(Note that hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the<unk> s). The<unk> s provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behavior down the group.
The<unk> s are all highly reactive and are rarely found in elemental form in nature. As a result, in the laboratory they are stored under mineral oil. They also tarnish easily and have low melting points and densities. Potassium and rubidium possess a weak radioactive characteristic (harmless) due to the presence of long duration radioactive isotopes.
The<unk> s are silver-colored (cesium has a golden tinge), soft, low-density metals, which react readily with halogens to form ionic salts, and with water to form strongly<unk> ne (basic) hydroxides. These elements all have one electron in their outermost shell, so the energetically preferred state of achieving a filled electron shell is to lose one electron to form a singly charged positive ion, or cation.
Hydrogen, with a solitary electron, is usually placed at the top of Group 1 of the periodic table, but it is not considered an<unk> ; rather it exists naturally as a diatomic gas. Removal of its single electron requires considerably more energy than removal of the outer electron for the<unk> s.  As in the halogens, only one additional electron is required to fill in the outermost shell of the hydrogen atom, so hydrogen can in some circumstances behave like a halogen, forming the negative hydride ion. Binary compounds of hydride with the<unk> s and some transition metals have been prepared. Under extremely high pressure, such as is found at the core of Jupiter, hydrogen does become metallic and behaves like an<unk> ; see metallic hydrogen.<unk> s have the lowest ionization potentials in their respective periods, as removing the single electron from the outermost shell gives them the stable inert gas configuration. But their second ionization potentials are very high, as removing an electron from a species having a noble gas configuration is very difficult.<unk> s are famous for their vigorous reactions with water, and these reactions become increasingly violent as one moves down the group.  The reaction with water is as follows:<unk>  + water →<unk>  hydroxide + hydrogen gas
With potassium as an example:<unk> s dissolve in liquid ammonia to give blue solutions that are paramagnetic.
As the solution approaches saturation, it becomes deep purple, then metallic.
Because the solution contains free electrons, it occupies more space than the sum of the volumes of the metal and ammonia. The presence of free electrons also makes these solutions very good reducing agents and good electrical conductors. Since they are easier to handle than the metals themselves they are sometimes used as substitutes.
The solution is not stable over long periods, and the dissolved<unk>  will react to form the corresponding amide.  This reaction, when accelerated with a catalyst (usually iron(III) nitrate), is used for the production of sodium amide:
The amide can be extracted, or it can be converted to sodium azide by bubbling nitrous oxide through the ammonia solution:
Alkali metal
Metal
Halide
Organophosphorus
Terpene
Phytochemical
Organotin
Organic compound