Alkali metal
Alkali metals | |||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|
| |||||||||||
↓ Period | |||||||||||
2 | Lithium (Li) 3 | ||||||||||
3 | Sodium (Na) 11 | ||||||||||
4 | Potassium (K) 19 | ||||||||||
5 | Rubidium (Rb) 37 | ||||||||||
6 | Caesium (Cs) 55 | ||||||||||
7 | Francium (Fr) 87 | ||||||||||
Legend
|
The alkali metals consist of the
The alkali metals are all shiny,
All of the discovered alkali metals occur in nature as their compounds: in order of
Most alkali metals have many different applications. One of the best-known applications of the pure elements is the use of rubidium and caesium in
History
Sodium compounds have been known since ancient times; salt (
Pure potassium was first isolated in 1807 in England by
Rubidium and caesium were the first elements to be discovered using the
Around 1865
After 1869,
There were at least four erroneous and incomplete discoveries
- 227
89Ac
223
87Fr
223
88Ra
219
86Rn
The next element below francium (eka-francium) in the periodic table would be ununennium (Uue), element 119.[37]: 1729–1730 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 the Lawrence Berkeley National Laboratory in Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.[38][39]
It is highly unlikely
Occurrence
In the Solar System
The Oddo–Harkins rule holds that elements with even atomic numbers are more common that those with odd atomic numbers, with the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability.[45][46][47] All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the noble gases and the alkaline earth metals) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones as the alkali metals from rubidium onward can only be synthesised in supernovae and not in stellar nucleosynthesis. Lithium is also much less abundant than sodium and potassium as it is poorly synthesised in both Big Bang nucleosynthesis and in stars: the Big Bang could only produce trace quantities of lithium, beryllium and boron due to the absence of a stable nucleus with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this bottleneck by the triple-alpha process, fusing three helium nuclei to form carbon, and skipping over those three elements.[44]
On Earth
The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the
The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are
Sodium and potassium are very abundant on Earth, both being among the ten
Despite its chemical similarity, lithium typically does not occur together with sodium or potassium due to its smaller size.
Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite,[56] although none of these contain only rubidium and no other alkali metals.[11]: 70 Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.[57]
Properties
Physical and chemical
The physical and chemical properties of the alkali metals can be readily explained by their having an ns1 valence
Name | Lithium | Sodium | Potassium | Rubidium | Caesium | Francium |
---|---|---|---|---|---|---|
Atomic number | 3 | 11 | 19 | 37 | 55 | 87 |
Standard atomic weight[note 7][58][59] | 6.94(1)[note 8] | 22.98976928(2) | 39.0983(1) | 85.4678(3) | 132.9054519(2) | [223][note 9] |
Electron configuration | [He] 2s1 | [Ne] 3s1 | [Ar] 4s1 | [Kr] 5s1 | [Xe] 6s1 | [Rn] 7s1 |
Melting point (°C) | 180.54 | 97.72 | 63.38 | 39.31 | 28.44 | ? |
Boiling point (°C) | 1342 | 883 | 759 | 688 | 671 | ? |
Density (g·cm−3) | 0.534 | 0.968 | 0.89 | 1.532 | 1.93 | ? |
Heat of fusion (kJ·mol−1)
|
3.00 | 2.60 | 2.321 | 2.19 | 2.09 | ? |
Heat of vaporisation (kJ·mol−1)
|
136 | 97.42 | 79.1 | 69 | 66.1 | ? |
Heat of formation of monatomic gas (kJ·mol−1)
|
162 | 108 | 89.6 | 82.0 | 78.2 | ? |
Electrical resistivity at 25 °C (nΩ ·cm)
|
94.7 | 48.8 | 73.9 | 131 | 208 | ? |
pm )
|
152 | 186 | 227 | 248 | 265 | ? |
Ionic radius of hexacoordinate M+ ion (pm) | 76 | 102 | 138 | 152 | 167 | ? |
First kJ·mol−1 )
|
520.2 | 495.8 | 418.8 | 403.0 | 375.7 | 392.8[68] |
Electron affinity (kJ·mol−1) | 59.62 | 52.87 | 48.38 | 46.89 | 45.51 | ? |
Enthalpy of dissociation of M2 (kJ·mol−1) | 106.5 | 73.6 | 57.3 | 45.6 | 44.77 | ? |
Pauling electronegativity | 0.98 | 0.93 | 0.82 | 0.82 | 0.79 | ?[note 10] |
Allen electronegativity | 0.91 | 0.87 | 0.73 | 0.71 | 0.66 | 0.67 |
Standard electrode potential (E°(M+→M0); V)[71] | −3.04 | −2.71 | −2.93 | −2.98 | −3.03 | ? |
nm )
|
Crimson 670.8 |
Yellow 589.2 |
Violet 766.5 |
Red-violet 780.0 |
Blue 455.5 |
? |
The alkali metals are more similar to each other than the elements in any other group are to each other.[5] Indeed, the similarity is so great that it is quite difficult to separate potassium, rubidium, and caesium, due to their similar ionic radii; lithium and sodium are more distinct. For instance, when moving down the table, all known alkali metals show increasing atomic radius,[72] decreasing electronegativity,[72] increasing reactivity,[5] and decreasing melting and boiling points[72] as well as heats of fusion and vaporisation.[11]: 75 In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium.[72] One of the very few properties of the alkali metals that does not display a very smooth trend is their reduction potentials: lithium's value is anomalous, being more negative than the others.[11]: 75 This is because the Li+ ion has a very high hydration energy in the gas phase: though the lithium ion disrupts the structure of water significantly, causing a higher change in entropy, this high hydration energy is enough to make the reduction potentials indicate it as being the most electropositive alkali metal, despite the difficulty of ionising it in the gas phase.[11]: 75
The stable alkali metals are all silver-coloured metals except for caesium, which has a pale golden tint:[73] it is one of only three metals that are clearly coloured (the other two being copper and gold).[11]: 74 Additionally, the heavy alkaline earth metals calcium, strontium, and barium, as well as the divalent lanthanides europium and ytterbium, are pale yellow, though the colour is much less prominent than it is for caesium.[11]: 74 Their lustre tarnishes rapidly in air due to oxidation.[5]
All the alkali metals are highly reactive and are never found in elemental forms in nature.[21] Because of this, they are usually stored in mineral oil or kerosene (paraffin oil).[74] They react aggressively with the halogens to form the alkali metal halides, which are white ionic crystalline compounds that are all soluble in water except lithium fluoride (LiF).[5] The alkali metals 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 if the same number of moles of each metal is used.[5][75][57] The alkali metals have the lowest first ionisation energies in their respective periods of the periodic table[65] because of their low effective nuclear charge[5] and the ability to attain a noble gas configuration by losing just one electron.[5] Not only do the alkali metals react with water, but also with proton donors like alcohols and phenols, gaseous ammonia, and alkynes, the last demonstrating the phenomenal degree of their reactivity. Their great power as reducing agents makes them very useful in liberating other metals from their oxides or halides.[11]: 76
The second ionisation energy of all of the alkali metals is very high
In aqueous solution, the alkali metal ions form
Lithium
The chemistry of lithium shows several differences from that of the rest of the group as the small Li+ cation
