molecular solid

Examples of molecular solid in the following topics:

• Molecular Crystals

  • Molecules held together by van der Waals forces form molecular solids.
  • Whereas the characteristic melting point of metals and ionic solids is ~1000 °C, most molecular solids melt well below ~300 °C.
  • Molecular solids also have relatively low density and hardness.
  • The term "molecular solid" may refer not to a certain chemical composition, but to a specific form of a material.
  • Conductivity of molecular solids can be induced by "doping" fullerenes (e.g., C60).

• Crystal Structure: Closest Packing

  • This section considers how the packing of atoms within unit cells contributes to a crystalline solid's lattice structure.
  • The three dimensional structure of a solid crystalline material is established through the periodic patterning of the atoms that make up the crystal.
  • The arrangement of the atoms in a crystalline solid affects atomic coordination numbers, interatomic distances, and the types and strengths of bonding that occur within the solid.

• Three States of Matter

  • A solid's particles are packed closely together.
  • A solid can transform into a liquid through melting, and a liquid can transform into a solid through freezing.
  • It can also exist in equilibrium with a liquid (or solid), in which case the gas pressure equals the vapor pressure of the liquid (or solid).
  • What does a phase change look like at the molecular level?
  • This video takes a look at the molecular structure of solids, liquids, and gases and examines how the kinetic energy of the particles changes.

• The Kinetic Molecular Theory of Matter

  • The kinetic molecular theory of matter explains how matter can change among the phases of solid, liquid, and gas.
  • The kinetic molecular theory of matter offers a description of the microscopic properties of atoms (or molecules) and their interactions, leading to observable macroscopic properties (such as pressure, volume, temperature).
  • This in turn determines whether the substance exists in the solid, liquid, or gaseous state.
  • Molecules in the solid phase have the least amount of energy, while gas particles have the greatest amount of energy.
  • A description of what happens during the process of dissolution of a solid in a liquid and diffusion.

• Substances that Exist as Gases

  • Gas is one of the three classical states of matter (the others being liquid and solid).
  • Near absolute zero, a substance exists as a solid.
  • All particles have energy, but the energy varies depending on whether the substance is a solid, liquid, or gas; solid particles have the least amount of energy and gas particles the most.
  • Explore the structure of a gas at the molecular level.
  • The process of a solid converting to a liquid is known as "melting"; liquid to a gas is "vaporization"; and gas back to a solid is "deposition."

• Boiling & Melting Points

  • Experience shows that many compounds exist normally as liquids and solids; and that even low-density gases, such as hydrogen and helium, can be liquified at sufficiently low temperature and high pressure.
  • First there is molecular size.
  • Molecular shape is also important, as the second group of compounds illustrate.
  • Halogens also form polar bonds to carbon, but they also increase the molecular mass, making it difficult to distinguish among these factors.
  • The melting points of crystalline solids cannot be categorized in as simple a fashion as boiling points.

• Ionic Crystals

  • It is a good conductor of electricity when molten (melted state), but very poor in the solid state.
  • As in all ionic structures, there are no distinguishable "molecular" units that correspond to the NaCl simplest formula.
  • The exothermicity of such reactions results in the stability of ionic solids.
  • That is, energy is needed to break apart the ionic solid into its constituent elements.
  • The most energetically stable arrangement of solids made up of identical molecular units are generally those in which there is a minimum of empty space.

• Conductors

  • In solid-state physics, the band structure of a solid describes those ranges of energy, called energy bands, that an electron within the solid may have ("allowed bands") and ranges of energy called band gaps ("forbidden bands"), which it may not have.
  • It successfully uses a material's band structure to explain many physical properties of solids.
  • Bands may also be viewed as the large-scale limit of molecular orbital theory.
  • If several atoms are brought together into a molecule, their atomic orbitals split into separate molecular orbitals, each with a different energy.
  • This produces a number of molecular orbitals proportional to the number of valence electrons.

• Crystalline Solids

  • This behavior is shown in the diagram below, with the green segment representing the solid phase, light blue the liquid, and red the temperature invariant liquid/solid equilibrium.
  • Below the temperature of the isothermal line ced, the mixture is entirely solid, consisting of a conglomerate of solid A and solid B.
  • An interesting but less common mixed system involves molecular components that form a tight complex or molecular compound, capable of existing as a discrete species in equilibrium with a liquid of the same composition.
  • An example of such a system is shown below, the molecular compound being represented as A:B or C.
  • It has a rigid flat molecular structure, and in dilute solution has a light yellow color.

• Molecular, Ionic, and Complete Ionic Equations

  • Precipitation reactions can be written as molecular, ionic, or complete ionic equations.
  • In the molecular equation, electrolytes are written as salts followed by (aq) to indicate that the electrolytes are completely dissociated into their constituent ions; the (aq) designation indicates that the ions are in aqueous solution.
  • On the right hand side of the equation, the precipitant (AgCl) is written in its full formula and designated as a solid, since this is the precipitate that is formed in the reaction.
  • Recognize whether a chemical equation is written in molecular, ionic, or complete ionic form.