Alkali metal

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Not to be confused with Alkaline earth metal.

Alkali metals
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
noble gases ←    → alkaline earth metals
IUPAC group number 1
Name by element lithium group
Trivial name alkali metals
CAS group number
(US, pattern A-B-A)
 IA
old IUPAC number
(Europe, pattern A-B)
 IA
↓ Period
2
Image: Lithium metal stored under paraffin
Lithium (Li)
3
3
Image: Sodium metal
Sodium (Na)
11
4
Image: Potassium metal
Potassium (K)
19
5
Image: Rubidium metal in a glass ampoule
Rubidium (Rb)
37
6
Image: Caesium metal in a glass ampoule
Caesium (Cs)
55
7 Francium (Fr)
87
Legend
primordial
element by radioactive decay

The alkali metals consist of the

s-orbital: this shared electron configuration results in their having very similar characteristic properties.[note 4] Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour.[5]
This family of elements is also known as the lithium family after its leading element.

The alkali metals are all shiny,

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

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

electrolytes
, and although the other alkali metals are not essential, they also have various effects on the body, both beneficial and harmful.

History

A sample of petalite
Petalite, the lithium mineral from which lithium was first isolated

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,[6] and Henri-Louis Duhamel du Monceau was able to prove this difference in 1736.[7] 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.[8][9]

Pure potassium was first isolated in 1807 in England by

caustic soda (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different.[8][9][12][13]

Johann Wolfgang Döbereiner was among the first to notice similarities between what are now known as the alkali metals.

detected the presence of a new element while analysing petalite ore.[17][18] 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.[19] 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.[20][15][18] 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.[21]

A sample of lepidolite
Lepidolite, the rubidium mineral from which rubidium was first isolated

Rubidium and caesium were the first elements to be discovered using the

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).[24][25]

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.[26][27] 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.[21]

Dmitri Mendeleev's periodic system proposed in 1871 showing hydrogen and the alkali metals as part of his group I, along with copper, silver, and gold

After 1869,

d-block, while alkali metals were left in group IA. Later the group's name was changed to group 1 in 1988.[4] 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.[5]

There were at least four erroneous and incomplete discoveries

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.[34] 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%.[35]

227
89Ac
 α (1.38%)→21.77 y 223
87Fr
 β−→22 min 223
88Ra
 α→11.4 d219
86Rn

The next element below francium (eka-francium) in the periodic table would be ununennium (Uue), element 119.[36]: 1729–1730  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.[37][38]

254
99Es
 + 48
20Ca
 → 302
119Uue
* → no atoms[note 6]

It is highly unlikely

drip instabilities, that only the lower period 8 elements, up to around element 128, are physically possible.[41][42] 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.[36]
: 1737–1739 

Occurrence

In the Solar System

Estimated abundances of the chemical elements in the Solar System. Hydrogen and helium are most common, from the Big Bang. The next three elements (lithium, beryllium, and boron) are rare because they are poorly synthesised in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers, and (2) a general decrease in abundance, as elements become heavier. Iron is especially common because it represents the minimum-energy nuclide that can be made by fusion of helium in supernovae.[43]

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.[44][45][46] 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.[43]

On Earth

Spodumene, an important lithium mineral

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 5.98×1024 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.[47]

The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are

ionic radii.[48]

Sodium and potassium are very abundant on Earth, both being among the ten most common elements in Earth's crust;[49][50] sodium makes up approximately 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element overall[51] and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.[51] 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.[51] 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.[10]: 69  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.[10]: 69 

Despite its chemical similarity, lithium typically does not occur together with sodium or potassium due to its smaller size.

ppm, comparable to that of gallium and niobium. Commercially, the most important lithium mineral is spodumene, which occurs in large deposits worldwide.[10]
: 69 

Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite,[55] although none of these contain only rubidium and no other alkali metals.[10]: 70  Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.[56]

Francium-223, the only naturally occurring isotope of francium,[57][58] is the product of the alpha decay of actinium-227 and can be found in trace amounts in uranium minerals.[59] In a given sample of uranium, there is estimated to be only one francium atom for every 1018 uranium atoms.[60][61] 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.[62][63]

Properties

Physical and chemical

The physical and chemical properties of the alkali metals can be readily explained by their having an ns1 valence

radioactivity;[5] 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.[64]