Lithium fluoride is the only alkali metal halide that is poorly soluble in water,
Francium
Francium is also predicted to show some differences due to its high
Nuclear
Z |
Alkali metal |
Stable |
Decays |
unstable: italics odd–odd isotopes coloured pink
| ||
---|---|---|---|---|---|---|
3 | lithium | 2 | — | 7 Li |
6 Li |
|
11 | sodium | 1 | — | 23 Na |
||
19 | potassium | 2 | 1 | 39 K |
41 K |
40 K |
37 | rubidium | 1 | 1 | 85 Rb |
87 Rb |
|
55 | caesium | 1 | — | 133 Cs |
||
87 | francium | — | — | No primordial isotopes (223 Fr is a radiogenic nuclide) | ||
Radioactive: 40K, t1/2 1.25 × 109 years; 87Rb, t1/2 4.9 × 1010 years; 223Fr, t1/2 22.0 min. |
All the alkali metals have odd atomic numbers; hence, their isotopes must be either
Due to the great rarity of odd–odd nuclei, almost all the primordial isotopes of the alkali metals are odd–even (the exceptions being the light stable isotope lithium-6 and the long-lived
All of the alkali metals except lithium and caesium have at least one naturally occurring
Periodic trends
The alkali metals are more similar to each other than the elements in any other group are to each other.[5] For instance, when moving down the table, all known alkali metals show increasing atomic radius,[72] decreasing electronegativity,[72] increasing reactivity,[5] and decreasing melting and boiling points[72] as well as heats of fusion and vaporisation.[11]: 75 In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium.[72]
Atomic and ionic radii
The
The ionic radii of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different electron shell than the inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the effective nuclear charge has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases.[5]
First ionisation energy
The first
The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled electron shell and is thus difficult to remove.[5]
Reactivity
The
Electronegativity
Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example,
Melting and boiling points
The
Density
The alkali metals all have the same
Compounds
The alkali metals form complete series of compounds with all usually encountered anions, which well illustrate group trends. These compounds can be described as involving the alkali metals losing electrons to acceptor species and forming monopositive ions.
Hydroxides
External videos | |
---|---|
Reactions of the alkali metals with water, conducted by The Open University |
All the alkali metals react vigorously or explosively with cold water, producing an aqueous solution of a strongly basic alkali metal hydroxide and releasing hydrogen gas.[101] This reaction becomes more vigorous going down the group: lithium reacts steadily with effervescence, but sodium and potassium can ignite, and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers.[5] When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing hydrogen gas; this takes place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water).[75] The alkali metal hydroxides are the most basic known hydroxides.[11]: 87
Recent research has suggested that the explosive behavior of alkali metals in water is driven by a Coulomb explosion rather than solely by rapid generation of hydrogen itself.[106] All alkali metals melt as a part of the reaction with water. Water molecules ionise the bare metallic surface of the liquid metal, leaving a positively charged metal surface and negatively charged water ions. The attraction between the charged metal and water ions will rapidly increase the surface area, causing an exponential increase of ionisation. When the repulsive forces within the liquid metal surface exceeds the forces of the surface tension, it vigorously explodes.[106]
The hydroxides themselves are the most basic hydroxides known, reacting with acids to give salts and with alcohols to give
Intermetallic compounds
The alkali metals form many
Compounds with the group 13 elements
The intermetallic compounds of the alkali metals with the heavier group 13 elements (aluminium, gallium, indium, and thallium), such as NaTl, are poor conductors or semiconductors, unlike the normal alloys with the preceding elements, implying that the alkali metal involved has lost an electron to the Zintl anions involved.[109] Nevertheless, while the elements in group 14 and beyond tend to form discrete anionic clusters, group 13 elements tend to form polymeric ions with the alkali metal cations located between the giant ionic lattice. For example, NaTl consists of a polymeric anion (—Tl−—)n with a covalent diamond cubic structure with Na+ ions located between the anionic lattice. The larger alkali metals cannot fit similarly into an anionic lattice and tend to force the heavier group 13 elements to form anionic clusters.[110]
Compounds with the group 14 elements
Lithium and sodium react with carbon to form acetylides, Li2C2 and Na2C2, which can also be obtained by reaction of the metal with acetylene. Potassium, rubidium, and caesium react with graphite; their atoms are intercalated between the hexagonal graphite layers, forming graphite intercalation compounds of formulae MC60 (dark grey, almost black), MC48 (dark grey, almost black), MC36 (blue), MC24 (steel blue), and MC8 (bronze) (M = K, Rb, or Cs). These compounds are over 200 times more electrically conductive than pure graphite, suggesting that the valence electron of the alkali metal is transferred to the graphite layers (e.g. M+C−8).[66] Upon heating of KC8, the elimination of potassium atoms results in the conversion in sequence to KC24, KC36, KC48 and finally KC60. KC8 is a very strong reducing agent and is pyrophoric and explodes on contact with water.[114][115] While the larger alkali metals (K, Rb, and Cs) initially form MC8, the smaller ones initially form MC6, and indeed they require reaction of the metals with graphite at high temperatures around 500 °C to form.[116] Apart from this, the alkali metals are such strong reducing agents that they can even reduce buckminsterfullerene to produce solid fullerides MnC60; sodium, potassium, rubidium, and caesium can form fullerides where n = 2, 3, 4, or 6, and rubidium and caesium additionally can achieve n = 1.[11]: 285
When the alkali metals react with the heavier elements in the
Nitrides and pnictides
Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with
All the alkali metals react readily with
Oxides and chalcogenides
All the alkali metals react vigorously with
The smaller alkali metals tend to polarise the larger anions (the peroxide and superoxide) due to their small size. This attracts the electrons in the more complex anions towards one of its constituent oxygen atoms, forming an oxide ion and an oxygen atom. This causes lithium to form the oxide exclusively on reaction with oxygen at room temperature. This effect becomes drastically weaker for the larger sodium and potassium, allowing them to form the less stable peroxides. Rubidium and caesium, at the bottom of the group, are so large that even the least stable superoxides can form. Because the superoxide releases the most energy when formed, the superoxide is preferentially formed for the larger alkali metals where the more complex anions are not polarised. The oxides and peroxides for these alkali metals do exist, but do not form upon direct reaction of the metal with oxygen at standard conditions.
