alkali metal

Any of a group of soft, white, low-density, low-melting, highly reactive metallic elements, including lithium, sodium, potassium, rubidium, cesium, and francium.

Concept

Group 1 of the periodic table of elements consists of hydrogen, and below it the six alkali metals: lithium, sodium, potassium, rubidium, cesium, and francium. The last three are extremely rare, and have little to do with everyday life; on the other hand, it is hard to spend a day without encountering at least one of the first three—particularly sodium, found in table salt. Along with potassium, sodium is an important component of the human diet, and in compounds with other substances, it has an almost endless array of uses. Lithium does not have as many applications, but to many people who have received it as a medication for bipolar disorder, it is quite literally a life-saver.

How It Works

Electron Configuration of the Alkali Metals

In the essay on Families of Elements, there is a lengthy discussion concerning the relationship between electron configuration and the definition of a particular collection of elements as a "family." Here that subject will only be touched upon lightly, inasmuch as it relates to the alkali metals.

All members of Group 1 on the periodic table of elements have a valence electron configuration of s¹. This means that a single electron is involved in chemical bonding, and that this single electron moves through an orbital, or range of probabilities, roughly corresponding to a sphere.

Most elements bond according to what is known as the octet rule, meaning that when two or more atoms are bonded, each has (or shares) eight valence electrons. It is for this reason that the noble gases, at the opposite side of the periodic table from the alkali metals, almost never bond with other elements: they already have eight valence electrons.

The alkali metals, on the other hand, are quite likely to find "willing partners," since they each have just one valence electron. This brings up one of the reasons why hydrogen, though it is also part of Group 1, is not included as an alkali metal. First and most obviously, it is not a metal; additionally, it bonds according to what is called the duet rule, such that it shares two electrons with another element.

Chemical and Physical Characteristics of the Alkali Metals

The term "alkali" (essentially the opposite of an acid) refers to a substance that forms the negatively charged hydroxide ion (OH−) in contact with water. On their own, however, alkali metals almost always form positive ions, or cations, with a charge of +1.

When alkali metals react with water, one hydrogen atom splits off from the water molecule to form hydrogen gas, while the other hydrogen atom joins the oxygen to form hydroxide. Where the heavier members of the alkali metal family are concerned, reactions can often be so vigorous that the result is combustion or even explosion. Alkali metals also react with oxygen to produce either an oxide, peroxide, or superoxide, depending on the particular member of the alkali metal family involved.

Shiny and soft enough to be cut with a knife, the alkali metals are usually white (though cesium is more of a yellowish white). When placed in a flame, most of these substances produce characteristic colors: lithium, for instance, glows bright red, and sodium an intense yellow. Heated potassium produces a violet color, rubidium a dark red, and cesium a light blue. This makes it possible to identify the metals, when heated, by color—a useful trait, since they are so often inclined to be bonded with other elements.

Melting and Boiling Points

As one moves down the rows or periods of the periodic table, the mass of atoms increases, as does the energy each atom possesses. Yet the amount of energy required to turn a solid alkali metal into a liquid, or to vaporize a liquid alkali metal, actually decreases with higher atomic number. In other words, the higher the atomic number, the lower the boiling and melting points.

The list below lists the six alkali metals in order of atomic number, along with chemical symbol, atomic mass, melting point, and boiling point. Note that for francium, which is radioactive, the figures given are for its most stable isotope, francium-223 (²²³Fr)—which has a half-life of only about 21 minutes.