Properties of the alkali metals[10]: 75 [65]
Name Lithium Sodium Potassium Rubidium Caesium Francium
Atomic number 3 11 19 37 55 87
Standard atomic weight[note 7][57][58] 6.94(1)[note 8] 22.98976928(2) 39.0983(1) 85.4678(3) 132.9054519(2) [223][note 9]
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 ?
pm
) 		152 186 227 248 265 ?
Ionic radius of hexacoordinate M+ ion (pm) 76 102 138 152 167 ?
First
kJ·mol−1
) 		520.2 495.8 418.8 403.0 375.7 392.8[67]
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 ?[note 10]
Allen electronegativity 0.91 0.87 0.73 0.71 0.66 0.67
Standard electrode potential (E°(M+→M0); V)[70] −3.04 −2.71 −2.93 −2.98 −3.03 ?
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

densities increase when moving down the table, with the exception that potassium is less dense than sodium.[71] 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.[10]: 75  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.[10]
: 75 

The stable alkali metals are all silver-coloured metals except for caesium, which has a pale golden tint:[72] it is one of only three metals that are clearly coloured (the other two being copper and gold).[10]: 74  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.[10]: 74  Their lustre tarnishes rapidly in air due to oxidation.[5]

Potassium reacts violently with water at room temperature
Caesium reacts explosively with water even at low temperatures

All the alkali metals are highly reactive and are never found in elemental forms in nature.

ionisation energies in their respective periods of the periodic table[64] because of their low effective nuclear charge[5] and the ability to attain a noble gas configuration by losing just one electron.[5] 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.[10]
: 76 

The second ionisation energy of all of the alkali metals is very high

metastable or stable.[79][80]

In aqueous solution, the alkali metal ions form

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.[65][82]: 126–127  While it was previously thought that the heavier alkali metals also formed octahedral hexaaqua ions, it has since been found that potassium and rubidium probably form the [K(H2O)8]+ and [Rb(H2O)8]+ ions, which have the square antiprismatic structure, and that caesium forms the 12-coordinate [Cs(H2O)12]+ ion.[83]

Lithium

The chemistry of lithium shows several differences from that of the rest of the group as the small Li+ cation

organometallic compounds with significant covalent character (e.g. LiMe and MgMe2).[85]

Lithium fluoride is the only alkali metal halide that is poorly soluble in water,

air-conditioners.[10]
: 76 

Francium

Francium is also predicted to show some differences due to its high

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.[67][36]: 1729 [86] 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.[86] Some of the few properties of francium that have been predicted taking relativity into account are the electron affinity (47.2 kJ/mol)[87] and the enthalpy of dissociation of the Fr2 molecule (42.1 kJ/mol).[88] 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.[86] 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.[86]

Nuclear

Primordial isotopes of the alkali metals
Z
 Alkali metal
Stable
Decays

 unstable: italics
odd–odd isotopes coloured pink
3 lithium 2 — 7
Li
 6
Li
  
11 sodium 1 — 23
Na
    
19 potassium 2 1 39
K
 41
K
 40
K
37 rubidium 1 1 85
Rb
 87
Rb
  
55 caesium 1 — 133
Cs
    
87 francium — — No primordial isotopes
(223
Fr
 is a radiogenic nuclide)
Radioactive: 40K, t1/2 1.25 × 109 years; 87Rb, t1/2 4.9 × 1010 years; 223Fr, t1/2 22.0 min.

All the alkali metals have odd atomic numbers; hence, their isotopes must be either

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.[89]

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.[89]

All of the alkali metals except lithium and caesium have at least one naturally occurring

radioactive.[91] Caesium was also thought to be radioactive in the early 20th century,[92][93] although it has no naturally occurring radioisotopes.[91] (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,[94] 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.[30][31] Natural rubidium is similarly slightly radioactive, with 27.83% being the long-lived radioisotope rubidium-87.[10]
: 74 

Chernobyl nuclear power plant.[97] 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.[98]
: 114 

The alkali metals are more similar to each other than the elements in any other

densities increase when moving down the table, with the exception that potassium is less dense than sodium.[71]

Atomic and ionic radii

Effective nuclear charge on an atomic electron

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.[71]

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.[5]

First ionisation energy

Periodic trend for ionisation energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases. Predicted values are used for elements beyond 104.

The first

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.[71] 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.[36]
: 1729 

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.[5]

Reactivity

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.[100]

Electronegativity

main groups of the periodic table from the second to the sixth period
.

