Alkali metals are the elements in group 1 of the periodic table, excluding hydrogen from the alkali metal classification. The alkali metals are lithium, sodium, potassium, rubidium, caesium and francium. They are characterised by having one valence electron in their outer s orbital and by a strong general tendency to lose that electron and form singly positive ions with an oxidation state of plus 1. They are all metals and are among the most reactive metallic elements.
The name alkali comes from the ability of their compounds, particularly their hydroxides, to produce alkaline solutions in water. Alkali metals are generally soft, relatively low density metals with comparatively low melting points, although the exact physical properties vary down the group. Their high reactivity means that they are rarely found as free elements in nature and instead occur mainly in ionic compounds and minerals.
The electronic configuration of alkali metals explains much of their chemistry. Lithium has the approximate ground state configuration 1s2 2s1, sodium has a filled neon core followed by 3s1, and the heavier members similarly possess one electron outside a filled inner shell configuration. Removing this single outer electron produces a cation with a stable noble gas like electron configuration.
The first ionisation energy generally decreases down the group because the valence electron occupies a shell farther from the nucleus and experiences greater shielding. Although nuclear charge increases down the group, the increased distance and shielding make the outer electron easier to remove. This trend contributes to the general increase in chemical reactivity down the group, although lithium has several distinctive properties because its small ionic radius and high charge density give it unusual behaviour compared with the heavier alkali metals.
Alkali metals react vigorously with water, producing hydrogen gas and an alkaline metal hydroxide. For sodium, the simplified reaction is represented as 2Na plus 2H2O producing 2NaOH plus H2. Potassium and the heavier alkali metals generally react even more vigorously under comparable conditions. The reaction involves oxidation of the metal from oxidation state zero to plus 1 while water is reduced to hydrogen. The resulting hydroxide remains in solution, increasing hydroxide ion concentration and making the solution alkaline.
The reactions become increasingly vigorous down the group under comparable conditions, partly because the outer electron is easier to remove. Lithium is comparatively less vigorous and can float on water because of its low density, while sodium can melt during the reaction because the reaction releases enough heat to raise the metal temperature substantially. These reactions illustrate why alkali metals must be handled carefully and generally stored under conditions that limit contact with air and moisture.
Alkali metals react with other non metals to form a wide range of ionic compounds. Sodium reacts with chlorine to produce sodium chloride, potassium reacts with oxygen under appropriate conditions to produce oxygen containing compounds, and lithium forms compounds with nitrogen that are unusual among the group because lithium can react directly with atmospheric nitrogen. Their oxides, peroxides and superoxides show systematic differences down the group, reflecting changes in ionic size and lattice and hydration energies.
Sodium commonly forms sodium oxide and sodium peroxide under different conditions, while potassium, rubidium and caesium can form increasingly stable superoxides. The chemistry of these oxygen compounds demonstrates that periodic trends are influenced not only by simple electron loss but also by the relative sizes and stabilisation energies of the resulting ions and anions.
Alkali metals have distinctive physical properties. They are generally soft enough to be cut with a knife, have relatively low densities and have lower melting points than many transition metals. Lithium, sodium and potassium are less dense than water, although their vigorous reactions with water make immersion in water an unsuitable means of demonstrating their density. Their fresh surfaces can appear shiny but quickly tarnish in air because of reactions with oxygen and moisture.
Electrical conductivity is relatively high because metallic bonding provides mobile electrons. Down the group, atomic and ionic radii increase, while ionisation energy generally decreases. Their relatively low melting points and softness are connected with comparatively weak metallic bonding relative to many metals with more extensive valence electron participation, although detailed explanations depend on crystal structure and electronic band properties.
Alkali metals have important industrial and scientific uses. Lithium is central to many rechargeable battery technologies because lithium ions are small and have favourable electrochemical properties. Sodium compounds are extensively used in glass manufacture, chemical production, food processing and many other industries. Potassium compounds are major agricultural fertilisers because potassium is an essential plant nutrient.
Rubidium and caesium have specialised uses in atomic clocks, research and certain electronic technologies, while francium is extremely rare and radioactive and has no substantial commercial application. Sodium and potassium ions are also biologically important because they help establish electrical gradients across cell membranes and participate in nerve and muscle function. Their biological importance does not mean elemental alkali metals are safe in biological environments; the reactive elemental forms differ fundamentally from the stable ions found in salts and aqueous solutions.
The alkali metals show important trends and exceptions that illustrate periodic chemistry. Reactivity generally increases down the group because the valence electron becomes easier to remove. Atomic radius and ionic radius increase, while first ionisation energy generally decreases. Melting points generally decrease down the group, although detailed values depend on crystal structure and bonding. Lithium differs significantly from the heavier members because its small size gives it unusually high polarising power and leads to distinctive compounds and reactions.
Lithium also shows a diagonal relationship with magnesium in some chemical properties. This phenomenon, called the diagonal relationship, demonstrates that periodic trends can involve more than simple vertical comparisons. Sodium and potassium also have different oxygen chemistry, illustrating that group trends can change the stability of different compound types rather than producing a single uniform progression.
Alkali metals are easily confused with hydrogen because hydrogen is placed above group 1 in the periodic table and has one electron, but hydrogen is not classified as an alkali metal. Hydrogen is a non metal under ordinary conditions and has chemical behaviour that differs profoundly from lithium, sodium and the other group 1 metals. Alkali metals should also be distinguished from alkaline earth metals in group 2, which have two valence electrons and commonly form plus 2 ions.
The extreme reactivity of alkali metals is a consequence of their atomic structure rather than an indication that all group 1 compounds are equally reactive. Once an alkali metal forms a stable ion within a compound, its behaviour can be completely different from that of the elemental metal. This distinction between elemental properties and compound properties is essential for understanding why sodium metal is highly reactive while sodium chloride is a stable crystalline salt used widely in everyday life.