Rubidium and caesium can form a great variety of suboxides with the metals in formal oxidation states below +1.[11]: 85 Rubidium can form Rb6O and Rb9O2 (copper-coloured) upon oxidation in air, while caesium forms an immense variety of oxides, such as the ozonide CsO3[126][127] and several brightly coloured suboxides,[128] such as Cs7O (bronze), Cs4O (red-violet), Cs11O3 (violet), Cs3O (dark green),[129] CsO, Cs3O2,[130] as well as Cs7O2.[131][132] The last of these may be heated under vacuum to generate Cs2O.[57]
The alkali metals can also react analogously with the heavier chalcogens (sulfur, selenium, tellurium, and polonium), and all the alkali metal chalcogenides are known (with the exception of francium's). Reaction with an excess of the chalcogen can similarly result in lower chalcogenides, with chalcogen ions containing chains of the chalcogen atoms in question. For example, sodium can react with sulfur to form the sulfide (Na2S) and various polysulfides with the formula Na2Sx (x from 2 to 6), containing the S2−
x ions.[66] Due to the basicity of the Se2− and Te2− ions, the alkali metal selenides and tellurides are alkaline in solution; when reacted directly with selenium and tellurium, alkali metal polyselenides and polytellurides are formed along with the selenides and tellurides with the Se2−
x and Te2−
x ions.[133] They may be obtained directly from the elements in liquid ammonia or when air is not present, and are colourless, water-soluble compounds that air oxidises quickly back to selenium or tellurium.[11]: 766 The alkali metal polonides are all ionic compounds containing the Po2− ion; they are very chemically stable and can be produced by direct reaction of the elements at around 300–400 °C.[11]: 766 [134][135]
Halides, hydrides, and pseudohalides
The alkali metals are among the most
The alkali metals also react similarly with hydrogen to form ionic alkali metal hydrides, where the hydride anion acts as a pseudohalide: these are often used as reducing agents, producing hydrides, complex metal hydrides, or hydrogen gas.[11]: 83 [66] Other pseudohalides are also known, notably the cyanides. These are isostructural to the respective halides except for lithium cyanide, indicating that the cyanide ions may rotate freely.[11]: 322 Ternary alkali metal halide oxides, such as Na3ClO, K3BrO (yellow), Na4Br2O, Na4I2O, and K4Br2O, are also known.[11]: 83 The polyhalides are rather unstable, although those of rubidium and caesium are greatly stabilised by the feeble polarising power of these extremely large cations.[11]: 835
Coordination complexes
Alkali metal cations do not usually form
Ammonia solutions
The alkali metals dissolve slowly in liquid
Organometallic
Organolithium
Being the smallest alkali metal, lithium forms the widest variety of and most stable
Alkyllithiums and aryllithiums may also react with N,N-disubstituted amides to give aldehydes and ketones, and symmetrical ketones by reacting with carbon monoxide. They thermally decompose to eliminate a β-hydrogen, producing alkenes and lithium hydride: another route is the reaction of ethers with alkyl- and aryllithiums that act as strong bases.[11]: 105 In non-polar solvents, aryllithiums react as the carbanions they effectively are, turning carbon dioxide to aromatic carboxylic acids (ArCO2H) and aryl ketones to tertiary carbinols (Ar'2C(Ar)OH). Finally, they may be used to synthesise other organometallic compounds through metal-halogen exchange.[11]: 106
Heavier alkali metals
Unlike the organolithium compounds, the organometallic compounds of the heavier alkali metals are predominantly ionic. The application of
Alkyl and aryl derivatives of sodium and potassium tend to react with air. They cause the cleavage of
- RM + R'X → R–R' + MX
As such, they have to be made by reacting
The alkali metals and their hydrides react with acidic hydrocarbons, for example
Representative reactions of alkali metals
Reaction with oxygen
Upon reacting with oxygen, alkali metals form oxides, peroxides, superoxides and suboxides. However, the first three are more common. The table below[144] shows the types of compounds formed in reaction with oxygen. The compound in brackets represents the minor product of combustion.
Alkali metal | Oxide | Peroxide | Superoxide |
Li | Li2O | (Li2O2) | |
Na | (Na2O) | Na2O2 | |
K | KO2 | ||
Rb | RbO2 | ||
Cs | CsO2 |
The alkali metal peroxides are ionic compounds that are unstable in water. The peroxide anion is weakly bound to the cation, and it is hydrolysed, forming stronger covalent bonds.
- Na2O2 + 2H2O → 2NaOH + H2O2
The other oxygen compounds are also unstable in water.
- 2KO2 + 2H2O → 2KOH + H2O2 + O2[145]
- Li2O + H2O → 2LiOH
Reaction with sulfur
With sulfur, they form sulfides and polysulfides.[146]
- 2Na + 1/8S8 → Na2S + 1/8S8 → Na2S2...Na2S7
Because alkali metal sulfides are essentially salts of a weak acid and a strong base, they form basic solutions.
- S2- + H2O → HS− + HO−
- HS− + H2O → H2S + HO−
Reaction with nitrogen
Lithium is the only metal that combines directly with nitrogen at room temperature.
- 3Li + 1/2N2 → Li3N
Li3N can react with water to liberate ammonia.
- Li3N + 3H2O → 3LiOH + NH3
Reaction with hydrogen
With hydrogen, alkali metals form saline hydrides that hydrolyse in water.
Reaction with carbon
Lithium is the only metal that reacts directly with carbon to give
Reaction with water
On reaction with water, they generate hydroxide ions and hydrogen gas. This reaction is vigorous and highly exothermic and the hydrogen resulted may ignite in air or even explode in the case of Rb and Cs.[144]
- Na + H2O → NaOH + 1/2H2
Reaction with other salts
The alkali metals are very good reducing agents. They can reduce metal cations that are less electropositive. Titanium is produced industrially by the reduction of titanium tetrachloride with Na at 400 °C (van Arkel–de Boer process).
- TiCl4 + 4Na → 4NaCl + Ti
Reaction with organohalide compounds
Alkali metals react with halogen derivatives to generate hydrocarbon via the Wurtz reaction.
- 2CH3-Cl + 2Na → H3C-CH3 + 2NaCl
Alkali metals in liquid ammonia
Alkali metals dissolve in liquid ammonia or other donor solvents like aliphatic amines or hexamethylphosphoramide to give blue solutions. These solutions are believed to contain free electrons.[144]
- Na + xNH3 → Na+ + e(NH3)x−
Due to the presence of solvated electrons, these solutions are very powerful reducing agents used in organic synthesis.
