Periodic table, main group elements

The main groups of the periodic table are groups 1,2 and 13 through 18. Elements in these groups are collectively know as main group or representative elements

Group I

Group	1
Period
2	3 Li
3	11 Na
4	19 K
5	37 Rb
6	55 Cs
7	87 Fr

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 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 vigourous reactions with water, and these reactions become increasingly violent as you move down the periods. The reaction with water is as follows:

Alkali metal + water → Alkali metal hydroxide + hydrogen

With potassium as an example:

${K}_{(s)}+{H_{2}O}_{(l)}\to {KOH}_{(aq)}+{\begin{matrix}{\frac {1}{2}}\end{matrix}}{H_{2}}_{(g)}$

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

Explanation of above periodic table slice

Alkali metals

atomic number in black are solids

solid borders are primordial elements (older than the Earth)

dashed borders are naturally radioactive elements with no isotopes older than the Earth

See also

External

• Science aid:Alkali metals A simple look at alkali metals
• Doc Brown, Alkali metals

Wiki

• Lithium
• Sodium
• Potassium
• Rubidium
• Caesium
• Francium

Group 2

Group	2
Period
2	4 Be
3	12 Mg
4	20 Ca
5	38 Sr
6	56 Ba
7	88 Ra

The alkaline earth metals are the series of elements in Group 2 (IUPAC style) of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) (though radium is not always considered an alkaline earth due to its radioactivity).

The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These were named alkaline earths because of their intermediate nature between the alkalis (oxides of the alkali metals) and the rare earths (oxides of rare earth metals). The classification of some apparently inert substances as 'earths' is millennia old. The earliest known system used by the ancient Greeks consisted of four elements, including earth. This system was later refined by philosophers and alchemists such as Aristotle (4th century B.C.E.), Paracelsus (first half of 16th century), John Becher (mid 17th century) and Georg Stahl (late 17th century), with later thinkers subdividing 'earth' into three or more types. The realization that 'earths' were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them Substances simples salifiables terreuses, or salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths.

The alkaline earth metals are silvery colored, soft, low-density metals, which react readily with halogens to form ionic salts, and with water, though not as rapidly as the alkali metals, to form strongly alkaline (basic) hydroxides. Beryllium is an exception: It does not react with water or steam, and its halides are covalent. For example, where sodium and potassium react with water at room temperature, magnesium reacts only with steam and calcium with hot water. These elements all have two electrons in their outermost shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions.

Explanation of above periodic table slice:

Alkaline earth metals

atomic number in black are solids

solid borders are primordial elements (older than the Earth)

dashed borders are naturally radioactive elements with no isotopes older than the Earth

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