Atomic Number, Mass, and Melting and Boiling Points of the Alkali Metals:

• 3. Lithium (Li), 6.941; Melting Point: 356.9°F (180.5°C); Boiling Point: 2,457°F (1,347°C)
• 11. Sodium (Na), 22.99; Melting Point: 208°F (97.8°C); Boiling Point: 1,621.4°F (883°C)
• 19. Potassium (K), 39.10; Melting Point: 145.9°F (63.28°C); Boiling Point: 1,398.2°F (759°C)
• 37. Rubidium (Rb), 85.47; Melting Point: 102.8°F (39.31°C); Boiling Point: 1,270.4°F (688°C)
• 55. Cesium (Cs), 132.9; Melting Point: 83.12°F (28.4°C); Boiling Point: 1,239.8°F (671°C)
• 87. Francium (Fr), (223); Melting Point: 80.6°F (27°C); Boiling Point: 1,250.6°F (677°C)

Abundance of Alkali Metals

Sodium and potassium are, respectively, the sixth and seventh most abundant elements on Earth, comprising 2.6% and 2.4% of the planet's known elemental mass. This may not seem like much, but considering the fact that just two elements—oxygen and silicon—make up about 75%, and that just 16 elements make up most of the remainder, it is an impressive share.

Lithium, on the other hand, is much less abundant, and therefore, figures for its part of Earth's known elemental mass are measured in parts per million (ppm). The total lithium in Earth's crust is about 17 ppm. Surprisingly, rubidium is more abundant, at 60 ppm; less surprisingly, cesium, with just 3 ppm, is very rare. Almost no francium is found naturally, except in very small quantities within uranium ores.

Real-Life Applications

Lithium

Swedish chemist Johan August Arfvedson (1792-1841) discovered lithium in 1817, and named it after the Greek word for "stone." Four years later, another scientist named W. T. Brande succeeded in isolating the highly reactive metal. Most of the lithium available on Earth's crust is bound up with aluminum and silica in minerals.

Since the time of its discovery, lithium has been used in lubricants, glass, and in alloys of lead, aluminum, and magnesium. In glass, it acts as a strengthening agent; likewise, metal alloys that contain lithium tend to be stronger, yet less dense. In 1994, physicist Jeff Dahn of Simon Fraser University in British Columbia, Canada, developed a lithium battery. Not only was the battery cheaper to produce than the traditional variety, Dahn and his colleagues announced, but the disposal of used lithium batteries presented less danger to the environment.

One of the most striking uses of lithium occurred in 1932, when English physicist John D. Cockcroft (1897-1967) and Irish physicist Ernest Walton (1903-1995) built the first particle accelerator. By bombarding lithium atoms, they produced highly energized alpha particles. This was the first nuclear reaction brought about by the use of artificially accelerated particles—in other words, without the need for radioactive materials such as uranium-235. Cockcroft's and Walton's experiment with lithium thus proved pivotal to the later creation of the atomic bomb.

Lithium in Psychiatric Treatment

The most important application of lithium, however, is in treatment for the psychiatric condition once known as manic depression, today identified as bipolar disorder. Persons suffering from bipolar disorder tend toward mood swings: during some periods the patient is giddy ("manic," or in a condition of "mania"), and during others the person is suicidal. Indeed, prior to the development of lithium as a treatment for bipolar disorder, as many as one in five patients with this condition committed suicide.

Doctors do not know exactly how lithium does what it does, but it obviously works: between 70% and 80% of patients with the bipolar condition respond well to treatment, and are able to go on with their lives in such a way that their condition is no longer outwardly evident. Lithium is also administered to patients who suffer unipolar depression and some forms of schizophrenia.

Early Medicinal Uses of Lithium

It is said that the great Greco-Roman physician Galen (129-c. 199) counseled patients suffering from "mania" to bathe in, and even drink the water from, alkaline springs. If so, he was nearly 2,000 years ahead of his time. Even in the 1840s, not long after lithium was discovered, the mineral—mixed with carbonate or citrate—was touted as a cure for insomnia, gout, epilepsy, diabetes, and even cancer.

None of these alleged cures proved a success; nor did a lithium chloride treatment administered in the 1940s as a salt substitute for patients on low-sodium diets. As it turned out, when not enough sodium is present, the body experiences a buildup of sodium's sister element, lithium. The result was poisoning, which in some cases proved fatal.

Cade's Breakthrough

Then in 1949, Australian psychiatrist John Cade discovered the value of lithium for psychiatric treatment. He approached the problem from an entirely different angle, experimenting with uric acid, which he believed to be a cause of manic behavior. In administering the acid to guinea pigs, he added lithium salts merely to keep the uric acid soluble—and was very surprised by what he discovered. The uric acid did not make the guinea pigs manic, as he had expected; instead, they became exceedingly calm.

