The alkali metals are so reactive that they are generally found in nature combined with other elements. Simple minerals, such as halite ( sodium chloride, NaCl), sylvite (potassium chloride, KCl), and carnallite (a potassium-magnesium chloride, KCl · MgCl₂· 6H₂O), are soluble in water and therefore are easily extracted and purified. More complex, water-insoluble minerals are, however, far more abundant in Earth's crust. A very dilute gas of atomic sodium (about 1,000 atoms per cubic cm 【about 16,000 atoms per cubic inch】) is produced in Earth's mesosphere (altitude about 90 km 【60 miles】) by ablation of meteors (meteor and meteoroid). Subsequent reaction of sodium with ozone and atomic oxygen produces excited sodium atoms that emit the light we see as the “tail” of a meteor as well as the more diffuse atmospheric nightglow. Smaller amounts of lithium and potassium are also present.

The alkali metals have the silver-like lustre, high ductility, and excellent conductivity of electricity and heat generally associated with metals. Lithium is the lightest metallic element. The alkali metals have low melting points, ranging from a high of 179 °C (354 °F) for lithium to a low of 28.5 °C (83.3 °F) for cesium. Alloys (alloy) of alkali metals exist that melt as low as −78 °C (−109 °F).

The alkali metals react readily with atmospheric oxygen and water vapour. (Lithium also reacts with nitrogen.) They react vigorously, and often violently, with water to release hydrogen and form strong caustic solutions. Most common nonmetallic substances such as halogens (halogen element), halogen acids, sulfur, and phosphorus react with the alkali metals. The alkali metals themselves react with many organic compounds, particularly those containing a halogen or a readily replaceable hydrogen atom.

Sodium is by far the most important alkali metal in terms of industrial use. The metal is employed in the reduction of organic compounds and in the preparation of many commercial compounds. As a free metal, it is used as a heat-transfer fluid in some nuclear reactors (nuclear reactor). Hundreds of thousands of tons of commercial compounds that contain sodium are used annually, including common salt (NaCl), baking soda (NaHCO₃), sodium carbonate (Na₂CO₃), and caustic soda (NaOH). Potassium has considerably less use than sodium as a free metal. Potassium salts, however, are consumed in considerable tonnages in the manufacture of fertilizers (fertilizer). Lithium metal is used in certain light-metal alloys and as a reactant in organic syntheses. An important use of lithium is in the construction of lightweight batteries (battery). Primary lithium batteries (battery) (not rechargeable) are widely used in many devices such as cameras, cellular telephones, and pacemakers. Rechargeable lithium storage batteries (battery) that could be suitable for vehicle propulsion or energy storage are the subject of intensive research. Rubidium and cesium and their compounds have limited use, but cesium metal vapour is used in atomic clocks (atomic clock), which are so accurate that they are used as time standards.

Alkali metal salts were known to the ancients. The Old Testament refers to a salt called neter (sodium carbonate), which was extracted from the ash of vegetable matter. saltpetre (potassium nitrate) was used in gunpowder, which was invented in China about the 9th century AD and had been introduced into Europe by the 13th century.

In October 1807 the English chemist Sir Humphry Davy (Davy, Sir Humphry, Baronet) isolated potassium and then sodium. The name sodium is derived from the Italian soda, a term applied in the Middle Ages to all alkalies; potassium comes from the French potasse, a name used for the residue left in the evaporation of aqueous solutions derived from wood ashes.

While the German chemists Robert Bunsen (Bunsen, Robert Wilhelm) and Gustav Kirchhoff (Kirchhoff, Gustav Robert) were investigating the mineral waters in the Palatinate in 1860, they obtained a filtrate that was characterized by two lines in the blue region of its spectrum (the light emitted when the sample was inserted into a flame). They suggested the presence of a new alkali element and called it cesium, derived from the Latin caesius, used to designate the blue of the sky. The same researchers, on extracting the alkalies from the mineral lepidolite, separated another solution, which yielded two spectral lines of red colour. They proposed the name rubidium for the element in this solution from the Latin rubidus, which was used for the darkest red colour. Francium was not discovered until 1939 by Marguerite Perey of the Radium Institute in Paris.

