Periodic Trends in Bonding and Structure

• What is a semi-metal/metalloid?

  Elements near to the metal/non-metal divide (e.g. boron, silicon, germanium, arsenic, antimony)

  They can show properties in between metals and non-metals
• Metallic bonding, structures and properties?


  At room temp - all metals except mercury are solids

  Some are strong and hard (e.g. tungsten)

  Some are soft (e.g. lead)

  Some are light (e.g. aluminium)

  Some are heavy (e.g. osmium = 2x as dense as lead)

  One constant property= ability to conduct electricity

  remarkable property for a solid, as charge must be able to move within a rigid structure for conduction to take place
• Metallic Bonding?

  In a solid metal structure - each atom has donated its negative outer-shell electrons to a shared pool of electrons - delocalised (spread out) throughout whole structure

  Positive ions (cations) left behind consist of the nucleus and the inner electron shells of the metal atoms

  The cations are fixed in place - maintaining structure and shape of the metal

  Delocalised electrons are mobile and are able to move throughout the structure
• What is metallic bonding?

  The strong electrostatic attraction between cations (positive ions) and delocalised electrons
• What is a giant metallic lattice?

  In a metal structure, billions of metal atoms are held together by metallic bonding
• Properties of metals?

  Strong metallic bonds
• Properties of metals?

  Strong metallic bonds - attraction between positive ions and delocalised electrons

  high electrical conductivity

  high melting and boiling points
• Electrical conductivity?

  Metals conduct electricity in solid and liquid states. When a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge
• Melting and Boiling points?

  Most metals have high melting and boiling points. Melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice.

  For most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons. This strong attraction results in most metals having high melting and boiling points.
• Solubility?

  Metals do not dissolve. It might be expected that there would be some interaction between polar solvents and the charges in a metallic lattice, as with ionic compounds, but any interactions would lead to a reaction, rather than dissolving (e.g. sodium and water)
• Simple molecular structures?

  Many non-metallic elements exist as simple covalently bonded molecules. In the solid state, these molecules form a simple molecular lattice structure held together by weak intermolecular forces. These structures therefore have low melting and boiling points.
• Giant covalent structures?

  Non-metals boron, carbon and silicon have very different lattice structures. Instead of small molecules and intermolecular forces, many billions of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice.
• Carbon and silicon?

  Group 4 and so their atoms have 4 electrons in the outer shells. Carbon (in its diamond form) and silicon use these four electrons to form covalent bonds to other carbon or silicon atoms. The result is a tetrahedral structure.

  Bond angles 109.5 by electron -pair repulsion
• Properties of giant covalent structures?

  Properties are dominated by the strong covalent bonds, which make for very stable structures that are very difficult to break down
• Melting points and boiling points (GCS)?

  High melting and boiling points because covalent bonds are strong and high temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds.
• Solubility (GCS)?

  Insoluble in almost all solvents because the covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents
• Electrical conductivity (GCS)?

  Non-conductors - only exceptions are graphene and graphite which are forms of carbon

  In carbon (diamond) and silicon, all four outer-shell electrons are involved in covalent bonding, so non are available for conducting electricity.

  Carbon is special in forming several structures in which one of the electrons is available for conductivity e.g. graphene and graphite are able to conduct
• Graphene and Graphite?

  Apart form diamond, carbon forms giant covalent structures based on planar hexagonal layers

  Only 3 of outer 4 electrons used in covalent bonding

  Remaining electron is released to a pool of delocalised electrons shared by all atoms in the structure.

  Structures of carbon containing planar hexagonal layers are therefore good electrical conductors.

  Graphene and Graphite are both giant covalent structures of carbon based on planar hexagonal layers with bond angles 120 by elctron-pair repulsion.
• Graphene?

  Single layer of graphite

  Composed of hexagonally arranged carbon atoms, linked by strong covalent bonds

  Same electrical conductivity as copper, and is the thinnest and strongest material ever made
• Graphite?

  Composed of parallel layers of hexagonally arranged carbon atoms, like a stack of graphene layers.

  Layers bonded by weak london forces.

  Bonding in hexagonal layers only uses three of carbon's 4 outer-shell electrons. Spare electron is delocalised between the layers, so electrically can be conducted as in metals.
• Across period 2 and 3

  1. Melting point increases from group 1 to group 4

  2. Sharp decrease in melting point between group 4 and 5

  3. Melting points are comparatively low from group 5 to group 0
• Sharp decrease in melting point between group 4 and 5


  Marks a change from giant to simple molecular structures
• Giant structures have strong forces to overcome

  Have high melting points
• Simple molecular structures have weak forces to overcome

  Have much lower melting points