus the electrons are attracted more strongly towards
the nucleus and the ionic radius decreases.


First ionisation energy
=========================
The first ionisation energy of an element or molecule is the energy
required to move the most loosely held electron from one mole of
gaseous atoms of the element or molecules to form one mole of gaseous
ions with electric charge +1. The factors affecting the first
ionisation energy are the nuclear charge, the amount of shielding by
the inner electrons and the distance from the most loosely held
electron from the nucleus, which is always an outer electron in main
group elements. The first two factors change the effective nuclear
charge the most loosely held electron feels. Since the outermost
electron of alkali metals always feels the same effective nuclear
charge (+1), the only factor which affects the first ionisation energy
is the distance from the outermost electron to the nucleus. Since this
distance increases down the group, the outermost electron feels less
attraction from the nucleus and thus the first ionisation energy
decreases. This trend is broken in francium due to the relativistic
stabilisation and contraction of the 7s orbital, bringing francium's
valence electron closer to the nucleus than would be expected from
non-relativistic calculations. This makes francium's outermost
electron feel more attraction from the nucleus, increasing its first
ionisation energy slightly beyond that of caesium.

The second ionisation energy of the alkali metals is much higher than
the first as the second-most loosely held electron is part of a fully
filled electron shell and is thus difficult to remove.


Reactivity
============
The reactivities of the alkali metals increase going down the group.
This is the result of a combination of two factors: the first
ionisation energies and atomisation energies of the alkali metals.
Because the first ionisation energy of the alkali metals decreases
down the group, it is easier for the outermost electron to be removed
from the atom and participate in chemical reactions, thus increasing
reactivity down the group. The atomisation energy measures the
strength of the metallic bond of an element, which falls down the
group as the atoms increase in radius and thus the metallic bond must
increase in length, making the delocalised electrons further away from
the attraction of the nuclei of the heavier alkali metals. Adding the
atomisation and first ionisation energies gives a quantity closely
related to (but not equal to) the activation energy of the reaction of
an alkali metal with another substance. This quantity decreases going
down the group, and so does the activation energy; thus, chemical
reactions can occur faster and the reactivity increases down the
group.


Electronegativity
===================
Electronegativity is a chemical property that describes the tendency
of an atom or a functional group to attract electrons (or electron
density) towards itself. If the bond between sodium and chlorine in
sodium chloride were covalent, the pair of shared electrons would be
attracted to the chlorine because the effective nuclear charge on the
outer electrons is +7 in chlorine but is only +1 in sodium. The
electron pair is attracted so close to the chlorine atom that they are
practically transferred to the chlorine atom (an ionic bond). However,
if the sodium atom was replaced by a lithium atom, the electrons will
not be attracted as close to the chlorine atom as before because the
lithium atom is smaller, making the electron pair more strongly
attracted to the closer effective nuclear charge from lithium. Hence,
the larger alkali metal atoms (further down the group) will be less
electronegative as the bonding pair is less strongly attracted towards
them. As mentioned previously, francium is expected to be an
exception.

Because of the higher electronegativity of lithium, some of its
compounds have a more covalent character. For example, lithium iodide
(LiI) will dissolve in organic solvents, a property of most covalent
compounds. Lithium fluoride (LiF) is the only alkali halide that is
not soluble in water, and lithium hydroxide (LiOH) is the only alkali
metal hydroxide that is not deliquescent.


Melting and boiling points
============================
The melting point of a substance is the point where it changes state
from solid to liquid while the boiling point of a substance (in liquid
state) is the point where the vapour pressure of the liquid equals the
environmental pressure surrounding the liquid and all the liquid
changes state to gas. As a metal is heated to its melting point, the
metallic bonds keeping the atoms in place weaken so that the atoms can
move around, and the metallic bonds eventually break completely at the
metal's boiling point. Therefore, the falling melting and boiling
points of the alkali metals indicate that the strength of the metallic
bonds of the alkali metals decreases down the group. This is because
metal atoms are held together by the electromagnetic attraction from
the positive ions to the delocalised electrons. As the atoms increase
in size going down the group (because their atomic radius increases),
the nuclei of the ions move further away from the delocalised
electrons and hence the metallic bond becomes weaker so that the metal
can more easily melt and boil, thus lowering the melting and boiling
points. The increased nuclear charge is not a relevant factor due to
the shielding effect.


Density
=========
The alkali metals all have the same crystal structure (body-centred
cubic) and thus the only relevant factors are the number of atoms that
can fit into a certain volume and the mass of one of the atoms, since
density is defined as mass per unit volume. The first factor depends
on the volume of the atom and thus the atomic radius, which increases
going down the group; thus, the volume of an alkali metal atom
increases going down the group. The mass of an alkali metal atom also
increases going down the group. Thus, the trend for the densities of
the alkali metals depends on their atomic weights and atomic radii; if
figures for these two factors are known, the ratios between the
densities of the alkali metals can then be calculated. The resultant
trend is that the densities of the alkali metals increase down the
table, with an exception at potassium. Due to having the lowest atomic
weight and the largest atomic radius of all the elements in their
periods, the alkali metals are the least dense metals in the periodic
table. Lithium, sodium, and potassium are the only three metals in the
periodic table that are less dense than water: in fact, lithium is the
least dense known solid at room temperature.


                             Compounds
======================================================================
The alkali metals form complete series of compounds with all usually
encountered anions, which well illustrate group trends. These
compounds can be described as involving the alkali metals losing
electrons to acceptor species and forming monopositive ions. This
description is most accurate for alkali halides and becomes less and
less accurate as cationic and anionic charge increase, and as the
anion becomes larger and more polarisable. For instance, ionic bonding
gives way to metallic bonding along the series NaCl, Na2O, Na2S, Na3P,
Na3As, Na3Sb, Na3Bi, Na.


[[Hydroxides]]
================
All the alkali metals react vigorously or explosively with cold water,
producing an aqueous solution of a strongly basic alkali metal
hydroxide and releasing hydrogen gas. This reaction becomes more
vigorous going down the group: lithium reacts steadily with
effervescence, but sodium and potassium can ignite, and rubidium and
caesium sink in water and generate hydrogen gas so rapidly that shock
waves form in the water that may shatter glass containers. When an
alkali metal is dropped into water, it produces an explosion, of which
there are two separate stages. The metal reacts with the water first,
breaking the hydrogen bonds in the water and producing hydrogen gas;
this takes place faster for the more reactive heavier alkali metals.
Second, the heat generated by the first part of the reaction often
ignites the hydrogen gas, causing it to burn explosively into the
surrounding air. This secondary hydrogen gas explosion produces the
visible flame above the bowl of water, lake or other body of water,
not the initial reaction of the metal with water (which tends to
happen mostly under water). The alkali metal hydroxides are the most
basic known hydroxides.

Recent research has suggested that the explosive behavior of alkali
metals in water is driven by a Coulomb explosion rather than solely by
rapid generation of hydrogen itself. All alkali metals melt as a part
of the reaction with water. Water molecules ionise the bare metallic
surface of the liquid metal, leaving a positively charged metal
surface and negatively charged water ions. The attraction between the
charged metal and water ions will rapidly increase the surface area,
causing an exponential increase of ionisation. When the repulsive
forces within the liquid metal surface exceeds the forces of the
surface tension, it vigorously explodes.

The hydroxides themselves are the most basic hydroxides known,
reacting with acids to give salts and with alcohols to give oligomeric
alkoxides. They easily react with carbon dioxide to form carbonates or
bicarbonates, or with hydrogen sulfide to form sulfides or bisulfides,
and may be used to separate thiols from petroleum. They react with
amphoteric oxides: for example, the oxides of aluminium, zinc, tin,
and lead react with the alkali metal hydroxides to give aluminates,
zincates, stannates, and plumbates. Silicon dioxide is acidic, and
thus the alkali metal hydroxides can also attack silicate glass.