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.[71]

Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example,

deliquescent.[5]

Melting and boiling points

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.[71][104] 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.[71] This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons.[71][104] 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.[71] The increased nuclear charge is not a relevant factor due to the shielding effect.[71]

Density

The alkali metals all have the same

body-centred cubic)[10] 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.[71] Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water:[5] in fact, lithium is the least dense known solid at room temperature.[10]
: 75 

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.

metallic bonding along the series NaCl, Na2O, Na2S, Na3P, Na3As, Na3Sb, Na3Bi, Na.[10]
: 81 

Hydroxides

External videos
video icon Alkali Metals – 20 Reactions of the alkali metals with water, conducted by
The Royal Society of Chemistry
A large orange-yellow explosion
A reaction of 3 pounds (≈ 1.4 kg) of sodium with water

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.[100] 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.[5] 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).[74] The alkali metal hydroxides are the most basic known hydroxides.[10]: 87 

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.[105] 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.[105]

The hydroxides themselves are the most basic hydroxides known, reacting with acids to give salts and with alcohols to give

silicate glass.[10]
: 87 

Intermetallic compounds

Liquid NaK alloy at room temperature

The alkali metals form many

eutectic mixture melts at −12.6 °C.[107] An alloy of 41% caesium, 47% sodium, and 12% potassium has the lowest known melting point of any metal or alloy, −78 °C.[22]

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.[108] 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.[109]

Wade's rules to forming Zintl anions like the rest of group 13.[112]

Compounds with the group 14 elements

Side (left) and top (right) views of the graphite intercalation compound KC8

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. M+C−8).[65] 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.[113][114] 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.[115] Apart from this, the alkali metals are such strong reducing agents that they can even reduce buckminsterfullerene to produce solid fullerides MnC60; sodium, potassium, rubidium, and caesium can form fullerides where n = 2, 3, 4, or 6, and rubidium and caesium additionally can achieve n = 1.[10]: 285 

When the alkali metals react with the heavier elements in the

Zintl) ions such as Ge2−4, Ge4−9, Ge2−9, and [(Ge9)2]6−, is largely analogous to that of the corresponding silicides.[10]: 393  Alkali metal stannides are mostly ionic, sometimes with the stannide ion (Sn4−),[109] and sometimes with more complex Zintl ions such as Sn4−9, which appears in tetrapotassium nonastannide (K4Sn9).[116] 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 Pb4−9. These alkali metal germanides, stannides, and plumbides may be produced by reducing germanium, tin, and lead with sodium metal in liquid ammonia.[10]
: 394 

Nitrides and pnictides

tetrahedral structure would be expected, but that would be geometrically impossible: thus lithium nitride takes on this unique crystal structure.[10]
: 76 

Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with

endothermic, so they do not form nitrides at standard conditions.[84] 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.[118][119] Steric hindrance forbids the existence of rubidium or caesium nitride.[10]: 417  However, sodium and potassium form colourless azide salts involving the linear N−3 anion; due to the large size of the alkali metal cations, they are thermally stable enough to be able to melt before decomposing.[10]
: 417 

All the alkali metals react readily with

intermetallic compounds containing partially metallic and partially ionic bonds.[123]

Oxides and chalcogenides

See also:
Alkali metal oxide
octahedra connected to each other by one face
The ball-and-stick diagram shows three regular octahedra where each octahedron is connected to both of the others by one face each. All three octahedra have one edge in common. All eleven vertices of the structure are violet spheres representing caesium, and at the centre of each octahedron is a small red sphere representing oxygen.
Cs11O3 cluster, composed of three regular octahedra where each octahedron is connected to both of the others by one face each. All three octahedra have one edge in common.

All the alkali metals react vigorously with

pyrophoric (spontaneously catch fire in air).[84]

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.

water vapour, naturally present in breath, makes the removal of carbon dioxide by potassium superoxide even more efficient.[65][124] 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.[10]: 85  Potassium, rubidium, and caesium also form sesquioxides M2O3, which may be better considered peroxide disuperoxides, [(M+)4(O2−2)(O−2)2].[10]
: 85 

Rubidium and caesium can form a great variety of suboxides with the metals in formal oxidation states below +1.[10]: 85  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[125][126] and several brightly coloured suboxides,[127] such as Cs7O (bronze), Cs4O (red-violet), Cs11O3 (violet), Cs3O (dark green),[128] CsO, Cs3O2,[129] as well as Cs7O2.[130][131] The last of these may be heated under vacuum to generate Cs2O.[56]

The alkali metals can also react analogously with the heavier chalcogens (

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 Se2−
x and Te2−
x ions.[132] 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.[10]: 766  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.[10]: 766 [133][134]

