alkali metal

Group →	1
↓ Period
2	3 Li
3	11 Na
4	19 K
5	37 Rb
6	55 Cs
7	87 Fr
Legend Alkali metal Primordial From decay

The alkali metals are a group of chemical elements in the periodic table with very similar properties: they are all shiny, soft, silvery, highly reactive metals at standard temperature and pressure^[1] and readily lose their outermost electron to form cations with charge +1.^[2]:28 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.^[1] In the modern IUPAC nomenclature, the alkali metals comprise the group 1 elements,^{[note 1]} excluding hydrogen (H), which is nominally a group 1 element^[4][5] but not normally considered to be an alkali metal^[6][7] as it rarely exhibits behaviour comparable to that of the alkali metals.^[8] All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.^[1][9]

The alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).^[4] This group lies in the s-block of the periodic table^[10] as all alkali metals have their outermost electron in an s-orbital.^[1][11][12] The alkali metals provide the best example of group trends in properties in the periodic table,^[1] with elements exhibiting well-characterized homologous behaviour.^[1]

All the discovered alkali metals occur in nature.^[13][14] Experiments have been conducted to attempt the synthesis of ununennium (Uue), which is likely to be the next member of the group, but they have all met with failure.^[15] 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.^[16]

Most alkali metals have many different applications. Two of the most well-known applications of the pure elements are rubidium and caesium atomic clocks,^[17] of which caesium atomic clocks are the most accurate representation of time known as of 2012.^[18][19] A common application of the compounds of sodium is the sodium vapour lamp, which emits very efficient light.^[20][21] Table salt, or sodium chloride, has been used since antiquity.

Characteristics

Chemical

Series of alkali metals, stored in mineral oil ("Natrium" is sodium.)

Like other groups, the members of this family show patterns in its electronic configuration, especially the outermost shells, resulting in trends in chemical behavior:

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 radioactivity,^[1] and ununennium has not yet been discovered; thus, the presentation of their properties here is limited. All the alkali metals are all highly reactive and are never found in elemental forms in nature.^[23] Because of this, they are usually stored in mineral oil or kerosene (paraffin oil).^[24]

Caesium reacts explosively with water even at low temperatures

The alkali metals are all silver-coloured except for metallic caesium, which can have a golden tint.^[25] All are soft and have low densities,^[1] melting points,^[1] and boiling points.^[1] In chemical terms, all of the alkali metals 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).^[1] 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.^[1][9][18] The alkali metals have the lowest first ionisation energies in their respective periods of the periodic table^[12] because of their low effective nuclear charge^[1] and the ability to attain a noble gas configuration by losing just one electron. The second ionisation energy of all of the alkali metals is very high^[1][12] as it is in a full shell that is also closer to the nucleus;^[1] thus, they almost always lose a single electron, forming cations.^[2]:28

The chemistry of lithium is anomalous as the small Li⁺ cation polarises anions and gives its compounds a more covalent character.^[1] Lithium and magnesium have a diagonal relationship.^[1] Lithium fluoride is the only alkali metal halide that is not soluble in water,^[1] and lithium hydroxide is the only alkali metal hydroxide that is not deliquescent.^[1]

The chemistry of ununennium, the undiscovered seventh alkali metal, is predicted to be closer to that of potassium^[26] or rubidium^{[22]:1729–1730} instead of caesium or francium. This is unusual as periodic trends would predict ununennium to be even more reactive than caesium and francium. This lowered reactivity is due to the energetic properties^{[clarification needed]} of ununennium's valence electron, increasing ununennium's ionisation energy and decreasing the metallic and ionic radii.^[26] It may also show the +3 oxidation state,^{[22]:1729-1730} which is not seen in any other alkali metal,^[2]:28 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.^{[2]:28[22]:1729–1730} This assumes that ununennium will behave chemically as an alkali metal, which may not be true due to relativistic effects.^[16]

Compounds and reactions

Reaction with water

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. 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).^[9]

Physical and atomic

The table below is a summary of the key physical and atomic properties of the alkali metals. Data marked with question marks are either uncertain or are estimations partially based on periodic trends rather than observations.

