Structure & Physical Properties (OCR AS Chemistry)

• Metal atoms are tightly packed together in lattice structures
• When the metal atoms are in lattice structures, the electrons in their outer shells are free to move throughout the structure
• The free-moving electrons are called delocalised electrons and they are not bound to their atom
• When the electrons are delocalised, the metal atoms become positively charged ions
• The positive charges repel each other and keep the neatly arranged lattice in place
• There are very strong forces between the positive metal centres and the ‘sea’ of delocalised electrons

Covalent lattices

• Covalent bonds are bonds between nonmetals where there is a shared pair of electrons between the atoms
• In some cases, it is not possible to satisfy the bonding capacity of a substance in the form of a molecule
  • The bonds between atoms continue indefinitely, and a large lattice is formed
  • There are no individual molecules and covalent bonding exists between all adjacent atoms
• Such substances are called giant covalent substances
• The most important examples are the carbon allotropes graphite, diamond and graphene as well as silicon(IV) oxide

Diamond

• Diamond is a giant covalent lattice (or macromolecule) of carbon atoms
• Each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of 109.5^o
• The result is a giant lattice structure with strong bonds in all directions
• Diamond is the hardest substance known
  • For this reason, it is used in drills and glass-cutting tools

The structure of diamond

Graphite

• In graphite, each carbon atom is bonded to three others in a layered structure
• The layers are made of hexagons with a bond angle of 120^o
• The spare electrons are delocalised and occupy the space between the layers
• All atoms in the same layer are held together by strong covalent bonds
• However, the layers are held together by weak intermolecular forces
  • These weak intermolecular forces allow the layers to slide over each other

The structure of graphite

Graphene

• Some substances contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers.
  • Graphene is an example
• Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons
• Graphene is one million times thinner than paper; so thin that it is actually considered two dimensional

The structure of graphene

Silicon(IV) oxide

• Silicon(IV) oxide is also known as silicon dioxide, but you will be more familiar with it as the white stuff on beaches!
• Silicon(IV) oxide adopts the same structure as diamond -  a giant covalent lattice / macromolecular structure made of tetrahedral units all bonded by strong covalent bonds
• Each silicon is shared by four oxygens and each oxygen is shared by two silicons
• This gives an empirical formula of SiO₂

The structure of silicon dioxide

• Different types of structure and bonding have different effects on the physical properties of substances such as their melting and boiling points, electrical conductivity and solubility

Properties of metallic substances

• Due to the delocalised ‘sea’ of electrons, metallic structures have some characteristic properties:
• High melting and boiling point: as a lot of energy is required to overcome the strong electrostatic forces of attraction between positive ions and the 'sea' of delocalised electrons
• Solubility: metals do not dissolve. There is some interaction between polar solvents and charges in the metallic lattice but these lead to reactions, rather than dissolving e.g. sodium and water
• Electrical conductivity: conduct electricity in both solid and liquid states. This is due to the delocalised electrons which are free to move / carry charge around the structure

Properties of giant covalent substances

• Giant covalent lattices have very high melting and boiling points
  • These compounds have a large number of covalent bonds linking the whole structure
  • A lot of energy is required to break the lattice

• The compounds can be hard or soft
  • Graphite is soft as the intermolecular forces between the carbon layers are weak
  • Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds
  • Graphene is strong, flexible and transparent, which makes it potentially a very useful material

• Most compounds are insoluble with water
• Most covalent substances do not conduct electricity
  • For example, diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom is involved in a covalent bond , so there are no free electrons available
• There are some covalent substances that are exceptions because they do conduct electricity 
  • Graphite has delocalised electrons between the carbon layers, which can move along the layers when a voltage is applied
  • Graphene is an excellent conductor of electricity due to the delocalised electrons

Periodic trend in melting points

• Across Period 2 and Period 3,
  • Melting point increases from Group 1 to Group 4 (14) 
    • Groups 1 to 3 (13) have metallic bonding which increases in strength due to increased forces of attraction between more electrons in the outer shell that are released to the sea of electrons and a smaller positive ion
    • Group 4 (14) has a giant covalent structure with many strong covalent bonds requiring a lot of energy to overcome
  • A sharp decrease in melting point from Group 4 (14) to Group 5 (15)
    • Groups 5 (15) to 0 (18) have simple molecular structures with weak London forces between molecules requiring little energy to overcome