The Group 2 Elements of the modern periodic table are also known as the alkaline earth metals. The group 2 elements are Beryllium(Be), Magnesium(Mg), Calcium(Ca), Strontium(Sr), Barium(Br), and Radium(Ra).
Beryllium was known since the time of the Ptolemaic Kingdom in Egypt, in the mineral, beryl. Attempts were made over the time to isolate beryllium, but were unsuccessful. In 1898, French chemist Paul Lebeau performed an electrolysis of a mixture of beryllium fluoride and sodium fluoride, and was successful in isolating and extracting large pure samples of beryllium.
Magnesium was produced in 1808 in England by Humphry Davy using electrolysis of a mixture of magnesia and mercuric oxide.
Calcium was used in the form of lime, as a material for building since 7000 BCE. Calcium has been used in other forms over the time. It was not until 1808 that calcium was isolated. The discovery was done by Humphry Davy, using electrolysis on a mixture of lime and mercuric oxide.
Strontium was first discovered as an ore and named as stronite by Adair Crawford in 1790. It was eventually isolated by Humphry Davy in 1808 by performing electrolysis on strontium chloride and mercuric oxide.
Barium was first recognised in the form of the ore barite in 1774 by Carl Scheele. Many attempts were made to isolate barium, but were unsuccessful. Barium was finally isolated by Humphry Davy in 1808 using electrolysis on molten salts.
Radium was discovered in December 1898 by Marie and Pierre Curie, when they discovered, while studying uraninite, that even after uranium had decayed, the material was still radioactive.
The Group 2 Elements or Alkaline-Earth Metals
The alkaline-earth metals have very similar properties. These elements are shiny, silvery-white and reactive metals at standard temperature and pressure. All discovered alkaline earth metals occur in nature, except for radium, which occurs only through decay of uranium and thorium.
|Valence Electronic Configuration||1s22s2||[Ne]3s2||[Ar]4s2||[Kr]5s2||[Xe]6s2||[Rn]7s2|
|Melting Point/Boiling Point(K)||1560/2742||923/1363||1115/1757||1050/1650||1000/2118||973/2010|
|First Ionization Energy/kJ mol-1||899.5||737.7||589.8||549.5||502.9||509.3|
|Common Oxidation State(s)||0, +1, +2||+1, +2||+1, +2||+1, +2||+1, +2||+2|
Trends in Group 2 Elements
- Electronic Configuration: The general electronic configuration of group 2 elements is ns2. Since these metals have a full s-orbital in their respective valence shells, they tend to readily lose two electrons to form cations with a charge of +2. Thus, the most common oxidation state exhibited by the alkaline earth metals is +2.
- Atomic Radii: Atomic radius increases down the group due to increase in the number of orbits.
- Ionisation Energy: With an increase in atomic size, the valence electron gets shielded by the inner electrons and becomes easily removable with less energy requirement. Hence the ionization energy decreases down the group with an increasing atomic number or atomic size.
- Melting and Boiling Points: Down the group, except for magnesium, the melting and boiling points decrease regularly.
Anomalous Behaviour of Beryllium
Beryllium, compared to the other alkaline-earth metals, has more covalent nature due to its smallest size, high ionization energy, high electro-positive nature and strongest polarizing nature. Beryllium differs from other alkaline earth metals, with the following properties:
- The hardest metal among alkaline earth metals.
- Does not react with water.
- Highest melting and boiling point of beryllium.
- It does not react directly with hydrogen to form hydride.
- Does not liberate hydrogen from acid because of higher electrode potential.
- Beryllium oxide and hydroxide are amphoteric. Dissolves in acids to form salts and in bases to form beryllate.
- Beryllium forms carbide of a different formula and yields methane and not acetylene like other metal on reaction with water.
- Beryllium nitride is volatile.
- Does not react with atmospheric nitrogen and oxygen.
Diagonal Relationship of Beryllium with Aluminium
A diagonal relationship exists between beryllium and aluminium. Beryllium of group 2 resembles more with aluminum of group 13:
- They occur together in the mineral, “Beryl” 3BeO Al2O3 6SiO2.
- They do not react with atmospheric oxygen and nitrogen.
- They do not react with water even at high temperatures.
- They do not liberate hydrogen from acid. They become passive on treatment with concentrated nitric acid.
- They form polyvalent bridged hydrides of covalent nature.
- Halides of both the elements are polyvalent, bridged, and of low melting points. Halides are Lewis acids.
- Water hydrolyzes both nitrides liberate ammonia.
Applications of Group 2 Elements
- Beryllium is mostly used in military applications.
- It is also used as a p-type dopant in semiconductors.
- Magnesium is used in products that benefit from being lightweight, such as car seats, luggage, laptops, cameras and power tools.
- Magnesium is also added to molten iron and steel to remove sulfur.
- Magnesium is used in flares, fireworks and sparklers, due to its property of easy flammability and burning with a bright light.
- Calcium metal is used as a reducing agent in preparing other metals such as thorium and uranium.
- It is also used as an alloying agent for aluminium, beryllium, copper, lead and magnesium alloys.
- Strontium is used as strontium carbonate in the manufacture of red fireworks.
- Pure strontium is used in the study of neurotransmitter release in neurons.
- Barium is used in vacuum tubes to remove gases.
- Radium’s uses have been stopped due to its high radioactivity.
The group 2 elements are called alkaline-earth metals because their respective oxides and hydroxides are strongly alkaline in nature and these metal oxides are found in earth crust.
All alkaline-earth metals are in solid state at room temperature.