Reaction 1) is known as Birch reduction. Other reductions[144] that can be carried by these solutions are:
- S8 + 2e− → S82-
- Fe(CO)5 + 2e− → Fe(CO)42- + CO
Extensions
Although francium is the heaviest alkali metal that has been discovered, there has been some theoretical work predicting the physical and chemical characteristics of hypothetical heavier alkali metals. Being the first
The stabilisation of ununennium's valence electron and thus the contraction of the 8s orbital cause its atomic radius to be lowered to 240
Not as much work has been done predicting the properties of the alkali metals beyond ununennium. Although a simple extrapolation of the periodic table (by the
The probable properties of further alkali metals beyond unsepttrium have not been explored yet as of 2019, and they may or may not be able to exist.[148] In periods 8 and above of the periodic table, relativistic and shell-structure effects become so strong that extrapolations from lighter congeners become completely inaccurate. In addition, the relativistic and shell-structure effects (which stabilise the s-orbitals and destabilise and expand the d-, f-, and g-orbitals of higher shells) have opposite effects, causing even larger difference between relativistic and non-relativistic calculations of the properties of elements with such high atomic numbers.[37]: 1732–1733 Interest in the chemical properties of ununennium, unhexpentium, and unsepttrium stems from the fact that they are located close to the expected locations of islands of stability, centered at elements 122 (306Ubb) and 164 (482Uhq).[154][155][156]
Pseudo-alkali metals
Many other substances are similar to the alkali metals in their tendency to form monopositive cations. Analogously to the pseudohalogens, they have sometimes been called "pseudo-alkali metals". These substances include some elements and many more polyatomic ions; the polyatomic ions are especially similar to the alkali metals in their large size and weak polarising power.[157]
Hydrogen
The element
Hydrogen, like the alkali metals, has one valence electron[123] and reacts easily with the halogens,[123] but the similarities mostly end there because of the small size of a bare proton H+ compared to the alkali metal cations.[123] Its placement above lithium is primarily due to its electron configuration.[158] It is sometimes placed above fluorine due to their similar chemical properties, though the resemblance is likewise not absolute.[162]
The first ionisation energy of hydrogen (1312.0
The 1s1 electron configuration of hydrogen, while analogous to that of the alkali metals (ns1), is unique because there is no 1p subshell. Hence it can lose an electron to form the
Ammonium and derivatives
The ammonium ion (NH+4) has very similar properties to the heavier alkali metals, acting as an alkali metal intermediate between potassium and rubidium,[157][168] and is often considered a close relative.[169][170][171] For example, most alkali metal salts are soluble in water, a property which ammonium salts share.[172] Ammonium is expected to behave stably as a metal (NH+4 ions in a sea of delocalised electrons) at very high pressures (though less than the typical pressure where transitions from insulating to metallic behaviour occur around, 100 GPa), and could possibly occur inside the ice giants Uranus and Neptune, which may have significant impacts on their interior magnetic fields.[170][171] It has been estimated that the transition from a mixture of ammonia and dihydrogen molecules to metallic ammonium may occur at pressures just below 25 GPa.[170] Under standard conditions, ammonium can form a metallic amalgam with mercury.[173]
Other "pseudo-alkali metals" include the alkylammonium cations, in which some of the hydrogen atoms in the ammonium cation are replaced by alkyl or aryl groups. In particular, the quaternary ammonium cations (NR+4) are very useful since they are permanently charged, and they are often used as an alternative to the expensive Cs+ to stabilise very large and very easily polarisable anions such as HI−2.[11]: 812–9 Tetraalkylammonium hydroxides, like alkali metal hydroxides, are very strong bases that react with atmospheric carbon dioxide to form carbonates.[123]: 256 Furthermore, the nitrogen atom may be replaced by a phosphorus, arsenic, or antimony atom (the heavier nonmetallic pnictogens), creating a phosphonium (PH+4) or arsonium (AsH+4) cation that can itself be substituted similarly; while stibonium (SbH+4) itself is not known, some of its organic derivatives are characterised.[157]
Cobaltocene and derivatives
Thallium
Copper, silver, and gold
The
In Mendeleev's 1871 periodic table, copper, silver, and gold are listed twice, once under group VIII (with the
The coinage metals were traditionally regarded as a subdivision of the alkali metal group, due to them sharing the characteristic s1 electron configuration of the alkali metals (group 1: p6s1; group 11: d10s1). However, the similarities are largely confined to the
Production and isolation
The production of pure alkali metals is somewhat complicated due to their extreme reactivity with commonly used substances, such as water.[5][66] From their silicate ores, all the stable alkali metals may be obtained the same way: sulfuric acid is first used to dissolve the desired alkali metal ion and aluminium(III) ions from the ore (leaching), whereupon basic precipitation removes aluminium ions from the mixture by precipitating it as the hydroxide. The remaining insoluble alkali metal carbonate is then precipitated selectively; the salt is then dissolved in hydrochloric acid to produce the chloride. The result is then left to evaporate and the alkali metal can then be isolated.[66] Lithium and sodium are typically isolated through electrolysis from their liquid chlorides, with calcium chloride typically added to lower the melting point of the mixture. The heavier alkali metals, however, are more typically isolated in a different way, where a reducing agent (typically sodium for potassium and magnesium or calcium for the heaviest alkali metals) is used to reduce the alkali metal chloride. The liquid or gaseous product (the alkali metal) then undergoes fractional distillation for purification.[66] Most routes to the pure alkali metals require the use of electrolysis due to their high reactivity; one of the few which does not is the pyrolysis of the corresponding alkali metal azide, which yields the metal for sodium, potassium, rubidium, and caesium and the nitride for lithium.[123]: 77
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.[184]
Sodium occurs mostly in seawater and dried seabed,[5] but 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.[185][186] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.[187] Potassium occurs in many minerals, such as sylvite (potassium chloride).[5] Previously, potassium was generally made from the electrolysis of potassium chloride or potassium hydroxide,[188] 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.[189] It can also be produced from seawater.[5] However, these methods are problematic because the potassium metal tends to dissolve in its molten chloride and vaporises significantly at the operating temperatures, potentially forming the explosive superoxide. As a result, pure potassium metal is now produced by reducing molten potassium chloride with sodium metal at 850 °C.[11]: 74
- Na (g) + KCl (l) ⇌ NaCl (l) + K (g)
Although sodium is less reactive than potassium, this process works because at such high temperatures potassium is more volatile than sodium and can easily be distilled off, so that the equilibrium shifts towards the right to produce more potassium gas and proceeds almost to completion.[11]: 74