Intermetallic compounds
=========================
The alkali metals form many intermetallic compounds with each other
and the elements from groups 2 to 13 in the periodic table of varying
stoichiometries, such as the sodium amalgams with mercury, including
Na5Hg8 and Na3Hg. Some of these have ionic characteristics: taking the
alloys with gold, the most electronegative of metals, as an example,
NaAu and KAu are metallic, but RbAu and CsAu are semiconductors. NaK
is an alloy of sodium and potassium that is very useful because it is
liquid at room temperature, although precautions must be taken due to
its extreme reactivity towards water and air. The eutectic mixture
melts at −12.6 °C. An alloy of 41% caesium, 47% sodium, and 12%
potassium has the lowest known melting point of any metal or alloy,
−78 °C.


Compounds with the group 13 elements
======================================
The intermetallic compounds of the alkali metals with the heavier
group 13 elements (aluminium, gallium, indium, and thallium), such as
NaTl, are poor conductors or semiconductors, unlike the normal alloys
with the preceding elements, implying that the alkali metal involved
has lost an electron to the Zintl anions involved. Nevertheless, while
the elements in group 14 and beyond tend to form discrete anionic
clusters, group 13 elements tend to form polymeric ions with the
alkali metal cations located between the giant ionic lattice. For
example, NaTl consists of a polymeric anion (--Tl−--)n with a covalent
diamond cubic structure with Na+ ions located between the anionic
lattice. The larger alkali metals cannot fit similarly into an anionic
lattice and tend to force the heavier group 13 elements to form
anionic clusters.

Boron is a special case, being the only nonmetal in group 13. The
alkali metal borides tend to be boron-rich, involving appreciable
boron-boron bonding involving deltahedral structures, and are
thermally unstable due to the alkali metals having a very high vapour
pressure at elevated temperatures. This makes direct synthesis
problematic because the alkali metals do not react with boron below
700 °C, and thus this must be accomplished in sealed containers with
the alkali metal in excess. Furthermore, exceptionally in this group,
reactivity with boron decreases down the group: lithium reacts
completely at 700 °C, but sodium at 900 °C and potassium not until
1200 °C, and the reaction is instantaneous for lithium but takes hours
for potassium. Rubidium and caesium borides have not even been
characterised. Various phases are known, such as LiB10, NaB6, NaB15,
and KB6. Under high pressure the boron-boron bonding in the lithium
borides changes from following Wade's rules to forming Zintl anions
like the rest of group 13.


Compounds with the group 14 elements
======================================
Lithium and sodium react with carbon to form acetylides, Li2C2 and
Na2C2, which can also be obtained by reaction of the metal with
acetylene. Potassium, rubidium, and caesium react with graphite; their
atoms are intercalated between the hexagonal graphite layers, forming
graphite intercalation compounds of formulae MC60 (dark grey, almost
black), MC48 (dark grey, almost black), MC36 (blue), MC24 (steel
blue), and MC8 (bronze) (M = K, Rb, or Cs). These compounds are over
200 times more electrically conductive than pure graphite, suggesting
that the valence electron of the alkali metal is transferred to the
graphite layers (e.g. ). Upon heating of KC8, the elimination of
potassium atoms results in the conversion in sequence to KC24, KC36,
KC48 and finally KC60. KC8 is a very strong reducing agent and is
pyrophoric and explodes on contact with water. While the larger alkali
metals (K, Rb, and Cs) initially form MC8, the smaller ones initially
form MC6, and indeed they require reaction of the metals with graphite
at high temperatures around 500 °C to form. Apart from this, the
alkali metals are such strong reducing agents that they can even
reduce buckminsterfullerene to produce solid fullerides M'n'C60;
sodium, potassium, rubidium, and caesium can form fullerides where 'n'
= 2, 3, 4, or 6, and rubidium and caesium additionally can achieve 'n'
= 1.

When the alkali metals react with the heavier elements in the carbon
group (silicon, germanium, tin, and lead), ionic substances with
cage-like structures are formed, such as the silicides M4Si4 (M = K,
Rb, or Cs), which contains M+ and tetrahedral  ions. The chemistry of
alkali metal germanides, involving the germanide ion Ge4− and other
cluster (Zintl) ions such as , , , and [(Ge9)2]6−, is largely
analogous to that of the corresponding silicides. Alkali metal
stannides are mostly ionic, sometimes with the stannide ion (Sn4−),
and sometimes with more complex Zintl ions such as , which appears in
tetrapotassium nonastannide (K4Sn9). The monatomic plumbide ion (Pb4−)
is unknown, and indeed its formation is predicted to be energetically
unfavourable; alkali metal plumbides have complex Zintl ions, such as
. These alkali metal germanides, stannides, and plumbides may be
produced by reducing germanium, tin, and lead with sodium metal in
liquid ammonia.


Nitrides and pnictides
========================
Lithium, the lightest of the alkali metals, is the only alkali metal
which reacts with nitrogen at standard conditions, and its nitride is
the only stable alkali metal nitride. Nitrogen is an unreactive gas
because breaking the strong triple bond in the dinitrogen molecule
(N2) requires a lot of energy. The formation of an alkali metal
nitride would consume the ionisation energy of the alkali metal
(forming M+ ions), the energy required to break the triple bond in N2
and the formation of N3− ions, and all the energy released from the
formation of an alkali metal nitride is from the lattice energy of the
alkali metal nitride. The lattice energy is maximised with small,
highly charged ions; the alkali metals do not form highly charged
ions, only forming ions with a charge of +1, so only lithium, the
smallest alkali metal, can release enough lattice energy to make the
reaction with nitrogen exothermic, forming lithium nitride. The
reactions of the other alkali metals with nitrogen would not release
enough lattice energy and would thus be endothermic, so they do not
form nitrides at standard conditions. Sodium nitride (Na3N) and
potassium nitride (K3N), while existing, are extremely unstable, being
prone to decomposing back into their constituent elements, and cannot
be produced by reacting the elements with each other at standard
conditions. Steric hindrance forbids the existence of rubidium or
caesium nitride. However, sodium and potassium form colourless azide
salts involving the linear  anion; due to the large size of the alkali
metal cations, they are thermally stable enough to be able to melt
before decomposing.

All the alkali metals react readily with phosphorus and arsenic to
form phosphides and arsenides with the formula M3Pn (where M
represents an alkali metal and Pn represents a pnictogen - phosphorus,
arsenic, antimony, or bismuth). This is due to the greater size of the
P3− and As3− ions, so that less lattice energy needs to be released
for the salts to form. These are not the only phosphides and arsenides
of the alkali metals: for example, potassium has nine different known
phosphides, with formulae K3P, K4P3, K5P4, KP, K4P6, K3P7, K3P11,
KP10.3, and KP15. While most metals form arsenides, only the alkali
and alkaline earth metals form mostly ionic arsenides. The structure
of Na3As is complex with unusually short Na-Na distances of 328-330 pm
which are shorter than in sodium metal, and this indicates that even
with these electropositive metals the bonding cannot be
straightforwardly ionic. Other alkali metal arsenides not conforming
to the formula M3As are known, such as LiAs, which has a metallic
lustre and electrical conductivity indicating the presence of some
metallic bonding. The antimonides are unstable and reactive as the
Sb3− ion is a strong reducing agent; reaction of them with acids form
the toxic and unstable gas stibine (SbH3). Indeed, they have some
metallic properties, and the alkali metal antimonides of stoichiometry
MSb involve antimony atoms bonded in a spiral Zintl structure.
Bismuthides are not even wholly ionic; they are intermetallic
compounds containing partially metallic and partially ionic bonds.