Cade changed the focus of his research, and tested lithium treatment on ten manic patients. Again, the results were astounding: one patient who had suffered from an acute bipolar disorder (as it is now known) for five years was released from the hospital after three months of lithium treatment, and went on to lead a healthy, normal life.

Encouraged by the changes he had seen in patients who received lithium, Cade published a report on his findings in the Medical Journal of Australia, but his work had little impact at the time. Nor did the idea of lithium treatment meet with an enthusiastic reception on the other side of the Pacific: in the aftermath of the failed experiments with lithium as a sodium substitute in the 1940s, stories of lithium poisoning were widespread in the United States.

Lithium Today

Were it not for the efforts of Danish physician Mogens Schou, lithium might never have taken hold in the medical community. During the 1950s and 1960s, Schou campaigned tirelessly for recognition of lithium as a treatment for manic-depressive illness. Finally during the 1960s, the U.S. Food and Drug Administration began conducting trials of lithium, and approved its use in 1974. Today some 200,000 Americans receive lithium treatments.

A non-addictive and non-sedating medication, lithium—as evidenced by the failed experiment in the 1940s—may still be dangerous in large quantities. It is absorbed quickly into the bloodstream and carried to all tissues in the brain and body before passing through the kidneys. Both lithium and sodium are excreted through the kidneys, and since sodium affects lithium excretion, it is necessary to maintain a proper quantity of sodium in the body. For this reason, patients on lithium are cautioned to avoid a low-salt diet.

Sodium

Sodium compounds had been known for some time prior to 1807, when English chemist Sir Humphry Davy (1778-1829) succeeded in isolating sodium itself. The element is represented by a chemical symbol (Na), reflecting its Latin name, natrium. In its pure form, sodium has a bright, shiny surface, but in order to preserve this appearance, it must be stored in oil: sodium reacts quickly with oxygen, forming a white crust of sodium oxide.

Pure sodium never occurs in nature; instead, it combines readily with other substances to form compounds, many of which are among the most widely used chemicals in industry. It is also highly soluble: thus whereas sodium and potassium occur in crystal rocks at about the same ratio, sodium is about 30 times more abundant in sea-water than its sister element.

Obtaining Sodium Chloride

Though the extraction of sodium involves the use of a special process, the metal is plentiful in the form of sodium chloride—better known as table salt. In fact, the term salt in chemistry refers generally to any combination of a metal with a nonmetal. More specifically, salts are (along with water) the product of reactions between acids and bases.

Sodium chloride is so easy to obtain, and therefore so cheap, that most industries making other sodium compounds use it, simply separating out the chloride (as described below) before adding other elements. The United States is the world's largest producer of sodium chloride, obtained primarily from brine, a term used to describe any solution of sodium chloride in water. Brine comes from seawater, subterranean wells, and desert lakes, such as the Great Salt Lake in Utah. Another source of sodium chloride is rock salt, created underground by the evaporation of long-buried saltwater seas.

Other top sodium-chloride-producing nations include China, Germany, Great Britain, France, India, and various countries in the former Soviet Union. Salt may be cheap and plentiful for the world in general, but there are places where it is a precious commodity. One such place is the Sahara Desert, where salt caravans ply a brisk trade today, much as they have since ancient times.

Isolating Sodium

Modern methods for the production of sodium represent an improvement in the technique Davy used in 1807, although the basic principle is the same. Though several decades passed before electricity came into widespread public use, scientists had been studying its properties for years, and Davy applied it in a process called electrolysis.

Electrolysis is the use of an electric current to produce a chemical reaction—in this case, to separate sodium from the other element or elements with which it is combined. Davy first fused or melted a sample of sodium chloride, then electrolyzed it. Using an electrode, a device that conducts electricity and is used to emit or collect electric charge, he separated the sodium chloride in such a way that liquid sodium metal collected on the cathode, or negatively charged end. Meanwhile, the gaseous chlorine was released through the anode, or the positively charged end.