Metals like sodium are obtained by electrolysis of molten salts. Rb & Cs obtained mainly as by products of Li processing. To make pure caesium, ores of caesium and rubidium are crushed and heated to 650 °C with sodium metal, generating an alloy that can then be separated via a fractional distillation technique. Because metallic caesium is too reactive to handle, it is normally offered as caesium azide (CsN3). Caesium hydroxide is formed when caesium interacts aggressively with water and ice (CsOH).[190]
Rubidium is the 16th most abundant element in the earth's crust; however, it is quite rare. Some minerals found in North America, South Africa, Russia, and Canada contain rubidium. Some potassium minerals (lepidolites, biotites, feldspar, carnallite) contain it, together with caesium. Pollucite, carnallite, leucite, and lepidolite are all minerals that contain rubidium. As a by-product of lithium extraction, it is commercially obtained from lepidolite. Rubidium is also found in potassium rocks and brines, which is a commercial supply. The majority of rubidium is now obtained as a byproduct of refining lithium. Rubidium is used in vacuum tubes as a getter, a material that combines with and removes trace gases from vacuum tubes.[191][192]
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.[193] Today the largest producers of caesium, for example the Tanco Mine in Manitoba, Canada, produce rubidium as by-product from pollucite.[194] Today, a common method for separating rubidium from potassium and caesium is the fractional crystallisation of a rubidium and caesium alum (Cs, Rb)Al(SO4)2·12H2O, which yields pure rubidium alum after approximately 30 recrystallisations.[194][195] The limited applications and the lack of a mineral rich in rubidium limit the production of rubidium compounds to 2 to 4 tonnes per year.[194] 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.[194][196] Both metals are produced as by-products of lithium production: after 1958, when interest in lithium's thermonuclear properties increased sharply, the production of rubidium and caesium also increased correspondingly.[11]: 71 Pure rubidium and caesium metals are produced by reducing their chlorides with calcium metal at 750 °C and low pressure.[11]: 74
As a result of its extreme rarity in nature,
Applications
Lithium, sodium, and potassium have many applications, while rubidium and caesium are very useful in academic contexts but do not have many applications yet.[11]: 68 Lithium is often used in lithium-ion batteries, and lithium oxide can help process silica. Lithium stearate is a thickener and can be used to make lubricating greases; it is produced from lithium hydroxide, which is also used to absorb carbon dioxide in space capsules and submarines.[11]: 70 Lithium chloride is used as a brazing alloy for aluminium parts.[200] Metallic lithium is used in alloys with magnesium and aluminium to give very tough and light alloys.[11]: 70
Sodium compounds have many applications, the most well-known being sodium chloride as
Potassium compounds are often used as
Rubidium and caesium are often used in
Francium has no commercial applications,
Biological role and precautions
Metals
Pure alkali metals are dangerously reactive with air and water and must be kept away from heat, fire, oxidising agents, acids, most organic compounds, halocarbons, plastics, and moisture. They also react with carbon dioxide and carbon tetrachloride, so that normal fire extinguishers are counterproductive when used on alkali metal fires.[216] Some Class D dry powder extinguishers designed for metal fires are effective, depriving the fire of oxygen and cooling the alkali metal.[217]
Experiments are usually conducted using only small quantities of a few grams in a
Ions
The bioinorganic chemistry of the alkali metal ions has been extensively reviewed.[221] Solid state crystal structures have been determined for many complexes of alkali metal ions in small peptides, nucleic acid constituents, carbohydrates and ionophore complexes.[222]
Lithium naturally only occurs in traces in biological systems and has no known biological role, but does have effects on the body when ingested.
Sodium and potassium occur in all known biological systems, generally functioning as
Potassium is the major
Due to their similar atomic radii, rubidium and caesium in the body mimic potassium and are taken up similarly. Rubidium has no known biological role, but may help stimulate metabolism,[238][239][240] and, similarly to caesium,[238][241] replace potassium in the body causing potassium deficiency.[238][240] Partial substitution is quite possible and rather non-toxic: a 70 kg person contains on average 0.36 g of rubidium, and an increase in this value by 50 to 100 times did not show negative effects in test persons.[242] Rats can survive up to 50% substitution of potassium by rubidium.[240][243] Rubidium (and to a much lesser extent caesium) can function as temporary cures for hypokalemia; while rubidium can adequately physiologically substitute potassium in some systems, caesium is never able to do so.[239] There is only very limited evidence in the form of deficiency symptoms for rubidium being possibly essential in goats; even if this is true, the trace amounts usually present in food are more than enough.[244][245]
Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic. Like rubidium, caesium tends to substitute potassium in the body, but is significantly larger and is therefore a poorer substitute.
Notes
- ^ The symbols Na and K for sodium and potassium are derived from their Latin names, natrium and kalium; these are still the origins of the names for the elements in some languages, such as German and Russian.
- ^ 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.
- Roman numeral).[4]
- ^ While hydrogen also has this electron configuration, it is not considered an alkali metal as it has very different behaviour owing to the lack of valence p-orbitals in period 1 elements.
- ^ In the 1869 version of Mendeleev's periodic table, copper and silver were placed in their own group, aligned with hydrogen and mercury, while gold was tentatively placed under uranium and the undiscovered eka-aluminium in the boron group.
- ^ The asterisk denotes an excited state.
- least significant figure(s) of the number prior to the parenthesised value (ie. counting from rightmost digit to left). For instance, 1.00794(7) stands for 1.00794±0.00007, while 1.00794(72) stands for 1.00794±0.00072.[67]
- ^ The value listed is the conventional value suitable for trade and commerce; the actual value may range from 6.938 to 6.997 depending on the isotopic composition of the sample.[59]
- ^ The element does not have any stable nuclides, and a value in brackets indicates the mass number of the longest-lived isotope of the element.[58][59]
- Pauling scale, the same as caesium;[69] the value for caesium has since been refined to 0.79, although there are no experimental data to allow a refinement of the value for francium.[70] Francium has a slightly higher ionisation energy than caesium,[68] 392.811(4) kJ/mol as opposed to 375.7041(2) kJ/mol for caesium, as would be expected from relativistic effects, and this would imply that caesium is the less electronegative of the two.
References
- ISBN 0-85404-438-8. pp. 248–49. Electronic version..
- ISBN 978-0-8412-3999-9.
- (PDF) from the original on 9 October 2022.
- ^ (PDF) from the original on 9 October 2022. Retrieved 24 March 2012.
- ^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah Royal Society of Chemistry. "Visual Elements: Group 1 – The Alkali Metals". Visual Elements. Royal Society of Chemistry. Archived from the original on 5 August 2012. Retrieved 13 January 2012.
