Saturday, April 30, 2011

Alkaline Earth Metals

The elements of Group 2, the Alkaline Earth Metals, are:

  symbol electron configuration
beryllium Be [He]2s2
magnesium Mg [Ne]3s2
calcium Ca [Ar]4s2
strontium Sr [Kr]5s2
barium Ba [Xe]6s2
radium Ra [Rn]7s2

The last element, radium, is radioactive and will not be considered here.
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, but not in the elemental form as they are so reactive. Instead, they are widely distributed in rock structures. The main minerals in which magnesium is found are carnellite, magnesite and dolomite. Calcium is found in chalk, limestone, gypsum and anhydrite. Magnesium is the eighth most abundant element in the Earth’s crust, and calcium is the fifth.
Of the elements in this Group only magnesium is produced 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 hydroxide is then converted to the chloride, which is electrolysed in a Downs cell to extract magnesium metal.
Physical Properties
The metals of Group 2 are harder and denser than sodium and potassium, and have higher melting points. These properties are due largely to the presence of two valence electrons on each atom, which leads to stronger metallic bonding than occurs in Group 1.
Three of these elements give characteristic colours when heated in a flame:

Mg brilliant white   Ca brick-red   Sr crimson   Ba apple green
Atomic and ionic radii increase smoothly down the Group. The ionic radii are all much smaller than the corresponding atomic radii. This is because the atom contains two electrons in an s level relatively far from the nucleus, and it is these electrons which are removed to form the ion. Remaining electrons are thus in levels closer to the nucleus, and in addition the increased effective nuclear charge attracts the electrons towards the nucleus and decreases the size of the ion.
Chemical Properties
The chemical properties of Group 2 elements are dominated by the strong reducing power of the metals. The elements become increasingly electropositive on descending the Group.
Once started, the reactions with oxygen and chlorine are vigorous:
2Mg(s) + O2(g) ® 2MgO(s)
Ca(s) + Cl2(g) ® CaCl2(s)
All the metals except beryllium form oxides in air at room temperature which dulls the surface of the metal. Barium is so reactive it is stored under oil.
All the metals except beryllium reduce water and dilute acids to hydrogen:
Mg(s) + 2H+(aq) ® Mg(aq) + H2(g)
Magnesium reacts only slowly with water unless the water is boiling, but calcium 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) ® CaH2(s)
The hot metals are also sufficiently strong reducing agents to reduce nitrogen gas and form nitrides:
3Mg(s) + N2(g) ® Mg3N2(s)
Magnesium can reduce, and burn in, carbon dioxide:
2Mg(s) + CO2(g) ® 2MgO(s) + C(s)
This means that magnesium fires cannot be extinguished using carbon dioxide fire extinguishers.
The oxides of alkaline earth metals have the general formula MO and are basic. They are normally prepared by heating the hydroxide or carbonate to release carbon dioxide gas. They have high lattice enthalpies and melting points. Peroxides, MO2, are known for all these elements except beryllium, as the Be2+ cation is too small to accommodate the peroxide anion.
Calcium, strontium and barium oxides react with water to form hydroxides:
CaO(s) + H2O(l) ® 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 lime water which is used to test for the acidic gas carbon dioxide.
The Group 2 halides are normally found in the hydrated form. They are all ionic except beryllium chloride. Anhydrous calcium chloride has such a strong affinity for water it is used as a drying agent.
Oxidation States and lonisation Energies
In all their compounds these metals have an oxidation number of +2 and, with few exceptions, their compounds are ionic. The reason for this can be seen by examination of the electron configuration, which always has two electrons in an outer quantum level. These electrons are relatively easy to remove, but removing the third electron is much more difficult, as it is close to the nucleus and in a filled quantum shell. This results in the formation of M2+. The ionisation energies reflect this electron arrangement. The first two ionisation energies are relatively low, and the third very much higher.
Industrial Information
Magnesium is the only Group 2 element used on a large scale. It is used in flares, tracer bullets and incendiary bombs as it burns with a brilliant white light. It is also alloyed with aluminium to produce a low-density, strong material used in aircraft. Magnesium oxide has such a high melting point it is used to line furnaces.
Further Information
For further information look up the individual elements.

  Atomic Number Relative Atomic Mass Melting Point/K Density/kg m-3
Be 4   9.012   1551   1847.7  
Mg 12   24.31   922   1738  
Ca 20   40.08   1112   1550  
Sr 38   87.62   1042   2540  
Ba 56   137.33   1002   3594  

Ionisation Energies/kJ mol-1

  1st 2nd 3rd
Be 899.4 1757.1 14848
Mg 737.7 1450.7 7732.6
Ca 589.7 1145 4910
Sr 549.5 1064.2 4210
Ba 502.8 965.1 3600

  Atomic Radius/nm Ionic Radius/nm (M2+) Standard Electrode Potentials/V
Be 0.113 0.034 -1.85
Mg 0.160 0.078 -2.36
Ca 0.197 0.106 -2.87
Sr 0.215 0.127 -2.89
Ba 0.217 0.143 -2.90


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