Oxides and chalcogenides
==========================
All the alkali metals react vigorously with oxygen at standard
conditions. They form various types of oxides, such as simple oxides
(containing the O2− ion), peroxides (containing the  ion, where there
is a single bond between the two oxygen atoms), superoxides
(containing the  ion), and many others. Lithium burns in air to form
lithium oxide, but sodium reacts with oxygen to form a mixture of
sodium oxide and sodium peroxide. Potassium forms a mixture of
potassium peroxide and potassium superoxide, while rubidium and
caesium form the superoxide exclusively. Their reactivity increases
going down the group: while lithium, sodium and potassium merely burn
in air, rubidium and caesium are pyrophoric (spontaneously catch fire
in air).

The smaller alkali metals tend to polarise the larger anions (the
peroxide and superoxide) due to their small size. This attracts the
electrons in the more complex anions towards one of its constituent
oxygen atoms, forming an oxide ion and an oxygen atom. This causes
lithium to form the oxide exclusively on reaction with oxygen at room
temperature. This effect becomes drastically weaker for the larger
sodium and potassium, allowing them to form the less stable peroxides.
Rubidium and caesium, at the bottom of the group, are so large that
even the least stable superoxides can form. Because the superoxide
releases the most energy when formed, the superoxide is preferentially
formed for the larger alkali metals where the more complex anions are
not polarised. The oxides and peroxides for these alkali metals do
exist, but do not form upon direct reaction of the metal with oxygen
at standard conditions. In addition, the small size of the Li+ and O2−
ions contributes to their forming a stable ionic lattice structure.
Under controlled conditions, however, all the alkali metals, with the
exception of francium, are known to form their oxides, peroxides, and
superoxides. The alkali metal peroxides and superoxides are powerful
oxidising agents. Sodium peroxide and potassium superoxide react with
carbon dioxide to form the alkali metal carbonate and oxygen gas,
which allows them to be used in submarine air purifiers; the presence
of water vapour, naturally present in breath, makes the removal of
carbon dioxide by potassium superoxide even more efficient. All the
stable alkali metals except lithium can form red ozonides (MO3)
through low-temperature reaction of the powdered anhydrous hydroxide
with ozone: the ozonides may be then extracted using liquid ammonia.
They slowly decompose at standard conditions to the superoxides and
oxygen, and hydrolyse immediately to the hydroxides when in contact
with water. Potassium, rubidium, and caesium also form sesquioxides
M2O3, which may be better considered peroxide disuperoxides, .

Rubidium and caesium can form a great variety of suboxides with the
metals in formal oxidation states below +1. Rubidium can form Rb6O and
Rb9O2 (copper-coloured) upon oxidation in air, while caesium forms an
immense variety of oxides, such as the ozonide CsO3 and several
brightly coloured suboxides, such as Cs7O (bronze), Cs4O (red-violet),
Cs11O3 (violet), Cs3O (dark green), CsO, Cs3O2, as well as Cs7O2. The
last of these may be heated under vacuum to generate Cs2O.

The alkali metals can also react analogously with the heavier
chalcogens (sulfur, selenium, tellurium, and polonium), and all the
alkali metal chalcogenides are known (with the exception of
francium's). Reaction with an excess of the chalcogen can similarly
result in lower chalcogenides, with chalcogen ions containing chains
of the chalcogen atoms in question. For example, sodium can react with
sulfur to form the sulfide (Na2S) and various polysulfides with the
formula Na2S'x' ('x' from 2 to 6), containing the  ions. Due to the
basicity of the Se2− and Te2− ions, the alkali metal selenides and
tellurides are alkaline in solution; when reacted directly with
selenium and tellurium, alkali metal polyselenides and polytellurides
are formed along with the selenides and tellurides with the  and
ions. They may be obtained directly from the elements in liquid
ammonia or when air is not present, and are colourless, water-soluble
compounds that air oxidises quickly back to selenium or tellurium. The
alkali metal polonides are all ionic compounds containing the Po2−
ion; they are very chemically stable and can be produced by direct
reaction of the elements at around 300-400 °C.


Halides, hydrides, and pseudohalides
======================================
The alkali metals are among the most electropositive elements on the
periodic table and thus tend to bond ionically to the most
electronegative elements on the periodic table, the halogens
(fluorine, chlorine, bromine, iodine, and astatine), forming salts
known as the alkali metal halides. The reaction is very vigorous and
can sometimes result in explosions. All twenty stable alkali metal
halides are known; the unstable ones are not known, with the exception
of sodium astatide, because of the great instability and rarity of
astatine and francium. The most well-known of the twenty is certainly
sodium chloride, otherwise known as common salt. All of the stable
alkali metal halides have the formula MX where M is an alkali metal
and X is a halogen. They are all white ionic crystalline solids that
have high melting points. All the alkali metal halides are soluble in
water except for lithium fluoride (LiF), which is insoluble in water
due to its very high lattice enthalpy. The high lattice enthalpy of
lithium fluoride is due to the small sizes of the Li+ and F− ions,
causing the electrostatic interactions between them to be strong: a
similar effect occurs for magnesium fluoride, consistent with the
diagonal relationship between lithium and magnesium.

The alkali metals also react similarly with hydrogen to form ionic
alkali metal hydrides, where the hydride anion acts as a pseudohalide:
these are often used as reducing agents, producing hydrides, complex
metal hydrides, or hydrogen gas. Other pseudohalides are also known,
notably the cyanides. These are isostructural to the respective
halides except for lithium cyanide, indicating that the cyanide ions
may rotate freely. Ternary alkali metal halide oxides, such as Na3ClO,
K3BrO (yellow), Na4Br2O, Na4I2O, and K4Br2O, are also known. The
polyhalides are rather unstable, although those of rubidium and
caesium are greatly stabilised by the feeble polarising power of these
extremely large cations.


Coordination complexes
========================
Alkali metal cations do not usually form coordination complexes with
simple Lewis bases due to their low charge of just +1 and their
relatively large size; thus the Li+ ion forms most complexes and the
heavier alkali metal ions form less and less (though exceptions occur
for weak complexes). Lithium in particular has a very rich
coordination chemistry in which it exhibits coordination numbers from
1 to 12, although octahedral hexacoordination is its preferred mode.
In aqueous solution, the alkali metal ions exist as octahedral
hexahydrate complexes [M(H2O)6]+, with the exception of the lithium
ion, which due to its small size forms tetrahedral tetrahydrate
complexes [Li(H2O)4]+; the alkali metals form these complexes because
their ions are attracted by electrostatic forces of attraction to the
polar water molecules. Because of this, anhydrous salts containing
alkali metal cations are often used as desiccants. Alkali metals also
readily form complexes with crown ethers (e.g. 12-crown-4 for Li+,
15-crown-5 for Na+, 18-crown-6 for K+, and 21-crown-7 for Rb+) and
cryptands due to electrostatic attraction.


Ammonia solutions
===================
The alkali metals dissolve slowly in liquid ammonia, forming
ammoniacal solutions of solvated metal cation M+ and solvated electron
e−, which react to form hydrogen gas and the alkali metal amide (MNH2,
where M represents an alkali metal): this was first noted by Humphry
Davy in 1809 and rediscovered by W. Weyl in 1864. The process may be
speeded up by a catalyst. Similar solutions are formed by the heavy
divalent alkaline earth metals calcium, strontium, barium, as well as
the divalent lanthanides, europium and ytterbium. The amide salt is
quite insoluble and readily precipitates out of solution, leaving
intensely coloured ammonia solutions of the alkali metals. In 1907,
Charles A. Kraus identified the colour as being due to the presence of
solvated electrons, which contribute to the high electrical
conductivity of these solutions. At low concentrations (below 3 M),
the solution is dark blue and has ten times the conductivity of
aqueous sodium chloride; at higher concentrations (above 3 M), the
solution is copper-coloured and has approximately the conductivity of
liquid metals like mercury. In addition to the alkali metal amide salt
and solvated electrons, such ammonia solutions also contain the alkali
metal cation (M+), the neutral alkali metal atom (M), diatomic alkali
metal molecules (M2) and alkali metal anions (M−). These are unstable
and eventually become the more thermodynamically stable alkali metal
amide and hydrogen gas. Solvated electrons are powerful reducing
agents and are often used in chemical synthesis.