The apparatus used for sodium separation today is known as the Downs cell, after its inventor, J. C. Downs. In a Downs cell, sodium chloride and calcium chloride are combined in a molten mixture in which the presence of calcium chloride lowers the melting point of the sodium chloride by more than 30%. When an electric current is passed through the mixture, sodium ions move to the cathode, where they pick up electrons to become sodium atoms. At the same time, ions of chlorine migrate to the anode, losing electrons to become chlorine atoms.

Sodium is a low-density material that floats on water, and in the Downs cell, the molten sodium rises to the top, where it is drawn off. The chlorine gas is allowed to escape through a vent at the top of the anode end of the cell, and the resulting sodium metal—that is, the elemental form of sodium—is about 99.8% pure.

Uses for Sodium Chloride

As indicated earlier, sodium chloride is by far the most widely known and commonly used sodium compound—and this in itself is a distinction, given the fact that so many sodium compounds are a part of daily life. Today people think of salt primarily as a seasoning to enhance the taste of food, but prior to the development of refrigeration, it was vital as a preservative because it kept microbes away from otherwise perishable food items.

Salt does not merely improve the taste of food; it is an essential nutrient. Sodium compounds regulate transmission of signals through the nervous system, alter the permeability of membranes, and perform a number of other life-preserving functions. On the other hand, too much salt can aggravate high blood pressure. Thus, since the 1970s and 1980s, food manufacturers have increasingly offered products low in sodium, a major selling point for health-conscious consumers.

Other Sodium Compounds

In addition to its widespread use in consumer goods, sodium chloride is the principal source of sodium used in making other sodium compounds. These include sodium hydroxide, for manufacturing cellulose products such as film, rayon, soaps, and paper, and for refining petroleum. In its application as a cleaning solution, sodium hydroxide is known as caustic soda or lye.

Another widely used sodium compound is sodium carbonate or, soda ash, applied in glass-making, paper production, textile manufacturing, and other areas, such as the production of soaps and detergents. Sodium also can be combined with carbon to produce sodium bicarbonate, or baking soda. Sodium sulfate, sometimes known as salt cake, is used for making cardboard and kraft paper. Yet another widely used sodium compound is sodium silicate, or "water glass," used in the production of soaps, detergents, and adhesives; in water treatment; and in bleaching and sizing of textiles.

Still other sodium compounds used by industry and/or consumers include sodium borate, or borax; sodium tartrate, or sal tartar; the explosive sodium nitrate, or Chilean salt-peter; and the food additive monosodium glutamate (MSG). Perhaps ironically, there are few uses for pure metallic sodium. Once applied as an "anti-knock" additive in leaded gasoline, before those products were phased out for environmental reasons, metallic sodium is now used as a heat-exchange medium in nuclear reactors. But its widest application is in the production of the many other sodium compounds used around the world.

Potassium

In some ways, potassium is a strange substance, as evidenced by its behavior in response to water. As everyone knows, water tends to put out a fire, and most explosives, when exposed to sufficient quantities of water, become ineffective. Potassium, on the other hand, explodes in contact with water and reacts violently with ice at temperatures as low as −148°F (−100°C). In a complete reversal of the procedures normally followed for most substances, potassium is stored in kerosene, because it might burst into flames if exposed to moist air!

Many aspects of potassium mirror those already covered with regard to sodium. The two have a number of the same applications, and in certain situations, potassium is used as a sodium substitute. Like sodium, potassium is never found alone in nature; instead, it comes primarily from sylvinite and carnalite, two ores containing potassium chloride. Also, like sodium, potassium was first isolated in 1807 by Davy, using the process of electrolysis described above. A few years later, a German chemist dubbed the newly isolated element "kalium," apparently a derivation of the Arabic qali, for "alkali"; hence the use of K as the chemical symbol for potassium.

