Which Of The Compounds Below Is An Example Of A Network Solid?

To understand the correlation between bonding and the properties of solids. To classify solids as ionic, molecular, covalent (network), or metallic, where the general order of increasing strength of interactions.

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Crystalline solids fall into one of four categories. All four categories involve packing discrete molecules or atoms into a lattice or repeating array, though network solids are a special case. The categories are distinguished by the nature of the interactions holding the discrete molecules or atoms together. Based on the nature of the forces that hold the component atoms, molecules, or ions together, solids may be formally classified as ionic, molecular, covalent (network), or metallic. The variation in the relative strengths of these four types of interactions correlates nicely with their wide variation in properties.

Table (PageIndex{1}): Solids may be formally classified as ionic, molecular, covalent (network), or metallic

Type of SolidInteractionPropertiesExamples
Ionic Ionic High Melting Point, Brittle, Hard NaCl, MgO
Molecular Hydrogen Bonding, Dipole-Dipole, London Dispersion Low Melting Point, Nonconducting H2, CO2
Metallic Metallic Bonding Variable Hardness and Melting Point (depending upon strength of metallic bonding), Conducting Fe, Mg
Network Covalent Bonding High Melting Point, Hard, Nonconducting C (diamond), SiO2 (quartz)

In ionic and molecular solids, there are no brianowens.tvical bonds between the molecules, atoms, or ions. The solid consists of discrete brianowens.tvical species held together by intermolecular forces that are electrostatic or Coulombic in nature. This behavior is most obvious for an ionic solid such as (NaCl), where the positively charged Na+ ions are attracted to the negatively charged (Cl^-) ions. Even in the absence of ions, however, electrostatic forces are operational. For polar molecules such as (CH_2Cl_2), the positively charged region of one molecular is attracted to the negatively charged region of another molecule (dipole-dipole interactions). For a nonpolar molecule such as (CO_2), which has no permanent dipole moment, the random motion of electrons gives rise to temporary polarity (a temporary dipole moment). Electrostatic attractions between two temporarily polarized molecules are called London Dispersion Forces.

Hydrogen bonding is a term describing an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation. (See the IUPAC Provisional Recommendation on the definition of a hydrogen bond.) Dots are employed to indicate the presence of a hydrogen bond: X–H•••Y. The attractive interaction in a hydrogen bond typically has a strong electrostatic contribution, but dispersion forces and weak covalent bonding are also present.

In metallic solids and network solids, however, brianowens.tvical bonds hold the individual brianowens.tvical subunits together. The crystal is essential a single, macroscopic molecule with continuous brianowens.tvical bonding throughout the entire structure. In metallic solids, the valence electrons are no longer exclusively associated with a single atom. Instead these electrons exist in molecular orbitals that are delocalized over many atoms, producing an electronic band structure. The metallic crystal essentially consists of a set of metal cations in a sea of electrons. This type of brianowens.tvical bonding is called metallic bonding.

Ionic Solids

You learned previously that an ionic solid consists of positively and negatively charged ions held together by electrostatic forces. The strength of the attractive forces depends on the charge and size of the ions that compose the lattice and determines many of the physical properties of the crystal.

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The lattice energy (i.e., the energy required to separate 1 mol of a crystalline ionic solid into its component ions in the gas phase) is directly proportional to the product of the ionic charges and inversely proportional to the sum of the radii of the ions. For example, NaF and CaO both crystallize in the face-centered cubic (fcc) sodium chloride structure, and the sizes of their component ions are about the same: Na+ (102 pm) versus Ca2+ (100 pm), and F− (133 pm) versus O2− (140 pm). Because of the higher charge on the ions in CaO, however, the lattice energy of CaO is almost four times greater than that of NaF (3401 kJ/mol versus 923 kJ/mol). The forces that hold Ca and O together in CaO are much stronger than those that hold Na and F together in NaF, so the heat of fusion of CaO is almost twice that of NaF (59 kJ/mol versus 33.4 kJ/mol), and the melting point of CaO is 2927°C versus 996°C for NaF. In both cases, however, the values are large; that is, simple ionic compounds have high melting points and are relatively hard (and brittle) solids.

Molecular Solids

Molecular solids consist of atoms or molecules held to each other by dipole–dipole interactions, London dispersion forces, or hydrogen bonds, or any combination of these. The arrangement of the molecules in solid benzene is as follows:

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