# Alkali Metals

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

The alkali metals are a series of chemical elements in the periodic table. In the modern IUPAC nomenclature, the alkali metals are called the Group 1 elements.  The alkali metals include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs)[note 1] and francium (Fr).[4] Hydrogen (H), although nominally also a member of Group 1,[5] very rarely exhibits behaviour comparable to the alkali metals.

This group lies in the s-block of the periodic table, which means that all its elements have their outermost electron in an s-orbital. The s-block also includes alkaline earth metals, plus hydrogen and helium. The alkali metals provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behaviour down the group.

All of the alkali metals discovered, as of 2011, are naturally occurring, although francium is the second-rarest naturally occurring element. These metals share similar chemical properties: they are all highly reactive metals under standard conditions. Experiments have been conducted to attempt the synthesis of the next member of the group, which is probably ununennium (Uue), but these have all met with failure, but since ununennium is the first period 8 element and only the first element on the periodic table that has not been discovered, it is likely to be found in the near future. Note that ununennium may not be the next alkali metal due to relativistic effects.

## Characteristics

### Chemistry

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

Z Element No. of electrons/shell Electron configuration
1 hydrogen 1 1s1
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

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). They also tarnish easily and have low melting points and densities.[6]

Physically, the alkali metals are mostly silver-colored, except for metallic caesium, which can have a golden tint. These elements are all soft metals of low density. Chemically, all of the alkali metals react aggressively with the halogens to form ionic salts. They all react with water to form strongly alkaline hydroxides. The vigor of reaction increases down the group. All of the atoms of alkali metals have one electron in their valence shells, hence their only way for achieving the equivalent of filled outmost electron shells is to give up one electron to an element with high electronegativity, and hence to become singly charged positive ions, i.e. cations.

Caesium reacts explosively with water even at low temperatures

The alkali metals have the lowest ionization enthalpies in their periods of the periodic table, because the removal of their single electrons from their outmost electron shells gives them the stable electron configuration of inert gases. Another way of stating this is that they all have a high electropositivity. The “second ionization potential” of all of the alkali metals is very high, since removing any electron from an atom having a noble gas configuration is difficult to do.

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

All of the alkali metals are notable for their vigorous reactions with water, and these reactions become increasingly vigorous when going down their column in the periodic table towards the heaviest alkali metals, such as caesium. Their chemical reactions with water are as follows:

Alkali metal + water → Alkali metal hydroxide + hydrogen gas

For a typical example (M represents an alkali metal):

2 M (s) + 2 H2O (l) → 2 MOH (aq) + H2 (g)

Alkali metals form a very wide range of amalgams.[7] They tend to form ionically bonded salts with most electronegative elements on the periodic table, like caesium fluoride and sodium chloride.

### Physics

Potassium and rubidium are naturally weakly radioactive as they have naturally occurring radioisotopes (40
K
and 87
Rb
) and francium‘s only naturally occurring isotope (223
Fr
) is also radioactive with a half-life of only 22.0 minutes. Caesium was also thought to have natural radioisotopes before the discovery of francium, which made the group strange compared to the others.[citation needed]

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. Their densities generally increase, with the notable exception that potassium is less dense than sodium, and the possible exception of francium being less dense than caesium (the highly radioactive element francium only exists in microscopic quantities, so its physical properties have not been measured).

Below, hydrogen is included for comparison.

Group 1 element Standard atomic weight (u) Melting point (K) Melting point (°C) Boiling point (K) Boiling point (°C) Density (g/cm3) Electronegativity (Pauling)
Hydrogen 1.00794 14.2 −258.8 20.3 −252.7 0.00008988 2.20
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

†Estimation[8]

## Hydrogen

The element hydrogen, with its solitary one electron per atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not counted as an alkali metal. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule.

The removal of the single electron of hydrogen requires considerably more energy than removal of the outer electron from the atoms of 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 hydrogen with the alkali metals and some transition metals have been produced in the laboratory, but these are only laboratory curiosities without much practical use. Under extremely high pressures and low temperatures, such as those found at the cores of the planets Jupiter and Saturn, hydrogen does become metallic, and it behaves like an alkali metal. This type of hydrogen is known as metallic hydrogen.