- ^ Harper, Douglas. "salary". Online Etymology Dictionary.
- ^ Marggraf, Andreas Siegmund (1761). Chymische Schriften (in German). p. 167.
- ^ du Monceau, H. L. D. (1736). "Sur la Base de Sel Marine". Mémoires de l'Académie Royale des Sciences (in French): 65–68.
- ^ .
- ^ S2CID 38152048.
- ^ ISBN 978-0-08-037941-8.
- ISBN 978-3-527-30666-4.
- .
- S2CID 96141217.
- ^ Ralph, Jolyon; Chau, Ida (24 August 2011). "Petalite: Petalite mineral information and data". Retrieved 27 November 2011.
- ^ a b Winter, Mark. "WebElements Periodic Table of the Elements | Lithium | historical information". Retrieved 27 November 2011.
- ISBN 978-0-7661-3872-8.
- ^ "Johan Arfwedson". Archived from the original on 5 June 2008. Retrieved 10 August 2009.
- ^ a b van der Krogt, Peter. "Lithium". Elementymology & Elements Multidict. Retrieved 5 October 2010.
- ^ Clark, Jim (2005). "Compounds of the Group 1 Elements". chemguide. Retrieved 10 August 2009.
- ^ ISBN 978-0-313-33438-2.
- ^ a b c d Leach, Mark R. (1999–2012). "The Internet Database of Periodic Tables". meta-synthesis.com. Retrieved 6 April 2012.
- ^ a b Kaner, Richard (2003). "C&EN: It's Elemental: The Periodic Table – Cesium". American Chemical Society. Retrieved 25 February 2010.
- (PDF) from the original on 9 October 2022.
- .
- ^ "caesium". Oxford English Dictionary (2nd ed.). Oxford University Press. (Subscription or participating institution membership required.)
- ^ Newlands, John A. R. (20 August 1864). "On Relations Among the Equivalents". Chemical News. 10: 94–95. Archived from the original on 1 January 2011. Retrieved 25 November 2013.
- ^ Newlands, John A. R. (18 August 1865). "On the Law of Octaves". Chemical News. 12: 83. Archived from the original on 1 January 2011. Retrieved 25 November 2013.
- ^ Mendelejew, Dimitri (1869). "Über die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente". Zeitschrift für Chemie (in German): 405–406.
- ^ doi:10.1021/ed080p952. Archived from the original(PDF) on 11 June 2010. Retrieved 6 May 2012.
- ^ a b Fontani, Marco (10 September 2005). "The Twilight of the Naturally-Occurring Elements: Moldavium (Ml), Sequanium (Sq) and Dor (Do)". International Conference on the History of Chemistry. Lisbon. pp. 1–8. Archived from the original on 24 February 2006. Retrieved 8 April 2007.
- ^ a b Van der Krogt, Peter (10 January 2006). "Francium". Elementymology & Elements Multidict. Retrieved 8 April 2007.
- ^ "Education: Alabamine & Virginium". Time. 15 February 1932. Archived from the original on 30 September 2007. Retrieved 1 April 2007.
- .
- ^ Adloff, Jean-Pierre; Kaufman, George B. (25 September 2005). Francium (Atomic Number 87), the Last Discovered Natural Element Archived 4 June 2013 at the Wayback Machine. The Chemical Educator 10 (5). Retrieved 26 March 2007.
- ISBN 978-0-07-913665-7.
- ^ ISBN 978-1-4020-3555-5.
- ^ PMID 9953034.
- ^ van der Krogt, Peter. "Ununennium". Elementymology & Elements Multidict. Retrieved 14 February 2011.
- .
- ^ "Hunt for element 119 set to begin". Chemistry World. 12 September 2017. Retrieved 9 January 2018.
- ^ a b c Seaborg, G. T. (c. 2006). "transuranium element (chemical element)". Encyclopædia Britannica. Retrieved 16 March 2010.
- ISBN 978-0-19-960563-7.
- ^ doi:10.1086/375492.
- .
- .
- ISBN 978-0-226-59441-5.
- PMID 16592930.
- ISBN 978-0-521-89148-6.
- ^ "Abundance in Earth's Crust". WebElements.com. Retrieved 14 April 2007.
- ^ "List of Periodic Table Elements Sorted by Abundance in Earth's crust". Israel Science and Technology Directory.
- ^ ISBN 0-8493-0486-5.
- ^ "Lithium Occurrence". Institute of Ocean Energy, Saga University, Japan. Archived from the original on 2 May 2009. Retrieved 13 March 2009.
- ^ "Some Facts about Lithium". ENC Labs. Archived from the original on 10 July 2011. Retrieved 15 October 2010.
- S2CID 93866412.
- S2CID 140585007.
- ^ a b c d e Butterman, William C.; Brooks, William E.; Reese, Robert G. Jr. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Archived from the original (PDF) on 22 November 2009. Retrieved 27 December 2009.
- ^ (PDF) from the original on 9 October 2022. Retrieved 7 February 2012.
- ^ (PDF) from the original on 9 October 2022. Retrieved 11 February 2012.
- ISBN 978-0-8493-0474-3.
- ^ ISBN 978-0-19-850341-5.
- ^ a b Gagnon, Steve. "Francium". Jefferson Science Associates, LLC. Archived from the original on 31 March 2007. Retrieved 1 April 2007.
- ^ a b Winter, Mark. "Geological information". Francium. The University of Sheffield. Retrieved 26 March 2007.
- ^ "It's Elemental — The Periodic Table of Elements". Jefferson Lab. Archived from the original on 29 April 2007. Retrieved 14 April 2007.
- ^ a b c Lide, D. R., ed. (2003). CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, FL: CRC Press.
- ^ ISBN 978-0-8053-3799-0. Retrieved 24 June 2013.
- CODATA reference. National Institute of Standards and Technology. Retrieved 26 September 2011.
- ^ PMID 10035190.
- ISBN 978-0-8014-0333-0.
- .
- ^ Vanýsek, Petr (2011). “Electrochemical Series”, in Handbook of Chemistry and Physics: 92nd Edition Archived 24 July 2017 at the Wayback Machine (Chemical Rubber Company).
- ^ a b c d e f g h i j k l m n o p q r s Clark, Jim (2005). "Atomic and Physical Properties of the Group 1 Elements". chemguide. Retrieved 30 January 2012.
- ^ Gray, Theodore. "Facts, pictures, stories about the element Cesium in the Periodic Table". The Wooden Periodic Table Table. Retrieved 13 January 2012.
