Appearance
The Group 2 elements are all metals with a shiny, silvery white colour

General Reactivity
The Alkaline Earth Metals are high in the reactivity series of metals, but not as high as the Alkali Metals of Group 1.

Occurrence and Extraction
These elements are all found in the Earth's crust, widely distributed in rock structures in their non-elemental forms. Only Magnesium is extracted on a large scale. It is extracted from sea water by the addition of Calcium hydroxide, which precipitates out the less soluble Magnesium hydroxide. This is then converted to Magnesium chloride, which is then electrolysed in a Downs Cell.

Physical Properties
The metals of Group 2 are harder and denser than Sodium and Potassium, and have higher melting points. These properties are due to the presence of two valence electrons on each atom, which gives stronger metallic bonding than occurs in Group 1.

These elements give characteristic colours when heated in a flame:

Magnesium - Brilliant White
Calcium - Brick Red
Strontium - Crimson
Barium - Green
Radium - Red

Atomic and ionic radii increase down the Group, but ionic radii are smaller than the corresponding atomic radii. This is because the atoms lose their two outer electrons to form ions. The remaining electrons occupy closer levels, and the increased effective nuclear charge attracts the electrons towards the nucleus.

Chemical Properties
The chemical properties of Group2 elements are dominated by the strong reducing power of the metals. Once started, the reactions with Oxygen and Chlorine are vigorous:

2Mg(s) + O2(g) è2MgO(s)
Ca(s) + Cl2(g) è CaCl2(s)


All the metals except Beryllium form oxides in air at room temperature which dulls the surface of the metal. Beryllium has to be stored under oil.

All the metals, except Beryllium, reduce water and dilute acids to Hydrogen:

Mg(s) + 2H+(aq) è Mg(aq) + H2(g)


Magnesium reacts only slowly with water unless the water is boiling, but Clacium reacts rapidly even at room temperature, and forms a cloudy white suspension of sparingly soluble Calcium hydroxide.

Calcium, Strontium and Barium can reduce Hydrogen gas when heated, forming the hydride:

Ca(s) + H2(g) è CaH2(s)


The hot metals are also sufficiently strong reducing agents to reduce Nitrogen gas and form nitrides:

3Mg(s) + N2(g) è Mg3N2(s)


Magnesium can reduce, and burn in, Carbon dioxide:


2Mg(s) + CO2(g) è 2MgO(s) + C(s)


This means that Magnesium fires cannot be extinguished using Carbon dioxide fire extinguishers.

Oxides
The oxides of Group 2 metals have the general formula MO and are basic. They are normally prepared by heating the hydroxide or carbonate. They have high melting points. Peroxides, MO2, are known for all these elements except Beryllium

Hydroxides
Calcium, Strontium and Barium oxides react with water to form hydroxides;

CaO(s) + H2O(l) è Ca(OH)2(s)

Calcium hydroxide is known as Slaked Lime. It is sparingly soluble in water and the resulting mildly alkaline solution is known as Limewater which is used to test for the gas Carbon dioxide.

Halides
The Group 2 halides are normally found in the hydrated form. They are all ionic except for Beryllium chloride. Anhydrous Calcium chloride has such a strong affinity for water it is used as a drying agent.