Organolithium
===============
|title= Structures of Classical Reagents in Chemical Synthesis:
(nBuLi)6, (tBuLi)4, and the Metastable (tBuLi · Et2O)2
|last=T. Kottke|first=D. Stalke
|journal= Angew. Chem. Int. Ed. Engl.
|date= September 1993
|volume= 32
|issue= 4
|pages= 580-582
|doi= 10.1002/anie.199305801

Being the smallest alkali metal, lithium forms the widest variety of
and most stable organometallic compounds, which are bonded covalently.
Organolithium compounds are electrically non-conducting volatile
solids or liquids that melt at low temperatures, and tend to form
oligomers with the structure (RLi)'x' where R is the organic group. As
the electropositive nature of lithium puts most of the charge density
of the bond on the carbon atom, effectively creating a carbanion,
organolithium compounds are extremely powerful bases and nucleophiles.
For use as bases, butyllithiums are often used and are commercially
available. An example of an organolithium compound is methyllithium
((CH3Li)'x'), which exists in tetrameric ('x' = 4, tetrahedral) and
hexameric ('x' = 6, octahedral) forms. Organolithium compounds,
especially 'n'-butyllithium, are useful reagents in organic synthesis,
as might be expected given lithium's diagonal relationship with
magnesium, which plays an important role in the Grignard reaction. For
example, alkyllithiums and aryllithiums may be used to synthesise
aldehydes and ketones by reaction with metal carbonyls. The reaction
with nickel tetracarbonyl, for example, proceeds through an unstable
acyl nickel carbonyl complex which then undergoes electrophilic
substitution to give the desired aldehyde (using H+ as the
electrophile) or ketone (using an alkyl halide) product.

:LiR \ + \ Ni(CO)4 \ \longrightarrow Li^{+}[RCONi(CO)3]^{-}
:Li^{+}[RCONi(CO)3]^{-}->[\ce{H^{+}}][\ce{solvent}] \ Li^{+} \ + \
RCHO \ + \ [(solvent)Ni(CO)3]
:Li^{+}[RCONi(CO)3]^{-}->[\ce{R^{'}Br}][\ce{solvent}] \ Li^{+} \ +
\ RR^{'}CO \ + \ [(solvent)Ni(CO)3]

Alkyllithiums and aryllithiums may also react with
'N','N'-disubstituted amides to give aldehydes and ketones, and
symmetrical ketones by reacting with carbon monoxide. They thermally
decompose to eliminate a β-hydrogen, producing alkenes and lithium
hydride: another route is the reaction of ethers with alkyl- and
aryllithiums that act as strong bases. In non-polar solvents,
aryllithiums react as the carbanions they effectively are, turning
carbon dioxide to aromatic carboxylic acids (ArCO2H) and aryl ketones
to tertiary carbinols (Ar'2C(Ar)OH). Finally, they may be used to
synthesise other organometallic compounds through metal-halogen
exchange.


Heavier alkali metals
=======================
Unlike the organolithium compounds, the organometallic compounds of
the heavier alkali metals are predominantly ionic. The application of
organosodium compounds in chemistry is limited in part due to
competition from organolithium compounds, which are commercially
available and exhibit more convenient reactivity. The principal
organosodium compound of commercial importance is sodium
cyclopentadienide. Sodium tetraphenylborate can also be classified as
an organosodium compound since in the solid state sodium is bound to
the aryl groups. Organometallic compounds of the higher alkali metals
are even more reactive than organosodium compounds and of limited
utility. A notable reagent is Schlosser's base, a mixture of
'n'-butyllithium and potassium 'tert'-butoxide. This reagent reacts
with propene to form the compound allylpotassium (KCH2CHCH2).
'cis'-2-Butene and 'trans'-2-butene equilibrate when in contact with
alkali metals. Whereas isomerisation is fast with lithium and sodium,
it is slow with the heavier alkali metals. The heavier alkali metals
also favour the sterically congested conformation. Several crystal
structures of organopotassium compounds have been reported,
establishing that they, like the sodium compounds, are polymeric.
Organosodium, organopotassium, organorubidium and organocaesium
compounds are all mostly ionic and are insoluble (or nearly so) in
nonpolar solvents.

Alkyl and aryl derivatives of sodium and potassium tend to react with
air. They cause the cleavage of ethers, generating alkoxides. Unlike
alkyllithium compounds, alkylsodiums and alkylpotassiums cannot be
made by reacting the metals with alkyl halides because Wurtz coupling
occurs:
:RM + R'X → R-R' + MX

As such, they have to be made by reacting alkylmercury compounds with
sodium or potassium metal in inert hydrocarbon solvents. While
methylsodium forms tetramers like methyllithium, methylpotassium is
more ionic and has the nickel arsenide structure with discrete methyl
anions and potassium cations.

The alkali metals and their hydrides react with acidic hydrocarbons,
for example cyclopentadienes and terminal alkynes, to give salts.
Liquid ammonia, ether, or hydrocarbon solvents are used, the most
common of which being tetrahydrofuran. The most important of these
compounds is sodium cyclopentadienide, NaC5H5, an important precursor
to many transition metal cyclopentadienyl derivatives. Similarly, the
alkali metals react with cyclooctatetraene in tetrahydrofuran to give
alkali metal cyclooctatetraenides; for example, dipotassium
cyclooctatetraenide (K2C8H8) is an important precursor to many metal
cyclooctatetraenyl derivatives, such as uranocene. The large and very
weakly polarising alkali metal cations can stabilise large, aromatic,
polarisable radical anions, such as the dark-green sodium
naphthalenide, Na+[C10H8•]−, a strong reducing agent.


Reaction with oxygen
======================
Upon reacting with oxygen, alkali metals form oxides, peroxides,
superoxides and suboxides. However, the first three are more common.
The table below shows the types of compounds formed in reaction with
oxygen. The compound in brackets represents the minor product of
combustion.
|**Alkali metal**||**Oxide**||**Peroxide**||**Superoxide**
|Li     Li2O    (Li2O2)
|Na     (Na2O)  Na2O2
|K                      KO2
|Rb                     RbO2
|Cs                     CsO2
The alkali metal peroxides are ionic compounds that are unstable in
water. The peroxide anion is weakly bound to the cation, and it is
hydrolysed, forming stronger covalent bonds.
:Na2O2 + 2H2O → 2NaOH + H2O2
The other oxygen compounds are also unstable in water.
:2KO2 + 2H2O → 2KOH + H2O2 + O2
:Li2O + H2O → 2LiOH


Reaction with sulfur
======================
With sulfur, they form sulfides and polysulfides.
:2Na + 1/8S8 → Na2S + 1/8S8 → Na2S2...Na2S7
Because alkali metal sulfides are essentially salts of a weak acid and
a strong base, they form basic solutions.
:S2- + H2O → HS− + HO−
:HS− + H2O → H2S + HO−


Reaction with nitrogen
========================
Lithium is the only metal that combines directly with nitrogen at room
temperature.
:3Li + 1/2N2 → Li3N
Li3N can react with water to liberate ammonia.
:Li3N + 3H2O → 3LiOH + NH3