Halides, hydrides, and pseudohalides

Main article: Alkali metal halide

The alkali metals are among the most

electrostatic interactions between them to be strong:[5] a similar effect occurs for magnesium fluoride, consistent with the diagonal relationship between lithium and magnesium.[10]
: 76 

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.[10]: 83 [65] 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.[10]: 322  Ternary alkali metal halide oxides, such as Na3ClO, K3BrO (yellow), Na4Br2O, Na4I2O, and K4Br2O, are also known.[10]: 83  The polyhalides are rather unstable, although those of rubidium and caesium are greatly stabilised by the feeble polarising power of these extremely large cations.[10]: 835 

Coordination complexes

18-crown-6 coordinating a potassium ion
Structure of 2.2.2-Cryptand encapsulating a potassium cation (purple). At crystalline state, obtained with an X-ray diffraction.[135]

Alkali metal cations do not usually form

21-crown-7 for Rb+) and cryptands due to electrostatic attraction.[65]

Ammonia solutions

The alkali metals dissolve slowly in liquid

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.[65]

Organometallic

Organolithium

Main article: Organolithium reagent
n-butyllithium hexamer, (C4H9Li)6.[137] The aggregates are held together by delocalised covalent bonds between lithium and the terminal carbon of the butyl chain.[138] There is no direct lithium–lithium bonding in any organolithium compound.[122]
: 264 
Solid phenyllithium forms monoclinic crystals that can be described as consisting of dimeric Li2(C6H5)2 subunits. The lithium atoms and the ipso carbons of the phenyl rings form a planar four-membered ring. The plane of the phenyl groups is perpendicular to the plane of this Li2C2 ring. Additional strong intermolecular bonding occurs between these phenyllithium dimers and the π electrons of the phenyl groups in the adjacent dimers, resulting in an infinite polymeric ladder structure.[139]

Being the smallest alkali metal, lithium forms the widest variety of and most stable

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.[10]
: 105 

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.[10]: 105  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.[10]: 106 

Heavier alkali metals

Unlike the organolithium compounds, the organometallic compounds of the heavier alkali metals are predominantly ionic. The application of

sterically congested conformation.[141] Several crystal structures of organopotassium compounds have been reported, establishing that they, like the sodium compounds, are polymeric.[142] Organosodium, organopotassium, organorubidium and organocaesium compounds are all mostly ionic and are insoluble (or nearly so) in nonpolar solvents.[65]

Alkyl and aryl derivatives of sodium and potassium tend to react with air. They cause the cleavage of

Wurtz coupling occurs:[122]
: 265 

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.[122]
: 265 

The alkali metals and their hydrides react with acidic hydrocarbons, for example

sodium naphthalenide, Na+[C10H8•]−, a strong reducing agent.[122]
: 266 

Representative reactions of alkali metals

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[143] 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[144]
Li2O + H2O → 2LiOH

Reaction with sulfur

With sulfur, they form sulfides and polysulfides.[145]

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.

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.[146]

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.[143]

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.[143]

Na + xNH3 → Na+ + e(NH3)x−

Due to the presence of solvated electrons, these solutions are very powerful reducing agents used in organic synthesis.

Reduction reactions using sodium in liquid ammonia

Reaction 1) is known as Birch reduction. Other reductions[143] that can be carried by these solutions are:

S8 + 2e− → S82-
Fe(CO)5 + 2e− → Fe(CO)42- + CO

Extensions

ninth period, measured in angstroms[36]: 1730 [147]

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

polarisability of ununennium.[36]: 1729–1730  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.[36]
: 1724 

electron volts[36]: 1730 [147]

The stabilisation of ununennium's valence electron and thus the contraction of the 8s orbital cause its atomic radius to be lowered to 240 

covalent character, due to the involvement of the 7p3/2 electrons in the bonding.[86]

Empirical (Na–Fr, Mg–Ra) and predicted (Uue–Uhp, Ubn–Uhh) ionisation energy of the alkali and alkaline earth metals from the third to the ninth period, measured in electron volts[36]: 1730 [147]

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

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).[36]: 1729–1733  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.[151][152]

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.[147] 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.[36]: 1732–1733  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).[153][154][155]

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.[156]

Hydrogen

The element

diatomic gas consisting of two atoms per molecule (H2);[159] however, the alkali metals form diatomic molecules (such as dilithium, Li2) only at high temperatures, when they are in the gaseous state.[160]

Hydrogen, like the alkali metals, has one valence electron[122] and reacts easily with the halogens,[122] but the similarities mostly end there because of the small size of a bare proton H+ compared to the alkali metal cations.[122] Its placement above lithium is primarily due to its electron configuration.[157] It is sometimes placed above fluorine due to their similar chemical properties, though the resemblance is likewise not absolute.[161]

The first ionisation energy of hydrogen (1312.0

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.[165]

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

electrostatic, and electron correlative contributing phenomena.[161] While analogous lithium bonds are also known, they are mostly electrostatic.[161] 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).[166]

Ammonium and derivatives

Similarly to the alkali metals, ammonia reacts with hydrochloric acid to form the salt ammonium chloride.

