| << Chapter < Page | Chapter >> Page > |
Originally known as eka-caesium, francium was discovered in 1939 by Marguerite Perey ( [link] ) 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 was caused by a previously unidentified element, which exhibited chemical properties of an alkali metal. This led Perey to believe that it was element 87, caused by the alpha decay of actinium-227. Perey named the new isotope actinium-K (now referred to as francium-223) and in 1946, she proposed the name catium for her newly discovered element, however, she subsequently suggested francium, after France.
The abundance of the alkali metals is given in [link] . Sodium’s high abundance is due mainly to large underground deposits of rock salt (NaCl). However, sodium is also more abundant in seawater (10,800 ppm) as compared to potassium (380 ppm).
| Element | Terrestrial abundance (ppm) |
| Li | 20 (Earth’s crust), 40 (soil), 0.17 (sea water) |
| Na | 23,000 (Earth’s crust), 10,500 (sea water) |
| K | 21,000 (Earth’s crust), 14 (soil), 380 (sea water) |
| Rb | 90 (Earth’s crust), 30 - 250 (soil), 0.1 (sea water) |
| Cs | 3 (Earth’s crust), 0.0001 (soil), 0.0003 (sea water) |
| Fr | Essentially nil |
The naturally abundant isotopes of the alkali metals are listed in [link] . All of the isotopes of francium are radioactive. Lithium-7 and sodium-23 are both useful NMR nucleus having I = 1 / 2 .
| Isotope | Natural abundance (%) |
| Lithium-6 | 7.5 |
| Lithium-7 | 92.5 |
| Sodium-23 | 100 |
| Potassium-39 | 93 |
| Potassium-40 | 0.0118 |
| Potassium-41 | 6.9 |
| Caesium-133 | 100 |
Potassium has three isotopes ( [link] ), of which potassium-40 is radioactive and provides the basis for the determination of the age of rocks between 10 5 and 10 9 years old, i.e., those formed in proterozoic and cenozoic periods of geological time. The decay of potassium-40 occurs with a half life of 1.31 x 10 9 years, by two routes. That associated with a beta particle decay accounts for 89% of the decay:
While that associated with an electron capture and by positron emission decay accounts for 11% of the decay to give argon-40. Since many rocks contain potassium containing minerals the decay of potassium-40 after solidification of the rock will result in the formation of argon trapped in the rock. The argon-40 content is determined by mass spectrometry, while the potassium content is determined by flame spectrophotometry. The ratio of the two, [link] , will allow for the determination of the elapsed time since the rock solidified.
Caesium has at least 39 known isotopes (more than any other element except francium) ranging from caesium-112 caesium-151; however, caesium-133 is the only naturally occurring stable isotope. The other isotopes have half-lives from a few days to fractions of a second. The radiogenic isotope caesium-137 is produced from the detonation of nuclear weapons and is produced in nuclear power plants, and was released to the atmosphere most notably from the 1986 Chernobyl accident.
Many of the physical properties of the alkali metals ( [link] ) are typical of metals, e.g., thermal and electrical conductivity. However, due to the relatively weak inter-atomic forces (weak M-M bonding) they are soft and readily cut with a knife.
| Element | Mp (°C) | Bp (°C) | Density (g/cm 3 ) | Electrical resistivity (Ω·cm) |
| Li | 453 | 1615 | 0.534 | 12.17 @ 86 °C |
| Na | 370 | 1156 | 0.968 | 5.23 @ 29 °C |
| K | 336 | 1032 | 0.89 | 7.01 @ 22.8 °C |
| Rb | 312 | 961 | 1.532 | 12.53 @ 53 °C |
| Cs | 301 | 944 | 1.93 | 37.38 @ 28.1 °C |
All the alkali metals are highly reactive and are as a consequence of the stability of the M + ion are strong reducing agents ( [link] ). The metals react readily with hydrogen and oxygen.
| Reduction | Reduction potential (V) |
| Li + + e - → Li | -3.045 |
| Na + + e - → Na | -2.7109 |
| K + + e - → K | -2.924 |
| Rb + + e - → Rb | -2.925 |
| Cs + + e - → Cs | -2.923 |
All of the alkali metals react with water to liberate hydrogen.
A similar, but less violent, reaction is also observed with ammonia when catalyzed by transition metal ions.
In the absence of a catalyst, the Group 1 metals dissolve in liquid ammonia to form solutions with characteristic properties.
As an example, the dissolution of sodium in liquid ammonia results in the formation of solvated Na + cations and electrons.
The solvated electrons are stable in liquid ammonia and form a complex: [e - (NH 3 ) 6 ]. It is this solvated electron that gives the strong reducing properties of the solution as well as the characteristic signal in the ESR spectrum associated with a single unpaired electron. The blue color of the solution is often ascribed to these solvated electrons; however, their absorption is in the far infra-red region of the spectrum. A second species, Na - (solv) , is actually responsible for the blue color of the solution.
The formation of the sodium anion is confirmed by complexation of the cation with a cryptan ligand (“C”) such as a crown ether.
The resulting complex is found to be isostructural to the iodide analog in the solid state.
All the alkali metals form M 2 dimers in the vapor phase in an analogous manner to hydrogen. As with dihydrogen the bonding is associated with the molecular orbital combination of the two valence s-orbitals ( [link] ).
Sodium vapor is commonly used for lighting in a gas discharge lamp, which uses sodium in an excited state to produce light ( [link] ). There are two varieties of such lamps: low pressure and high pressure.
Notification Switch
Would you like to follow the 'Chemistry of the main group elements' conversation and receive update notifications?
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|