- ^ The OpenLearn team (2012). "Alkali metals". OpenLearn. The Open University. Retrieved 9 July 2012.
- ^ a b Gray, Theodore. "Alkali Metal Bangs". Theodore Gray. Retrieved 13 May 2012.
- .
- .
- .
- PMID 12590555.
- ^ PMID 12022811.
- ^ (PDF) from the original on 9 October 2022.
- ISBN 978-0-85312-027-8.
- ^ ISBN 978-0-471-97058-3.
- (PDF) from the original on 9 October 2022. Retrieved 23 August 2014.
- ^ a b c d e Clark, Jim (2005). "Reaction of the Group 1 Elements with Oxygen and Chlorine". chemguide. Retrieved 27 June 2012.
- ]
- ^ ISBN 978-1-4020-9974-8.
- .
- ISBN 978-3-527-33541-1.
- ^ OCLC 179976746. Archived from the originalon 24 July 2017. Retrieved 23 May 2008.
- ^ "Universal Nuclide Chart". Nucleonica. Institute for Transuranium Elements. 2007–2012. Retrieved 17 April 2011.
- ^ a b c Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brookhaven National Laboratory. Archived from the original on 21 July 2011. Retrieved 4 October 2012.
- .
- .
- ^ "Potassium-40" (PDF). Human Health Fact Sheet. Argonne National Laboratory, Environmental Science Division. August 2005. Archived (PDF) from the original on 9 October 2022. Retrieved 7 February 2012.
- ^ National Institute of Standards and Technology (6 September 2009). "Radionuclide Half-Life Measurements". Archived from the original on 12 August 2016. Retrieved 7 November 2011.
- ^ Radioisotope Brief: Cesium-137 (Cs-137). U.S. National Center for Environmental Health
- ^ IAEA. 1988.
- ISBN 978-1-870965-87-3.
- ISBN 978-0-13-061142-0.
- ^ a b Clark, Jim (2005). "Reaction of the Group 1 Elements with Water". chemguide. Retrieved 18 June 2012.
- ISBN 978-0-07-023684-4. Section 17.43, page 321
- ISBN 978-1-56670-495-3.
- ^ a b Clark, Jim (2000). "Metallic Bonding". chemguide. Retrieved 23 March 2012.
- ^ PMID 25698335.
- ISBN 978-0-471-93623-7.
- ^ "Sodium-Potassium Alloy (NaK)" (PDF). BASF. December 2004. Archived from the original (PDF) on 27 September 2007.
- ISBN 978-0-471-49315-0.
- ^ ISBN 0-471-93620-0
- ISBN 978-0-470-14549-4.
- ISBN 978-3-642-66620-9.
- PMID 23066852. Archived from the original(PDF) on 27 September 2020. Retrieved 21 August 2016.
- ISBN 978-0-13-175553-6.
- ^ NIST Ionizing Radiation Division 2001 – Technical Highlights. physics.nist.gov
- PMID 27878015.
- PMID 11932511.
- .
- PMID 19750706.
- Chemical & Engineering News 80 No. 20 (20 May 2002)
- ISBN 0-471-93620-0
- ISBN 978-1-4097-6995-8.
- ^ ISBN 978-0-471-18602-1.
- ^ "Welcome to Arthur Mar's Research Group". University of Alberta. 1999–2013. Archived from the original on 4 December 2012. Retrieved 24 June 2013.
- .
- .
- S2CID 250883291.
- .
- .
- S2CID 96084147.
- .
- .
- ISBN 978-0-12-356786-4.
- doi:10.2172/4367751. TID-5221.
- ISBN 978-0-12-023604-6.
- PMID 11457025.
- ISBN 978-0-471-17560-5.
- .
- ISBN 3-527-29390-6.
- .
- .
- (PDF) from the original on 9 October 2022.
- .
- ^ a b c d "Inorganic Chemistry" by Gary L. Miessler and Donald A. Tar, 6th edition, Pearson
- ISBN 978-8122413847.
- ^ "The chemistry of the Elements" by Greenwood and Earnshaw, 2nd edition, Elsevier
- ^ "Inorganic Chemistry" by Cotton and Wilkinson
- ^ PMID 20967377.
- ^ Gäggeler, Heinz W. (5–7 November 2007). "Gas Phase Chemistry of Superheavy Elements" (PDF). Lecture Course Texas A&M. Archived from the original (PDF) on 20 February 2012. Retrieved 26 February 2012.
- . Retrieved 16 November 2022.
- ISBN 978-3-540-07109-9. Retrieved 4 October 2013.
- doi:10.1016/0092-640X(77)90010-9. Archived from the original(PDF) on 22 March 2016. Retrieved 25 February 2016.
- ^ Kul'sha, A. V. "Есть ли граница у таблицы Менделеева?" [Is there a boundary to the Mendeleev table?] (PDF). www.primefan.ru (in Russian). Archived (PDF) from the original on 9 October 2022. Retrieved 8 September 2018.
- ^ Kratz, J. V. (5 September 2011). The Impact of Superheavy Elements on the Chemical and Physical Sciences (PDF). 4th International Conference on the Chemistry and Physics of the Transactinide Elements. Archived (PDF) from the original on 9 October 2022. Retrieved 27 August 2013.
- ^ Nuclear scientists eye future landfall on a second 'island of stability'. EurekAlert! (2008-04-06). Retrieved on 2016-11-25.
- S2CID 120251297.
- ^ PMID 17441140.
- ^ a b "International Union of Pure and Applied Chemistry > Periodic Table of the Elements". IUPAC. Retrieved 1 May 2011.
- ^ Folden, Cody (31 January 2009). "The Heaviest Elements in the Universe" (PDF). Saturday Morning Physics at Texas A&M. Archived from the original (PDF) on 10 August 2014. Retrieved 9 March 2012.
- ^ Emsley, J. (1989). The Elements. Oxford: Clarendon Press. pp. 22–23.
- ISBN 0-19-855694-2
- ^ (PDF) from the original on 9 October 2022.
- ^ Huheey, J.E.; Keiter, E.A. and Keiter, R.L. (1993) Inorganic Chemistry: Principles of Structure and Reactivity, 4th edition, HarperCollins, New York, USA.
- ^ James, A.M. and Lord, M.P. (1992) Macmillan's Chemical and Physical Data, Macmillan, London, UK.
- .
- S2CID 97261131.
- PMID 24217230. Retrieved 7 August 2014.
- ^ Leach, Mark R. "2002 Inorganic Chemist's Periodic Table". Retrieved 16 October 2012.
- ISBN 0-12-352651-5
- ^ S2CID 4199721.
- ^ .
- ^ "Solubility Rules!". chem.sc.edu.
- .
