Covalent Network Solids (College Board AP® Chemistry): Revision Note

Diamond

• Covalent network solids consist of atoms held together in large networks by covalent bonds

• Covalent bonds are much stronger than intermolecular forces

  • This means that covalent solids are much harder and have higher melting points than molecular solids

• Examples of covalent network solids include:

  • Silicon

  • Germanium

  • Silicon dioxide (quartz)

  • Silicon carbide

• Covalent network solids form only from nonmetals and metalloids, either as elements (diamond, graphite, silicon) or as binary compounds (silicon dioxide, silicon carbide)

• Two familiar examples of covalent network solids are diamond and graphite

  • These are allotropes of carbon

• In diamond, each carbon atom is surrounded by a tetrahedral arrangement of other carbon atoms to form a giant three-dimensional network

  • These carbon atoms are sp³ hybridized and held together by strong carbon-carbon single covalent bonds

Structure of diamond

Diagram showing the tetrahedral structure of diamond

• Industrial grade diamonds are used as cutting tools because they are very hard

  • This is due to the number, strength and directionality of the covalent bonds

• Diamond also has a very high melting point because many strong covalent bonds must be broken to melt it

  • Its rigid, strongly bonded network also makes diamond a good conductor of heat

  • However due to a lack of mobile valence electrons, diamond is a poor conductor of electricity

Graphite

• In graphite, carbon atoms are arranged in layers of six-membered (hexagonal) rings where each carbon atom forms covalent bonds to three other carbon atoms

  • Hence, the carbon atoms are sp² hybridized and have one unhybridized 2p orbital

• Graphite is a good conductor of electricity, unlike diamond

  • The unhybridized 2p orbitals on neighboring carbon atoms overlap to form a delocalized π system across each layer, and these delocalized electrons can move along the layer

• Graphite is soft and slippery, so it is used as a lubricant

  • This is because the layers of hexagonal carbon rings are held together only by weak London dispersion forces, so adjacent layers can slide past each other relatively easily

  • These weak forces mean that the layers can slide past one another, an advantage in lubrication because the sliding layers allow for movement

• The enormous differences in physical properties of graphite and diamond — both of which are pure carbon — arise from differences in their three-dimensional structure and bonding

Structure of graphite

Diagram showing the bonding and structure of graphite

Silicon dioxide

• Silicon dioxide, SiO₂, is also known as quartz or sand

• It is another example of a naturally occurring covalent network solid with a similar structure to diamond

  • Like diamond, it has a tetrahedral structure

  • However, each silicon atom forms covalent bonds with four oxygen atoms while each oxygen atom forms covalent bonds with two silicon atoms

• The strong covalent bonds in silicon dioxide are responsible for its hardness and high melting point

  • The high melting point is due to the large amount of energy required to break a large number of strong covalent bonds in the solid

  • It also explains its hardness and use as an abrasive and in the manufacture of glass

• Like diamond, silicon dioxide is unable to conduct electricity

  • This is because all its valence electrons are held in localized covalent bonds or lone pairs, so there are no mobile charge carriers

Structure of silicon dioxide

Diagram showing the tetrahedral structure of silicon dioxide. The red atoms are oxygen and the blue atoms are silicon

Silicon carbide

• Silicon carbide, SiC, is another example of a covalent network solid consisting of covalently bonded silicon and carbon atoms

• Silicon carbide has a tetrahedral crystalline structure consisting of four carbon atoms covalently bonded to a single silicon atom at the center

• Unlike other examples of covalent network solids which are naturally occurring, silicon carbide is mostly made synthetically and exists naturally only in rare forms

Properties of silicon carbide

• Pure silicon carbide behaves as an insulator

  • This is because there are no free electrons which can act as mobile charge carriers

  • But, it can exhibit the electrical properties of a semiconductor when impurities are added

• Silicon carbide is very hard with a hardness close to that of diamond

  • Like in diamond, the hardness of silicon carbide is derived from the tetrahedral structure of silicon and carbon atoms which are held together by strong covalent bonds

  • This makes it useful for cutting tools, bearings and mechanical seals

• Silicon carbide is resistant to high temperature

  • Due to the strong silicon-carbon covalent bonds, silicon carbide expands very little when heated and resists high temperatures

  • This means that it is used in the manufacture of fire bricks and other heat-resistant materials

Structure of silicon carbide

Diagram showing the compact crystal structure of silicon carbide. The black atoms are carbon and the blue atoms are silicon