Alkali metal	Standard atomic weight (u)[note 3][28][29]	Melting point (K)	Melting point (°C)	Boiling point (K)[12]	Boiling point (°C)[12]	Density (g/cm3)	Electronegativity (Pauling)	First ionisation energy (kJ·mol−1)	Covalent radius (pm)[30]	Flame test colour
Lithium	6.941(2)[note 4]	453.69	180.54	1615	1342	0.534	0.98	520.2	145	Red[1][31]
Sodium	22.98976928(2)	370.87	97.72	1156	883	0.968	0.93	495.8	180	Strong persistent orange or yellow[1][31]
Potassium	39.0983(1)	336.53	63.38	1032	759	0.89	0.82	418.8	220	Lilac or pink[1][31]
Rubidium	85.4678(3)	312.46	39.31	961	688	1.532	0.82	403.0	235	Red or reddish-violet[1][31]
Caesium	132.9054519(2)	301.59	28.44	944	671	1.93	0.79	375.7	260	Blue or violet[1][31]
Francium	[223][note 5]	? 300	? 27	? 950[32]	? 677[32]	? 1.87	? 0.7	380	?	Unknown

Radioactivity

All of the alkali metals except lithium and caesium have at least one naturally occurring radioisotope: sodium-22 and sodium-24 are trace radioisotopes, 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,^[33][34] although it has no naturally occurring radioisotopes.^{[citation needed]} (Francium had not been discovered yet at that time.)^{[citation needed]}

The natural radioisotope of potassium, potassium-40, makes up about 0.012% of natural potassium,^[35] and thus natural potassium is weakly radioactive. The Soviet scientist D. K. Dobroserdov observed this weak radioactivity in a sample of potassium in 1925 and incorrectly assumed that eka-caesium (currently known to be francium) was contaminating the sample.^[36] He then claimed to have discovered eka-caesium and predicted its properties, naming it russium after his home country.^[37] Shortly thereafter, Dobroserdov began to focus on his teaching career at the Polytechnic Institute of Odessa, and he did not pursue the element further.^[36]

Periodic trends

The alkali metals are more similar to each other than the elements in any other group are to each other.^[1] For instance, when moving down the table, all alkali metals show increasing atomic radius,^[38] decreasing electronegativity,^[38] increasing reactivity,^[1] and decreasing melting and boiling points.^[38] In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium.^[38]

Atomic and ionic radii

Effective nuclear charge on an atomic electron

The atomic radii of the alkali metals increase going down the group.^[38] 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.^[39] 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.^[38]

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.^[1]

First ionisation energy and reactivity

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 feel 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.^[38] (This trend is broken in francium due to relativistic effects.)^{[citation needed]} Therefore, it is easier for the outer electron to be removed from the atom and participate in chemical reactions, thus increasing reactivity down the group.^{[citation needed]}

Electronegativity

The variation of Pauling electronegativity (y-axis) as one descends the main groups of the periodic table from the second period to the sixth period

Electronegativity is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself.^[45] 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.^[38]

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.^[38] Lithium fluoride (LiF) is the only alkali halide that is not soluble in water,^[1] and lithium hydroxide (LiOH) is the only alkali metal hydroxide that is not deliquescent.^[1]

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 vapor pressure of the liquid equals the environmental pressure surrounding the liquid^[47][48] 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.^[38][49] 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.^[38] This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons.^[38][49] 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.^[38] (The increased nuclear charge is not a relevant factor due to the shielding effect.)^[38]

Density

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Other substances similar to the alkali metals

Hydrogen

This section requires expansion.

Hydrogen gas glowing in a discharge tube

The element hydrogen, with one electron per neutral atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not normally considered to be an alkali metal;^[6] when it is considered to be an alkali metal, it is because of its atomic properties and not its chemical properties.^[7] Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule (H₂).^[50]

Hydrogen, like the alkali metals, has one valence electron^[8] and reacts easily with the halogens,^[8] but the similarities end there.^[8] Its placement above lithium is primarily due to its electron configuration and not its chemical properties.^[6][8] It is sometimes placed above carbon due to their similar electronegativities^[51] or fluorine due to their similar chemical properties.^[8][51]

The first ionisation energy of hydrogen (1312.0 kJ/mol) is much higher than that of the alkali metals.^[40][41] 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 sometimes considered to be a halogen.^[8] The alkali metals can also form negative ions, known as alkalides. Under extremely high pressures, such as those found at the cores of Jupiter and Saturn, hydrogen does become metallic^[52] and behaves like an alkali metal; in this phase, it is known as metallic hydrogen.

