The four types of solids are metallic, ionic, covalent-network, and molecular. Metallic solids are held together by collectively shared delocalized valence electrons. This type of bonding is responsible for the fact that most metals are strong without the characteristic of a brittle consistency. It also allows metals to conduct electricity. Metallic solids can have a wide range of melting points. Ionic solids are held together by the mutual attraction between cations and anions. NaCl, for example, is an ionic solid that is an extended network of ions held together by ion-ion interactions. Ionic solids are conductive in aqueous solution, but not in solid form, and they tend to have high melting points due to the difference in charge between the ions. Covalent-network solids are held together by an extended network of covalent bonds, which generally results in materials that are extremely hard. These solids are non-conductive, and have high boiling points. This type of bonding also accounts for the unique properties of semiconductors. Last but not least, molecular solids are held together by intermolecular forces including dispersion forces, dipole-dipole interactions, and hydrogen bonds. Molecular solids tend to be soft with low melting points due to the fact that these forces are relatively weak. Most molecular solids are not soluble in water, and they are not conductive.
In the lab, we used the properties of conductivity, solubility in both polar and non-polar solvents, and melting point to classify the solids. Based on our definitions, metallic solids will conduct, while covalent-network, molecular, and ionic solids will not. All solids out of these four will have high melting points except for molecular solids.
Our results were as follows:
CaCl2: Ionic Sand and Charcol: Network Covalent Citric Acid: Molecular Zinc: Metallic
I think that the closeness of the bonds
- What microscopic differences do you think account (i.e., as shown by particle diagrams) for the property differences observed?