- PMID 11848774.
- .
- .
- doi:10.1107/S0567739476001551. Archived from the original(PDF) on 17 March 2020. Retrieved 4 September 2019.
- ^ .
- ISBN 0-85404-438-8. pp. 51. Electronic version..
- .
- ISBN 0-471-64952-X
- ^ Deming HG (1940) Fundamental Chemistry, John Wiley & Sons, New York, pp. 705–7
- ISBN 1-57215-291-5
- ^ Ober, Joyce A. "Lithium" (PDF). United States Geological Survey. pp. 77–78. Archived (PDF) from the original on 11 July 2007. Retrieved 19 August 2007.
- ^ Pauling, Linus. General Chemistry (1970 ed.). Dover Publications.
- ^ "Los Alamos National Laboratory – Sodium". Retrieved 8 June 2007.
- ^ Merck Index, 9th ed., monograph 8325
- ^ Winter, Mark. "WebElements Periodic Table of the Elements | Potassium | Essential information". Webelements. Retrieved 27 November 2011.
- ISBN 978-0-471-23896-6.
- ^ "Cesium | Cs (Element) – PubChem". pubchem.ncbi.nlm.nih.gov. Retrieved 18 December 2021.
- ^ "WebElements Periodic Table » Rubidium » geological information". www.webelements.com. Retrieved 18 December 2021.
- ^ Liu, Jinlian & Yin, Zhoulan & Li, Xinhai & Hu, Qiyang & Liu, Wei. (2019). A novel process for the selective precipitation of valuable metals from lepidolite. Minerals Engineering. 135. 29–36. 10.1016/j.mineng.2018.11.046.
- .
- ^ a b c d Butterman, William C.; Brooks, William E.; Reese, Robert G. Jr. (2003). "Mineral Commodity Profile: Rubidium" (PDF). United States Geological Survey. Archived (PDF) from the original on 9 October 2022. Retrieved 4 December 2010.
- ^ bulletin 585. United States. Bureau of Mines. 1995.
- ISBN 978-0-471-48494-3.
- .
- ^ .
- ^ a b Price, Andy (20 December 2004). "Francium". Retrieved 19 February 2012.
- ^ USGS (2011). "Lithium" (PDF). Archived (PDF) from the original on 9 October 2022. Retrieved 4 December 2011.
- ^ "Soaps & Detergents: Chemistry". Retrieved 20 July 2015.
- ISBN 978-0-88173-212-2.
- ISBN 978-0-88173-351-8.
- ^ Stampers, National Association of Drop Forgers and (1957). Metal treatment and drop forging.
- ^ Harris, Jay C (1949). Metal cleaning bibliographical abstracts. p. 76.
- ^ Cordel, Oskar (1868). Die Stassfurter Kalisalze in der Landwirthschalt: Eine Besprechung ... (in German). L. Schnock.
- ISBN 978-0-313-32579-3.
- ISBN 978-3-527-30673-2.
- ^ "Cesium Atoms at Work". Time Service Department—U.S. Naval Observatory—Department of the Navy. Archived from the original on 23 February 2015. Retrieved 20 December 2009.
- ^ "The NIST reference on Constants, Units, and Uncertainty". National Institute of Standards and Technology. 5 February 2015.
- ^ Koch, E.-C. (2002). "Special Materials in Pyrotechnics, Part II: Application of Caesium and Rubidium Compounds in Pyrotechnics". Journal Pyrotechnics. 15: 9–24. Archived from the original on 13 July 2011. Retrieved 3 November 2011.
- ISBN 978-0-8306-3015-8.
- ^ Winter, Mark. "Uses". Francium. The University of Sheffield. Archived from the original on 31 March 2007. Retrieved 25 March 2007.
- S2CID 15917603.
- (PDF) from the original on 9 October 2022. Retrieved 11 September 2009.
- ^ a b c Lerner, Michael M. (2013). "Standard Operating Procedure: Storage and Handling of Alkali Metals". Oregon State University. Retrieved 26 August 2016.
- ISBN 978-0-87765-472-8.
- ISBN 978-0-935702-48-4.
- ISBN 978-1-903996-65-2.
- ^ Wray, Thomas K. "Danger: peroxidazable chemicals" (PDF). Environmental Health & Public Safety (North Carolina State University). Archived from the original (PDF) on 8 June 2011.
- S2CID 5983458.
- PMID 26860299.
- ^ a b c d e Winter, Mark. "WebElements Periodic Table of the Elements | Lithium | biological information". Webelements. Retrieved 15 February 2011.
- ^ Gray, Theodore. "Facts, pictures, stories about the element Lithium in the Periodic Table". theodoregray.com. Retrieved 9 January 2012.
- PMID 17848039.
- PMID 21301855.
- ^ a b c Winter, Mark. "WebElements Periodic Table of the Elements | Potassium | biological information". WebElements. Retrieved 13 January 2012.
- ^ a b Winter, Mark. "WebElements Periodic Table of the Elements | Sodium | biological information". WebElements. Retrieved 13 January 2012.
- ^ "Sodium" (PDF). Northwestern University. Archived from the original (PDF) on 23 August 2011. Retrieved 21 November 2011.
- ^ "Sodium and Potassium Quick Health Facts". Retrieved 7 November 2011.
- United States National Academies. 11 February 2004. Archived from the originalon 6 October 2011. Retrieved 23 November 2011.
- OCLC 738512922. Archived from the original(PDF) on 27 October 2011. Retrieved 23 November 2011.
- PMID 15369026. Archived from the originalon 1 August 2018. Retrieved 30 August 2017.
- S2CID 19315480. Archived from the original(PDF) on 28 January 2012.
- ^ PMID 16253415.
- ^ ISBN 978-0-7817-2845-4.
- ISBN 978-1-56053-503-4.
- ^ a b c Winter, Mark. "WebElements Periodic Table of the Elements | Rubidium | biological information". Webelements. Retrieved 15 February 2011.
- ^ PMID 13409924.
- ^ S2CID 2574742. Archived from the originalon 9 July 2012.
- ^ a b Winter, Mark. "WebElements Periodic Table of the Elements | Caesium | biological information". WebElements. Retrieved 13 January 2012.
- S2CID 33738527.
- doi:10.1152/ajplegacy.1943.138.2.246.)
{{cite journal}}
: CS1 maint: numeric names: authors list (link - ISBN 978-0-7872-7680-5.
- ISBN 978-0-7637-0765-1.
- S2CID 19186683.
- ^ .
- PMID 1154391.
- S2CID 11947121.
- ^ Wood, Leonie. "'Cured' cancer patients died, court told". The Sydney Morning Herald. 20 November 2010.