Reaction with hydrogen
========================
With hydrogen, alkali metals form saline hydrides that hydrolyse in
water.
:2 Na \ + H2 \ ->[\ce{\Delta}] \ 2 NaH
:2 NaH \ + \ 2 H2O \ \longrightarrow \ 2 NaOH \ + \ H2 \uparrow


Reaction with carbon
======================
Lithium is the only metal that reacts directly with carbon to give
dilithium acetylide. Na and K can react with acetylene to give
acetylides.
:2 Li \ + \ 2 C \ \longrightarrow \ Li2C2
: 2 Na \ + \ 2 C2H2 \ ->[\ce{150 \ ^{o}C}] \ 2 NaC2H \ + \ H2
: 2 Na \ + \ 2 NaC2H \ ->[\ce{220 \ ^{o}C}] \ 2 Na2C2 \ + \ H2


Reaction with water
=====================
On reaction with water, they generate hydroxide ions and hydrogen gas.
This reaction is vigorous and highly exothermic and the hydrogen
resulted may ignite in air or even explode in the case of Rb and Cs.
:Na + H2O → NaOH + 1/2H2


Reaction with other salts
===========================
The alkali metals are very good reducing agents. They can reduce metal
cations that are less electropositive. Titanium is produced
industrially by the reduction of titanium tetrachloride with Na at 400
°C (van Arkel-de Boer process).
:TiCl4 + 4Na → 4NaCl + Ti


Reaction with organohalide compounds
======================================
Alkali metals react with halogen derivatives to generate hydrocarbon
via the Wurtz reaction.
:2CH3-Cl + 2Na → H3C-CH3 + 2NaCl


Alkali metals in liquid ammonia
=================================
Alkali metals dissolve in liquid ammonia or other donor solvents like
aliphatic amines or hexamethylphosphoramide to give blue solutions.
These solutions are believed to contain free electrons.
:Na + xNH3 → Na+ + e(NH3)x−
Due to the presence of solvated electrons, these solutions are very
powerful reducing agents used in organic synthesis.


Reaction 1) is known as Birch reduction.
Other reductions that can be carried by these solutions are:
:S8 + 2e− → S82-
:Fe(CO)5 + 2e− → Fe(CO)42- + CO


                             Extensions
======================================================================
Although francium is the heaviest alkali metal that has been
discovered, there has been some theoretical work predicting the
physical and chemical characteristics of hypothetical heavier alkali
metals. Being the first period 8 element, the undiscovered element
ununennium (element 119) is predicted to be the next alkali metal
after francium and behave much like their lighter congeners; however,
it is also predicted to differ from the lighter alkali metals in some
properties. Its chemistry is predicted to be closer to that of
potassium or rubidium instead of caesium or francium. This is unusual
as periodic trends, ignoring relativistic effects would predict
ununennium to be even more reactive than caesium and francium. This
lowered reactivity is due to the relativistic stabilisation of
ununennium's valence electron, increasing ununennium's first
ionisation energy and decreasing the metallic and ionic radii; this
effect is already seen for francium. This assumes that ununennium will
behave chemically as an alkali metal, which, although likely, may not
be true due to relativistic effects. The relativistic stabilisation of
the 8s orbital also increases ununennium's electron affinity far
beyond that of caesium and francium; indeed, ununennium is expected to
have an electron affinity higher than all the alkali metals lighter
than it. Relativistic effects also cause a very large drop in the
polarisability of ununennium. On the other hand, ununennium is
predicted to continue the trend of melting points decreasing going
down the group, being expected to have a melting point between 0 °C
and 30 °C.

The stabilisation of ununennium's valence electron and thus the
contraction of the 8s orbital cause its atomic radius to be lowered to
240 pm, very close to that of rubidium (247 pm), so that the chemistry
of ununennium in the +1 oxidation state should be more similar to the
chemistry of rubidium than to that of francium. On the other hand, the
ionic radius of the Uue+ ion is predicted to be larger than that of
Rb+, because the 7p orbitals are destabilised and are thus larger than
the p-orbitals of the lower shells. Ununennium may also show the +3
and +5 oxidation states, which are not seen in any other alkali metal,
in addition to the +1 oxidation state that is characteristic of the
other alkali metals and is also the main oxidation state of all the
known alkali metals: this is because of the destabilisation and
expansion of the 7p3/2 spinor, causing its outermost electrons to have
a lower ionisation energy than what would otherwise be expected.
Indeed, many ununennium compounds are expected to have a large
covalent character, due to the involvement of the 7p3/2 electrons in
the bonding.

Not as much work has been done predicting the properties of the alkali
metals beyond ununennium. Although a simple extrapolation of the
periodic table (by the Aufbau principle) would put element 169,
unhexennium, under ununennium, Dirac-Fock calculations predict that
the next element after ununennium with alkali-metal-like properties
may be element 165, unhexpentium, which is predicted to have the
electron configuration [Og] 5g18 6f14 7d10 8s2 8p1/22 9s1. This
element would be intermediate in properties between an alkali metal
and a group 11 element, and while its physical and atomic properties
would be closer to the former, its chemistry may be closer to that of
the latter. Further calculations show that unhexpentium would follow
the trend of increasing ionisation energy beyond caesium, having an
ionisation energy comparable to that of sodium, and that it should
also continue the trend of decreasing atomic radii beyond caesium,
having an atomic radius comparable to that of potassium. However, the
7d electrons of unhexpentium may also be able to participate in
chemical reactions along with the 9s electron, possibly allowing
oxidation states beyond +1, whence the likely transition metal
behaviour of unhexpentium. Due to the alkali and alkaline earth metals
both being s-block elements, these predictions for the trends and
properties of ununennium and unhexpentium also mostly hold quite
similarly for the corresponding alkaline earth metals unbinilium (Ubn)
and unhexhexium (Uhh). Unsepttrium, element 173, may be an even better
heavier homologue of ununennium; with a predicted electron
configuration of [Usb] 6g1, it returns to the alkali-metal-like
situation of having one easily removed electron far above a closed
p-shell in energy, and is expected to be even more reactive than
caesium.

The probable properties of further alkali metals beyond unsepttrium
have not been explored yet as of 2019, and they may or may not be able
to exist. In periods 8 and above of the periodic table, relativistic
and shell-structure effects become so strong that extrapolations from
lighter congeners become completely inaccurate. In addition, the
relativistic and shell-structure effects (which stabilise the
s-orbitals and destabilise and expand the d-, f-, and g-orbitals of
higher shells) have opposite effects, causing even larger difference
between relativistic and non-relativistic calculations of the
properties of elements with such high atomic numbers. Interest in the
chemical properties of ununennium, unhexpentium, and unsepttrium stems
from the fact that they are located close to the expected locations of
islands of stability, centered at elements 122 (306Ubb) and 164
(482Uhq).


                        Pseudo-alkali metals
======================================================================
Many other substances are similar to the alkali metals in their
tendency to form monopositive cations. Analogously to the
pseudohalogens, they have sometimes been called "pseudo-alkali
metals". These substances include some elements and many more
polyatomic ions; the polyatomic ions are especially similar to the
alkali metals in their large size and weak polarising power.


Hydrogen
==========
The element hydrogen, with one electron per neutral atom, is usually
placed at the top of Group 1 of the periodic table because of its
electron configuration. But hydrogen is not normally considered to be
an alkali metal.  Metallic hydrogen, which only exists at very high
pressures, is known for its electrical and magnetic properties, not
its chemical properties. Under typical conditions, pure hydrogen
exists as a diatomic gas consisting of two atoms per molecule (H2);
however, the alkali metals form diatomic molecules (such as dilithium,
Li2) only at high temperatures, when they are in the gaseous state.

Hydrogen, like the alkali metals, has one valence electron and reacts
easily with the halogens, but the similarities mostly end there
because of the small size of a bare proton H+ compared to the alkali
metal cations. Its placement above lithium is primarily due to its
electron configuration. It is sometimes placed above fluorine due to
their similar chemical properties, though the resemblance is likewise
not absolute.