Ammonium

The ammonium ion (NH+
4) has very similar properties to the heavier alkali metals and is often considered a close relative.^[53][54][55] For example, all alkali metal salts are soluble in water, a property which ammonium salts share.^{[citation needed]} Ammonium is expected to behave as a metal (NH+
4 ions in a sea of electrons) at very high pressures, such as inside the ice giants Uranus and Neptune.^[54][55]

Thallium

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

Mendeleev's 1869 periodic table shows thallium as an alkali metal.

Because thallium displays the +1 oxidation state^[2]:28 that all the known alkali metals display,^[2]:28 and thallium compounds with thallium in its +1 oxidation state closely resemble the corresponding potassium or silver compounds,^{[citation needed]} thallium was sometimes considered an alkali metal in continental Europe (but not in England) in the years immediately following its discovery,^[56]:126 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.^[57] (Mendeleev's 1871 periodic table and Meyer's 1870 periodic table put thallium in its current position in the boron group and leave the space below caesium blank.)^[57] However, thallium also displays the oxidation state +3,^[2]:28 which no known alkali metal displays^[2]:28 (although ununennium, the undiscovered seventh alkali metal, is predicted to possibly display the +3 oxidation state).^{[22]:1729–1730} The sixth alkali metal is now considered to be francium.^[4]

History

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Etymology

This section requires expansion.

The alkali metals are so called because their hydroxides are all strong alkalis when dissolved in water.^[1]

Discovery

Lithium

Petalite, the lithium mineral from which lithium was first isolated

Petalite (LiAlSi₄O₁₀) was discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada e Silva in a mine on the island of Utö, Sweden.^[58][59][60] However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jöns Jakob Berzelius, detected the presence of a new element while analyzing petalite ore.^[61][62] This new element formed 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.^[63] 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".^[23][59][62]

Sodium

Caustic soda (sodium hydroxide), the sodium compound from which sodium was first isolated

Sodium compounds have been known since ancient times; salt (sodium chloride) has been an important commodity in human activities, as testified by the English word salary, referring to salarium, the wafers of salt sometimes given to Roman soldiers along with their other wages.^{[citation needed]} In medieval Europe a compound of sodium^{[clarification needed]} with the Latin name of sodanum was used as a headache remedy.^{[citation needed]} Pure sodium was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda (now called sodium hydroxide),^[64] a very similar method to the one used to isolate potassium earlier that year.

Potassium

Caustic potash (potassium hydroxide), the potassium compound from which potassium was first isolated

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,^[65] and Henri Louis Duhamel du Monceau was able to prove this difference in 1736.^[66] 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 include the alkali in his list of chemical elements in 1789.^[67][68] Pure potassium was first isolated in 1807 in England by Sir 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. Potassium was the first metal that was isolated by electrolysis.^[69] 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.^{[64][67][68][70]}

Rubidium

Lepidolite, the rubidium mineral from which rubidium was first isolated

Rubidium was discovered in 1861 in Heidelberg, Germany by Robert Bunsen and Gustav Kirchhoff, the first people to suggest finding new elements by spectrum analysis, in the mineral lepidolite through the use of a spectroscope. Because of the bright red lines in its emission spectrum, they chose a name derived from the Latin word rubidus, meaning dark red.^[71][72] Rubidium's discovery succeeded that of caesium, also discovered by Bunsen and Kirchhoff through spectroscopy.^[73]

Caesium

In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in the mineral water from Dürkheim, Germany. Due to the bright-blue lines in its emission spectrum, they chose a name derived from the Latin word caesius, meaning sky-blue.^{[71][note 11][74]} Caesium was the first element to be discovered spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.^[73]

Francium

There were at least four erroneous and incomplete discoveries^{[36][37][75][76]} 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.^[14] 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%.^[77]

Eka-francium

The next element below francium (eka-francium) is likely to be ununennium (Uue), element 119,^{[citation needed]} although this is not certain due to relativistic effects.^[16] 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 Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.^[15][78]

^{[note 12]}

It is highly unlikely^[15] 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 ²⁵⁴Es 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. 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.^{[citation needed]} Currently, none of the period 8 elements have been discovered yet. It is also possible, due to drip instabilities, that only the lower period 8 elements are physically possible.^[26]

