1. 3.2.2 Group 2, the Alkaline Earth Metals

2. The physical properties of the Group 2 elements, magnesium to barium
Atomic radii
Going down the group from magnesium to barium the atomic radii increase. This is because each element has an extra filled electron shell compared with the element above it and so the outermost electrons become progressively further away from the nucleus.

3. First ionisation energy
Going down the group from magnesium to barium the first ionisation energy decreases as it takes less energy to remove one of the outermost electrons. This is because:
the atomic radius increases, increasing the distance between the outermost electrons and the nucleus
and
the outermost electrons become increasingly shielded from the positive charge of the nucleus. 

4. Melting points
Group 2 elements are metals with high melting points. Going down the group from calcium to barium the strength of the metallic bond decreases. This is because the metal ions become larger and so the electrons in the “sea” of delocalised electrons become further away from the positive nuclei. Thus the force of attraction between the positive nuclei and negative delocalised electrons decreases and less energy is required to overcome it.
Magnesium does not follow the trend. It has the lowest melting point as it has a different crystal structure to the other Group 2 elements.

5. Melting points
Magnesium does not follow the trend. It has the lowest melting point as it has a different crystal structure to the other Group 2 elements.

6. The chemical reactions of the Group 2 elements, magnesium to barium
Oxidation is loss of electrons. In all their reactions the Group 2 metals are oxidised. The metals go from oxidation state 0 to oxidation state +2. These are redox reactions.

7. Reaction with water
The metals get more reactive going down the group.
The general reaction is:
M (s) + 2H2O (l)  M(OH)2 (aq) + H2 (g)
Oxidation state:
(M) (H) (M) (H)

8. Mg(s) + H2O (g)  MgO(s) + H2(g)
Reaction with water
Magnesium reacts very slowly with cold water.
Magnesium reacts rapidly with steam to form an alkaline oxide and hydrogen:
Mg(s) + H2O (g)  MgO(s) + H2(g)
The other Group 2 metals (calcium, strontium and barium) react increasingly vigorously with cold water.

9. The solubilities of hydroxides and sulfates
Going down Group 2 the metal hydroxides become more soluble:
magnesium hydroxide is almost insoluble, calcium hydroxide is sparingly soluble, strontium hydroxide is more soluble and barium hydroxide is very soluble.
Going down Group 2 the metal sulfates become less soluble.

10. Uses of Group 2 metals and their compounds
Magnesium is used in the extraction of titanium. Titanium cannot be extracted from its oxide by reaction with carbon because the metal reacts with carbon to form titanium carbide which is a brittle compound.

11. Magnesium is used in the extraction of titanium.
Titanium oxide is first reacted with chlorine and carbon to form titanium chloride and carbon monoxide and then titanium chloride is reduced to titanium by reaction with magnesium.
TiO2(s) + 2C(s) + 2Cl2(g)  TiCl4(l) + 2CO(g)
TiCl4 (l) + 2Mg (s)  2MgCl2 (s) + Ti (s)

12. Magnesium hydroxide is milk of magnesia
Magnesium hydroxide is milk of magnesia. It is used as an antacid to neutralise excess stomach acid which causes heartburn and indigestion.

13. Calcium hydroxide is sometimes called slaked lime
Calcium hydroxide is sometimes called slaked lime. It is used to reduce soil acidity so that a wider range of crops can be grown and to provide calcium ions which are essential for plant growth.

14. Calcium carbonate and calcium oxide are used in the process of flue gas desulfurisation. Calcium oxide can be produced by heating calcium carbonate. As basic compounds they act to neutralise acidic sulphur dioxide.
CaO + SO2  CaSO3
CaCO3 + SO2  CaSO3 + CO2

15. The product of the reaction of calcium oxide with sulphur dioxide is calcium sulphite which can be oxidised and hydrated to produce hydrated calcium sulfate (gypsum). CaSO4·2H2O.
Gypsum is used to make plasterboard for the building industry.

16. Hydrated calcium sulfate (gypsum) CaSO4·2H2O
also occurs naturally and, in the right conditions, can form very large crystals.

17. CO32-(aq) + 2H+(aq)  H2O(l) +CO2(g)
Barium chloride solution is used to test for sulfate ions. It must be acidified first so that any contaminating carbonate ions can be removed from the test solution such that they do not interfere with the test (by producing a precipitate).
Removal of carbonate ions:
CO32-(aq) + 2H+(aq)  H2O(l) +CO2(g)

18. Ba2+(aq) + SO42-(aq)  BaSO4(s)
Barium chloride solution is used to test for sulfate ions.
Test for sulfate ions:
Ba2+(aq) + SO42-(aq)  BaSO4(s)
The barium sulfate is insoluble, so if sulfate ions are present a white precipitate will be formed when the two colourless solutions are mixed.

19. Barium sulfate is used as a contrast agent when taking X rays of soft tissues. It is taken by mouth as a barium meal or given rectally as a barium enema in order to outline the gastrointestinal tract. The heavy barium atom is very good at absorbing x-rays (it appears white on a processed x-ray).

20. Barium sulfate in X rays of soft tissues.
Due to the fact that barium sulfate is insoluble the procedure is safe, despite barium compounds being highly toxic. The barium sulfate does not get absorbed by the body.