The first ionisation energy of hydrogen (1312.0 kJ/mol) is much higher
than that of the alkali metals. As only one additional electron is
required to fill in the outermost shell of the hydrogen atom, hydrogen
often behaves like a halogen, forming the negative hydride ion, and is
very occasionally considered to be a halogen on that basis. (The
alkali metals can also form negative ions, known as alkalides, but
these are little more than laboratory curiosities, being unstable.) An
argument against this placement is that formation of hydride from
hydrogen is endothermic, unlike the exothermic formation of halides
from halogens. The radius of the H− anion also does not fit the trend
of increasing size going down the halogens: indeed, H− is very diffuse
because its single proton cannot easily control both electrons. It was
expected for some time that liquid hydrogen would show metallic
properties; while this has been shown to not be the case, under
extremely high pressures, such as those found at the cores of Jupiter
and Saturn, hydrogen does become metallic and behaves like an alkali
metal; in this phase, it is known as metallic hydrogen. The electrical
resistivity of liquid metallic hydrogen at 3000 K is approximately
equal to that of liquid rubidium and caesium at 2000 K at the
respective pressures when they undergo a nonmetal-to-metal transition.

The 1s1 electron configuration of hydrogen, while analogous to that of
the alkali metals (ns1), is unique because there is no 1p subshell.
Hence it can lose an electron to form the hydron H+, or gain one to
form the hydride ion H−. In the former case it resembles superficially
the alkali metals; in the latter case, the halogens, but the
differences due to the lack of a 1p subshell are important enough that
neither group fits the properties of hydrogen well. Group 14 is also a
good fit in terms of thermodynamic properties such as ionisation
energy and electron affinity, but hydrogen cannot be tetravalent. Thus
none of the three placements are entirely satisfactory, although group
1 is the most common placement (if one is chosen) because of the
electron configuration and the fact that the hydron is by far the most
important of all monatomic hydrogen species, being the foundation of
acid-base chemistry. As an example of hydrogen's unorthodox properties
stemming from its unusual electron configuration and small size, the
hydrogen ion is very small (radius around 150 fm compared to the
50-220 pm size of most other atoms and ions) and so is nonexistent in
condensed systems other than in association with other atoms or
molecules. Indeed, transferring of protons between chemicals is the
basis of acid-base chemistry. Also unique is hydrogen's ability to
form hydrogen bonds, which are an effect of charge-transfer,
electrostatic, and electron correlative contributing phenomena. While
analogous lithium bonds are also known, they are mostly electrostatic.
Nevertheless, hydrogen can take on the same structural role as the
alkali metals in some molecular crystals, and has a close relationship
with the lightest alkali metals (especially lithium).


Ammonium and derivatives
==========================
The ammonium ion () has very similar properties to the heavier alkali
metals, acting as an alkali metal intermediate between potassium and
rubidium, and is often considered a close relative. For example, most
alkali metal salts are soluble in water, a property which ammonium
salts share. Ammonium is expected to behave stably as a metal ( ions
in a sea of delocalised electrons) at very high pressures (though less
than the typical pressure where transitions from insulating to
metallic behaviour occur around, 100 GPa), and could possibly occur
inside the ice giants Uranus and Neptune, which may have significant
impacts on their interior magnetic fields. It has been estimated that
the transition from a mixture of ammonia and dihydrogen molecules to
metallic ammonium may occur at pressures just below 25 GPa. Under
standard conditions, ammonium can form a metallic amalgam with
mercury.

Other "pseudo-alkali metals" include the alkylammonium cations, in
which some of the hydrogen atoms in the ammonium cation are replaced
by alkyl or aryl groups. In particular, the quaternary ammonium
cations () are very useful since they are permanently charged, and
they are often used as an alternative to the expensive Cs+ to
stabilise very large and very easily polarisable anions such as .
Tetraalkylammonium hydroxides, like alkali metal hydroxides, are very
strong bases that react with atmospheric carbon dioxide to form
carbonates. Furthermore, the nitrogen atom may be replaced by a
phosphorus, arsenic, or antimony atom (the heavier nonmetallic
pnictogens), creating a phosphonium () or arsonium () cation that can
itself be substituted similarly; while stibonium () itself is not
known, some of its organic derivatives are characterised.


Cobaltocene and derivatives
=============================
Cobaltocene, Co(C5H5)2, is a metallocene, the cobalt analogue of
ferrocene. It is a dark purple solid. Cobaltocene has 19 valence
electrons, one more than usually found in organotransition metal
complexes, such as its very stable relative, ferrocene, in accordance
with the 18-electron rule. This additional electron occupies an
orbital that is antibonding with respect to the Co-C bonds.
Consequently, many chemical reactions of Co(C5H5)2 are characterized
by its tendency to lose this "extra" electron, yielding a very stable
18-electron cation known as cobaltocenium. Many cobaltocenium salts
coprecipitate with caesium salts, and cobaltocenium hydroxide is a
strong base that absorbs atmospheric carbon dioxide to form
cobaltocenium carbonate. Like the alkali metals, cobaltocene is a
strong reducing agent, and decamethylcobaltocene is stronger still due
to the combined inductive effect of the ten methyl groups. Cobalt may
be substituted by its heavier congener rhodium to give rhodocene, an
even stronger reducing agent. Iridocene (involving iridium) would
presumably be still more potent, but is not very well-studied due to
its instability.


Thallium
==========
Thallium is the heaviest stable element in group 13 of the periodic
table. At the bottom of the periodic table, the inert-pair effect is
quite strong, because of the relativistic stabilisation of the 6s
orbital and the decreasing bond energy as the atoms increase in size
so that the amount of energy released in forming two more bonds is not
worth the high ionisation energies of the 6s electrons. It displays
the +1 oxidation state that all the known alkali metals display, and
thallium compounds with thallium in its +1 oxidation state closely
resemble the corresponding potassium or silver compounds
stoichiometrically due to the similar ionic radii of the Tl+ (164 pm),
K+ (152 pm) and Ag+ (129 pm) ions. It was sometimes considered an
alkali metal in continental Europe (but not in England) in the years
immediately following its discovery, and was placed just after caesium
as the sixth alkali metal in Dmitri Mendeleev's 1869 periodic table
and Julius Lothar Meyer's 1868 periodic table. Mendeleev's 1871
periodic table and Meyer's 1870 periodic table put thallium in its
current position in the boron group and left the space below caesium
blank. However, thallium also displays the oxidation state +3, which
no known alkali metal displays (although ununennium, the undiscovered
seventh alkali metal, is predicted to possibly display the +3
oxidation state). The sixth alkali metal is now considered to be
francium. While Tl+ is stabilised by the inert-pair effect, this inert
pair of 6s electrons is still able to participate chemically, so that
these electrons are stereochemically active in aqueous solution.
Additionally, the thallium halides (except TlF) are quite insoluble in
water, and TlI has an unusual structure because of the presence of the
stereochemically active inert pair in thallium.


Copper, silver, and gold
==========================
The group 11 metals (or coinage metals), copper, silver, and gold, are
typically categorised as transition metals given they can form ions
with incomplete d-shells. Physically, they have the relatively low
melting points and high electronegativity values associated with
post-transition metals. "The filled 'd' subshell and free 's' electron
of Cu, Ag, and Au contribute to their high electrical and thermal
conductivity. Transition metals to the left of group 11 experience
interactions between 's' electrons and the partially filled 'd'
subshell that lower electron mobility." Chemically, the group 11
metals behave like main-group metals in their +1 valence states, and
are hence somewhat related to the alkali metals: this is one reason
for their previously being labelled as "group IB", paralleling the
alkali metals' "group IA". They are occasionally classified as
post-transition metals. Their spectra are analogous to those of the
alkali metals. Their monopositive ions are paramagnetic and contribute
no colour to their salts, like those of the alkali metals.