The

soluble in water, a property which ammonium salts share.[172] Ammonium is expected to behave stably as a metal (NH+4 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.[170][171] 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.[170] Under standard conditions, ammonium can form a metallic amalgam with mercury.[173]

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 (NR+4) 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 HI−2.[10]: 812–9  Tetraalkylammonium hydroxides, like alkali metal hydroxides, are very strong bases that react with atmospheric carbon dioxide to form carbonates.[122]: 256  Furthermore, the nitrogen atom may be replaced by a phosphorus, arsenic, or antimony atom (the heavier nonmetallic pnictogens), creating a phosphonium (PH+4) or arsonium (AsH+4) cation that can itself be substituted similarly; while stibonium (SbH+4) itself is not known, some of its organic derivatives are characterised.[156]

Cobaltocene and derivatives

Iridocene (involving iridium) would presumably be still more potent, but is not very well-studied due to its instability.[176]

Thallium

Very pure thallium pieces in a glass ampoule, stored under argon gas

TlI has an unusual structure because of the presence of the stereochemically active inert pair in thallium.[180]

Copper, silver, and gold

A crystal of a coppery-colored metal mineral of standing on a white surface
Copper
A crystal of a silvery metal crystal lying on a grey surface
Silver
A crystal of a yellow metal lying on a white surface
Gold

The

paramagnetic and contribute no colour to their salts, like those of the alkali metals.[183]

In Mendeleev's 1871 periodic table, copper, silver, and gold are listed twice, once under group VIII (with the

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.[29] 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.[29] 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.[29]

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

electrochemical series, whereas the coinage metals are almost at the very bottom.[10]: 1177  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.[10]: 1177  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.[10]
: 1177 

Production and isolation

alt1
alt2
Salt flats are rich in lithium, such as these in Salar del Hombre Muerto, Argentina (left) and Uyuni, Bolivia (right). The lithium-rich brine is concentrated by pumping it into solar evaporation ponds (visible in Argentina image).

The production of pure alkali metals is somewhat complicated due to their extreme reactivity with commonly used substances, such as water.[5][65] 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.[65] 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.[65] 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.[122]: 77 

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.[184]

Sodium occurs mostly in seawater and dried seabed,[5] 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.[185][186] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.[187] Potassium occurs in many minerals, such as sylvite (potassium chloride).[5] Previously, potassium was generally made from the electrolysis of potassium chloride or potassium hydroxide,[188] 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.[189] It can also be produced from seawater.[5] 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.[10]: 74 

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.[10]: 74 

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).[190]

Rubidium is the 16th most abundant element in the earth's crust; however, it is quite rare. Some minerals found in North America, South Africa, Russia, and Canada contain rubidium. Some potassium minerals (lepidolites, biotites, feldspar, carnallite) contain it, together with caesium. Pollucite, carnallite, leucite, and lepidolite are all minerals that contain rubidium. As a by-product of lithium extraction, it is commercially obtained from lepidolite. Rubidium is also found in potassium rocks and brines, which is a commercial supply. The majority of rubidium is now obtained as a byproduct of refining lithium. Rubidium is used in vacuum tubes as a getter, a material that combines with and removes trace gases from vacuum tubes.[191][192]

A shiny gray 5-centimeter piece of matter with a rough surface.
This sample of uraninite contains about 100,000 atoms (3.3×10−20 g) of francium-223 at any given time.[60]

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.[193] Today the largest producers of caesium, for example the Tanco Mine in Manitoba, Canada, produce rubidium as by-product from pollucite.[194] 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.[194][195] 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.[194] 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.[194][196] 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.[10]: 71  Pure rubidium and caesium metals are produced by reducing their chlorides with calcium metal at 750 °C and low pressure.[10]: 74 

As a result of its extreme rarity in nature,

francium-211.[197] The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,[198] which were synthesised using the nuclear reaction given above.[198] 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.[199]

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.[10]: 68  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.[10]: 70  Lithium chloride is used as a brazing alloy for aluminium parts.[200] 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.[10]: 70 