In the 19th century the only use for the alkali metals was the employment of sodium as a reagent in the manufacture of aluminum. When the electrolytic (electrolysis) process for aluminum purification was established, it appeared that large-scale use of sodium would cease. Subsequent improvements in the electrolytic production of sodium, however, reduced the cost of this element to such an extent that it can be employed economically to manufacture gasoline additives, reagents for chemical industry, herbicides (herbicide), insecticides (insecticide), nylon, pharmaceuticals, and reagents for metal refining. The continuous electrolysis of sodium hydroxide, a technique called the Castner process, was replaced in 1926 by the Downs cell process. This process, in which a molten sodium chloride–calcium chloride mixture (to reduce the melting point) is electrolyzed, produces both sodium metal and chlorine.

The alkali metals have the high thermal and electrical conductivity, lustre, ductility, and malleability that are characteristic of metals. Each alkali metal atom has a single electron in its outermost shell. This valence electron is much more weakly bound than those in inner shells. As a result, the alkali metals tend to form singly charged positive ions (ion) (cations (cation)) when they react with nonmetals. The compounds that result have high melting points (melting point) and are hard crystals (crystal) that are held together by ionic bonds (ionic bond) (resulting from mutually attractive forces that exist between positive and negative electrical charges). In the metallic state, either pure or in alloys (alloy) with other alkali metals, the valence electrons become delocalized and mobile as they interact to form a half-filled valence band. As with other metals, such a partially filled valence band is a conduction band and is responsible for the valence properties typical of metals (metal). In passing from lithium to francium, the single electron tends to be less strongly held. Generally, the energy necessary to remove the outermost electron from the atoms (atom) of an element, the ionization energy, decreases in the periodic table (periodic law) toward the left and downward in each vertical file, with the result that the most easily ionizable element in the entire table is francium, followed closely by cesium. The alkali metals, which make up the extreme left-hand file, have ionization energies ranging from 124.3 kilocalories per mole (kcal/mole) in lithium to 89.7 kcal/mole in cesium (omitting the rare radioactive element francium). The alkaline-earth metals (alkaline-earth metal), the next group to the right, have higher ionization energies ranging from 214.9 in beryllium to 120.1 kcal/mole in barium.

The electronegativity scale of the elements compares the ability of the atoms of the various elements to attract electrons to themselves. In the periodic table the electronegativities range from 0.7 for cesium, the least electronegative of the elements, to 4.0 for fluorine, the most electronegative. Metals are ordinarily considered to be those elements having values less than 2.0 on the electronegativity scale. As a group the alkali metals are the least electronegative of the elements, ranging from 0.7 to 1.0 on the scale, while the alkaline earths, the next group on the table, have electronegativities ranging from about 0.9 to 1.5.

Vapour-pressure (vapour pressure) data for the alkali metals and for two alloys formed between elements of the group show that the vapour pressures increase in regular fashion with increasing atomic weight. Cesium is the most volatile of the alkali metals, with a boiling point of 671 °C (1,240 °F). The boiling points of the alkali metals decrease in regular fashion as the atomic numbers (atomic number) increase, with the highest, 1,317 °C (2,403 °F), being that of lithium.

The melting points (melting point) of the alkali metals as a group are lower than those of any other nongaseous group of the periodic table, ranging between 179 °C (354 °F) for lithium and 28.5 °C (83.3 °F) for cesium. Among the metallic elements, only mercury has a lower melting point (−38.9 °C, or −38.02 °F) than cesium. The low melting points of the alkali metals are a direct result of the large interatomic distances in their crystals and the weak bond energies associated with such loose arrays. These same factors are responsible for the low densities, low heats of fusion, and small changes in volume upon fusion of the metals. Lithium, sodium, and potassium are less dense than water.