In Mendeleev's 1871 periodic table, copper, silver, and gold are
listed twice, once under group VIII (with the iron triad and platinum
group metals), and once under group IB. Group IB was nonetheless
parenthesised to note that it was tentative. Mendeleev's main
criterion for group assignment was the maximum oxidation state of an
element: on that basis, the group 11 elements could not be classified
in group IB, due to the existence of copper(II) and gold(III)
compounds being known at that time. However, eliminating group IB
would make group I the only main group (group VIII was labelled a
transition group) to lack an A-B bifurcation. Soon afterward, a
majority of chemists chose to classify these elements in group IB and
remove them from group VIII for the resulting symmetry: this was the
predominant classification until the rise of the modern medium-long
18-column periodic table, which separated the alkali metals and group
11 metals.

The coinage metals were traditionally regarded as a subdivision of the
alkali metal group, due to them sharing the characteristic s1 electron
configuration of the alkali metals (group 1: p6s1; group 11: d10s1).
However, the similarities are largely confined to the stoichiometries
of the +1 compounds of both groups, and not their chemical properties.
This stems from the filled d subshell providing a much weaker
shielding effect on the outermost s electron than the filled p
subshell, so that the coinage metals have much higher first ionisation
energies and smaller ionic radii than do the corresponding alkali
metals. Furthermore, they have higher melting points, hardnesses, and
densities, and lower reactivities and solubilities in liquid ammonia,
as well as having more covalent character in their compounds. Finally,
the alkali metals are at the top of the electrochemical series,
whereas the coinage metals are almost at the very bottom. The coinage
metals' filled d shell is much more easily disrupted than the alkali
metals' filled p shell, so that the second and third ionisation
energies are lower, enabling higher oxidation states than +1 and a
richer coordination chemistry, thus giving the group 11 metals clear
transition metal character. Particularly noteworthy is gold forming
ionic compounds with rubidium and caesium, in which it forms the
auride ion (Au−) which also occurs in solvated form in liquid ammonia
solution: here gold behaves as a pseudohalogen because its 5d106s1
configuration has one electron less than the quasi-closed shell
5d106s2 configuration of mercury.


                      Production and isolation
======================================================================
The production of pure alkali metals is somewhat complicated due to
their extreme reactivity with commonly used substances, such as water.
From their silicate ores, all the stable alkali metals may be obtained
the same way: sulfuric acid is first used to dissolve the desired
alkali metal ion and aluminium(III) ions from the ore (leaching),
whereupon basic precipitation removes aluminium ions from the mixture
by precipitating it as the hydroxide. The remaining insoluble alkali
metal carbonate is then precipitated selectively; the salt is then
dissolved in hydrochloric acid to produce the chloride. The result is
then left to evaporate and the alkali metal can then be isolated.
Lithium and sodium are typically isolated through electrolysis from
their liquid chlorides, with calcium chloride typically added to lower
the melting point of the mixture. The heavier alkali metals, however,
are more typically isolated in a different way, where a reducing agent
(typically sodium for potassium and magnesium or calcium for the
heaviest alkali metals) is used to reduce the alkali metal chloride.
The liquid or gaseous product (the alkali metal) then undergoes
fractional distillation for purification. Most routes to the pure
alkali metals require the use of electrolysis due to their high
reactivity; one of the few which does not is the pyrolysis of the
corresponding alkali metal azide, which yields the metal for sodium,
potassium, rubidium, and caesium and the nitride for lithium.

Lithium salts have to be extracted from the water of mineral springs,
brine pools, and brine deposits. The metal is produced
electrolytically from a mixture of fused lithium chloride and
potassium chloride.

Sodium occurs mostly in seawater and dried seabed, but is now produced
through electrolysis of sodium chloride by lowering the melting point
of the substance to below 700 °C through the use of a Downs cell.
Extremely pure sodium can be produced through the thermal
decomposition of sodium azide. Potassium occurs in many minerals, such
as sylvite (potassium chloride). Previously, potassium was generally
made from the electrolysis of potassium chloride or potassium
hydroxide, found extensively in places such as Canada, Russia,
Belarus, Germany, Israel, United States, and Jordan, in a method
similar to how sodium was produced in the late 1800s and early 1900s.
It can also be produced from seawater. However, these methods are
problematic because the potassium metal tends to dissolve in its
molten chloride and vaporises significantly at the operating
temperatures, potentially forming the explosive superoxide. As a
result, pure potassium metal is now produced by reducing molten
potassium chloride with sodium metal at 850 °C.
:Na (g) + KCl (l)  NaCl (l) + K (g)
Although sodium is less reactive than potassium, this process works
because at such high temperatures potassium is more volatile than
sodium and can easily be distilled off, so that the equilibrium shifts
towards the right to produce more potassium gas and proceeds almost to
completion.

Metals like sodium are obtained by electrolysis of molten salts. Rb
& Cs obtained mainly as by products of Li processing. To make pure
caesium, ores of caesium and rubidium are crushed and heated to 650 °C
with sodium metal, generating an alloy that can then be separated via
a fractional distillation technique. Because metallic caesium is too
reactive to handle, it is normally offered as caesium azide (CsN3).
Caesium hydroxide is formed when caesium interacts aggressively with
water and ice (CsOH).

For several years in the 1950s and 1960s, a by-product of the
potassium production called Alkarb was a main source for rubidium.
Alkarb contained 21% rubidium while the rest was potassium and a small
fraction of caesium. Today the largest producers of caesium, for
example the Tanco Mine in Manitoba, Canada, produce rubidium as
by-product from pollucite. Today, a common method for separating
rubidium from potassium and caesium is the fractional crystallisation
of a rubidium and caesium alum (Cs, Rb)Al(SO4)2·12H2O, which yields
pure rubidium alum after approximately 30 recrystallisations. The
limited applications and the lack of a mineral rich in rubidium limit
the production of rubidium compounds to 2 to 4 tonnes per year.
Caesium, however, is not produced from the above reaction. Instead,
the mining of pollucite ore is the main method of obtaining pure
caesium, extracted from the ore mainly by three methods: acid
digestion, alkaline decomposition, and direct reduction. Both metals
are produced as by-products of lithium production: after 1958, when
interest in lithium's thermonuclear properties increased sharply, the
production of rubidium and caesium also increased correspondingly.
Pure rubidium and caesium metals are produced by reducing their
chlorides with calcium metal at 750 °C and low pressure.

As a result of its extreme rarity in nature, most francium is
synthesised in the nuclear reaction 197Au + 18O → 210Fr + 5 n,
yielding francium-209, francium-210, and francium-211. The greatest
quantity of francium ever assembled to date is about 300,000 neutral
atoms, which were synthesised using the nuclear reaction given above.
When the only natural isotope francium-223 is specifically required,
it is produced as the alpha daughter of actinium-227, itself produced
synthetically from the neutron irradiation of natural radium-226, one
of the daughters of natural uranium-238.


                            Applications
======================================================================
Lithium, sodium, and potassium have many useful applications, while
rubidium and caesium are very notable in academic contexts but do not
have many applications yet. Lithium is the key ingredient for a range
of lithium-based batteries, and lithium oxide can help process silica.
Lithium stearate is a thickener and can be used to make lubricating
greases; it is produced from lithium hydroxide, which is also used to
absorb carbon dioxide in space capsules and submarines. Lithium
chloride is used as a brazing alloy for aluminium parts. In medicine,
some lithium salts are used as mood-stabilising pharmaceuticals.
Metallic lithium is used in alloys with magnesium and aluminium to
give very tough and light alloys.