Occurrence

Spodumene, an important lithium mineral

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.^{[citation needed]} Potassium, rubidium and caesium are also incompatible elements due to their low ionic radii.^[79] Sodium and potassium are among the ten most common elements in Earth's crust.^[13][80]

Lithium, due to its relatively low reactivity, can be found in seawater in large amounts; it is estimated that seawater is approximately 0.14 to 0.25 parts per million (ppm)^[81][82] or 25 micromolar.^[83]

Sodium and potassium are very abundant in earth; sodium makes up approximately 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element overall^[84] and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.^[84] 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.^[84]

Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.^[18] Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite.^[85]

Francium-223, the only naturally occurring isotope of francium,^[28][29] is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals.^[86] In a given sample of uranium, there is estimated to be only one francium atom for every 1 × 10¹⁸ uranium atoms.^[87][88] It has been calculated that there is at most 30 g of francium in the earth's crust at any time, due to its extremely short half-life of 22 minutes.^[89][90]

Production

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 difficult due to their extreme reactivity with commonly used substances, such as water.^{[citation needed]} The alkali metals are so reactive that they cannot be displaced by other elements and must be isolated through electrolysis.^[1]

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.^[91]

Potassium occurs in many minerals, such as sylvite (potassium chloride).^[1] It is occasionally produced through separating the potassium from the chlorine in potassium chloride, but is more often produced through electrolysis of potassium hydroxide,^[92] 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.^[93] It can also be produced from seawater. Sodium occurs mostly in seawater and dried seabed,^[1] 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.^[94][95] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.^[96]

Pollucite ore, from which caesium can be extracted

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.^[97] Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite.^[98] Today, a common method for separating rubidium from potassium and caesium is the fractional crystallization of a rubidium and caesium alum (Cs,Rb)Al(SO₄)₂·12H₂O, which yields pure rubidium alum after approximately 30 different reactions.^[98][99] The limited applications and the lack of a mineral rich in rubidium limits the production of rubidium compounds to 2 to 4 tonnes per year.^[98] 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.^[98][100]

This sample of uraninite contains about 100,000 atoms (3.3×10⁻²⁰ g) of francium-223 at any given time.^[87]

Francium-223, the only naturally occurring isotope of francium,^[28][29] is produced naturally as the product of the alpha decay of actinium-227. Francium can be found in trace amounts in uranium and thorium minerals;^[86] it has been calculated that at most there are 30 g of francium in the earth's crust at any given time.^[89] As a result of its extreme rarity in nature, most francium is synthesized in the nuclear reaction ¹⁹⁷Au + ¹⁸O → ²¹⁰Fr + 5 n, yielding francium-209, francium-210, and francium-211.^[101] The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,^[102] which were synthesized using the nuclear reaction given above.^[102]

Applications

All of the discovered alkali metals excluding francium have many applications. Lithium is often used in batteries, and lithium oxide can help process silica. Lithium can also be used to make lubricating greases, air treatment, and aluminium production.^[103]

Pure sodium has many applications, including use in sodium vapour lamps, which produce very efficient light compared to other types of lighting,^[20][21] and can help smooth the surface of other metals.^[104][105] Sodium compounds have many applications as well, the most well-known compound being table salt.^{[citation needed]} Sodium is also used in soap as salts of fatty acids.^{[citation needed]}

Potassium compounds are often used as fertilisers^[2]:73[106] as potassium is an important element for plant nutrition. Other potassium ions are often used to hold anions.^{[citation needed][clarification needed]} Potassium hydroxide is a very strong base, and is used to control the pH of various substances.^[107][108]

FOCS 1, a caesium atomic clock in Switzerland

Rubidium and caesium are often used in atomic clocks.^[17] 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 two seconds.^[18] For that reason, caesium atoms are used as the definition of the second.^[19] Rubidium ions are often used in purple fireworks,^[109] and caesium is often used in drilling fluids in the petroleum industry.^[18][110]

Francium has no commercial applications,^{[87][88][111]} 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 between subatomic particles.^[112] 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.^[113]

Biological role and precautions

Lithium carbonate

Lithium does not occur naturally in biological systems and has no biological role, but does have effects on the body when ingested.^[114] 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.^[114] Excessive ingestion of lithium causes drowsiness, slurred speech and vomiting, among other symptoms,^[114] and poisons the central nervous system,^[114] which is dangerous as the required dosage of lithium to treat bipolar disorder is only slightly lower than the toxic dosage.^[114][115]