The large size of an alkali metal atom (and the resulting low density of the metal) results from the presence of only one, weakly bound electron in the large outer s-type orbital. Upon going from the noble-gas configuration of argon (atomic number 18) to potassium (atomic number 19), the added electron goes into the large 4s orbital rather than the smaller 3p orbital. When, however, potassium, rubidium, or cesium metals are subjected to increasing pressure (up to one-half million atmospheres or more), a number of phase transitions occur. Ultimately, occupation of a d-type orbital becomes preferred over that of the s-orbital, with the result that these alkali metals resemble transition metals (transition element). Under such circumstances, alloys with transition metals (such as iron) can form, a result that does not occur at low pressures. It has been proposed that the lower-than-expected density of Earth's core may be the result of the formation of a potassium-iron alloy under the extreme pressures that occur there.

The alkali metals have played an important role in quantum physics. Some alkali metal isotopes, such as rubidium-87, are bosons (boson). Dilute atomic gases of such alkali metal isotopes, confined by magnetic fields (magnetic field) or “ laser mirrors” and cooled to temperatures near absolute zero, form Bose-Einstein condensates. In this state, the cluster of atoms is in a single quantum state and exhibits macroscopic behaviours normally seen only with atomic-sized particles. These include interference effects and coherent motion of the entire “cloud” of atoms.

Since the alkali metals are the most electropositive (the least electronegative) of elements, they react with a great variety of nonmetals. In its chemical reactivity, lithium more closely resembles Group 2 (IIa) of the periodic table (periodic law) than it does the other metals of its own group. It is less reactive than the other alkali metals with water, oxygen, and halogens (halogen element) and more reactive with nitrogen, carbon, and hydrogen.

The alkali metals tend to form ionic solids in which the alkali metal has an oxidation number of +1. Therefore, neutral compounds with oxygen can be readily classified according to the nature of the oxygen species involved. Ionic oxygen species include the oxide, O^2-, peroxide, O₂^2-, superoxide, O₂^-, and ozonide O₃^-. Compounds that can be prepared that contain an alkali metal, M, and oxygen are therefore the monoxide, M₂O, peroxide, M₂O₂, superoxide, MO₂, and ozonide, MO₃. Rubidium and cesium and, possibly, potassium also form the sesquioxide, M₄O₆, which contains two peroxide anions and one superoxide anion per formula unit. Lithium forms only the monoxide and the peroxide.

All the alkali metals react directly with oxygen; lithium and sodium form monoxides, Li₂O and Na₂O, and the heavier alkali metals form superoxides, MO₂. The rate of reaction with oxygen, or with air, depends upon whether the metals are in the solid or liquid state, as well as upon the degree of mixing of the metals with the oxygen or air. In the liquid state, alkali metals can be ignited in air with ease, generating copious quantities of heat and a dense choking smoke of the oxide.

The free energy of formation (a measure of stability) of the alkali metal oxides at 25 °C (77 °F) varies widely from a high of −133 kcal/mole for lithium oxide to −63 kcal/mole for cesium oxide. The close approach of the small lithium ion to the oxygen atom results in the unusually high free energy of formation of the oxide. The peroxides (Li₂O₂and Na₂O₂) can be made by passing oxygen through a liquid- ammonia solution of the alkali metal, although sodium peroxide is made commercially by oxidation of sodium monoxide with oxygen. Sodium superoxide (NaO₂) can be prepared with high oxygen pressures, whereas the superoxides of rubidium, potassium, and cesium can be prepared directly by combustion in air. By contrast, no superoxides have been isolated in pure form in the case of lithium or the alkaline-earth metals, although the heavier members of that group can be oxidized to the peroxide state. The cyanides of potassium, rubidium, and cesium, which are less stable than the lower oxides, can be prepared by the reaction of the superoxides with ozone.

The alkali metals all react violently with water according to M + H₂O → MOH + 1/2 H₂. The rate of the reaction depends on the degree of metal surface presented to the liquid. With small metal droplets or thin films of alkali metal, the reaction can be explosive. The rate of the reaction of water with the alkali metals increases with increasing atomic weight of the metal. With the heavier alkali metals, the hydroxides (hydroxide) are highly soluble; thus, they are removed readily from the reacting surface, and the reaction can proceed with unabated vigour. The reaction involves equimolar mixtures (that is, equal numbers of atoms or molecules (molecule)) of the alkali metal and water to form a mole (an amount equal to that of the reactants) of alkali metal hydroxide and half a mole of hydrogen gas. These reactions are highly exothermic (give off heat), and the hydrogen that is generated can react with oxygen to increase further the heat that is generated.