Sodium compounds have many applications, the most well-known being
sodium chloride as table salt. Sodium salts of fatty acids are used as
soap. Pure sodium metal also has many applications, including use in
sodium-vapour lamps, which produce very efficient light compared to
other types of lighting, and can help smooth the surface of other
metals. Being a strong reducing agent, it is often used to reduce many
other metals, such as titanium and zirconium, from their chlorides.
Furthermore, it is very useful as a heat-exchange liquid in fast
breeder nuclear reactors due to its low melting point, viscosity, and
cross-section towards neutron absorption.  Sodium-ion batteries may
provide cheaper alternatives to their equivalent lithium-based cells.
Both sodium and potassium are commonly used as GRAS counterions to
create more water-soluble and hence more bioavailable salt forms of
acidic pharmaceuticals.

Potassium compounds are often used as fertilisers as potassium is an
important element for plant nutrition. Potassium hydroxide is a very
strong base, and is used to control the pH of various substances.
Potassium nitrate and potassium permanganate are often used as
powerful oxidising agents. Potassium superoxide is used in breathing
masks, as it reacts with carbon dioxide to give potassium carbonate
and oxygen gas. Pure potassium metal is not often used, but its alloys
with sodium may substitute for pure sodium in fast breeder nuclear
reactors.

Rubidium and caesium are often used in atomic clocks. Caesium atomic
clocks are extraordinarily accurate; if a clock had been made at the
time of the dinosaurs, it would be off by less than four seconds
(after 80 million years). For that reason, caesium atoms are used as
the definition of the second. Rubidium ions are often used in purple
fireworks, and caesium is often used in drilling fluids in the
petroleum industry.

Francium has no commercial applications, but because of francium's
relatively simple atomic structure, among other things, it has been
used in spectroscopy experiments, leading to more information
regarding energy levels and the coupling constants of the weak
interaction. Studies on the light emitted by laser-trapped
francium-210 ions have provided accurate data on transitions between
atomic energy levels, similar to those predicted by quantum theory.


Metals
========
Pure alkali metals are dangerously reactive with air and water and
must be kept away from heat, fire, oxidising agents, acids, most
organic compounds, halocarbons, plastics, and moisture. They also
react with carbon dioxide and carbon tetrachloride, so that normal
fire extinguishers are counterproductive when used on alkali metal
fires. Some Class D dry powder extinguishers designed for metal fires
are effective, depriving the fire of oxygen and cooling the alkali
metal.

Experiments are usually conducted using only small quantities of a few
grams in a fume hood. Small quantities of lithium may be disposed of
by reaction with cool water, but the heavier alkali metals should be
dissolved in the less reactive isopropanol. The alkali metals must be
stored under mineral oil or an inert atmosphere. The inert atmosphere
used may be argon or nitrogen gas, except for lithium, which reacts
with nitrogen. Rubidium and caesium must be kept away from air, even
under oil, because even a small amount of air diffused into the oil
may trigger formation of the dangerously explosive peroxide; for the
same reason, potassium should not be stored under oil in an
oxygen-containing atmosphere for longer than 6 months.


Ions
======
The bioinorganic chemistry of the alkali metal ions has been
extensively reviewed.
Solid state crystal structures have been determined for many complexes
of alkali metal ions in small peptides, nucleic acid constituents,
carbohydrates and ionophore complexes.

Lithium naturally only occurs in traces in biological systems and has
no known biological role, but does have effects on the body when
ingested. Lithium carbonate is used as a mood stabiliser in psychiatry
to treat bipolar disorder (manic-depression) in daily doses of about
0.5 to 2 grams, although there are side-effects. Excessive ingestion
of lithium causes drowsiness, slurred speech and vomiting, among other
symptoms, and poisons the central nervous system, which is dangerous
as the required dosage of lithium to treat bipolar disorder is only
slightly lower than the toxic dosage. Its biochemistry, the way it is
handled by the human body and studies using rats and goats suggest
that it is an essential trace element, although the natural biological
function of lithium in humans has yet to be identified.

Sodium and potassium occur in all known biological systems, generally
functioning as electrolytes inside and outside cells. Sodium is an
essential nutrient that regulates blood volume, blood pressure,
osmotic equilibrium and pH; the minimum physiological requirement for
sodium is 500 milligrams per day. Sodium chloride (also known as
common salt) is the principal source of sodium in the diet, and is
used as seasoning and preservative, such as for pickling and jerky;
most of it comes from processed foods. The Dietary Reference Intake
for sodium is 1.5 grams per day, but most people in the United States
consume more than 2.3 grams per day, the minimum amount that promotes
hypertension; this in turn causes 7.6 million premature deaths
worldwide.

Potassium is the major cation (positive ion) inside animal cells,
while sodium is the major cation outside animal cells. The
concentration differences of these charged particles causes a
difference in electric potential between the inside and outside of
cells, known as the membrane potential. The balance between potassium
and sodium is maintained by ion transporter proteins in the cell
membrane. The cell membrane potential created by potassium and sodium
ions allows the cell to generate an action potential--a "spike" of
electrical discharge. The ability of cells to produce electrical
discharge is critical for body functions such as neurotransmission,
muscle contraction, and heart function. Disruption of this balance may
thus be fatal: for example, ingestion of large amounts of potassium
compounds can lead to hyperkalemia strongly influencing the
cardiovascular system. Potassium chloride is used in the United States
for lethal injection executions.

Due to their similar atomic radii, rubidium and caesium in the body
mimic potassium and are taken up similarly. Rubidium has no known
biological role, but may help stimulate metabolism, and, similarly to
caesium, replace potassium in the body causing potassium deficiency.
Partial substitution is quite possible and rather non-toxic: a 70 kg
person contains on average 0.36 g of rubidium, and an increase in this
value by 50 to 100 times did not show negative effects in test
persons. Rats can survive up to 50% substitution of potassium by
rubidium. Rubidium (and to a much lesser extent caesium) can function
as temporary cures for hypokalemia; while rubidium can adequately
physiologically substitute potassium in some systems, caesium is never
able to do so. There is only very limited evidence in the form of
deficiency symptoms for rubidium being possibly essential in goats;
even if this is true, the trace amounts usually present in food are
more than enough.

Caesium compounds are rarely encountered by most people, but most
caesium compounds are mildly toxic. Like rubidium, caesium tends to
substitute potassium in the body, but is significantly larger and is
therefore a poorer substitute. Excess caesium can lead to hypokalemia,
arrhythmia, and acute cardiac arrest, but such amounts would not
ordinarily be encountered in natural sources. As such, caesium is not
a major chemical environmental pollutant. The median lethal dose
(LD50) value for caesium chloride in mice is 2.3 g per kilogram, which
is comparable to the LD50 values of potassium chloride and sodium
chloride. Caesium chloride has been promoted as an alternative cancer
therapy, but has been linked to the deaths of over 50 patients, on
whom it was used as part of a scientifically unvalidated cancer
treatment.

Radioisotopes of caesium require special precautions: the improper
handling of caesium-137 gamma ray sources can lead to release of this
radioisotope and radiation injuries. Perhaps the best-known case is
the Goiânia accident of 1987, in which an improperly-disposed-of
radiation therapy system from an abandoned clinic in the city of
Goiânia, Brazil, was scavenged from a junkyard, and the glowing
caesium salt sold to curious, uneducated buyers. This led to four
deaths and serious injuries from radiation exposure. Together with
caesium-134, iodine-131, and strontium-90, caesium-137 was among the
isotopes distributed by the Chernobyl disaster which constitute the
greatest risk to health. Radioisotopes of francium would presumably be
dangerous as well due to their high decay energy and short half-life,
but none have been produced in large enough amounts to pose any
serious risk.


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