Sodium and potassium occur in all known biological systems, generally functioning as electrolytes inside and outside cells.^[116][117] 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.^[118] 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.^[119] The DRI for sodium is 1.5 grams per day,^[120] but most people in the United States consume more than 2.3 grams per day,^[121] the minimum amount that promotes hypertension;^[122] this in turn causes 7.6 million premature deaths worldwide.^[123]

Potassium is the major cation (positive ion) inside animal cells,^[116] while sodium is the major cation outside animal cells.^[116][117] 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 pumps in the cell membrane. All potassium ion channels are tetramers with several conserved secondary structural elements. The most recently resolved potassium ion channel is KirBac3.1, which gives a total of five potassium ion channels (KcsA, KirBac1.1, KirBac3.1, KvAP, MthK) with a determined structure.^{[124][clarification needed]} All five are from prokaryotic species.^[124] 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.^[124]

Rubidium has no known biological role, but may help stimulate metabolism,^{[125][126][127]} and, similarly to caesium,^[127][128] replace potassium in the body causing potassium deficiency.^[126][127]

Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium, allowing the caesium to replace the potassium in the body, causing potassium deficiency.^[128] Exposure to large amounts of caesium compounds can cause hyperirritability and spasms, but as such amounts would not ordinarily be encountered in natural sources, caesium is not a major chemical environmental pollutant.^[129] 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.^[130] Caesium chloride has been promoted as an alternative cancer therapy,^[131] but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.^[132]

Francium has no biological role^[133] and is most likely to be toxic due to its extreme radioactivity, causing radiation poisoning, but since the greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,^[102] it is unlikely that most people will ever encounter francium.

Notes

References

Further reading

• Bauer, Brent A., Robert Houlihan, Michael J. Ackerman, Katya Johnson, and Himeshkumar Vyas (2006). "Acquired Long QT Syndrome Secondary to Cesium Chloride Supplement". Journal of Alternative and Complementary Medicine 12 (10): 1011–1014. doi:10.1089/acm.2006.12.1011. PMID 17212573. 
• Campbell, Linda M., Aaron T. Fisk, Xianowa Wang, Gunter Kock, and Derek C. Muir (2005). "Evidence for Biomagnification of Rubidium in Freshwater and Marine Food Webs". Canadian Journal of Fisheries and Aquatic Sciences 62 (5): 1161–1167. doi:10.1139/f05-027. 
• Chang, Cheng-Hung, and Tian Y. Tsong (2005). "Stochastic Resonance of Na, K-Ion Pumps on the Red Cell Membrane". AIP Conference Proceedings: 18th International Conference on Noise and Fluctuations. 780. American Institute of Physics. pp. 587–590. doi:10.1063/1.2036821. ISBN 0-7354-0267-1. 
• Erermis, Serpil, Muge Tamar, Hatice Karasoy, Tezan Bildik, Eyup S. Ercan, and Ahmet Gockay (1997). "Double-Blind Randomised Trial of Modest Salt Restriction in Older People". Lancet 350 (9081): 850–854. doi:10.1016/S0140-6736(97)02264-2. PMID 9310603. 
• Joffe, Russell T., Stephen T. Sokolov and Anthony J. Levitt (2006). "Lithium and Triiodothyronine Augmentation of Antidepressants". Canadian Journal of Psychiatry 51 (12): 791–3. PMID 17168254. 
• Krachler, M, and E Rossipal (1999). "Trace Elements Transfer From Mother to the Newborn – Investigations on Triplets of Colostrum, Maternal and Umbilical Sera". European Journal of Clinical Nutrition 53 (6): 486–494. doi:10.1038/sj.ejcn.1600781. PMID 10403586. 
• Stein, Benjamin P., Stephen G. Benka, Phillip F. Schewe, and Bertram Schwarzhild (1996). "Physics Update". Physics Today 49 (6): 9. Bibcode 1996PhT....49f...9S. doi:10.1063/1.2807642. 

External links

• "Group 1: The Alkali Metals". Visual Elements. Royal Society of Chemistry. http://www.chemsoc.org/Viselements/pages/data/intro_groupi_data.html. Retrieved 8 December 2009. 
• Atomic and Physical Properties of the Group 1 Elements An in-depth look at alkali metals
• Alkali Metal Bangs Filmed reactions of five-gram samples of the alkali metals with water

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