In the solid state the compounds of the alkaline metals generally form ionic lattices. In fact (except for beryllium and to a lesser extend magnesium) the lattice parameters calculated from the ionic radii of Group 2 metals are within 1% of the experimentally determined values, indicating the highly ionic character.
Due to the increased ionic charge the cations of the alkaline earths are less than that of their alkali metal neighbors. Thus, the ionic radius of Ca2+ (0.99 Å) is less than that of K+ (1.33 Å), but more similar to Na+ (0.97 Å). This diagonal relationship is seen for the other metals in Group 2.
Oxides
Combustion of the Group 2 metals gives the monoxide, MO. In the case of SrO and BaO further reaction occurs by the absorption of oxygen under pressure to give the peroxides, MO2. The peroxide and superoxides are not stable for the lighter homologs because the smaller M2+ ions are more polarizing and cause the peroxide and superoxides to decompose to the monoxide. Calcium peroxide can be prepared by the reaction of the hydroxide with hydrogen peroxide.
As typical for the Group 2 oxides, calcium monoxide is basic and reacts with water to give the hydroxide, Equation 2. In fact, thin films of the Group 2 oxides of calcium, barium and strontium will readily absorb water and form the hydroxides.
Carbonate
Calcium carbonate (CaCO3) is a common substance found in rock in all parts of the world, and is the main component of shells of marine organisms, snails, pearls, and eggshells. Calcium carbonate is usually the principal cause of hard water. Most calcium carbonate used is extracted by quarrying or mining. However, pure calcium carbonate (e.g., for pharmaceutical use) can be produced from a pure quarried source (usually marble) or manufactured by the sequential reaction involving the thermal decomposition of the carbonate to the monoxide, Equation 3, followed by the reaction with water to give the hydroxide, Equation 2, and finally, the reaction with carbon dioxide to reform the carbonate, Equation 4.
Calcium carbonate crystallizes as a variety of mineral forms.
- Aragonite
- Calcite (Figure 1)
- Vaterite
- Chalk
- Travertine
- Limestone
- Marble
 |
Calcium carbonate is one of the most widely used mineral materials, the following represents a list of some of the main applications.
- Construction industry as a building material (marble) or as an ingredient of cement.
- Purification of iron from iron ore in a blast furnace.
- A drilling fluids in the oil industry.
- Filler material for latex gloves.
- Filler (extender) in paints.
- Filler in plastics.
- Babies' diapers.
- DIY adhesives, sealants, and decorating fillers.
- Whiting in ceramics/glazing application.
- The filler in glossy paper.
- The production of blackboard chalk (CaSO4).
- An abrasive in household cleaning products.
- Dietary calcium supplement.
- Inert filler for tablets and pharmaceuticals.
- Toothpaste.
Halides
Calcium chloride, bromide, and iodide are all ionic, water-soluble salts. In contrast, due to its high lattice energy for the small fluoride ion, CaF2 is only slightly soluble (Table 1). The fluorite structure typified for CaF2 (Figure 2) is found for most of the MX2 ionic solids.
| Compound | Solubility @ 20 °C (g/100 mL) | Solubility @ 100 °C (g/100 mL) |
| CaF2 | 0.0016 | 0.0017 |
| CaCl2 | 74.5 | 159 |
| CaBr2 | 142 | 312 |
| CaI2 | 209 | 426 |
 |