Of the alkali metals, only lithium reacts with nitrogen, and it forms a nitride (Li₃N). In this respect it is more similar to the alkaline-earth metals than to the Group 1 metals. Lithium also forms a relatively stable hydride, whereas the other alkali metals form hydrides that are more reactive. Lithium forms a carbide (Li₂C₂) similar to that of calcium. The other alkali metals do not form stable carbides, although they do react with the graphite form of carbon to give intercalation compounds (substances in which the metal atoms are inserted between layers of carbon atoms in the graphite structure).

The alkali metals can be burned in atmospheres of the various halogens to form the corresponding halides (halide mineral). The reactions are highly exothermic, producing up to 235 kcal/mole for lithium fluoride. The alkali metals react with nonmetals in Groups 15 and 16 (Va and VIa) of the periodic table. Sulfides (sulfide) can be formed by the direct reaction of the alkali metals with elemental sulfur, furnishing a variety of sulfides. phosphorus combines with the alkali metals to form phosphides with the general formula M₃P.

The characteristics of alloy behaviour in alkali metals can be evaluated in terms of the similarity of the elements participating in the alloy. Elements with similar atomic volumes form solid solutions (solid solution) (that is, mix completely in all proportions); some dissimilarity in atomic volumes results in eutectic-type (eutectic) systems (solutions formed over limited concentration ranges), and further dissimilarity results in totally immiscible systems. The high-pressure transition in potassium, rubidium, and cesium that converts these s-type metals to more transition metal-like d-type metals yields atomic volumes that are similar to those of many transition metals at the same pressure. This permits alloys or compounds to form between these alkali metals and such transition metals as nickel or iron.

The elements potassium, rubidium, and cesium, which have rather similar atomic volumes and ionization energies, form complete solid solutions and mixed crystals. Sodium, which is a significantly smaller atom than potassium and has a higher ionization energy, tends to form eutectic systems with potassium, rubidium, and cesium. Even greater dissimilarity exists in the atomic volumes of sodium and lithium, resulting in insolubilities of the liquid phases. The consolute temperature (the temperature at which the two liquids become completely miscible) increases on going from the lithium-sodium alloy system to the lithium-cesium system. Lithium and cesium can coexist as two separate liquid phases at temperatures up to at least 1,100 °C (2,000 °F).

There is only one example of solid miscibility in alkali–alkaline-earth-metal binaries—the lithium-magnesium system, in which the two elements are very similar. Sodium forms compounds only with barium in the alkaline-earth-metal series. The heavier alkali metals all tend to form immiscible liquid phases with the alkaline earth metals.

Several elements in Group 12 (IIb) of the periodic table ( zinc, cadmium, and mercury) react with the alkali metals to form compounds. Mercury forms at least six compounds, commonly termed amalgams, with each of the five alkali metals, and with the exception of the amalgam with lithium, the highest melting point compound in each series has the formula MHg₂. Lithium and sodium also form compounds with cadmium and zinc.

Until the late 1960s there were few complexes of the alkali metal cations (cation) with organic molecules. Specialized biological molecules (molecule) such as valinomycin were known to complex selectively the potassium cation K⁺ for transport across cell membranes (membrane), but synthetic ionophores (molecules that can form complexes with ions) were rare. All the alkali cations have a charge of +1 and, except for lithium, are chemically similar and rather inert. The only significant difference between one alkali cation and another is the size.