Uses for Potassium

Potassium has another similarity with sodium; although it was not isolated until the early nineteenth century, its compounds have been in use for many centuries. The Romans, for instance, used potassium carbonate, or potash, obtained from the ashes of burned wood, to make soap. During the Middle Ages, the Chinese applied a form of saltpeter, potassium nitrate, in making gunpowder. And in colonial America, potash went into the production of soap, glass, and other products.

The production of just one ton of potash required the burning of several acres' worth of trees—a wasteful practice in more ways than one. Though there was no environmentalist movement in those days, financial concerns never go out of style. In order to save the money lost by using up vast acres of timber, American industry in the nineteenth century sought another means of making potash. The many similarities between sodium and potassium provided a key, and the substitution of sodium carbonate for potassium carbonate saved millions of trees.

In 1847, German chemist Justus von Liebig (1803-1873) discovered potassium in living tissues. As a result, scientists became aware of the role this alkali metal plays in sustaining life: indeed, potassium is present in virtually all living cells. In the human body, potassium—which accounts for only 0.4% of the body's mass—is essential to the functioning of muscles. In larger quantities, however, it can be dangerous, causing a state of permanent relaxation known as potassium inhibition.

Since plants depend on potassium for growth, it was only logical that potassium, in the form of potassium chloride, was eventually applied as a fertilizer. This, at least, distinguishes it from its sister element: sodium, or sodium chloride, which can kill plants if administered to the soil in large enough quantities.

Another application of potassium is in the area pioneered by the Chinese about 800 years ago: the manufacture of fireworks and gunpowder from potassium nitrate. Like ammonium nitrate, made infamous by its use in the 1993 World Trade Center bombing and the Oklahoma City bombing in 1995, potassium nitrate doubles as a fertilizer.

Rubidium, Cesium, and Francium

The three heaviest alkali metals are hardly household names, though one of them, cesium, does have several applications in industry. Rubidium and cesium, discovered in 1860 by German chemist R. W. Bunsen (1811-1899) and German physicist Gustav Robert Kirchhoff (1824-1887), were the first elements ever found using a spectroscope. Matter emits electromagnetic radiation along various spectral lines, which can be recorded using a spectroscope and then analyzed to discern the particular "fingerprint" of the substance in question.

When Bunsen and Kirchhoff saw the bluish spectral lines emitted by one of the two elements, they named it cesium, after a Latin word meaning "sky blue." Cesium, which is very rare, appears primarily in compounds such as pollucite. It is used today in photoelectric cells, military infrared lamps, radio tubes, and video equipment. During the 1940s, American physicist Norman F. Ramsey, Jr. (1915-) built a highly accurate atomic clock based on the natural frequencies of cesium atoms.

Rubidium, by contrast, has far fewer applications, and those are primarily in areas of scientific research. On Earth it is found in pollucite, lepidolite, and carnallite. It is considerably more abundant than cesium, and vastly more so than francium. Indeed, it is estimated that if all the francium in Earth's crust were combined, it would have a mass of about 25 grams.

Francium was discovered in 1939 by French physicist Marguerite Perey (1909-1975), student of the famous French-Polish physicist and chemist Marie Curie (1867-1934). For about four decades, scientists had been searching for the mysterious Element 87, and while studying the decay products of an actinium isotope, actinium-227, Perey discovered that one out of 100 such atoms decayed to form the undiscovered element. She named it francium, after her home-land. Though the discovery of francium solved a mystery, the element has no known uses outside of its applications in research.

Where to Learn More

"Alkali Metals" (Web site). <http://www.midlink.com/~jfromm/elements/alkali.htm> (May 24, 2001).

"Alkali Metals" ChemicalElements.com (Web site). <http://www.chemicalelements.com/groups/alkali.html> (May 24, 2001).

"Hydrogen and the Alkali Metals." University of ColoradoDepartment of Physics (Web site). <http://www.colorado.edu/physics/2000/periodic_table/alkali metals.html> (May 24, 2001).

Kerrod, Robin. Matter and Materials. Illustrated by Terry Hadler. Tarrytown, N.Y.: Benchmark Books, 1996.

