Worksheet: Isostructural Network Solids

In this worksheet, we will practice comparing the properties of isostructural network solids, which exhibit different atoms in the same arrangement.

Q1:

Boron nitride contains alternating boron and nitrogen atoms in the same arrangement as the carbon atoms in graphite. How does the electrical conductivity of boron nitride compare with that of graphite?

  • ABoron nitride is more conductive because the empty p orbitals of boron allow for increased electron delocalization.
  • BBoron nitride and graphite are isoelectronic, so they have similar electrical conductivities.
  • CBoron nitride is more conductive because the partial charges of the boron and nitrogen atoms strengthen interactions between layers.
  • DBoron nitride is less conductive because electron density is shared less equally between neighboring atoms, resulting in reduced electron delocalization.
  • EBoron nitride is less conductive because the boron p orbitals are empty, resulting in fewer delocalized 𝜋 electrons.

Q2:

The crystal structure of elemental silicon at standard temperature and pressure is analogous to that of diamond. How does the hardness of silicon compare with that of diamond?

  • ASilicon is harder than diamond, as bonds between silicon atoms are strengthened by the overlap of d orbitals.
  • BSilicon is less hard than diamond, as the larger silicon atoms form weaker bonds that are easier to break.
  • CSilicon is less hard than diamond, as silicon atoms interact via metallic bonds, which are weaker than non-metallic covalent bonds.
  • DSilicon is harder than diamond, as bonds between the smaller carbon atoms are weakened by repulsion between electron lone pairs.
  • ESilicon and diamond are similarly hard, as there is only a small difference between the energies of silicon-silicon and carbon-carbon bonds.

Q3:

The crystal structure of elemental silicon at standard temperature and pressure is analogous to that of diamond. How does the electrical conductivity of silicon compare with that of diamond?

  • ASilicon is a semiconductor and diamond is an insulator: carbon atoms interact only via sp3 orbitals, but silicon atoms also exhibit overlap between unoccupied d orbitals, producing a narrower band gap.
  • BSilicon and diamond are conductors: the materials consist of planes of sp2-hybridized atoms, with delocalized electrons in a conjugated 𝜋 system.
  • CSilicon is a semiconductor and diamond is an insulator: the bonds between silicon atoms are weaker, so the energy difference between occupied and unoccupied orbitals is smaller, producing a narrower band gap.
  • DSilicon and diamond are both semiconductors, but the atoms in diamond are smaller and exhibit stronger orbital overlap, producing a narrower band gap.
  • ESilicon and diamond are insulators: the bonds between atoms are similarly strong, so the energy difference between occupied and unoccupied orbitals is large in both materials, producing wide band gaps.

Q4:

Silicon dioxide melts at approximately 1,700C, while silicon tetrafluoride melts at 90C. What is the main reason for the large difference between the melting points of the two materials?

  • ABoth SiO2 and SiF4 form 3D covalent networks, but SiF bonds are more ionic and less directional than SiO bonds, so the ordered structure is more easily lost.
  • BWeak interactions between molecules in the molecular crystal of SiF4 can be broken more easily than the covalent bonds in the 3D network of SiO2.
  • CWeak interactions between planar covalent SiF4 networks can be broken more easily than the covalent bonds in the 3D network of SiO2.
  • DBoth SiO2 and SiF4 form 3D covalent networks, but SiF bonds are weaker than SiO bonds, so less energy is needed to break up the lattice.
  • EThe higher coordination number of Si atoms in SiF4 results in steric crowding, destabilizing the ordered lattice.

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