In this worksheet, we will practice comparing the properties of isostructural network solids, which exhibit different atoms in the same arrangement.
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 less conductive because the boron p orbitals are empty, resulting in fewer delocalised electrons.
- BBoron nitride is more conductive because the empty p orbitals of boron allow for increased electron delocalisation.
- 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 neighbouring atoms, resulting in reduced electron delocalisation.
- EBoron nitride and graphite are isoelectronic so have similar electrical conductivities.
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 and diamond are similarly hard, as there is only a small difference between the energies of silicon-silicon and carbon-carbon bonds.
- BSilicon is harder than diamond, as bonds between silicon atoms are strengthened by the overlap of d orbitals.
- CSilicon is harder than diamond, as bonds between the smaller carbon atoms are weakened by repulsion between electron lone pairs.
- DSilicon is less hard than diamond, as the larger silicon atoms form weaker bonds that are easier to break.
- ESilicon is less hard than diamond, as silicon atoms interact via metallic bonds, which are weaker than non-metallic covalent bonds.
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 and diamond are conductors: the materials consist of planes of sp2-hybridized atoms, with delocalized electrons in a conjugated system.
- BSilicon 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.
- CSilicon 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.
- DSilicon 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.
- ESilicon and diamond are both semiconductors, but the atoms in diamond are smaller and exhibit stronger orbital overlap, producing a narrower band gap.
Silicon dioxide melts at approximately , while silicon tetrafluoride melts at . What is the main reason for the large difference between the melting points of the two materials?
- ABoth and form 3D covalent networks, but bonds are weaker than bonds so less energy is needed to break up the lattice.
- BWeak interactions between planar covalent networks can be broken more easily than the covalent bonds in the 3D network of .
- CThe higher coordination number of atoms in results in steric crowding, destabilising the ordered lattice.
- DWeak interactions between molecules in the molecular crystal of can be broken more easily than the covalent bonds in the 3D network of .
- EBoth and form 3D covalent networks, but bonds are more ionic and less directional than bonds so the ordered structure is more easily lost.