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 Natural Ratio

The alkali metals are a series of chemical elements in the periodic table. In the modern IUPAC nomenclature, the alkali metals comprise the group 1 elements, along with hydrogen. The alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).^[1] Hydrogen (H), although nominally also a member of Group 1,^[2] rarely exhibits behavior comparable to the alkali metals. This group lies in the s-block of the periodic table; all alkali metals' outermost electrons lie in an s-orbital. The alkali metals provide one of the best examples of group trends in properties in the periodic table, with elements exhibiting well characterized homologous behavior. For instance, when moving down the table, all alkali metals show decreasing electronegativity, increasing reactivity, and decreasing melting and boiling points. In general, their densities increase, with the notable exception that potassium is less dense than sodium.

All discovered alkali metals are naturally occurring, although francium is the second-rarest naturally occurring element, after astatine. All are highly reactive metals under standard conditions. Experiments have been conducted to attempt the synthesis of ununennium (Uue), likely the next member of the group, but have all met with failure, but because all elements with lower atomic numbers have been discovered, it is likely^[3] to be found in the near future. However, ununennium may not be an alkali metal due to relativistic effects.

Hydrogen

Main article: Hydrogen

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 usually considered as an alkali metal. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule (H₂).

Hydrogen, like the other alkali metals, has one valence electron; however, the similarities end there. In fact, its placement above lithium is primarily due to its electron configuration and not its chemical properties.^[2] It is sometimes placed above carbon due to similar electronegativity^[4] or fluorine due to similar chemical properties.^[4]

The first ionization energy of hydrogen is much higher than that of the alkali metals. As only one additional electron is required to fill in the outermost shell of the hydrogen atom, hydrogen often behaves like a halogen, forming the negative hydride ion. Binary compounds of hydrogen with the alkali metals and some transition metals have been produced in the laboratory^[5], but these are only laboratory curiosities without any practical use. Under extremely high pressures and low temperatures, however, such as those found at the cores of the planets Jupiter and Saturn, hydrogen does become metallic^[6] and behaves like an alkali metal; in this phase, it is known as metallic hydrogen.

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:

Z	Element	No. of electrons/shell	Electron configuration
3	lithium	2, 1	[He]2s1
11	sodium	2, 8, 1	[Ne]3s1
19	potassium	2, 8, 8, 1	[Ar]4s1
37	rubidium	2, 8, 18, 8, 1	[Kr]5s1
55	caesium	2, 8, 18, 18, 8, 1	[Xe]6s1
87	francium	2, 8, 18, 32, 18, 8, 1	[Rn]7s1

Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not very established and therefore the rest of this section does not deal with it. All the alkali metals are all highly reactive and are never found in elemental forms in nature. Because of this, they are usually stored in mineral oil or kerosene (paraffin oil).^[7]

H																	He
Li	Be											B	C	N	O	F	Ne
Na	Mg											Al	Si	P	S	Cl	Ar
K	Ca	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
Rb	Sr	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
Cs	Ba	*	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
Fr	Ra	**	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn	Uut	Uuq	Uup	Uuh	Uus	Uuo
		*	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
		**	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr

Alkali metals in the periodic table

Hydrogen in the periodic table

Caesium reacts explosively with water even at low temperatures

The alkali metals are all silver-colored except for metallic caesium, which can have a golden tint. All are soft and have low density, melting points, and boiling points. In chemical terms, all of the alkali metals react aggressively with the halogens to form ionic salts, and also react with water to form strongly alkaline hydroxides. The heavier alkali metals react more vigorously than the lighter ones; caesium produces much bigger explosions when dropped in water than potassium, for example. The alkali metals have the lowest first ionization energies in their respective periods of the periodic table because of their low effective nuclear charge and the ability to attain noble gas configuration by losing just one electron. The second ionization energy of all of the alkali metals is very high, thus they almost always lose a single electron, forming cations

Alkali metals form a very wide range of amalgams.^[8] They tend to form ionically bonded salts with most electronegative elements on the periodic table, for example caesium fluoride and sodium chloride.

The ammonium ion (NH+
4) has very similar properties to the heavier alkali metals and is often considered a close relative.^[9]

Physical and atomic

The table below is a summary of the key physical properties of the alkali metals. The density of Francium is an estimation partially based on periodic trends rather than observations.^[10]

Group 1 element	Standard atomic weight (u)	Melting point (K)	Melting point (°C)	Boiling point (K)	Boiling point (°C)	Density (g/cm3)	Electronegativity (Pauling)
Lithium	6.941	454	180.5	1615	1342	0.534	0.98
Sodium	22.98976928	370	97.8	1156	883	0.968	0.93
Potassium	39.0983	336	63.38	1032	759	0.89	0.82
Rubidium	85.4678	312	39.31	961	688	1.532	0.82
Caesium	132.9054519	301	28.44	944	671	1.93	0.79
Francium	(223)	300.15	27	950.15	677	1.87	0.70

History

Lithium

Petalite

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.^[11][12][13] 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.^[14][15][16] 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'.^[17] 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".^[12][16][18]

Sodium

Caustic soda

Although sodium has long been recognized in compounds, it was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda,^[19] a very similar method to the one used to isolate potassium earlier that year.

