Alkali metal

The alkali metals are the series of elements in Group 1 (IUPAC style) of the periodic table (excluding hydrogen in all but one rare circumstance): lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). They are all highly reactive and are never found in elemental form in nature. As a result, they are stored under mineral oil.

Introduction
The alkali metals are silver-colored (caesium has a golden tinge), soft, low-density metals, which react readily with halogens to form ionic salts, and with water to form strongly alkaline (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.

Hydrogen, with a solitary electron, is sometimes placed at the top of Group 1, but it is not an alkali metal (except under extreme circumstances as metallic hydrogen); 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 alkali metals. 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 alkali metals 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 alkali metal; see metallic hydrogen.

Alkali metals are highly reactive. They 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.

Reactions in water
Alkali metals are famous for their vigorous reactions with water, and these reactions become increasingly violent as you move down the groups. The reaction with water is as follows:

Alkali metal + water → Alkali metal hydroxide + hydrogen

With potassium as an example:


 * $$ 2{K}_{(s)} + 2{H_2O}_{(l)} \to 2{KOH}_{(aq)} + {H_2}_{(g)} $$

In this reaction, enough energy is produced to ignite the hydrogen, creating a lilac flame above the potassium.

Reaction in ammonia
Alkali metals dissolve in liquid ammonia to give blue solutions that are paramagnetic


 * $$ {K}_{(s)} + {NH_3}_{(l)} \to {K}^{+}  _{(solv)} + {e}^{-}   _{(solv)}$$

Because of the free electrons the solution occupies more space than the sum of the volumes of the metal and ammonia. The free electrons also makes these solutions very good reducing agents.