Alkali metal

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.^[3]

Most alkali metals have many different applications. Two of the most well-known applications of the pure elements are rubidium and caesium atomic clocks, of which caesium atomic clocks are the most accurate representation of time known as of 2012. A common application of the compounds of sodium is the sodium vapor lamp, which emit very efficient light. Table salt, or sodium chloride, has been used since antiquity. The alkali metals can also be used as fertilizers.

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

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; thus, the presentation of its properties here is limited. All the alkali metals are all highly reactive and are never found in elemental forms in nature.^[7] Because of this, they are usually stored in mineral oil or kerosene (paraffin oil).^[8]

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 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.^[9] 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.^[10] 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.^[11]

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.

History

Lithium

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.^[18][19][20] 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.^[21][22] 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.^[23] 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".^[7][19][22]

Sodium

Sodium has been known since ancient times; salt 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. In medieval Europe a compound of sodium with the Latin name of sodanum was used as a headache remedy. Pure sodium was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda,^[24] a very similar method to the one used to isolate potassium earlier that year.

Potassium

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,^[25] and Henri Louis Duhamel du Monceau was able to prove this difference in 1736.^[26] 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.^[27][28] Pure potassium 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.^[29] 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.^{[24][27][28][30]}

Rubidium

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.^[31][32] 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.^{[31][32][33][34]} Caesium was the first element to be discovered spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.^[35]

Francium

There were at least three erroneous and incomplete discoveries^{[36][37][38][39]} 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.^[40] 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%.^[41]

Eka-francium

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.^[3][42]

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

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.^[43]

Potassium is occasionally produced through separating the potassium from the chlorine in potassium chloride, but is more often produced through electrolysis of potassium hydroxide,^[44] 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.^[45] Sodium 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.^[46][47] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.^[48]

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.^[49] Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite.^[50] 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.^[50][51] 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.^[50] 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.^[50][52]

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;^[54] it has been calculated that at most there are 30 g of francium in the earth's crust at any given time.^[55] This makes it the second rarest element in the crust after astatine.^[53][56] 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.^[57]

As of 2012, all attempts at the synthesis of ununennium have been met with failure.^[3][42]

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)^[58][59] or 25 micromolar.^[60]

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^[61] and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.^[61] 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, soda niter, and zeolite.^[61]

Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.^[9] Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite.^[62]

Francium-223, the only naturally occurring isotope of francium,^[63] is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals^[54] In a given sample of uranium, there is estimated to be only one francium atom for every 1×10¹⁸ uranium atoms.^[53][64] 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,^[55] making it the second-rarest element in the crust after astatine.^[53][56]

Ununennium does not occur naturally in nature, and until it is synthesized in laboratories, it is not likely that any exists on earth.^[3][42]

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.^[65]

Pure sodium has many applications, including use in sodium-vapor lamps, which produce very efficient light compared to other types of lighting,^[66][67] and can help smooth the surface of other metals.^[68][69] 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,^[70] as it is an important element for plant nutrition. Other potassium ions are often used to hold anions. Potassium hydroxide is a very strong base, and is used to control pH of various substances^[71][72]

Rubidium and caesium are often used in atomic clocks.^[73] 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.^[9] For that reason, cesium atoms are used as the definition of the second.^[74] Rubidium ions are often used in purple fireworks,^[75] and caesium is often used in drilling fluids in the petroleum industry.^[9][76]

Francium has no commercial applications,^{[17][53][64][77]} 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.^[78] 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.^[79]

Biological occurrences

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.^[80]

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.^[81] Sodium chloride 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.^[82] The DRI for sodium is 1.5 grams per day,^[83] but most people in the United States consume more than 2.3 grams per day,^[84] the minimum amount that promotes hypertension;^[85] this in turn causes 7.6 million premature deaths worldwide.^[86]

Potassium is the major cation (positive ion) inside animal cells, while sodium is the major cation outside animal cells. The concentration differences of these charged particles causes a difference in electric potential between the inside and outside of cells, known as the membrane potential. The balance between potassium and sodium is maintained by ion 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.^[87] All five are from prokaryotic species. 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.^[87]

Rubidium has no known biological role, but may help stimulate metabolism,^[88][89] and, similarly to caesium, replace potassium in the body causing potassium deficiency.^[89]

Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium. 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.^[90] 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.^[91] Caesium chloride has been promoted as an alternative cancer therapy,^[92] but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.^[93]

Francium has no biological role^[94] and is most likely to be toxic due to its high radioactivity, although the greatest quantity of francium ever assembled to date is about 300,000 neutral atoms.^[95]

References

Bibliography

1. 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. 
2. 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. 
3. 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. 
4. 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. 
5. 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. 
6. 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. 
7. 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

1. "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. 
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

Links to related articles