Hydrogen is sometimes considered to be an alkali metal as it, like the other alkali metals, has one valence electron; however, hydrogen rarely acts like an alkali metal, as can be seen from the data above. In fact, it is sometimes placed over lithium (due to its electron configuration),[5] sometimes carbon (due to its electronegativity)[9] and sometimes fluorine (due to its chemical properties).[9]

## 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.[10][11][12] 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.[13][14][15] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.[16] Berzelius gave the alkaline 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 as “lithium“.[11][15][17]

### Sodium

Caustic soda

Although sodium (sometimes called “soda” in English) has long been recognized in compounds, it was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda.[18]

### 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 discovered voltaic pile. Potassium was the first metal that was isolated by electrolysis.[19] Later in the same year, Davy reported extraction of the metal sodium from a mineral derivative (caustic soda, NaOH, or lye), not a plant salt, by a similar technique, demonstrating the elements, and thus the salts, to be different.[18]

### Rubidium

Lepidolite

Robert Bunsen and Gustav Kirchhoff were the first to suggest finding new elements by spectrum analysis. Rubidium was discovered in 1861 by Bunsen and Kirchhoff, in Heidelberg, Germany, 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, dark red.[20][21]

### 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.[20][21][22][23] Caesium was the first element to be discovered spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.[24]

### Francium

Francium (then known as ekacaesium, because it is below caesium in the periodic table) was discovered in 1939 by Marguerite Perey of the Curie Institute in Paris, France when she purified 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 which 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.[25] 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 which she later revised to 1%.[26]

Before Perey’s discovery of francium, there were at least three erroneous and incomplete discoveries.[27][28][29][30]

### Eka-francium

The next element below francium is expected to be ununennium (Uue), element 119, although this is not certain. The synthesis of ununennium was 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.[31][32]

$\,^{254}_{99}\mathrm{Es} + \,^{48}_{20}\mathrm{Ca} \to \,^{302}_{119}\mathrm{Uue} ^{*} \to \ \ no\ atoms$
It is highly unlikely that this reaction will be useful 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, due to the rarity of the element, and extreme rarity of the isotope. 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. 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.

Regardless of what the next element below francium actually is, it will still be known as ekafrancium, as it is below francium in the periodic table.

## Production

### Lithium

Satellite images of the Salar del Hombre Muerto, Argentina (left), and Uyuni, Bolivia (right), salt flats are rich in lithium. The lithium-rich brine is concentrated by pumping it into solar evaporation ponds (visible in the left image).

Since the end of World War II lithium production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are 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.[33]

### Sodium

Sodium was first produced commercially in 1855 by thermal reduction of sodium carbonate with carbon at 1100 °C, in what is known as the Deville process.[34]

Na2CO3 (l) + 2 C (s) → 2 Na (g) + 3 CO (g)

A process based on the reduction of sodium hydroxide was developed in 1886.[34] Sodium is now produced commercially through the electrolysis of liquid sodium chloride, based on a process patented in 1924.[35][36] This is done in a Downs Cell in which the NaCl is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is less electropositive than sodium, no calcium will be formed at the anode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide. Very pure sodium can be isolated by the thermal decomposition of sodium azide.[37]

### Potassium

Pure potassium metal may be isolated by electrolysis of its hydroxide in a process that has changed little since Davy.[38] Thermal methods also are employed in potassium production, using potassium chloride.

Potassium salts such as carnallite, langbeinite, polyhalite, potash and sylvite form extensive deposits in ancient lake and seabeds,[citation needed] making extraction of potassium salts in these environments commercially viable.

### Rubidium

Although rubidium is more abundant in Earth’s crust than caesium 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.[39] There are several methods to separate potassium, rubidium and caesium. The fractional crystallization of a rubidium and caesium alum (Cs,Rb)Al(SO4)2·12H2O yields after 30 subsequent steps pure rubidium alum. Reports of two other methods are given in the literature the chlorostannate process and the ferrocyanide process.[39][40] 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.[41] Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite.[39]

### Caesium

Pollucite

The mining of pollucite ore is a selective process and is conducted on a small scale in comparison with most metal mining operations. The ore is crushed, hand-sorted, but not usually concentrated, and then ground. Caesium is then extracted from pollucite mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.[39][42]

### Francium

This sample of uraninite contains about 100,000 atoms (3.3×10−20 g) of francium-223 at any given time.[43]

Francium-223 is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals.[44] In a given sample of uranium, there is estimated to be only one francium atom for every 1×1018 uranium atoms.[43] It is also calculated that there is at most 30 g of francium in the earth’s crust at any time.[45] This makes it the second rarest element in the crust after astatine.[43][46]

