Group II

Group II Elements: Alkaline Earth Metals

1. General Characteristics and Trends

The Group II elements, also known as alkaline earth metals, share several common properties:

• They are all metals.
• They are good reducing agents.
• They typically form ionic compounds.
• Their oxides and hydroxides are generally basic.
• They liberate hydrogen when reacted with acids.

Summary of Properties (Table Data from Page 2):

As you descend Group II from Beryllium (Be) to Radium (Ra):

• Electronic Configuration: All end in ns2.
• Ionization Energy: Decreases (e.g., Be: 899 KJmol-1 to Ba: 500 KJmol-1), indicating easier removal of electrons.
• Ionic Radius: Increases (e.g., Be2+: 30 pm to Ba2+: 134 pm), due to added electron shells.
• Melting and Boiling Points: Do not vary regularly, primarily because the metals adopt different crystal structures. For example, Beryllium (M.Pt 1283°C, B.Pt 2477°C) is higher than Magnesium (M.Pt 650°C, B.Pt 1117°C), but then Calcium (M.Pt 850°C, B.Pt 1493°C) is higher again before generally decreasing.
• Radium (Ra) is radioactive.

2. Anomalous Behavior of Beryllium (Be)

Similar to Lithium in Group I, Beryllium is the first member of Group II and exhibits properties that are exceptions to the common pattern:

Reasons for Anomalous Behavior:

• Its electrons are not well shielded from its nucleus.
• It has a small ionic radius, resulting in a very dense center of positive charge and high polarizing power.

Consequences of Anomalous Behavior:

• Its compounds are more covalent than ionic.
• Its oxide (BeO) is amphoteric, not basic (unlike other Group II oxides).
• It forms complexes more readily than other members of the group.
• It dissolves in alkali, unlike other Group II metals.

3. Preparation of Group II Metals

Like Group I metals, Group II metals are very reactive and are prepared by electrolysis. For instance, magnesium is prepared by the electrolysis of fused magnesium chloride (MgCl2).

4. Formation of Group II Oxides

• Most Group II metals burn in air to form oxides (MO):
  2M + O2 → 2MO
  Example: 2Mg + O2 → 2MgO
• Beryllium is unreactive at room temperature. It can only react above 600°C, but its powder form is much more reactive.
• Group II metals also form peroxides down the group.
• Barium peroxide (BaO2) is formed with ease and increases in stability down the group:
  Ba(s) + O2(g) → BaO2(s)

5. Oxides and Hydroxides of Group II Elements

With the exception of beryllium oxide (BeO), which is insoluble in water, the oxides of Group II elements dissolve in water to give alkaline solutions.

• Example: CaO(s) + H2O(l) → Ca(OH)2(aq)
• Magnesium hydroxide (Mg(OH)2) is prepared by the reaction of magnesium chloride with sodium hydroxide:
  MgCl2(aq) + 2NaOH(aq) → Mg(OH)2(s) + 2NaCl(aq)
• Mg(OH)2 is weakly basic, which is why it's used as an antacid in indigestion tablets ('milk of magnesia').
• Be(OH)2 is amphoteric, while Mg, Ca, Sr, Ba hydroxides are basic.
• Basicity increases down the group.
• Ca(OH)2 and Ba(OH)2 are called lime water and baryta water, respectively, and are used to detect CO2.
• Ca(OH)2 and Ba(OH)2 are made by the reaction of their oxides with water.

6. Halides of Group II Elements

• The chlorides are the most common halides.
• The elements also form fluorides, bromides, and iodides.
• They are all soluble in water and easily form hydrates, e.g., CaCl2·6H2O.
• A single molecule of beryllium chloride (BeCl2) has a linear structure: Cl—Be—Cl
• However, the polymeric solid form of BeCl2 is a chain of linked molecules, indicating its covalent character.
• Ionic character in the chlorides increases down the group.
• Anhydrous calcium chloride (CaCl2) is deliquescent and is widely used as a drying agent.

7. Carbonates and Hydrogencarbonates

• Unlike Group I carbonates, Group II carbonates are slightly soluble in water.
• Solubility decreases down the group.
• On heating, they decompose to give the corresponding oxide and carbon dioxide:
  MgCO3(s) → MgO(s) + CO2(g)
• The ease of decomposition decreases down the group (i.e., carbonates become more thermally stable down the group).
• Unlike Group I metals which possess solid hydrogencarbonates, the hydrogencarbonates of Group II metals only exist in solution.

8. Sulphates of Group II Elements

• All Group II metals form sulphates (e.g., BeSO4, MgSO4, CaSO4, SrSO4, BaSO4).
• Solubility of their sulphates in water decreases down the group: BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4.
• BeSO4 and MgSO4 are soluble.
• CaSO4 is sparingly soluble.
• SrSO4 and BaSO4 are insoluble.
• The solubility trend is due to the high enthalpy of solvation of the smaller Be2+ and Mg2+ ions.
• Sulphates of Be, Mg, and Ca are found as hydrated crystals, e.g., BeSO4·4H2O, MgSO4·7H2O, CaSO4·2H2O.
• The stability of Group II sulphates depends on basicity; the more basic the sulphate, the more stable. Thermal decomposition temperatures increase down the group (e.g., BeSO4 decomposes at 500°C, MgSO4 at 895°C, CaSO4 at 1149°C, SrSO4 at 1374°C).

9. Some Uses of Group II Sulphates

• MgSO4·7H2O (Epsom salt) is used as a mild laxative.
• CaSO4·2H2O (gypsum) and CaSO4·H2O (plaster of Paris) are useful in the treatment of broken bones.
• BaSO4 is used in diagnosing stomach ulcers because it is insoluble in water and opaque to X-rays. It is used to provide a shadow of the stomach on an X-ray picture (known as a "barium meal").

10. Uses of Some Group II Elements

• Beryllium is used in nuclear reactors as a moderator.
• Magnesium is used in metal alloys for aircraft structures.
• Calcium is useful in the body for healthy bones.
• Strontium is useful in radiochemistry.
• Barium forms alloys with other metals such as lead and calcium. Its sulphate is used as "barium meal" in medicine.