Potassium

Caustic potash

Potassium metal was first isolated in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH) 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.^[20] Later that same year, Davy reported extraction of the metal sodium from the similar substance caustic soda (NaOH, or lye) by a similar technique, demonstrating the elements, and thus the salts, to be different.^{[19][21][22][23]}

Rubidium

Lepidolite

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.^[24][25] Rubidium's discovery succeeded that of caesium, also discovered by Bunsen and Kirchhoff through spectroscopy.

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.^{[24][25][26][27]} Caesium was the first element to be discovered spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.^[28]

Francium

There were at least three erroneous and incomplete discoveries^{[29][30][31][32]} before Marguerite Perey of the Curie Institute in Paris, France discoverd 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.^[33] 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%.^[34]

Eka-francium

See also: Extended periodic table

The next element below francium (Eka-francium) is likely to be ununennium (Uue), element 119, although this is not certain. 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.^[35][36]

$\,^{254}_{99}\mathrm{Es} + \,^{48}_{20}\mathrm{Ca} \to \,^{302}_{119}\mathrm{Uue} ^{*} \to \ \ no\ atoms$
It is highly unlikely^[3] 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.^[3] Currently, none of the period 8 elements have been discovered yet. It is also possible that, due to drip instabilities, only the lower period 8 elements are physically possible.

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. 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.^[37]

Potassium is occasionally produced through separating the potassium from the chlorine in potassium chloride, but is more often produced through electrolysis of potassium hydroxide,^[38], 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 in the late 1800s and early 1900s.^[39] Sodium is now produced through electrolysis of sodium chloride by lowering the melting point of the substance to below 700 ºC through the use of a Downs Cell.^[40][41] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.^[42]

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.^[43] Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite.^[44] 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.^[44][45] 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.^[44] Caesium, however, is not produced from the above reaction. Instead, the miming 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.^[44][46]

This sample of uraninite contains about 100,000 atoms (3.3×10⁻²⁰ g) of francium-223 at any given time.^[47]

Francium-223 is the only naturally occurring isotope of francium, produced naturally as the product of the alpha decay of actinium-227. Francium can be found in trace amounts in uranium and thorium minerals;^[48] it has been calculated that at most there are 30 g of francium in the earth's crust at any given time.^[49] This makes it the second rarest element in the crust after astatine.^[47][50] 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.^[51][52]

As of 2011, the synthesis of Ununennium has been met with failure.^[35][36]

Occurrence

The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. 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).^[53][54] or 25 micromolar;^[55]

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^[56]and the most abundant alkali metal, and Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.^[56]

Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.^[57] Rubidium is approximately as abundant as zinc and more abundant than copper.

Francium-223, the only naturally occurring isotope of francium,^[58] is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals^[48] In a given sample of uranium, there is estimated to be only one francium atom for every 1×10¹⁸ uranium atoms.^[47][59] 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,^[49] making it the second-rarest element in the crust after astatine.^[47][50]

Ununennium does not occur naturally in nature, and until it is synthesized in laboratories, it is not likely that any exists on earth.^[35][36]

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.^[60]

Pure sodium has many applications, including use in sodium-vapor lamps, which produce very efficient light compared to other types of lighting,^[61][62] and can help smooth the surface of other metals.^[63][64] Sodium compounds have many applications as well, the most well-known compound being table salt. Sodium is also used in soap as salts of fatty acids.

Potassium is often used as a fertilizer, ^[65] as it is an important element for plant nutrition. Other potassium ions are often used to hold anions. Potassium hydroxide is is a very strong base, and is used to control pH of various substances^[66][67]

Rubidium and caesium are often used in atomic clocks.^[68] Caesium atomic clocks are the most accurate clocks that humans have ever made; if a clock was made at the time of the dinosaurs, it would be off by less than two seconds.^[69] Rubidium ions are often used in purple fireworks,^[70] and caesium is often used in drilling fluids in the petroleum industry.^[69][71]

Francium has no commercial applications,^{[72][73][47][59]} 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.^[74] 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.^[75]

Biological occurrences

Lithium carbonate

The lighter alkali metals all have a role in the human body. Lithium carbonate is a mood stabilizer and is used to treat bipolar disorder (manic-depression), although there are side-effects. Excessive ingestion of lithium poisons the central nervous system, which is even more dangerous than the disorder itself, as the lethal dose is only just above the therapeutic dose.^[76] Sodium and potassium are essential elements for survival, as the cations Na⁺ and K⁺ are essential for nerve and heart function.

Rubidium and caesium have no known biological role, but rubidium may help stimulate metabolism. Despite this, rubidium, as well as caesium, can replace potassium in the body, especially in muscles. Caesium chloride has been promoted as an alternative cancer therapy,^[77] but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.^[78]

Francium has no biological role^[79] and is most likely extremely toxic due to its extreme radioactivity.

Precautions

The alkali metals all react violently with water and thus should be handled with great care. Caesium has the potential to replace potassium in the human body, causing potassium defiance.^[80] Francium's extreme radioactivity is a great hazard, although the greatest quantity of francium ever assembled to date is a sphere of radius 1 mm (about 300,000 neutral atoms).^[81]

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External links

1. "Group 1: The Alkali Metals". Visual Elements. Royal Society of Chemistry. http://www.chemsoc.org/Viselements/pages/data/intro_groupi_data.html. Retrieved 08 December 2009. 
2. Science aid: Alkali metals A simple look at alkali metals
3. Atomic and Physical Properties of the Group 1 Elements An in-depth look at alkali metals

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