Francium can also be synthesized in the nuclear reaction:

197Au + 18O → 210Fr + 5 n

This process, developed by Stony Brook Physics, yields francium isotopes with masses of 209, 210, and 211,[47] which are then isolated by the magneto-optical trap (MOT).[48]

### Eka-francium

Ununennium has not been produced yet as of 2011.[31][32]

## Occurrence

### Lithium

Although lithium is widely distributed on Earth, it does not naturally occur in elemental form due to its high reactivity.[17] The total lithium content of seawater is very large and is estimated as 230 billion tonnes, where the element exists at a relatively constant concentration of 0.14 to 0.25 parts per million (ppm),[49][50] or 25 micromolar;[51] higher concentrations approaching 7 ppm are found near hydrothermal vents.[50]

### Sodium

Owing to its high reactivity, sodium is found in nature only as a compound and never as the free element. Sodium makes up about 2.6% by weight of the Earth‘s crust, making it the sixth most abundant element overall[52] and the most abundant alkali metal. 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, as well as in solid deposits (halite). Others include amphibole, cryolite, soda niter and zeolite.[citation needed]

### Potassium

Potassium does not occur as the free element in nature due to its reactivity, but in compounds, makes up 1.5% of the weight of Earth’s crust and is the seventh most abundant chemical element.[citation needed]

### Rubidium

Rubidium is roughly as abundant as zinc and rather more common than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite and lepidolite.[citation needed]

### Caesium

Caesium is more abundant than antimony, cadmium, tin, tungsten, and much more abundant than mercury or silver, but 30 times less abundant than rubidium—with which it is so closely chemically associated.[53]

### Francium

Francium-223, the only naturally occurring isotope of francium,[54] is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals[44] In a given sample of uranium, there is estimated to be only one francium atom for every 1×1018 uranium atoms.[43][55] It is also calculated that there is at most 30 g of francium in the earth’s crust at any time.[45] This makes it the second rarest element in the crust after astatine.[43][46]

### Eka-francium

Ununennium has not been discovered as of 2011.[31][32]

## Applications

All of the alkali metals have very many applications, some of which are mentioned in the links below:

## Biological occurrences

Lithium carbonate

• Lithium has several biological effects. Lithium carbonate is used to treat bipolar disorder (manic-depression) disorders as it works as a mood stabiliser, although there are side effects, e.g. excessive ingestion of lithium poisons the central nervous system, which is even more dangerous as the lethal dose is only just above the therapeutic dose. Lithium is present in the human body at about 30 ppb.[56]
• Sodium is an essential element for survival as its cation, Na+, is important for nerve function. Sodium is present in the human body at about 0.14%.[57]
• Potassium is also an essential element for survival as its cation, K+, is essential for nerve and heart function. Potassium is present in the human body at about 0.2%.[58]
• Rubidium has no biological role but may help stimulate metabolism, and can accumulate ahead of potassium in muscle. Rubidium is present in the body at about 4.6 ppm.[59]
• Caesium has no biological role, but can replace potassium to some extent in the body due to having similar chemical properties. Thus, caesium compounds should be avoided when possible, especially those containing radioisotopes of caesium, such as 134Cs and 137Cs. Rats fed caesium instead of potassium die. Caesium chloride (non-radioactive) has also promoted as an alternative cancer therapy,[60] but has been linked to the deaths of over 50 patients, when it was used as part of a scientifically unvalidated cancer treatment.[61] Caesium is present in the body at about 20 ppb.[62]
• Francium has no biological role[63] and would most likely be extremely toxic due to its extreme radioactivity. If francium was not radioactive, based on periodic trends it would most likely behave like caesium and would presumably be similarly toxic.[citation needed]

## Precautions

The alkali metals all react with water ever more violently and thus should all be handled with great care. Of note is the toxicity of caesium because it replaces potassium, mentioned above. Also, francium’s extreme radioactivity is a great hazard, although fortunately the greatest quantity of francium ever assembled to date is a sphere of radius 1 mm (over 300,000 neutral atoms).[64]

## Notes

1. ^ Caesium is the spelling recommended by the International Union of Pure and Applied Chemistry (IUPAC).[1] The American Chemical Society (ACS) has used the spelling cesium since 1921,[2][3] following Webster’s Third New International Dictionary.

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