The crystal lattice of ionic solids consists of monoatomic or polyatomic ions held together by intense electrostatic interactions of the Coulomb type. These, by their nature, are not directional and therefore the ions of opposite charge are attracted independently of their spatial location: therefore, no single molecular units are recognizable.
In ionic solids, cations and anions are located on the nodes of the crystal lattice; these ions are joined together by intense Coulombian-type forces, therefore, the lattice energy of these crystals is very high and so is their melting point. All halides of alkali metals, zinc sulfide, calcium fluoride, lithium, calcium, titanium oxides, etc. form ionic solids.
Because of the strong attraction between opposite charges, it takes a lot of energy to overcome ionic bonds (the attractions between full charges are much larger than those between the partial charges in polar molecular compounds). This means that ionic compounds have very high melting points, often between 300÷1000 °C. The common characteristics of these substances are:
- high melting and boiling point: due to the strong electrostatic attractions between the ions of opposite sign, high energy is needed to separate them. The difference in behavior between the various ionic solids depends on the number of charges present in the ion. In fact, magnesium oxide has a higher melting point and boiling point than sodium chloride (the Mg2+ and O2- ions having charge +2 and -2 respectively, attract each other more than the ions Na+ and Cl– which have charge +1 and -1 respectively). Another factor influencing the melting and boiling temperature is given by the size of the ions: if the ions are very small they will be closer and the electrostatic attraction is greater. For example, rubidium iodide has lower melting and boiling temperatures than sodium chloride since both the Rb+ ion and the I– ion have significantly larger dimensions than Na+ and Cl– therefore the attractions between rubidium ion and iodide ion are low and less energy is required to separate them.
- transparency to visible radiation
- hardness and at the same time fragility given their easy flaking (they shatter rather than bend),
- solubility in polar solvents: many ionic solids are insoluble in almost all apolar or low-polar solvents due to the high reticular energy. When the energy released by the solvation of the ions exceeds the lattice energy these compounds can be soluble: this usually occurs in water. Most of the ionic solids are soluble in water and the positive ions are attracted by the solitary electronic doublet present on the oxygen with the formation of a dative bond. Furthermore, water molecules can form hydrogen bonds with negative ions. Many simple compounds formed by the reaction of a metallic element with a nonmetallic element are ionic;
- furthermore, since there are no free electrical charges, their conductivity is low; however, they do conduct when molten or dissolved because undergo electrolysis and their ions are free to move (the positive ion migrates towards the negative electrode while the negative ion migrates towards the positive electrode).
The ionic solids have a more complex structure than that of metals as they are made up of ions that have different ionic rays and also if on the one hand there is an attraction between ions of opposite charge on the other there is a repulsion between ions having the same charge. Furthermore, polyatomic ions such as, for example, the nitrate ion and the carbonate ion, cannot be assimilated to spheres.
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