Mebane, Robert C. and Thomas R. Rybolt. Metals. Illustrated by Anni Matsick. New York: Twenty-First Century Books, 1995.

Oxlade, Chris. Metal. Chicago, IL: Heinemann Library, 2001.

Snedden, Robert. Materials. Des Plaines, IL: Heinemann Library, 1999.

"Visual Elements: Group I—The Alkali Metals" (Web site). <http://www.chemsoc.org/viselements/pages/data/intro_groupi_data.html> (May 24, 2001).

The elements of group 1 in the periodic table (lithium, sodium, potassium, rubidium, cesium, francium). Of the alkali metals, lithium differs most from the rest of the group, and tends to resemble the alkaline-earth metals (group 2 of the periodic table) in many ways. In this respect lithium behaves as do many other elements that are the first members of groups in the periodic table; these tend to resemble the elements in the group to the right rather than those in the same group. Francium, the heaviest of the alkali-metal elements, has no stable isotopes and exists only in radioactive form.

In general, the alkali metals are soft, low-melting, reactive metals. This reactivity accounts for the fact that they are never found uncombined in nature but are always in chemical combination with other elements. This reactivity also accounts for the fact that they have no utility as structural metals (with the possible exception of lithium in alloys) and that they are used as chemical reactants in industry rather than as metals in the usual sense. The reactivity in the alkali-metal series increases in general with increase in atomic weight from lithium to cesium. See also Cesium; Electrochemical series; Francium; Lithium; Periodic table; Potassium; Rubidium; Sodium.

For more information on alkali metal, visit Britannica.com.

These are the elements at the left of the periodic table: lithium (Li, element 3), potassium (K, element 19), rubidium (Rb, element 37), cesium (Cs, element 55), francium (Fr, element 87), and sodium (Na, element 11). The alkali metals are sometimes called the sodium family of elements, or Group I elements. Because of their great chemical reactivity (they easily form positive ions), none exist in nature in the elemental state.

This article needs additional citations for verification. Please help improve this article by adding reliable references. Unsourced material may be challenged and removed. (November 2008)

Group →	1
↓ Period
1	1 H
2	3 Li
3	11 Na
4	19 K
5	37 Rb
6	55 Cs
7	87 Fr

The alkali metals are a series of chemical elements comprising Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals). The alkali metals provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behavior down the group.

Contents 1 Properties 2 Trends 3 Occurrence and production 4 Applications 5 Reference material 6 See also 7 External links

Properties

The alkali metals are all highly reactive and are never found in elemental form in nature. As a result, in the laboratory they are stored under mineral oil or parafin oil. They also tarnish easily and have low melting points and densities. Potassium and rubidium possess a weak radioactive characteristic due to the presence of long duration radioactive isotopes.

The alkali metals are silver-colored (caesium has a golden tinge), soft, low-density metals, which react readily with halogens to form ionic salts, and with water to form strongly alkaline (basic) hydroxides. These elements all have one electron in their outermost shell, so the energetically preferred state of achieving a filled electron shell is to lose one electron to form a singly charged positive ion, i.e. cation.

Hydrogen, with a solitary electron, is usually placed at the top of Group 1 of the periodic table, but it is not considered an alkali metal; rather it exists naturally as a diatomic gas. Removal of its single electron requires considerably more energy than removal of the outer electron for the alkali metals. As in the halogens, only one additional electron is required to fill in the outermost shell of the hydrogen atom, so hydrogen can in some circumstances behave like a halogen, forming the negative hydride ion. Binary compounds of hydride with the alkali metals and some transition metals have been prepared. Under extremely high pressure, such as is found at the core of Jupiter, hydrogen does become metallic and behaves like an alkali metal; see metallic hydrogen.

Alkali metals have the lowest ionization potentials in their respective periods, as removing the single electron from the outermost shell gives them the stable inert gas configuration. Their second ionization potentials are very high, as removing an electron from a species having a noble gas configuration is very difficult.