The synthesis of crown ethers (ether) by American chemist Charles J. Pedersen (Pedersen, Charles J.) in 1967 provided size-selective cyclic molecules consisting of ether oxygens forming a ring or “crown” that could complex a cation of the right size to fit into the hole in the centre of the molecule. In some cases two crown ether molecules can encapsulate a cation in a “sandwich” fashion. For example, K⁺ just fits into the centre of an 18-crown-6 ring (18 atoms in the ring, 12 of which are carbon atoms and 6 are ether oxygen atoms) to form a 1:1 complex (that is, 1 cation:1 crown ether), K⁺(18C6). Cs⁺ is too large to fit into the ring but can be complexed on one side to form the Cs⁺(18C6) complex or can be sandwiched between two 18-crown-6 molecules to form the 1:2 complex, Cs⁺(18C6)₂. Thus, the selectivity of a crown ether for a particular cation depends on the ring size. Common crown ethers are 12-crown-4, 15-crown-5, and 18-crown-6. These molecules are selective for Li⁺, Na⁺, and K⁺, respectively.

Even greater affinity for alkali cations was achieved by the synthesis of cryptands by French chemist Jean-Marie Lehn (Lehn, Jean-Marie) in 1968 and spherands by American chemist Donald Cram (Cram, Donald J.) in 1979. These are three-dimensional molecules with an internal cavity or crypt that can completely encapsulate the alkali cation. By synthesizing molecules with different cavity sizes, the selectivity for particular cations over those of the “wrong” size to fit in the cavity can be controlled. It should be noted, however, that these molecules are not rigid and that flexibility of the framework can alter the cavity size to accommodate alkali cations of different sizes, although with differences in the strength of complexation.

Since the initial syntheses of crown ethers and cryptands, thousands of complexants for cations of various sizes, charges, and geometries have been synthesized. This has led to an entirely new branch of chemistry called supramolecular chemistry.

Classical methods of separation and analysis of alkali metals are rather difficult and time consuming. For lithium they include such procedures as selective extraction of lithium chloride into organic solvents and the detection of lithium with azo dyes (azo dye) that give highly sensitive colour reactions in alkaline solutions. A modification of the uranyl acetate test (the precipitation of an insoluble sodium salt with uranyl acetate) has been used as a standard test for the presence of sodium. The use of a cobaltinitrite solution permits separation of potassium from sodium by precipitation of the insoluble potassium salt. There are essentially no satisfactory analytical methods for rubidium and cesium based on the use of reagents in solution.

Classical methods of separation of the alkali metals have been largely supplanted by chromatographic elution. Strongly acidic cation-exchange resins and aqueous acidic solutions are used. Generally the affinity increases with atomic weight so that the ions are eluted in the order Fr⁺ \\> Cs⁺ \\> Rb⁺ \\> K⁺ \\> Na⁺ \\> Li⁺, which is the order of decreasing size of the hydrated ions. Ion-exchange resins (ion-exchange resin) that are specific for lithium have been developed. Macrocyclic compounds such as crown ethers (ether) and cryptands that are selective for particular alkali metal ions (ion) have been synthesized. They form cationic complexes that can be dissolved in organic solvents such as chloroform (CHCl₃) with counterions such as picrate (C₆H₂【NO₂】₃O^-).

The characteristic flame colours of the alkali metals (red, yellow, violet, red, and blue for Li, Na, K, Rb, and Cs, respectively) are qualitative indicators of the modern analytical methods used to determine the concentrations of alkali-metal salts (salt) in aqueous solution. The intensities of the characteristic spectral lines in emission after excitation by a flame or ICP (inductively coupled plasma) give quantitative measures of the individual alkali metal concentration in the parts per million range or lower. Determination of one alkali metal in the presence of another, however, can result in interference, which can be reduced by using specially prepared standard solutions that contain known amounts of the interfering metals.

The analysis of the alkali-metal samples for the presence of nonmetallic elements, such as oxygen, carbon, hydrogen, and nitrogen, requires specialized techniques. The oxygen content of sodium and potassium samples can be determined by extraction of the free alkali metal with mercury, leaving behind mercury-insoluble oxides (oxide) and carbonates (carbonate), which can subsequently be analyzed by means of solution methods. The oxygen content of rubidium and cesium can be accurately determined by precise measurement of the freezing point of these two elements.

The carbon content of alkali metals can be analyzed by oxidation of the alkali metal in pure oxygen, followed by infrared measurement of the carbon dioxide generated during combustion. For the analysis of nitride in lithium, the nitride commonly is converted to ammonia, and the ammonia is measured by colorimetric analysis.