Sodium compounds have many applications, the most well-known being sodium chloride as

GRAS counterions to create more water-soluble and hence more bioavailable salt forms of acidic pharmaceuticals.[206]

Potassium compounds are often used as

fertilisers[10]: 73 [207] 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.[208][209] Potassium nitrate and potassium permanganate are often used as powerful oxidising agents.[10]: 73  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.[10]
: 74 

Rubidium and caesium are often used in

fireworks,[212] and caesium is often used in drilling fluids in the petroleum industry.[56][213]

Francium has no commercial applications,

atomic structure, among other things, it has been used in spectroscopy experiments, leading to more information regarding energy levels and the coupling constants between subatomic particles.[215] 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.[216]

Biological role and precautions

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.[217] Some Class D dry powder extinguishers designed for metal fires are effective, depriving the fire of oxygen and cooling the alkali metal.[218]

Experiments are usually conducted using only small quantities of a few grams in a

isopropanol.[217][219] 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.[217] 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.[220][221]

Ions

Lithium carbonate

The bioinorganic chemistry of the alkali metal ions has been extensively reviewed.[222] Solid state crystal structures have been determined for many complexes of alkali metal ions in small peptides, nucleic acid constituents, carbohydrates and ionophore complexes.[223]

Lithium naturally only occurs in traces in biological systems and has no known biological role, but does have effects on the body when ingested.

essential trace element, although the natural biological function of lithium in humans has yet to be identified.[226][227]

Sodium and potassium occur in all known biological systems, generally functioning as

electrolytes inside and outside cells.[228][229] 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.[230] 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.[231] The Dietary Reference Intake for sodium is 1.5 grams per day,[232] but most people in the United States consume more than 2.3 grams per day,[233] the minimum amount that promotes hypertension;[234] this in turn causes 7.6 million premature deaths worldwide.[235]

Potassium is the major

cation (positive ion) inside animal cells,[228] while sodium is the major cation outside animal cells.[228][229] 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.[236] 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.[236] 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.[237][238] Potassium chloride is used in the United States for lethal injection executions.[237]

A wheel type radiotherapy device which has a long collimator to focus the radiation into a narrow beam. The caesium-137 chloride radioactive source is the blue square, and gamma rays are represented by the beam emerging from the aperture. This was the radiation source involved in the Goiânia accident, containing about 93 grams of caesium-137 chloride.

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,[239][240][241] and, similarly to caesium,[239][242] replace potassium in the body causing potassium deficiency.[239][241] 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.[243] Rats can survive up to 50% substitution of potassium by rubidium.[241][244] 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.[240] 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.[245][246]

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.[242] Excess caesium can lead to hypokalemia, arrhythmia, and acute cardiac arrest,[247] but such amounts would not ordinarily be encountered in natural sources.[248] As such, caesium is not a major chemical environmental pollutant.[248] 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.[249] Caesium chloride has been promoted as an alternative cancer therapy,[250] but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.[251]

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.[97] 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.[199]

Notes

  1. ^ The symbols Na and K for sodium and potassium are derived from their Latin names, natrium and kalium; these are still the origins of the names for the elements in some languages, such as German and Russian.
  2. ^ Caesium is the spelling recommended by the International Union of Pure and Applied Chemistry (IUPAC).[1] The American Chemical Society (ACS) has used the spelling cesium since 1921,[2][3] following Webster's Third New International Dictionary.
  3. Roman numeral).[4]
  4. ^ While hydrogen also has this electron configuration, it is not considered an alkali metal as it has very different behaviour owing to the lack of valence p-orbitals in period 1 elements.
  5. ^ In the 1869 version of Mendeleev's periodic table, copper and silver were placed in their own group, aligned with hydrogen and mercury, while gold was tentatively placed under uranium and the undiscovered eka-aluminium in the boron group.
  6. ^ The asterisk denotes an excited state.
  7. least significant figure(s) of the number prior to the parenthesised value (ie. counting from rightmost digit to left). For instance, 1.00794(7) stands for 1.00794±0.00007, while 1.00794(72) stands for 1.00794±0.00072.[66]
  8. ^ The value listed is the conventional value suitable for trade and commerce; the actual value may range from 6.938 to 6.997 depending on the isotopic composition of the sample.[58]
  9. ^ The element does not have any stable nuclides, and a value in brackets indicates the mass number of the longest-lived isotope of the element.[57][58]
  10. relativistic effects
    , and this would imply that caesium is the less electronegative of the two.

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