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

Alkali metals are famous for their vigorous reactions with water, and these reactions become increasingly violent as one moves down the group. The reaction with water is as follows:

Alkali metal + water → Alkali metal hydroxide + hydrogen gas

With potassium as an example:

2K (s) + 2H₂O (l) → 2KOH (aq) + H₂ (g)

Trends

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

Z	Element	No. of electrons/shell
1	Hydrogen	1
3	Lithium	2, 1
11	Sodium	2, 8, 1
19	Potassium	2, 8, 8, 1
37	Rubidium	2, 8, 18, 8, 1
55	Caesium	2, 8, 18, 18, 8, 1
87	Francium	2, 8, 18, 32, 18, 8, 1

The alkali metals show a number of trends when moving down the group - for instance,decreasing electronegativity, increasing reactivity, and decreasing melting and boiling point. Density generally increases, with the notable exception of potassium being less dense than sodium, and the possible exception of francium being less dense than caesium.

Alkali metal	Standard Atomic Weight (u)	Melting Point (K)	Boiling Point (K)	Density (g·cm−3)	Electronegativity (Pauling)
Lithium	6.941	453	1615	0.534	0.98
Sodium	22.990	370	1156	0.968	0.93
Potassium	39.098	336	1032	0.89	0.82
Rubidium	85.468	312	961	1.532	0.82
Caesium	132.905	301	944	1.93	0.79
Francium	(223)	295	950	1.87	0.70

Occurrence and production

Applications

Reference material

• 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: 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". Noise and Fluctuations: 18th International Conference on Noise and Fluctuations. American Institute of Physics. 

• Sokolov, Stephen T., Russell T. Joffe, and Anthony J. Levitt (2006). "Lithium and Triiodothyronine Augmentation of Antidepressants". Canadian Journal of Psychiatry 51: 791–793. 

• Bauer, Brent A., Robert Houlihan, Michael J. Ackerman, Katya Johnson, and Himeshkumar Vyas (2006). "Acquired Long QT Syndrome Secondary to Cesium Chloride Supplement". The Journal of Alternative and Complementary Medicine 12: 1011–1014. doi:10.1089/acm.2006.12.1011. 

• 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: 850–854. doi:10.1016/S0140-6736(97)02264-2. 

• 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: 486–494. doi:10.1038/sj.ejcn.1600781. 

• "Physics Update." Physics Today June 1996: 9.

• "Visual Elements: Group 1 - The Alkali Metals." Visual Elements. Royal Society of Chemistry. <http://www.chemsoc.org/Viselements/pages/data/intro_groupi_data.html>.

See also

• Lithium
• Sodium
• Potassium
• Rubidium
• Caesium
• Francium

External links

• Science aid:Alkali metals A simple look at alkali metals
• Atomic and Physical Properties of the Group 1 Elements An in-depth look at alkali metals

Explanation of above periodic table slice:

Alkali metals

Atomic numbers in black are solids

Solid borders indicate primordial elements (older than the Earth)

Dashed borders indicate natural radioactive elements with no isotopes older than the Earth

v • d • e Periodic table
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	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
Fr	Ra	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo
Alkali metals Alkaline earth metals Lanthanoids Actinoids Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases

v • d • e Alkali metals     Lithium Li Atomic Number: 3 Atomic Weight: 6.941 Melting Point: 453.69 Boiling Point: 1615 Electronegativity: 0.98 Sodium Na Atomic Number: 11 Atomic Weight: 22.990 Melting Point: 370.87 Boiling Point: 1156 Electronegativity: 0.96 Potassium K Atomic Number: 19 Atomic Weight: 39.098 Melting Point: 336.58 Boiling Point: 1032 Electronegativity: 0.82 Rubidium Rb Atomic Number: 37 Atomic Weight: 85.468 Melting Point: 312.46 Boiling Point: 961 Electronegativity: 0.82 Caesium Cs Atomic Number: 55 Atomic Weight: 132.905 Melting Point: 301.59 Boiling Point: 944 Electronegativity: 0.79 Francium Fr Atomic Number: 87 Atomic Weight: (223) Melting Point: ?295 Boiling Point: ?950 Electronegativity: 0.7

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